Chemistry Notes

1

Scientific Units of Measurement (SI Units)

Length Meter m

Mass Kilogram (about the mass of a

standard paperclip)

kg

Time Second s

Temperature Kelvin K

Amount of substance Mole Mol

Volume Liters L or l

Energy Joules J

Uncertainty in scientific measurements- Significant Digits

There are 2 types of numbers

• Exact numbers (counted numbers or defined numbers such as 12in=1ft); no uncertainty; distinct

number of significant figures

• Inexact numbers (measured quantities)

• Measurements: attained with a measuring instrument; degree of uncertainty

Examples:

1. There are 37 chickens I counted – Exact number

2. I am 47.3 kg. I measured this on a scale – Inexact number

3. There are 12 inches in one foot – Exact number

How do I know what value to give inexact numbers?

1. Significant digits in inexact numbers

a. Numbers 1-9 are always significant (non-zero) numbers

b. Any zero between non-zeroes is always significant

c. Leading zeroes (place holders) are always NOT significant

i. Ex: 0.000304 – 3 significant figures

d. Tailing zeroes after decimal places are significant

i. Ex: 4.10 – 3 significant figures

e. Trailing zeroes without decimal places are not significant

i. Ex: 740 – 2 significant figures

ii. Ex: 740. – 3 significant figures

Weak Link Principle:

1. Based on the idea that a chain is only as strong as its weakest link

2. In calculations involving measured values, we can know the answer only as well as we know

the least well-known value in the calculation.

Math with significant figures

Addition or subtraction

1. Answer should be rounded to the same place as the number with the least number of significant

figures to the right of the decimal point → leftmost decimal place should be rounded

a. Ex: 114.12+3.137=117.257 → Rounded to 117.20

b.

Multiplication and division

1. Answer should have the same number of significant figures as the number with the least number

of significant figures.

a. 3.62 cm * 4.921e3 cm * 1.9 cm → Rounded to 3.4x10^4 cm3

Multistep Problems

1. Underline the last significant figure in each step of the calculation

2. Round only ONE time at the end of the problem

Chemistry Notes

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In-class problem: Consider 3 measured densities for the same sample of gold: 19.4 g/mL, 19.69 g/mL and

18.891 g/mL. What should be reported as the average density of gold, based on these measurements?

Work: (19.4+19.69+18.891)/3=19.327 g/mL=

Helpful Equation

Density=Mass/Volume

D=m/V

1 Angstrom = 1x10-10 m

Dimensional analysis problems:

1. Begin with a starting quantity (include units) usually found within the question

2. Multiply by one or more conversion factors to cancel unwanted units leaving only units required

in the solution

E.1-E.4

Quantification is the assignment of a number to some property of a substance or thing

SI units, their meanings and functions:

Quantity Unit symbol

Length Meter m

Mass Kilogram kg

Time Second s

Temperature Kelvin K

Amount of substance Mole mol

Electric current Ampere A

Luminous intensity Candela cd

Definitions of SI units:

Unit Measure of:

m The meter is a measure of length

kg The kilogram is a measure of mass

s The second is a measure of time

K The kelvin is a measure of temperature

Mass Mass is a quantity of the measurement of matter

Weight Weight is a measure of the gravitational pull on an

object

Temperature Temperature is the measurement of the average

kinetic energy of an object

Prefix Multipliers:

Prefix Symbol Multiplier Power of 10

Exa E 1,000,000,000,000,000,000 1018

Peta P 1,000,000,000,000,000 1015

Tera T 1,000,000,000,000 1012

Giga G 1,000,000,000 109

Mega M 1,000,000 106

Kilo k 1,000 103

Deci d 0.1 10-1

Centi c 0.01 10-2

Mili m 0.001 10-3

Micro µ 0.000001 10-6

Nano n 0.000000001 10-9

Pico p 0.000000000001 10-12

Femto f 0.000000000000001 10-15

Chemistry Notes

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Atto a 0.000000000000000001 10-18

Key equations and relationships:

Relationship between Kelvin (K) and Celsius temperature scales

K= ⁰C+273.15

Relationship between Celsius and Fahrenheit temperature scales

⁰C= (℉-32)/1.8

Relationship between density (d), mass (m), and volume (V)

d=m/V

E.5-E.9

• Density (d) is the ratio of something’s mass to that thing’s volume

• An Intensive Property is a property of a substance that is independent of the amount of substance

(such as: temperature, color, hardness, melting point, boiling point, pressure, molecular weight,

and density)

• An Extensive Property is a property of a substance that is dependent on the amount of substance

(such as: volume, mass, weight, and length)

• Energy is the capacity to do work

• Work is defined as the action of a force through a distance

• The total energy of an object is the sum of its kinetic energy, the energy associated with its

motion, and its potential energy, the energy associated with an object’s position or composition

• Thermal energy is the energy associated with the temperature of an object

Helpful Conversion Factors

1J=1kg m2/s2

1mol=NA=Avogadro’s number=6.022x1023

Joule (J)=unit of energy

Notes

• The total quantity of energy stays constant – law of conservation of energy (energy is neither

created nor destroyed)

• There is a tendency of systems with high potential energy to change in a way that lowers their

potential energy (Ex: gas tends to combust; rain tends to fall)

• The structure of a molecule; the way its protons and neutrons are arranged, determines the

potential energy of a molecule, which determines its properties

• Derived units are combinations of other units such as: g/cm3 or m/s or Nm/s, etc.

• Random error is error that has equal probability of being too high or too low

Density of Common substances at 20⁰C

Substance: Density (g/cm3)

Charcoal 0.57

Ethanol 0.789

Ice 0.917 (at 0⁰C)

Chemistry Notes

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Water 1.00 (at 4⁰C)

Sugar (sucrose) 1.58

Table salt (sodium chloride) 2.16

Glass 2.6

Aluminum 2.70

Titanium 4.51

Iron 7.86

Copper 8.96

Lead 11.4

Mercury 13.55

Gold 19.3

Platinum 21.4

Energy units:

An object of mass m, moving at velocity v, has a kinetic energy (KE) given by the equation:

KE= (mV2)/2

SI units for energy:

Mass: kg

Velocity: m/s

Energy: kg*m2/s2

Energy conversion factors:

1 calorie (cal)= 4.184 Joules

1 Calorie (Cal) or kilocalorie (kcal)= 1000 cal= 4184 J

1 kilowatt-hour (kWh)= 3.60x106 J

Quantifying changes in energy:

Energy can be transferred from one object or system to another such as: a weight being dropped or

gasoline combustion in a car

Illustration:

Falling weight (system) → Energy → Air, ground, etc. (surroundings)

System loses energy, change surroundings gain energy. Change in energy is positive

In energy is negative

• Chemical processes almost always involve chemical changes

• Processes in which a system loses energy are exothermic

• Processes in which a system gains energy are endothermic

• An exothermic process involves the transfer of energy from the system to the surroundings and

carries a positive sign (like a deposit into a checking account)

• An endothermic process involves the transfer of energy from the surroundings to the system and

carries a negative sign (like a withdrawal from a checking account)

Problem solving strategy for chemistry:

1. Identify the starting point (the GIVEN information)

2. Identify the ending point (what we must FIND)

3. Devise a way to arrive at the endpoint through what you already know or can look up (the

conceptual plan)

Chapter 1: Atoms

Chemistry Notes

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Chemistry- the science that seeks to understand the properties of matter by studying the structure of the

particles that compose it

Substance-a specific instance of matter (such as: air, water, or sand)

We can understand the particulate view of matter by classifying matter based on the particles that

compose it

Classification:

1 State of matter: depends on the relative positions of the particles and how strongly they interact

with one another (relative to temperature)

2 Composition of matter: depends on the types of particles something is composed of

1. States of matter:

Gas: particles are widely spaced, giving gas a fluid ability and making it compressible

Solid: molecules can only vibrate and stay closely packed to one another

Liquid: molecules are closely packed but can move past one another

2. Elements, compounds, and mixtures:

We can classify matter by its composition. We determine the types of particles in the matter and whether

there is: only one type, or more than one type. This is the difference between a pure substance and a

mixture.

Matter classification chart

Matter

One type of particle?

Pure substance

Mixture

Separable into simpler substances?

Uniform throughout?

Element

Compound Heterogeneous

Homogeneous

• A pure substance is made of only one type of particle (Ex: helium, water, sodium chloride)

• A mixture is a substance composed of two or more particles in proportion that can vary from one

sample to another (Ex: sweetened tea)

• An element is a substance that cannot be broken down into simpler substances

• A compound is a substance that is composed of two or more elements in fixed, definite

proportions (Ex: water)

• A heterogeneous mixture is a mixture in which the composition varies from one region of the

mixture to another (Ex: wet sand)

Chemistry Notes

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• A homogeneous mixture is a mixture which holds the same composition throughout. (Ex:

sweetened tea)

• A chemical reaction is a process in which one or more substances are converted into one or more

different substances.

• The law of definite proportions states: “when two elements (call them A and B) form two

different compounds, the masses of element B that combine with 1g of element A can be

expressed as a ratio of small whole numbers

Examples of matter classification:

Boba tea (HoM) H2 (el) concrete (HeM) oatmeal (HeM)

Air (HeM) Si (el) Gatorade (HoM) CH4 (c)

NaCl (c) water, H2O (c) NaCl auqueous (HoM)

Dalton’s atomic theory

1. Each element is composed of tiny, indestructible particles called atoms

2. All atoms of a given element have the same mass and other properties that distinguish them from

the atoms of other elements

3. Atoms combine in simple, whole-number ratios to form compounds

4. Atoms of one element cannot change into atoms of another element. In a chemical reaction,

atoms only change the way that they are bound together with other atoms.

The subatomic particles chart:

Particle Mass (kg) Mass (amu) charge

Proton 1.67x10-27 1 1+

Electron 9.11x10-31 0 1-

Neutron 1.67x10-27 1 0

An element is defined by the number of protons in the nucleus of one of its atoms

An isotope is an atom with a different number of neutrons in the nucleus than protons and can be

expressed in the following ways:

ZAXn+

X= identity of the atom

Z=atomic number= number of protons in the nucleus

A= atomic mass of the isotope

N+= charge of the atom (atoms with “-“have an excess of electrons)

Number of atoms also goes as a subscript at the bottom left in a chemical symbol

And

Xn+-A

X= identity of the atom

A= atomic mass of the isotope

N+= charge of the atom

Ions: losing and gaining electrons:

• Ions are charged particles that are a result of a chemical reaction in which an element has gained

or lost electrons. During chemical changes, atoms can lose or gain electrons and become charged

particles called ions

• The charge of an ion is indicated in the upper right corner of a chemical symbol. Ex: A Li+ ion

contains 3 protons and only two electrons. Lithium and fluorine occur in nature mostly as ions.

Chemistry Notes

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Positively charged ions such as Li+ are called cations and negatively charged ions such as F- are

called anions

• An average mass of each element is listed under each element in the periodic table and is called

the atomic mass. The atomic mass is represented by the average mass of all of the isotopes of an

element, weighed according to the natural abundance of each isotope.

• This process is derived using the equation:

o Atomic mass= ∑(𝑓𝑟𝑎𝑐𝑡𝑖𝑜𝑛 𝑜𝑓 𝑖𝑠𝑜𝑡𝑜𝑝𝑒 𝑛) ∗ (𝑚𝑎𝑠𝑠 𝑜𝑓 𝑖𝑠𝑜𝑡𝑜𝑝𝑒 𝑛) =

(fraction of isotope 1 x mass of isotope 1)

+ (fraction of isotope 2 x mass of isotope 2)

+ ……

+ (fraction of isotope n x mass of isotope n)

Mass Spectrometry

Mass spectrometry is a technique that helps scientists to separate particles according to their mass

Chapter 2: The Quantum-Mechanical Model of the Atom

The quantum-mechanical world is a model that explains the strange behavior of electrons. The model

describes electrons as the exist within atoms; later, we shall see how those electrons determine the

chemical and physical properties of elements.

2.2 The Nature of Light

Light and electrons have some things in common. Chief among these things is the wave-particle duality

of light. Certain properties of light are described best by thinking of light as a wave and other properties

of light are best described by thinking of light as a particle.

Light is electromagnetic radiation, a type of energy embodied in oscillating electric and magnetic fields.

A magnetic field is a region of space where a magnetic particle experiences a force

An electric field is a region of space where an electrically charged particle experiences a force. A proton,

for example, has an electric field around it (of electrons).

Electromagnetic radiation can be described as a wave oscillating electric and magnetic fields. The fields

oscillate in perpendicular planes.

A wave can be characterized by its amplitude and wavelength (λ).

The wavelength of light determines its placement on the electromagnetic spectrum.

Like all waves, light is also characterized by its frequency (ν), or the number of cycles that pass through a

stationary point in a given period. The units of frequency are (cycles/s) or simply s-1. An equivalent unit

of ν is the hertz (Hz), defined as 1 cycle/second.

The frequency of a wave is directly proportional to the speed at which the wave is traveling. Frequency is

also inversely proportional to the wavelength –the farther apart the crests are, the fewer will pass a fixed

location over a unit of time. For light, we can write:

ν=c/λ

Where the speed of light is c and the wavelength λ are both expressed in terms of the same unit of

distance.

The electromagnetic spectrum

Chemistry Notes

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Visible light makes up the smallest portion of the electromagnetic spectrum, which includes all

wavelengths of electromagnetic radiation. The shortest wavelengths of light contain the highest energy.

The shortest wavelengths are as follows:

• Gamma (γ) ray – gamma rays are dangerous to humans because they can damage biological cells

• X-rays – x-rays are also slightly harmful to humans but only a few exposures to the rays are

relatively harmless

• Ultraviolet (UV) radiation – while not as energetic as gamma or x-rays, UV light still carries

enough energy to damage biological molecules

• Visible light – visible light at low intensities does not damage biological cells. The light causes

certain molecules in our eyes to change shape, allowing us to see

• Infrared (IR) radiation – heat is infrared radiation. All warm objects, including human bodies,

emit infrared radiation

• Microwaves – although microwave radiation has longer wavelengths and therefore emits lower

energies than IR or visible light, it is efficiently absorbed by water and can therefore heat

substances that contain water

• Radio waves – radio waves are the longest waves. They are used to transmit signals responsible

for AM and FM radio, cellular telephone, television, and other forms of communication

Interference and Diffraction

Waves, including electromagnetic waves, interact with each other in a characteristic way called

interference: the waves cancel each other out or build each other up depending on their alignment.

• Constructive interference occurs if two waves are of equal amplitude and are in phase when they

interact – they align their overlapping crests

• If waves are completely out of phase – that is, they align so that the crest from one overlaps with

the trough from the other – the waves cancel by destructive interference.

Waves also exhibit a characteristic behavior called diffraction. When a wave encounters an obstacle or a

slit that is comparable in size to its wavelength, it bends, or diffracts around it. The diffraction of light

through two slits separated by a distance comparable to the wavelength of the light coupled with

interference, results in an interference pattern.

The particle nature of light

The photoelectric effect is the observation that many metals emit electrons when light shines upon them.

An electron’s binding energy is the energy with which an electron is bound to the metal. In 1905, Albert

Einstein (1879-1950) proposed a bold explanation for the photoelectric effect: light energy must come in

packets. According to Einstein, the amount of energy (E) in a light packet depends on its frequency ν

according to the following equation:

E=hν

Where h, called Planck’s constant, has the value h=6.626x10-34 J*s

A packet of light is called a photon or a quantum of light. Since ν=c/λ, the energy of a photon can also be

expressed in terms of wavelength:

E=hc/λ

Einstein’s idea that light is quantized elegantly explains the photoelectric effect.

2.3 Atomic Spectroscopy and the Bohr Model

The discovery of the particle nature of light began to break down the division that existed in nineteenth

century physics between electromagnetic radiation, which was thought of as a wave phenomenon and the

small particles that compose atoms. The photoelectric effect suggested a particle nature of light but other

things suggested the wave-like nature of light. The most important of these discoveries came from atomic

spectroscopy.

Chemistry Notes

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• When an atom absorbs energy in the form of heat, light, or electricity, it often re-emits that

energy as light

o Ex: when an electric current is run through a clear glass tube of neon, the neon emits a

blue light

• Atoms of each element emit characteristic colors

• Light emitted by atoms contains distinct wavelengths of light

• Shining this light through a prism reveals distinct wavelengths of light called an emission

spectrum

• Niels Bohr developed a model for the atom that attempted to explain this phenomenon

o Atoms exist at fixed distances around the atom (n=1, n=2,…)

o The energy of each Bohr orbit is fixed, or quantized

o Bohr called the orbits stationary states

o Bohr proposed that no radiation is emitted by an electron orbiting the nucleus. It is only

when an electron jumps, or makes a transition from one stationary state to another that

radiation is emitted or absorbed.

o The electron is never observed between states, only in one state or another

• The emission spectrum of an atom consists of discrete lines because the stationary states exist

only at specific, fixed energies. The energy of the photon emitted when an electron makes a

transition is the energy difference between the two stationary states.

• The emission spectrum of each atom is unique and can be used to identify the substance

• An absorption spectrum consists of dark lines on a bright background

• It is measured by passing white light through a sample and observing what wavelengths are

missing due to the absorption by the sample.

• The absorption lines are at the same wavelengths as the emission spectrum lines. This is because

the processes that produce them are mirror images

• In emission, an electron makes a transition from a higher-energy level to a lower energy level. In

absorption, the transition is between the same two energy levels but from the lower one to the

higher one

2.4 The Wave Nature of Matter: The de Broglie Wavelength, the Uncertainty Principle, and

Indeterminacy

• Single electron diffraction experiment shoots a single electron at two slits and detects where the

electron lands

• There are patterns of interference

• The interference pattern is not caused by pairs of electrons interfering with each other but rather

by single electrons interfering with themselves

• The wave nature of the electron is an inherent property of individual electrons

• The wavelength of an electron is given by the de Broglie relation:

λ=h/mv

• The act of (literally) observation forces an electron into a wave or particle state

• We can never see both the interference pattern and simultaneously determine which hole the

electron goes through

• “There is a limit to the fineness of our powers of observation and the smallness of the

accompanying disturbance – a limit which is inherent in the nature of things and can never be

surpassed by improved technique or increased skill on the part of the observer.

• The wave and particle-like nature of the electron are complementary properties

• We cannot simultaneously measure the electron’s position and velocity with infinite precision.

We can only predict where it is

Chemistry Notes

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• The behavior of an electron (impossible to predict) is called indeterminacy

2.5 Quantum Mechanics and the Atom

• An electron’s position is described in terms of an orbital, a probability distribution map showing

where the electron is likely to be found. Since chemical bonding often involves the sharing of

electrons between atoms, the spatial distribution of atomic electrons is important to bonding.

• The Schrodinger equation is the equation for finding the energies and orbitals of atoms. Its

general form is

Hψ=Eψ

• The symbol ψ is the wave function, a mathematical function that describes the wavelike nature of

the electron. A plot of the wave function squared (ψ2) represents an orbital.

• Each orbital is specified by three quantum numbers

• The principal quantum number, n

• The angular momentum quantum number, l

• The magnetic quantum number, ml

• A fourth quantum number, ms, specifies the spin orientation of the electron

• The principal quantum number, n, is an integer (n=1,2,3,…..) that determines the overall size and

energy of an orbital

o For the hydrogen atom, the energy of an electron in an orbital with quantum number n is

given by

o En= -2.18x10-18 J (1/n2)

o The energy is negative because the electron’s energy is lowered by its interaction with the

nucleus

• The angular momentum quantum number (l)

o The amq number is an integer that determines the shape of the orbital. The possible values for

l are 0,1,2,…,(n-1)

o In other words, l can be any integer up to (n-1)

o The values of l are often assigned letters as follows:

L=0 s

L=1 p

L=2 d

L=3 f

• The magnetic quantum number (ml)

o The mqn is an integer that specifies the orientation of the orbital

o The possible values of ml range from -l to l

• The spin quantum number (ms)

o The sqn determines the orientation of the spin of the electron

o Electron spin is a fundamental property of an electron (like its negative charge).

o All electrons have the same amount of spin

o The orientation of the electron’s spin is quantized, with only two possibilities that we can call

spin up (ms = +1/2) and spin down (ms = -1/2)

o Each specific combination of the first three quantum numbers specifies one specific atomic

orbital.

▪ Example: the orbital with n=1, l=0, and ml=0 is known as the 1s orbital. The 1 in 1s

is the value of n, and the s specifies that l=0. There is only one 1s orbital in an atom

and its ml value is zero

Atomic Spectroscopy Explained

Each wavelength in the emission spectrum of an atom corresponds to an electron transition between

quantum-mechanical orbitals. When an atom absorbs energy, an electron in a lower-energy orbital is

Chemistry Notes

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excited or promoted to a higher-energy orbital. In the new configuration, however, the atom is unstable,

and the electron quickly falls back or relaxes to a lower-energy orbital. As it does so, it releases a photon

of light containing an amount of energy precisely equal to the energy difference between the two energy

levels.

Excitation and Radiation

The energy from n=1 up increases

When an electron jumps up, the electron absorbs energy and is excited to an unstable energy level

Electron up → Energy absorbed

When an electron goes down, light is emitted as the electron goes to a lower-energy (PE) level

Electron down → Light Emitted; lower PE

2.6 The Shapes of Atomic Orbitals

The shapes of atomic orbitals are important because chemical bonds depend on the sharing of electrons

that occupy the orbitals. The shape of an atomic orbital is determined primarily by l, the angular

momentum quantum number. The value of l is assigned a corresponding letter. We call l= 0 , s orbitals,

l= 1 , p orbitals, l= 2 , d orbitals, and so on.

• s orbitals (l=0)

o The lowest energy orbital is the spherically symmetrical 1s orbital.

o An atomic orbital can also be represented by a geometrical shape that encompasses the

volume where the electron is likely to be found most frequently – typically, 90% of the

time

• P orbitals (l=1)

o The p orbitals are not spherically symmetric like the s orbitals

o P orbitals have two lobes of electron density on either side of the nucleus and a node

located at the nucleus. The three p orbitals differ only in their orientation and are

orthogonal to one another.

o A node is a point where the wave function (ψ), and therefore the probability density (ψ2)

and radical distribution function, all go through zero.

o The 3p, 4p, 5p, and higher p orbitals are all similar in shape to the 2p orbitals but they

have additional nodes.

• D orbitals (l=2)

Chemistry Notes

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o Each principal level with n=3 or greater contains 5 d orbitals. Four of the five 3d orbitals

have a cloverleaf shape, with four lobes of electron density around the nucleus and two

perpendicular nodal planes

• F orbitals (l=3)

o Each principal level with n=4 or greater contains seven f orbitals. The f orbitals have

more nodes and lobes than d orbitals.

The phase of orbitals

1. The sign of the amplitude of a wave is known as its phase. The phase of a wave determines how it

interferes with another wave

2. Just as a one-dimensional wave has a phase, so does a three-dimensional wave

The shape of atoms

1. If some orbitals are shaped like dumbbells and three-dimensional cloverleaves, and if most of the

volume of an atom is empty space, why do we often depict atoms as spheres?

2. Atoms are usually drawn as spheres because most atoms contain many electrons occupying

several different orbitals. Therefore, the shape of an atom is obtained by superimposing all its

orbitals.

3. If we superimpose the s, p, and d orbitals, we get a spherical shape

Quantum Numbers

A. Principal Quantum Number: n

a. Refer to these values as “energy levels” or “energy states”

b. Values: (1,2,3…, n)

c. Tells us approximate energy or distance from the nucleus

d. Large values of n correspond to large orbitals

B. Angular Momentum Quantum Number: l

a. Whole Number values: (0,1,2…,(n-1))

b. Referred to as “subshells”

c. Subshells are linked to the shapes of orbitals

d. l has numbers: s, p, f, d for l=0,1,2,3

e. When l=0, the orbital is a spherical shape (s-orbital)

f. When l=1, the orbital has two lobes of electron density (p-orbital)

g. When l=2, the orbital has a cloverleaf shape (d-orbital)

h. When l=3, the orbital has full lobes of electron density in each octant (f-orbital)

i. When l=4, l is referred to as a “g” subshell, then “h” for l=5 and so on

C. Magnetic Quantum Number: ml

a. Values: range → [-l, l]

b. Each value represents a different value with respect to the nucleus in space

c. Referred to as distinct “orbitals”

d. This number reflects the number of possible orbitals that could exist

D. Spin Quantum Number: ms

a. Values: +1/2, -1/2

b. Tells us which direction an electron is spinning

E. Generalizations

a. In any n level, there are n subshells

b. In any n level, there are n2 orbitals

i. Each orbital has the quantum numbers <n, l, ml>

F. Rules

a. Every electron in an atom must have a different set of quantum numbers

G. Orbital energies

a. As n increases, energy increases as well as distance from the nucleus (1<2<…n)

b. As l increases, energy also increases (s<p<d<f<….l)

Chemistry Notes

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c. A nucleus has a defined amount of attraction for e-

i. Infinite energy levels are possible

d. Energy levels (begins at n=3) overlap one another in space

e. To determine the actual energy of levels and subshells: use the following chart:

Or construct a chart as so:

This shows the filling order or orbitals

) Chapter 3: Periodic Properties of the Elements

I) Aluminum: Low Density Atoms Result in Low-Density Metal

A) Aluminum has a density of only 2.70 g/cm3

B) For comparison, Iron has a density of 7.86 g/cm3

C) Platinum has a density of 21.4 g/cm3

D) The density of aluminum is so low because the density of an aluminum atom is so low.

Although the arrangements of atoms in a solid must also be considered when evaluating the

density of the solid, the mass-to-volume ratio of the composite atoms is a very important

factor (definition of density). For this reason, the densities of the elements generally follow a

fairly well-defined trend: The density of the elements tends to increase as we move down a

column in the periodic table.

E) As we move down the column in the periodic table, the density of the elements increases

even though the radius generally increases as well.

(1) The mass of each successive atom increases more than its volume does

F) As we move down a column in the periodic table, the additional protons and neutrons add

more mass to the atoms.

(1) The increase in mass is greater than the increase of the radius of the atom

Chemistry Notes

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G) A periodic property is a property that is generally predictable based on an element’s position

in the periodic table.

H) Periodic Properties to consider in chapter 3

(1) Atomic radius

(2) Ionization energy

(3) Electron affinity

I) These periodic properties are explained by quantum-mechanical theory

(1) Quantum-mechanical theory explains the electronic structure of atoms – this in turn

determines the properties of those atoms.

J) Structure determines properties of atoms

(1) The arrangement of elements in the periodic table—originally based on similarities in the

properties of the elements—reflects how electrons fill quantum-mechanical orbitals.

II) The periodic Law and the Periodic Table – 3.2

A) Periodic Law: When the elements are arranged in order of increasing mass, certain sets of

properties recur periodically

B) The periodic table was arranged by Dmitri Mendeleev of elements in order of increasing

mass

(1) Mendeleev’s arrangement was a success allowing him to predict the existence and

properties of yet undiscovered elements such as eka-aluminum and eka-gallium (Eka

means the one beyond or the next one in a family of elements. So, eka-silicon means the

element beyond silicon in the same family as silicon)

C) The periodic table of elements is divided into two subgroups:

(1) Main-group elements

(2) Transition elements

D) Main group elements are elements whose properties tend to be largely predictable based on

their position in the periodic table

(1) Main group elements are in columns labeled with a number and the letter A

E) Transition elements (as well as inner transition elements or metals) are elements that tend to

be less predictable based simply on their position in the periodic table.

(1) Transition elements are in columns labeled with a number and the letter B

F) Each column within the main-group regions of the periodic table is a family or group of

elements. A family of elements has similar properties as observed by Mendeleev

III) Electron Configuration: How Electrons Occupy Orbitals – 3.3

(1) An electron configuration for an atom shows the particular orbitals that electrons occupy

for that atom.

(2) The ground state electron configuration of an atom is the lowest energy state of an atom.

This is the ground state configuration for the hydrogen atom:

(a) H → 1s1

(3) Electrons generally occupy the lowest energy orbitals available.

A) Electron spin and the Pauli Exclusion Principle

(1) We can represent the electron configuration of a hydrogen atom in a slightly different

way with an orbital diagram, which is similar to an electron configuration but symbolizes

the electron as an arrow and the orbital as a box

(a) H →

1s

(2) How do the spins of the two electrons in helium align relative to each other? The answer

to this question is addressed by the Pauli exclusion principle:

(a) Pauli exclusion principle: No two electrons in an atom can have the same four

quantum numbers

Chemistry Notes

15

(3) Because the two electrons occupying the same orbital have three identical quantum

numbers (n, l, and ml), they each must have a different spin quantum number. Since there

are only two possible spin quantum numbers, the Pauli Exclusion Principle implies that

each orbital can have a maximum of only two electrons, with opposing spins.

B) Sublevel Energy Splitting in Multi-Electron Atoms

(a) A major difference in the (approximate) solutions to the Schrodinger equation for

multi-electron atoms compared to the solutions for the hydrogen atom is the energy

ordering of the orbitals. In the hydrogen atom, the energy of an orbital depends only

on n, the principle quantum number. For example the 3s, 3p, and 3d orbitals (which

are empty for hydrogen in its lowest energy state) all have the same energy—we say

that they are degenerate.

(b) The orbitals within a principal level of a multi-electron atom, in contrast, are not

degenerate—their energy depends on the value of l. We say that the energies of the

sublevels are split.

(c) In general, the lower the value of l within a principle level, the lower the energy (E)

of the corresponding orbital. Thus, for a given value of n:

(i) E (s orbital)< E (p orbital)< E (d orbital)< E (f orbital)

(d) To understand why the sublevels split in this way, we must examine three key

concepts associated with the energy of an electron in the vicinity of a nucleus:

(i) Coulomb’s law, which describes the interactions between charged particles

(ii) Shielding, which describes how one electron can shield another electron from the

full charge of the nucleus

(iii) Penetration, lastly, which describes how one atomic orbital can overlap spatially

with another, thus penetrating into a region that is close to the nucleus (and

therefore is less shielded from nuclear charge)

(e) We then examine how these concepts, together with the spatial distributions of

electron probability for each orbital, result in the energy ordering presented above

(point 3a)

(2) Coulomb’s Law

(a) Coulomb’s law states that the potential energy (E) of two charged particles depends

on their charges (q1 and q2) and on their separation (r):

𝐸 =1

4𝜋 ∗ 𝜖0∗

𝑞1 ∗ 𝑞2

𝑟

(b) In this equation, ϵ0 is a constant (ϵ0 = 8.85 x 10-12 C2/J*m). The potential energy is

positive for the interaction of charges with the opposite sign (+ or -). The magnitude

of the potential energy depends inversely on the separation between the charged

particles.

(c) We can draw three important conclusions from Coulomb’s law:

(i) The potential energy (E) associated with the interaction of like charges is positive

but decreases as the particles get farther apart (as r increases).

1. Since systems tend towards lower potential energy, like charges that are

close together have high potential energy and tend to move away from each

other. Like charges therefore repel one another (like a magnet)

(ii) The potential energy (E) associated with the interaction of unlike charges is

negative and becomes more negative as the particles get closer together

1. Ibid.

(iii) The magnitude of the interaction between charged particles increases as the

charges of the particles increase.

Chemistry Notes

16

1. Consequently, an electron with a charge of 1 – is more strongly attracted to a

nucleus with a charge of 2 + than it is to a nucleus with a charge of 1 +.

(3) Shielding

(a) For multi-electron atoms, any single electron experiences both the positive charge of

the nucleus (which is attractive) and the negative charges of the other electrons

(which are repulsive)

(b) We can think of the repulsion of one electron by other electrons as screening or

shielding that electron from the full effects of nuclear charge.

(i) For example, consider a lithium ion (Li+). Because the lithium ion contains two

electrons, its electron configuration is identical to that of helium: Li+ 1s2

(ii) Now imagine bringing a third electron toward the lithium ion. When the third

electron is far from the nucleus, it experiences the 3+ charge of the nucleus

through the screen or shield of the 2 – charge of the two 1s electrons

(c) We can think of the third electron as experiencing an effective nuclear charge (Zeff)

of approximately 1+ (in a figure with the nucleus 3+ and two electrons around the

nucleus). The inner electrons in effect shield the outer electron from the full nuclear

charge.

(4) Penetration

(a) Now imagine that this third electron comes closer to the nucleus. As the electron

penetrates the electron cloud of the 1s electrons, it begins to experience the 3+ charge

(shown in a figure) of the nucleus more fully because the third electron is less

shielded by the intervening electrons. If the electron could somehow get closer to the

nucleus than the 1s electrons, it would experience the full 3+ charge.

(i) In other words, as the outer electron undergoes penetration into the region

occupied by the inner electrons, it experiences a greater nuclear charge and

therefore (according to Coulomb’s law), a lower energy.

(5) Electron Spatial Distributions and Sublevel Splitting

C) Electron Configurations for Multi-Electron Atoms

Chemistry Notes

17

(a) Now we know the energy ordering of orbitals in multi-electron atoms, we can

determine ground state electron configurations for the rest of the elements. Because

electrons occupy the lowest energy orbitals available when the atom is in its ground

state, and only two electrons are allowed in each orbital, we can systematically build

up the electron configurations for the elements. This pattern of orbital filling is

known as the Aufbau principle.

(b) The Aufbau principle is the principle that indicates the pattern of orbital filling in an

atom

(c) Hund’s rule state that when filling degenerate orbitals, electrons fill them singly first,

with parallel spins.

(2) Summarizing orbital filling:

(a) Electrons occupy orbitals to minimize the energy of the atom: therefore, lower-

energy orbitals fill before higher-energy orbitals. Orbitals fill in the following order:

(i) 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s

(b) Orbitals can hold no more than two electrons each. When two electrons occupy the

same orbital, their spins are opposite. This is another way of expressing the Pauli

Exclusion Principle

(c) When orbitals of identical energy are available, electrons first occupy these orbitals

singly with parallel spins rather than in pairs (Hund’s rule). Once the orbitals of equal

energy are half-full, the electrons start to pair.

(3) To write an electron configuration for an element, we first find its atomic number from

the periodic table—this number equals the number of electrons (in a neutral atom). Then

we use the order of filling to distribute the electrons in appropriate orbitals. Remember

that each orbital can hold a maximum of two electrons. Consequently,

(a) The s sublevel only has one orbital and can hold only two electrons

(b) The p sublevel only has three orbitals and can hold six electrons

(c) The d sublevel has five orbitals and can hold ten electrons

(d) The f sublevel has seven orbitals and can hold 14 electrons

Tabular display:

Sublevel Orbitals Electrons able to

hold

S 1 2

P 3 6

D 5 10

F 7 14

III) Electron Configurations, Valence Electrons, and the Periodic Table – 3.4

(a) Recall from section 3.2 that Mendeleev arranged the periodic table so that elements

with similar chemical properties lie in the same column. We can begin to make the

connection between an element’s properties and its electron configuration by

superimposing the electron configurations of the first 18 elements onto a partial

periodic table.

(b) As we move to the right across a row (a period), the orbitals fill in the correct order.

With each subsequent row, the highest principal quantum number increases by one.

(c) Notice that as we move down a column, the number of electrons in the outermost

principle energy level (highest n value) remains the same. The key connection

between the macroscopic world (an element’s chemical properties) and the

particulate world (an atom’s electronic structure) lies in these outermost electrons.

Chemistry Notes

18

(d) An

atom’s valence electrons are the most important in chemical bonding. For main group

elements, the valence electrons are those in the outermost principle energy level. For

transition elements, we also count the outermost d electrons among the valence

electrons (even though they are not in an outermost principle energy level).

(e) The chemical properties of an element depend on its valence electrons, which are

instrumental in bonding because they are held most loosely (and are therefore the

easiest to lose or share).

(f) We can now see why the elements in a column of the periodic table have similar

chemical properties: they have the same number of valence electrons

(g) We distinguish valence electrons from all the other electrons in an atom, which we

call core electrons. The core electrons are those in complete principle energy levels

and those in complete d and f sublevels.

(i) For example, silicon, with the electron configuration 1s22s22p63s23p2 has four

valence electrons (those in the n = 3 principle level) and ten core electrons:

1. Si 1s22s22p63s23p2

Core electrons Valence electrons

A) Orbital Blocks in the Periodic Table

(1) A pattern similar to what we saw for the first 18 elements exists for the entire periodic

table. Note that, because of the filling order of orbitals, the periodic table can be divided

into blocks representing the filling of particular sublevels.

(a) The first two columns on the left side of the periodic table constitute the s block, with

outer electron configurations of ns1 (group 1A) and the ns2 electron configurations

(b) The six columns on the right side of the periodic table constitute the p block

(i) Together, the s and p blocks constitute the main-group elements

(c) The transition elements constitute the d block

(d) The lanthanides and the actinides constitute the f block

(2) The number of columns in a block corresponds to the maximum number of electrons that

can occupy the particular sublevel of that block.

(a) The s block has two columns (corresponding to one s orbital)

(b) The p block has six columns (corresponding to three p orbitals)

(c) The d block has ten columns (corresponding to five d orbitals)

(d) The f block has 14 columns (corresponding to seven f orbitals)

(3) Except for helium, the number of valence electrons for any main-group element is equal

to its lettered group number.

(a) For example, we know that chlorine has seven valence electrons because it is in

group number 7A

Chemistry Notes

19

(4) Lastly, note that for main-group elements, the row number in the periodic table is equal

to the number (or n value) of the highest principle level.

(a) For example, because chlorine is in row 3, its highest principal level is the n = 3

level.

(5) Summarizing Periodic Table Organization

(a) The periodic table is divisible into four blocks corresponding to the filling of four

quantum sublevels (s, p, d, and f).

(b) The lettered group number of a main-group element is equal to the number of valence

electrons for that element.

(c) The row number of a main-group element is equal to the highest principal quantum

number of that element.

B) Writing an Electron Configuration for an Element from Its Position in the Periodic Table

(1) The organization of the periodic table allows us to write the electron configuration for

any element based on its position in the periodic table.

(2) For example, suppose we want to write an electron configuration for Cl.

(3) The inner electron configuration of Cl is that of the noble gas that precedes it in the

periodic table, Ne. So we represent the inner electron configuration of Cl with [Ne]. We

obtain the outer electron configuration—the configuration of the electrons beyond the

previous noble gas—by tracing the elements between Ne and Cl and assigning electrons

to the appropriate orbitals

(4) Cl [Ne] 3s23p5

C) The Transition and Inner Transition Elements

(1) The electron configurations of the transition elements (d block) and inner transition

elements (f block) exhibit trends that differ somewhat from those of the main-group

elements. As we move to the right across a row in the d block, the d orbitals fill as shown

here.

(2) Note that the principal quantum number of the d orbitals that fill across each row in the

transition series is equal to the row number minus one.

(a) The actual electron configurations are determined experimentally through atomic

spectroscopy and do not always conform to the general pattern. Nonetheless, the

patterns described allow us to accurately predict electron configurations for most of

the elements in the periodic table

(3) As we move across the f block (the inner transition series), the f orbitals fill. For these

elements, the principle quantum number of the f orbitals that fill across each row is the

row number minus two.

(4) In addition, within the inner transition series, the close energy spacing of the 5d and 4f

orbitals sometimes causes an electron to enter a 5d orbital instead of the expected 4f

orbital.

(a) For example, the electron configuration of gadolinium is [Xe] 6s24f75d1 (instead of

the expected) [Xe] 6s24f8

IV) Electron Configurations and Elemental Properties – 3.5

(a) The chemical properties of elements are largely determined by the number of valence

electrons the elements contain

Chemistry Notes

20

(b) Perhaps the most striking family in the periodic table is 8A, known as the noble gases

(i) Each noble gas has full outer electron shells.

(ii) This is why the noble gases are particularly stable and unreactive

(c) When a quantum level is completely full, the overall potential energy of the electrons

that occupy that level is particularly low

(2) We can explain a great deal of chemical behavior with this simple idea:

(a) Elements without a noble gas configuration react to attain a noble gas configuration

(3) First, we apply this idea to help differentiate between metals and nonmetals

(4) Next, we apply the idea to understand the properties of several individual families of

elements

(5) Last, we apply the idea to the formation of ions

B) Metals and Nonmetals

(1) Metals lie on the lower left side and middle of the periodic table and share some common

properties: they are good conductors of electricity; the can be pounded into flat sheets

(malleability); they can be drawn into wires (ductility); they are often shiny; and most

importantly, they tend to lose electrons when undergoing chemical changes.

(i) Na is among the most reactive metals. When we find sodium in nature, it is often

found as Na+, which has the electron configuration of neon.

(b) The other main group metals in the periodic table behave similarly: they tend to lose

their valence electrons in chemical changes to attain noble gas electron

configurations. The transition metals also tend to lose electrons in their chemical

changes, but they do not generally attain noble gas configurations

(2) Nonmetals lie on the upper right side of the periodic table. The division between metals

and nonmetals is the zigzag line running from boron to astatine.

(a) Nonmetals have varied properties—some are solids at room temperature, others are

liquids or gases—but as a whole, they tend to be poor conductors of electricity and

heat, and most importantly they all tend to gain electrons when they undergo

chemical changes.

(i) Chlorine is among the most reactive nonmetals. Chlorine is one electron short of

the noble gas configuration of argon

(b) Nonmetals tend to gain electrons in chemical changes to attain noble gas

configurations

(3) Many of the elements that lie along the zigzag diagonal line that divides metals and

nonmetals are metalloids and exhibit mixed properties.

(a) Several metalloids are classified as semiconductors because of their intermediate

(and highly temperature-dependent) electrical conductivity.

(b) This makes metalloids useful in the manufacture of devices central to computers and

phones and more

(4) Metals: lower left and middle of periodic table

(a) Good conductors of heat and electricity

(b) Tend to lose electrons when they undergo chemical reactions

(c) Ex: Na (sodium)

(5) Non-metals: upper right of the periodic table

(a) Poor conductors of heat and electricity

(b) Tend to gain electrons

(c) Ex: Cl (chlorine)

C) Families of Elements

(1) We can also understand the properties of families of elements based on their electron

configurations.

(2) Here is a tabular listing of: group or family, valence electron general format, and most

general charge taken on:

Chemistry Notes

21

Group Name General outer

electron configuration

Most common charge

taken on

1A Alkali metals ns1 1+

2A Alkaline earth metals ns2 2+

7A Halogens ns2np5 1 –

D) The Formation of Ions

(a) We learned that atoms can lose or gain electrons to form ions

(b) We also learned that metals tend to form

(i) Cations: positively charged ions

(c) And nonmetals tend to form

(i) Anions: negatively charged ions

(2) Notice on the periodic table that, for the main-group elements that form cations with

predictable charge, the charge is equal to the group number. For the main-group elements

that form anions with predictable charge, the charge is equal to the group number minus

eight.

(3) The tendency for many main-group elements to form ions with noble gas electron

configurations does not mean that the process is in itself energetically favorable. In fact,

forming cations always requires energy, and forming anions sometimes requires energy

as well.

(4) However, the energy cost of forming a cation or anion with a noble gas configuration is

often less than the energy payback that occurs when that cation or anion forms chemical

bonds (more on this topic in chapter 4).

(5)

V) Periodic Trends in Atomic Size and Effective Nuclear Charge – 3.6

(a) How do we define the size of an atom? One way to define atomic radii is to consider

the distance between the centers of adjacent atoms which can be determined from a

frozen solid’s density. An atomic radius determined in this way is the nonbonding

atomic radius or the van der Waals radius. Another way to define the size of an

atom, the bonding atomic radius or covalent radius, is defined differently for

nonmetals and metals as follows:

(i) Nonmetals: one-half the distance between two of the atoms bonded together

(ii) Metals: one-half the distance between two of the atoms next to each other in a

crystal of the metal

(b) Atomic Radius refers to a set of average bonding radii determined from

measurements on a large number of elements and compounds. The atomic radius

Chemistry Notes

22

represents the radius of an atoms when it is bonded to another atom and is always

smaller than the van der Waals radius.

(c) The general trends in the atomic radii of main-group elements, which are the same as

trends observed in van der Waals radii, are as follows:

(i) As we move down a family in the periodic table, the atomic radius increases

(ii) As we move to the right across a period in the periodic table, the atomic radius

decreases

(d) We can understand the observed trend in radius as we move down a column based on

the trends in the sizes of atomic orbitals. The atomic radius is largely determined by

the valence electrons, the electrons farthest from the nucleus. As we move down a

family in the periodic table, the highest principle quantum number n of the valence

electrons increases. Consequently, the valence electrons occupy larger orbitals,

resulting in larger atoms.

(e) The observed trend in atomic radius as we move to the right across a period,

however, is a bit more complex. To understand this trend, we now revisit some

concepts from 3.3, including effective nuclear charge and atomic shielding.

A) Effective Nuclear Charge

(1) The trend in atomic radius as we move across a period is determined by the inward pull

of the nucleus on the electrons in the outermost principle energy level (highest n value).

According to Coulomb’s law, the attraction between the nucleus and an electron

increases with increasing magnitude of nuclear charge.

(2) As shown in the figure below, even though the 2s orbital penetrates into the 1s orbital to

some degree, the majority of the 2s orbital is outside the 1s orbital. Therefore, the

electron in the 2s orbital is partially screened or shielded from the 3+ charge of the

nucleus by the 2- charge of the 1s (or core) electrons, reducing the net charge experienced

by the 2s atom.

(3) Recall that we define the average or net charge experienced by an electron as the effective

nuclear charge. The effective nuclear charge experienced by an electron in an atom is the

actual nuclear charge (Z) minus the charge shielded by other electrons (S):

Chemistry Notes

23

𝑍𝑒𝑓𝑓 = 𝑍 − 𝑆

(4) Consider the valence electrons in beryllium (Be) with atomic number 4. Its electron

configuration is: Be 1s22s2

(5) To estimate the effective nuclear charge experienced by the 2s electrons in beryllium, we

must distinguish between the two different types of shielding

(a) 1. The shielding of outermost electrons by the core electrons

(b) 2. The shielding of the outermost electrons by each other

(6) The key to understanding the trend in atomic radius is the difference between these two

types of shielding

(7) Core electrons efficiently shield the electrons in the outermost principle energy level from

nuclear charge, but outermost electrons do no efficiently shield one another from nuclear

charge.

(8) Summarizing Atomic Radii for Main-Group Elements:

(a) As we move down a column in the periodic table, the principle quantum number of

the electrons in the outermost principle energy level increases, resulting in larger

orbitals and therefore larger atomic radii.

(b) As we move to the right across a row in the periodic table, the effective nuclear

charge (Zeff) experienced by the electrons in the outermost principal energy level

increases, resulting in a stronger attraction between the outermost electrons and the

nucleus, and smaller atomic radii.

B) Atomic Radii and the Transition Elements

(1) Notice in the figure above that as we move down the first two rows in a column within

the transition metals, the elements follow the same general trend in atomic radii as the

main-group elements (the radii get larger moving down the first two rows in the transition

metals).

(2) In contrast, with the exception of the first couple of elements in each transition series, the

atomic radii of the transition elements do not follow the same trend as the main-group

elements as we move to the right across a row. Instead of decreasing in size, the radii of

the transition elements stay roughly constant across each row. Why?

(a) The difference is that, across a row of transition elements, the number of electrons in

the outermost principle energy level is nearly constant.

(i) As another proton is added to the nucleus of each successive element, another

electron is added as well, but the electron goes into an nhighest – 1 orbital (refer to

III.C.2). The number of outermost electrons stays constant, and the electrons

Chemistry Notes

24

experience a roughly constant effective nuclear charge, keeping the radius

approximately constant.

VI) Ions: Electron Configurations, Magnetic Properties, Radii, and Ionization Energy – 3.7

(1) Recall that ions are atoms that have lost or gained electrons. In this section, we examine

periodic trends in ionic electron configurations, magnetic properties, radii, and ionization

energies.

B) Electron Configurations and Magnetic Properties of Ions

(1) We can deduce the electron configuration of a main-group monoatomic ion from the

electron configuration of the neutral atom and the charge of the atom.

(a) For anions, we add the number of electrons indicated by the magnitude of the charge

of the anion

(b) For cations, we subtract the number of electrons indicated by the charge of the cation

in the reverse order of filling.

(c) For transition metal cations, the trend if different

(d) When writing the electron configuration of a transition metal cation, we remove the

electrons in the highest n-value orbitals first, even if this does not correspond to the

reverse order of filling.

(i) For example, the electron configuration of vanadium (V) is: V [Ar]4s23d3

(ii) The V2+ ion, however, has the following electron configuration: V [Ar]4s03d3

(e) Why this unexpected behavior of transition metals?

(f) The full answer to this is beyond our scope, but the following two factors contribute

to this phenomenon:

(i) As discussed, the ns and (n-1)d orbitals are extremely close in energy and,

depending on the exact configuration, can vary in relative energy ordering

(ii) As the (n-1)d orbitals begin to fill in the first transition series, the increasing

nuclear charge stabilizes the (n-1)d orbitals relative to the ns orbitals. This

happens because the (n-1)d orbitals are not the outermost (or highest n) orbitals

and are therefore not effectively shielded from the increasing nuclear charge by

the ns orbitals.

(g) The bottom is that we remove the ns electrons before the (n-1)d electrons when

writing electron configurations.

(h) An atom or ion that contains unpaired electrons is attracted to an external magnetic

field, and we say that the atom or ion is paramagnetic

(i) An atom or ion in which all electrons are paired is not attracted to an external

magnetic field—it is instead slightly repelled—and we say that the atom or ion is

diamagnetic.

(j) Observations in other transition metals confirm that the ns electrons are lost before

the (n-1)d electrons upon ionization.

C) Ionic Radii

(1) What happens to the radius of an atom when it becomes a cation? An anion?

(a) Cations are much smaller than their corresponding neutral atoms.

(b) Anions are much larger than their corresponding neutral atoms.

(c) We can observe an interesting trend in atomic size by examining the radii of an

isoelectronic series of ions—ions with the same number of electrons. Consider the

following ions and their radii:

(d) Isoelectronic series: a cohort of atoms/ions with the same number of electrons

S2— (184 pm) Cl— (181 pm) K+ (133 pm) Ca2+ (99 pm)

18 electrons 18 electrons 18 electrons 18 electrons

16 protons 17 protons 19 protons 20 protons

Chemistry Notes

25

(2) All of these ions have 18 electrons in exactly the same orbitals, but the radius of each ion

gets successively smaller. Why? The reason is the progressively greater number of

protons. As the protons increase in the atoms, there is a greater nuclear charge in the

atoms resulting in a smaller radius.

D) Ionization Energy

(2) The ionization energy (IE) of an atom or ion is the energy required to remove an electron

from the atom or ion in the gaseous state.

(3) Ionization energy is always positive because removing an electron always takes energy.

(The process is endothermic, which absorbs heat and therefore carries a positive sign)

(4) The energy required to remove the first electron is the first ionization energy (IE1). For

example, we represent the first ionization of sodium with the equation:

(a) Na(g) → Na+(g) + 1 e— IE1 = 496 kJ/mol

(5) The energy required to remove the second electron is the second ionization energy (IE2),

the energy required to remove the third electron is the third ionization energy (IE3) and so

on. We represent the second ionization energy of sodium as:

(a) Na+ (g) → Na2+ (g) + 1e— IE2 = 4560 kJ/mol

(6) Notice that the second ionization energy is not the energy required to remove two

electrons from sodium (which is the sum of the first and second ionization energies), but

rather the energy required to remove one electron from Na+.

E) Ionization Energy – Class Notes February 7, 2018

(2) Definition: amount of energy needed to remove an electron from an atom or ion in the

gaseous state (as opposed to the photoelectric effect which occurs in the solid state).

(3) Equation: X(g) → X+(g) + e- ΔE= 1st ionization energy (if you took another electron

away, you get X2+ and another electron. The energy it takes to complete this process is

known as second ionization energy…and so on)

(4) Successive ionization energies: amounts of energy needed to remove successive electrons

from atoms or ions

(a) These energies always increase for any atom or ion. It is harder to remove successive

electrons from an atom because there is a greater nuclear charge acting on the

electrons

(5) Valence electrons have relatively low ionization energies (they are relatively easily

removed compared to core electrons)

F) Trends in First Ionization Energy

(2) Here are two helpful figures to show trends in ionization energies

Chemistry Notes

26

(3) Based on what we have learned about electron configurations and effective nuclear

charge, how can we account for the observed trend? As we have seen, the principal

quantum number increases as we move down a family. For a given sublevel, orbitals with

higher principal quantum numbers are larger than orbitals with smaller principal quantum

numbers. Consequently, electrons in the outermost principal energy level are farther

away from the positively charged nucleus—and are therefore held less tightly—as we

Chemistry Notes

27

move down a column. This results in lower ionization energies as we move down a

column as shown in the figure above.

(4) What about the trend as we move to the right of a row? For example, does it take more

energy to remove an electron from Na or from Cl?

(a) We know that Na has an electron configuration of 3s1 and Cl has an outer electron

configuration of 3s23p5.

(b) The outermost electrons in chlorine experience a higher effective nuclear charge than

the outermost electrons in sodium (which is why chlorine has a smaller atomic radius

than sodium*). Consequently, we would expect chlorine to have a larger first

ionization energy than sodium, which is indeed the case.

(5) First ionization energy generally increases as we move to the right across a row in the

periodic table.

(6) Summarizing First Ionization Energy for Main-Group Elements

(a) First ionization energy generally decreases as we move down a column or family in

the periodic table because electrons in the outermost principle energy level are

increasingly farther away from the positively charged nucleus and are therefore held

less tightly.

(b) First ionization energy generally increases as we move to the right across a row (or

period) in the periodic table because electrons in the outermost principle energy level

generally experience a greater effective nuclear charge (Zeff).

G) Exceptions to Trends in First Ionization Energy

(1) Boron has a smaller ionization energy than beryllium even though it lies to the right of

beryllium in the same row

(a) This is caused by the change in going from the s block to the p block.

(i) Recall that the 2p orbital penetrates into the nuclear region less than the 2s

orbital.

(ii) Consequently, the 1s electrons shield the electron in the 2p orbital from nuclear

charge more than they shield the electrons in the 2s orbital. The result is that the

2p orbitals are higher in energy, and therefore the electron is easier to remove (it

has a lower first ionization energy). Similar exceptions occur for aluminum and

gallium, both directly below boron in group 3A.

(2) Another exception occurs between nitrogen and oxygen. Although oxygen is to the right

of nitrogen in the same row, is has a lower first ionization energy. This exception is

caused by the repulsion between electrons when they occupy the same orbital.

(a) Nitrogen has three electrons in three p orbitals, while oxygen has four. In nitrogen,

the 2p orbitals are half-filled (which makes the configuration particularly stable).

Oxygen’s fourth electron must pair with another electron, making it easier to remove

(and less stable). Exceptions for similar reasons occur for S and Se, directly below

oxygen in group 6A.

G) Trends in Second and Successive Ionization Energies

(1) Valence electrons are held much more loosely than core electrons therefore can be taken

much easier. It takes a tremendous amount of energy to remove a core electron, however.

(2) Here is a graph showing the successive values of ionization energies (this also shows how

many valence electrons an element has)

Chemistry Notes

28

VII) Electron affinities and metallic character – 3.8

(1) Electron affinity and metallic character also exhibit periodic trends. Electron affinity is a

measure of how easily an atom accepts an additional electron and is crucial to chemical

bonding because bonding involves the transfer of electrons.

(2) Electron affinity is the energy that is associated with the process of a gaseous atom

gaining an electron (units: kJ/mol).

(a) Equation X(g) + e- → X-(g) : ΔE=electron affinity

(b) Reflects how easily (or how hard it is) an atom gains electrons

(3) Metallic character is important because of the high proportion of metals in the periodic

table and the large role they play in our lives. Metallic character refers to the level of

reactivity of a metal. Metals tend to lose electrons in chemical reactions, as indicated by

their low ionization energies. Within a compound, metal atoms have relatively low

attraction for electrons, as indicated by their low electronegativities.

A) Electron Affinity

(1) Electron affinity (EA) of an atom or ion is the energy change associated with the gaining

of an electron by the atom in the gaseous state. Electron affinity is usually—though not

always—negative because an atom or ion usually releases energy when it gains an

electron (this process is exothermic and carries a negative sign).

(a) We can represent the electron affinity of chlorine with the equation:

(b) Cl (g) + 1 e— → Cl—- (g) EA = -349 kJ/mol

(2) This figure displays the electron affinities for a number of main-group elements. As we

can see from the figure, the trends in electron affinity are not as regular as trends in other

properties we have examined:

(5) Summarizing Electron Affinity for Main-Group Elements

(a) Most groups (columns) of the periodic table do not exhibit any definite trend in

electron affinity. Among the 1A metals, however, electron affinity becomes more

positive as we move down the column (adding an electron becomes less exothermic).

(b) Electron affinity becomes more negative (adding an electron becomes more

exothermic) as we move to the right across a period in the periodic table.

(c) Group 8A elements have the lowest electron affinities of all the elements

(d) Group 7A elements have the highest electron affinities of all the elements

(e) Similar trend to first ionization energy (except that this trend does not include the

noble gases and the trend in first ionization energy does).

Chemistry Notes

29

B) Metallic Character

(1) Metals are good conductors of heat and electricity; they can be pounded into flat sheets

(malleability); they can be drawn into wires (ductility); they are often shiny; and they

tend to lose electrons in chemical reactions

(2) Nonmetals have more varied physical properties; some are solids at room temperature,

others are gases, but in general nonmetals are typically poor conductors of heat and

electricity, and they all tend to gain electrons in chemical reactions.

(3) As we move to the right across a period in the periodic table, metallic character

decreases.

(4) As we move down a column in the periodic table, metallic character increases.

VIII) Periodic Trends Summary

(1) In this chapter, we have examined various trends in the periodic table that we can

understand in terms of electron configurations. Since electron configurations are a way of

specifying electronic structure, the trends in this chapter are a good example of the

overall theme of this book

(a) Structure determines properties

(2) We have seen how electronic structure determines the size, ionization energy, electron

affinity, and metallic character of atoms. Here is a summary table of these trends and the

reasons for them:

Property Trend moving

down a column

Reason for trend Trend moving

across a row

Reason for trend

Atomic Radii Increasing Size of outermost

occupied orbital

increases

Decreasing Effective nuclear

charge increases

First Ionization

Energy

Decreasing Outermost

electrons are

further away from

the nucleus (and

therefore are

easier to remove)

Increasing Effective nuclear

charge increases

Electron Affinity No definite trend Decreasing

(becoming more

negative)

Effective nuclear

charge increases

Metallic Character Increasing Ionization energy

decreases

Decreasing Ionization energy

increases

Chemistry Notes

30

) Molecules and Compounds – Chapter 4

I) Hydrogen, Oxygen, and Water – 4.1

A) Hydrogen gas is an explosive gas used as fuel in rocket engines. Oxygen gas is a natural

component of the air on Earth. Oxygen is not flammable but must be present for combustion.

Hydrogen and oxygen both have extremely low boiling points.

B) When hydrogen and oxygen come together to form H2O, a dramatically different substance

occurs. Water is a liquid, not gas, at room temperature. Water is not flammable but

responsible for extinguishing flames.

C) When two or more elements combine to form a compound, an entirely new substance results.

II) Types of Chemical Bonds – 4.2

A) A chemical compound is the force that holds atoms together in a compound.

(1) Chemical bonds form because they lower the potential energy of the charged particles

that compose atoms.

(2) The bond that forms between a metal and a non-metal is called an ionic bond

(a) Metals tend to lose electrons and nonmetals tend to gain them.

(b) When a metal interacts with a nonmetal, it can transfer one or more of its electrons to

the nonmetal. The metal atom becomes a cation and the nonmetal becomes an anion.

(c) These ions attract each other according to Coulomb’s law and form an ionic

compound, which in the solid state is composed of a lattice—a regular three-

dimensional array—of alternating cations and anions

(3) The bond that forms between two or more nonmetals is a covalent bond

(a) When a nonmetal bonds with another nonmetal, neither atom transfers electrons to

the other. Instead, the two atoms share some electrons. The shared electrons interact

with the nuclei of both bonding atoms, lowering their potential energy in accordance

with Coulomb’s law.

(i) Covalently bonded atoms form molecules, and the resulting compounds are

called molecular compounds

III) Representing Compounds: Chemical Formulas and Molecular Models – 4.3

A) The quickest and easiest way to represent a compound is with its chemical formula

(1) The chemical compound indicates the elements present in the compound and the relative

number of ions of each.

(a) the formula contains the symbol for each element and a subscript indicating the

relative number of atoms of the element.

(b) Chemical formulas usually list the more metallic element first, followed by the less

metallic (or more negatively charged element) element.

B) Types of Chemical Formulas

Chemistry Notes

31

(1) We categorize chemical formulas into three different types: empirical, molecular, and

structural

(a) An empirical formula indicates the relative number of atoms of each element in a

compound

(b) A molecular formula indicates the actual number of atoms of each element in a

molecule of a compound

(c) A structural formula uses lines to represent bonds and shows how atoms in a

molecule are bonded to each other.

(i) Two lines in a structural indicate a double bond which is generally stronger and

shorter than a single bond.

C) Molecular Models

(1) A molecular model is a more accurate and complete way to specify a compound

(a) A ball-and-stick molecular model represents atoms as balls and chemical bonds as

sticks; how the two connect reflects a molecule’s shape

(b) In a space-filling molecular model, atoms fill the space between each other to more

closely represent best estimates for how a molecule might appear if scaled to visible

size.

D) Noble gases have high 1st ionization energy and have very low electron affinity. Noble gases

have a stable electron arrangement. Except for helium, which has two valence electrons,

Noble gases have eight valence electrons.

(1) Need to know how to systematically name two types of compounds

(a) Ionic Compounds – compounds that form as a result of metals and non-metals

bonding

(b) Molecular Compounds – compounds which form as a result of covalent bonds

(sharing electrons)

(i) Molecular compounds come as a result of non-metals only.

IV) The Lewis Model: Representing Valence Electrons with Dots

A) Bonding theories are central to chemistry because they explain how atoms bond together to

form molecules. These theories explain why some atoms are stable and other atoms are not.

B) The Lewis model is introduced here named after the American Chemist G. N. Lewis (1875-

1946). In the Lewis model, we represent valence electrons as dots and we draw Lewis

electron-dot structures (or simply Lewis structures) to depict molecules.

C) In a Lewis symbol, we represent the valence electrons of main-group elements as dots

surrounding the abbreviation for the element

D) In the Lewis model, a chemical bond is the sharing of electrons to attain stable electron

configurations for the bonding atoms. Bonding atoms, whether ionic or covalent, obtain

stable electron configurations; since the stable electron configuration is usually eight

electrons in the outermost shell, this is known as the octet rule

E) The octet rule accounts for the success and longevity of the Lewis model because it uses a

practical approach that accurately predicts what we see in nature for a large number of

compounds.

V) Ionic Bonding: The Lewis Model and Lattice Energies

A) The basic unit of an ionic compound is the formula unit, the smallest, electrically neutral

collection of ions. NaCl is a formula unit for the molecules Na+ and Cl—.

B) Ionic Bonding and Electron Transfer

(1) The transfer of certain electrons to other atoms (such as the electron transferred between

K to Cl) gives all atoms an octet of valence electrons (in many cases).

C) Lattice Energy: The Rest of the Story

(1) The formation of an ionic compound from its constituent elements usually gives off quite

a bit of energy as heat.

Chemistry Notes

32

(a) For example, when one mole of sodium chloride forms from elemental sodium and

chlorine, 411 kJ of heat is evolved in the violent reaction.

(2) Where does the energy come from? In fact, the transfer of an electron from sodium to

chlorine—by itself—actually absorbs energy. The first ionization energy of sodium is

+496 kilojoules per mole, and the electron affinity of Cl is only -349 kilojoules per mole.

Based only on these energies, the reaction should absorb +147 kilojoules per mole. Why

is the reaction so exothermic?

(3) The answer lies in lattice energy—the energy associated with the formation of a

crystalline lattice of alternating cations and anions from the gaseous ions. Since the

sodium ions are positively charged and the chlorine ions are negatively charged, the

potential energy decreases—as described by Coulomb’s law—when these ions come

together to form a lattice.

D) Ionic Bonding: Models and Reality

(1) Ionic Bond – an electrostatic attraction (attraction between opposite charges)

(2) The Lewis model, when applied to ionic bonding, accounts, for the most part, for several

properties of atoms and molecules.

VI) Ionic Compounds: Formulas and Names – 4.6

A) Writing Formulas for Ionic Compounds

(1) Because ionic compounds are charge neutral and because many elements form only one

type of ion with a predictable charge, we can deduce the formulas for many ionic

compounds from their constituent elements.

(2) Summarizing Ionic Compound Formulas:

(a) Ionic compounds always contain positive and negative ions

(b) In a chemical formula, the sum of the charges of the positive ions (cations) must

equal the sum of the charges of the negative ions (anions)

(c) The formula of an ionic compound reflects the smallest whole-number ratio of ions.

(3) To write a formula for an ionic compound

(a) Write the symbol for the metal cation and its charge followed by the symbol for the

nonmetal anion and its charge. Determine charges from the element’s group number

in the periodic table

(b) Adjust the subscript on each cation and anion to balance the overall charge

(c) Check to make sure the sum of the charges of the cations equals the sum of the

charges of the anions

B) Naming Ionic Compounds

(1) Some ionic compounds have common names, which are nicknames of sorts learned by

familiarity. Chemists have also developed systematic names for different types of

compounds including ionic ones

(2) The first step in naming an ionic compound is identifying it as one. Remember, ionic

compounds are usually composed of metals and nonmetals.

(3) We categorize ionic compounds into two types, depending on the metal in the compound.

(a) The first type contains a metal whose charge is invariant from one compound to

another.

(b) Ionic compounds – metal and nonmetal (1. Metal forms only one type of ion), (2.

Metal forms more than one type ion)

Chemistry Notes

33

(c) The charges of the transition metals cannot be determined from their group number.

(d) The second type of ionic compound contains a metal with a charge that can differ in

different compounds. In other words, the metal in the second type of ionic compound

can form more than one type of cation. Therefore, we must specify its charge for a

given compound.

C) Naming Binary Ionic Compounds Containing

a Metal that Forms Only One Type of Cation

(1) Binary compounds contain only two

different elements. The names of binary

compounds take the form:

(a) [Name of cation (metal)][base name

of anion (nonmetal) + -“ide”]

D) Naming Binary Ionic Compounds Containing

a Metal That Forms More Than One Type of

Cation

(1) For these metals, the name of the cation

is followed by a roman numeral (in

parentheses), which indicates the charge

of the metal in that particular compound. For example, we distinguish between Fe2+ and

Fe3+ as follows: iron(II) and iron(III)

(2) The full names for compounds containing metals that form more than one kind of cation

have the form:

(a) [name of cation (metal)][charge of cation (metal) in roman numerals in

parentheses][base name of anion (nonmetal) + -“ide”]

(b) Example: CrBr3 chromium(III) bromide

Chemistry Notes

34

E) Naming Ionic Compounds Containing Polyatomic Ions

(1) Many common ionic compounds contain ions that are themselves composed of a group of

covalently bonded atoms with an overall charge. Hypochlorite is a polyatomic ion—an

ion composed of two or more atoms—with the formula ClO—.

(2) We name ionic compounds that contain a polyatomic ion in the same way that we name

other ionic compounds, except that we incorporate the name of the polyatomic ion

whenever it occurs.

(3) You should be able to recognize polyatomic ions in a chemical formula, so become

familiar with these ions in the table above. Most common polyatomic ions are oxyanions,

anions containing oxygen and another element.

(4) Notice that when a series of oxyanions contains different numbers of oxygen atoms, we

name them systematically according to the number of oxygen atoms in the ion. If there

are only two ions in the series, the one with more oxygen atoms has the ending -ate and

the one with fewer oxygen atoms has the ending -ite.

Chemistry Notes

35

(a) For example: NO3— is nitrate and NO2

— is nitrite

(5) Memorizing the “ate” ions:

(a) “Ate” suffix has a fixed number of oxygens with a charge (-2,-1,1, et al.)

(b) Nick the Camel ate a Clam Supper in Phoenix

(c) Number of vowels in acronym is the charge on the ion

(d) Number of consonants in acronymic word is number of oxygens

(i) N – nitrate

(ii) C – carbonate

(iii) Cl – chlorate

(iv) S – sulfate

(v) Ph – phosphate

(6) If there are more than two ions in the series, we use the prefixes hypo- meaning less than,

and per- meaning more than. So ClO— is hypochlorite—less oxygen than chlorite and

ClO4 is perchlorate—more oxygen than chlorate

(a) examples:

(i) ClO— hypochlorite

(ii) ClO2— chorite

(iii) ClO3— chlorate

(iv) ClO4— perchlorate

(7) Examples of naming ionic compounds:

(a) Copper(II) phosphide Cu2+ + P3- → Cu3P2

(b) Calcium cyanide Ca2+ + CN- → Ca(CN)2

(c) Vanadium(V) nitrite V5+ + NO2- → V(NO2)5

(d) Zinc fluoride Zn2+ + F- → ZnF2

(e) Nickel(III) hydroxide Ni3+ + OH- → Ni(OH)3

(f) Lead(IV) selenide Pb4+ + Se2- → PbSe2

(g) Barium sulfate Ba2+ + SO42- → BaSO4

(h) TiS2 titanium(IV) sulfide

(i) BaCrO4 Barium chromate

(j) NaC2H3O2 sodium acetate

(k) AlH3** [correct answer: aluminum hydride]

(l) (NH4)3PO3 ammonium phosphite

(m) Fe2(CO3)3 iron(III) carbonate

(8) Examples of naming binary molecular compounds

(a) CO2 carbon dioxide

(b) H2O water

(c) P4S10 tetraphosphorous decasulfide

(d) ClO2 chlorine dioxide

(e) N2H5 dinitrogen pentahydride

(f) Dinitrogen monoxide N2O

(g) Silicon tetraiodide SiI4

(h) Arsenic tribromide AsBr3

F) Hydrated Ionic Compounds

(1) Some ionic compounds called hydrates contain a specific number of water molecules

associated with each formula unit

(2) We name hydrates like we name other ionic compounds but we give them the additional

name “prefixhydrate” where prefix indicates the number of water molecules associated

with the formula unit.

G) Types of anions

(1) Monatomic – composed of only 1 type of element. The charge comes from the position

on the periodic table

Chemistry Notes

36

(a) Names include base name with -ide suffix + ion

(i) Example: Chloride ion (Cl-)

H) Formulas of Compounds and Calculations

(1) Compounds

(a) 1 mol = 6.022x1023 particles of a substance

(b) In naming compounds, you would say that you have a mol of molecules

(c) When one sums up all of the atomic masses of the elements in a compound,

something may be found which is called one of these things:

(i) Molar mass

(ii) Molecular weight

(iii) Formula weight

(d) For example, what is the molar mass (MM) of silicon tetraiodide?

(i) Solution: add MM Si and MM I x 4 = 535.69 amu or g/mol

VII) Covalent Bonding: Simple Lewis Structures

A) Single covalent bonds

(1) A shared pair of electrons is called a bonding pair, while a pair that is associated with

only one atom—and therefore is not involved in bonding—is called a lone pair. Lone pair

electrons are also called nonbonding electrons

(2) Common diatomic elements (elements that exist as diatomic molecules): H, N, O, F, Cl,

Br, I

(3) In naming diatomic elements, simply use the chemical name for the name. For example,

F2 should be named as “fluorine”

B) Double and Triple Covalent Bonds

(1) When two atoms share two electron pairs, the resulting bond is a double bond

(2) Triple bonds are shorter and stronger than double bonds

C) Covalent Bonding: Models and Reality

(1) The Lewis model accurately predicts some properties of bonds

VIII) Molecular Compounds: Formulas and Names

A) In contrast to anionic compound, the formula for a molecular compound cannot always be

determined from its constituent elements because the same combination of elements may

form several different molecular compounds, each with a different formula

B) The first step in naming a molecular compound is identifying it as one. Remember, molecular

compounds are composed of two or more nonmetals. In this section, we learn how to name

binary molecular compounds. Their names have the form:

(1) Prefix – name of first element – prefix – base name of 2nd element + -“ide”

C) When we write the name of a molecular compound, just as when we write its formula, the

first element is the more metal-like one (most toward the bottom and left of the periodic

table).

(1) Generally, we write the name of the element with the smallest group number first.

(2) If the two elements lie in the same group, we write the greatest row number first. I

(3) The prefixes given to each element indicate the number of atoms present:

Mono = 1 Tri = 3 Penta = 5 Hepta = 7 Nona = 9

Di = 2 Tetra = 4 Hexa = 6 Octa = 8 Deca = 10

IX) Skills needed for Exam #3

A) Given a molecular formula, you should be able to

(1) Generate a Lewis Structure and be able to obey the octet rule, minimize formal charges,

and know when it is acceptable to disobey the octet rule

(2) Identify and draw resonance structures

(3) Draw a three-dimensional representation with dashes and wedges as appropriate

(4) Identify electron geometry, molecular geometry, and hybridization for any atom in a

molecule

Chemistry Notes

37

(5) Predict any bond angle and distortion from ideal ____ to lone pair

(6) Identify and label bond polarities accurately

(7) Determine if the net dipole moment exists (vector sum of bond polarities)

(8) Identify σ vs. π bonds

B) The test will be from Chapter 5 through Chapter 6, Section 3

X) Lecture Notes for 10/31/17

A) Hybridization of central atoms in molecules

(1) Linear – sp – the hybridization of one s and one p orbital produce two hybrid orbitals

oriented 180° apart

(2) Trigonal planar – sp2 - the hybridization of one s and two p orbitals produce three hybrid

orbitals oriented 120° from each other all in the same plane.

(3) Tetrahedral – sp3 - the hybridization of one s and three p orbitals produce four hybrid

orbitals oriented toward the points of a regular tetrahedron, 109.5° apart.

(4) Trigonal bipyramidal – dsp3 or sp3d2 - the hybridization of one s , three p , and one d

orbitals produce five hybrid orbitals oriented in this weird shape: three equatorial hybrid

orbitals oriented 120° from each other all in the same plane and two axial orbitals

oriented 180° apart, orthogonal to the equatorial orbitals.

(5) Octahedral – d2sp3 or sp3d2 - the hybridization of one s , three p , and two d orbitals

produce six hybrid orbitals oriented toward the points of a regular octahedron 90° apart.

B) Chapter 7 – Intro to chemical reactions

(1) Predicting and balancing reactions, stoichiometry

(a) Stoichiometry: shows the molar ratio between chemical species in a balanced

chemical equation

(2) Yields and percent yields

(3) Chemical reaction – used to show a chemical change

(a) 2 Mg (s) + 1 O2 (g) → 2 MgO (s)

Stoichiometric coefficients

(4) Experiment performed*

(5) If I burn 0.13g Mg in excess O2, how much MgO should I generate?

(a) Mass Mg→mol Mg→mol MGO→mass MgO

(b) **theoretical yield calculated from above equation to get theoretical yield of 0.22g

MgO

(c) **magnesium burned and weighed to get 0.18g MgO

(d) Actual yield=0.18g

(e) % 𝑦𝑖𝑒𝑙𝑑 =𝑎𝑐𝑡𝑢𝑎𝑙 𝑦𝑖𝑒𝑙𝑑

(𝑡ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙 𝑦𝑖𝑒𝑙𝑑)𝑥 100%

(f) Applying above equation, 0.18g/0.22g x 100% = 82%

(6) Did propane and detergent bubbles experiment** grabbed bubbles and set them on fire

and fire came off of his hands but they did not burn

(a) Chemical equation to balance C3H8 (g) + 5O2 (g) → 3CO2 (g) + 4H2O (g)

(7) Predict a balanced chemical equation for the combustion of octane (C8H18)

(a) C8H18 (l) + 25/2 CO2 + 8H2O 9(g)

XI) Formula Mass and the Mole Concept for Compounds – 4.9

A) Did not write any notes for 4.9

) Chemical Bonding I: Drawing Lewis Structures and Determining Molecular Shapes – Chapter 5

I) Morphine: a molecular imposter – 5.1

A) Morphine is very effective against pain. Morphine is derived from the opium poppy. The

opium sap has been used for thousands of years. Scientists have discovered that humans have

natural morphine in their body called endorphins (short for endogenous morphine).

Chemistry Notes

38

B) Morphine works like a key and lock onto the ends of our nerve receptors allowing pain to be

blocked. In this chapter, we look at ways to predict and account for the shapes of molecules.

The molecules we examine are much smaller than the protein molecules we just discussed,

but the same principles apply to both.

C) The simple model that we examine to account for molecular shape is the valence shell

electron pair repulsion theory (VSEPR). We use this theory in conjunction with the Lewis

model to account for molecular shape.

II) Bonding Basics

A) Electronic stability

(1) Noble gases have stable electron arrangements; therefore,

(a) Noble gases have very high first Ionization Energy, and

(b) Electron affinities are very low

(c) They have 8 valence electrons (except He)

(d) The quantum number configuration for noble gases looks like ns2np6

(i) This is known as a stable octet

(2) Noble gases were not found in compounds until 1967

(3) Atoms bond to become electronically stable

B) Electronegativity

(1) Definition – the ability of an atom to attract electrons, specifically within a bond

(2) Measured from low values to high values of 0 – 4.0

C) Energy of bond formation

(1) Bond formation involves the release of energy (always the case)

(2) Atoms bond to form stable electron arrangements

D) Lewis Structures or Lewis Symbols

(1) X (nucleus + Core e-): valence electrons shown as dots around the symbol

(a) Example: B (three dots around Boron symbol)

III) Types of bonds

A) Metallic bond

(1) Hold metal atoms together

(a) Sea of electron theory – electrons aren’t specifically associated with one atom; they

can move freely; definition: valence electrons move amongst atoms; individual atoms

are stable for short periods of time

B) Ionic bonds

(1) Formed when electrons are transferred between atoms.

(2) One atom has very high electronegativity and one atom has low electronegativity

(3) Metals and nonmetals

(4) Example: NaCl (sodium becomes a cation while chlorine is an anion)

(5) Put Lewis symbol in square brackets ( [Na]+ ) and charge on the outside

C) Non-polar covalent bonds

(1) Occur when two atoms share electrons equally

(2) Formed when atoms have the same electronegativities

(a) Two non-metals of the same type

(b) Any diatomic elements have non-polar covalent bonds such as: I2, Br2, Cl2, F2, O2,

N2, and H2

(c) These types of bonds are not common

D) Polar covalent bond

(1) These types of bonds are the most prevalent on Earth

(2) In these types of bonds, atoms share electrons unequally

(3) These occur between non-metals of different types

(4) Example: water ( H2O)

Chemistry Notes

39

(a) In these types of bonds, atoms with larger electronegativities (Oxygen in this case)

attract more electrons in the bonds than the atoms with the lower electronegativities

(hydrogen in this case).

(b) Atom with high electronegativity develops partial negative charge

(i) For example, in a water molecule, an oxygen atom has partial negative charge

and hydrogen atom develops a partial positive charge.

(c) Polarity of bond is shown with a small case Greek letter delta and a “+” or “-“

symbol such as H—Oδ-

(d) Polar – charge separation exists positive and negative areas

(e) Dipole arrow shows polarity pointing from partially positive charge to more

electronegative atom and looks like the arrow below in IV.A

IV) Electronegativity and Bond Polarity – 5.2

A) From the Lewis model we know if an atom has a slight positive or negative charge in a

molecule. We use an arrow with a “+” sign at one end and a tail at the other end to represent

this: “+ ”

(1) A polar covalent bond is intermediate in nature between a pure covalent bond and an

ionic bond. In fact, the categories of pure covalent and ionic are really two extremes

within a broad continuum. Most covalent bonds between dissimilar atoms are actually

polar covalent, somewhere between the two extremes

B) Electronegativity

(1) The ability of an atom to attract electrons to itself in a chemical bond (which results in

polar and ionic bonds) is electronegativity.

(2) Here are some electronegativity general trends for main-group elements:

(a) Electronegativity generally increases across a period in the periodic table.

(b) Electronegativity generally decreases down a column in the periodic table.

(c) Fluorine is the most electronegative element.

(d) Francium is the least electronegative element (sometimes called the most

electropositive)

(3) Trends in electronegativity:

C) Bond Polarity, Dipole Moment, and Percent Ionic Character

(1) The degree of polarity in a chemical bond depends on the electronegativity difference

between the two bonding atoms. The greater the electronegativity difference, the more

polar the bond. If two atoms with identical electronegativities form a covalent bond, they

share the electrons equally and the bond is purely covalent or nonpolar

Chemistry Notes

40

V) Writing Lewis Structures for Molecular Compounds and Polyatomic Ions – 5.3

VI) Lewis Structures of molecules and polyatomic ions

A) Prediction of bonding

(1) Bond length and bond strength

(2) Reactive sites of a molecule

(3) Shapes of molecules and polarity of molecules

B) Steps in writing a Lewis Structure

(1) Determine total number of valence electrons

(a) Cation – subtract valence electrons based on the charge

(b) Anion – add valence electrons based on the charge

(2) Determine the skeletal structure of the molecule or ion

(a) The least electronegative atom goes in the center

(b) Usually written first in a chemical formula of a compound or ion

(3) Hydrogen may never be a central atom

(a) Hydrogen may only make one bond preventing it from being in the center

(4) Examples:

(a) NH3

(b) CH2O

(c) CN—

(5) Place single bonds in all locations where atoms are attached a single dash represents two

electrons

(6) Make exterior atoms stable with lone pairs of electrons

(7) Make the central atom stable

(a) Add lone pairs

(b) You may only do this if there are still electrons to account for

(8) Move electron pairs to create multiple bonds

(9) Polyatomic ions are enclosed in square brackets with the charge shown outside of the

square brackets

(10) Check to make sure that every atom is stable

(11) The number of electrons in the Lewis Structure must equal the total number of

valence electrons calculated

VII) Formal Charge

A) Formal charge is a tool used to evaluate Lewis Structures

B) Definition: HYPOTHETICAL, NOT TRUE, charge assigned to all atoms in a Lewis structure

(1) The difference between the number of valence electrons in an atom when it is free and

the number of electrons that you associate with that atom in a Lewis Structure

(2) This is under the assumption that all atoms are equally electronegative

C) Formula: Formal Charge = valence electrons – [all lone electrons + ½ bonds on an atom]

Chemistry Notes

41

D) Sum of all formal charges must equal the overall charge on the atom or ion for which a Lewis

Structure was made

E) The best Lewis Structure

(1) The best structure has as many atoms as possible which have formal charges FC=0

(2) Negative formal charge on the most electronegative atom

F) In Lewis Structures, include the formal charges for atoms that have nonzero formal charge

(put it near the atom it represents and circle it)

G) Here are the steps for writing a Lewis structure for a molecular compound

(1) Write the correct skeletal structure for the molecule.

(a) Hydrogen atoms are always terminal

(b) More electronegative atoms should be in the central positions

(2) Calculate the total number of electrons for the Lewis structure by summing the valence

electrons of each atom in the molecule

(a) If you are writing the Lewis structure for a polyatomic ion, you must consider the

charge of the ion when calculating the total number of electrons

(3) Distribute the electrons among the atoms, giving octets (or duets in the case of hydrogen)

to as many atoms as possible

(4) If any atoms lack an octet, form double or triple bonds as necessary to give them octets.

(5) Check

(a) Check for correct number of electrons

(b) Check that everything obeys the octet and duet rule

H) Writing Lewis Structures for Polyatomic Ions

(1) Write Lewis structures for polyatomic ions by following the same procedure, but pay

special attention to the charge of the ion when calculating the number of electrons for the

Lewis structure.

(2) Add one electron for a negative charge and subtract one for a positive charge

(3) Write the Lewis structure for a polyatomic ion within brackets with the charge of the ion

in the upper right-hand corner outside the bracket

VIII) Resonance and Formal Charge – 5.4

A) We need two additional concepts to write the best possible Lewis structures for a large

number of compounds. The concepts are resonance and formal charge

B) Resonance

(1) For some molecules we can write more than one valid Lewis structure.

(a) A resonance structure is one of two or more Lewis structures that have the same

skeletal formula but different electron arrangements. The actual structure of the

molecule is intermediate between the two (or more) resonance structures and is called

a resonance hybrid.

C) Formal Charge

(1) Formal charge is a fictitious charge assigned to each atom in a Lewis structure that helps

us to distinguish amongst Lewis structures.

(2) The formal charge of an atom in a Lewis structure is the charge it would have if all

bonding electrons were shared equally between the bonded atoms. In other words, formal

charge is the calculated charge for an atom in a molecule if we completely ignore the

effects of electronegativity.

(3) 𝐹𝑜𝑟𝑚𝑎𝑙 𝐶ℎ𝑎𝑟𝑔𝑒 = # 𝑣𝑎𝑙𝑒𝑛𝑐𝑒 𝑒𝑙. −(# 𝑛𝑜𝑛𝑏𝑜𝑛𝑑𝑖𝑛𝑔 𝑒𝑙. +1

2 # bonding el.)

(4) The concept of formal charge is useful because it can help us distinguish between

competing skeletal structures or competing resonance structures. In general, these four

rules apply:

(a) The sum of all formal charges in a neutral molecule must be zero

(b) The sum of all formal charges in an ion must equal the charge of the ion.

Chemistry Notes

42

(c) Small (or zero) formal charges on individual atoms are better than large ones.

(d) When formal charge cannot be avoided, negative formal charge should reside on the

most electronegative atom.

A) Bond Length

(1) Distance between two nuclei involved in covalent bonding

(2) More pairs of shared electrons between atoms results in shorter bonds

B) Bond Order – a concept used to characterize the length and strength of bonds

(1) Definition: the number of shared pairs of electrons between two atoms

(2) Bond order and bond length are inversely related

BOND ORDER LENGTH STRENGTH ENERGY

1 Single Longest Weakest Lowest

2 Double Shorter Stronger Higher

3 Triple Shortest Strongest Highest

C) Bond Strength

(1) Bond enthalpy – amount of energy that is needed to break covalent bonds and leave free

(unbonded) atoms at the end (enthalpy relates to energy).

(2) When atoms form bonds, energy is released (recall that bonds form to make atoms stable)

(3) When atoms break bonds, energy is gained

(4) Bond enthalpy is measured and averaged over many molecules in the gaseous state.

(5) Higher bond order results in greater bond enthalpy (which means stronger bonds)

(6) Bond enthalpy values are NOT additive quantities

(7) C – C 347 kJ/mol

(8) C double C is NOT 2(347 kJ/mol)

D) Resonance structures – equivalent Lewis Structures in terms of formal charge for a specific

molecule or ion.

(1) No one resonance structure accurately shows the molecule itself

(2) The actual molecule is a hybrid of all of the possible resonance structures

(3) To show resonance, we will use a double-sided arrow

(4) The actual molecular structure is a hybrid of the two resonance structures. It is useful to

think of the actual structure as an average of the two

(5) Resonance structures can be used to calculate the average bond order

(a) To do this:

(i) Pick a bond

(ii) Determine the order

(iii) Compare the equivalent Lewis Structures

(iv) Take the average of them to find bond order

(6) Resonance structures are only possible when you have double or triple bonds which can

be moved from one place to another

(a) Not all molecules with double bonds or triple bonds necessarily have resonance

(CH2O for example)

(7) Example: all of the bonds in ozone O3 are the same length and strength

(8)

II) Exceptions to the Octet Rule: Odd-Electron Species, Incomplete Octets, and Expanded Octets –

5.5

A) Exceptions to the octet rule

(1) H, He, Li, Be (Lithium and beryllium are stable with two electrons only. We don’t have

to worry about lithium or beryllium because they are metals)

(2) Boron is stable with six or eight electrons around it

(3) Radicals (or Free Radicals) are molecules which have an odd number of electrons and

thus cannot obey the octet rule.

(4) Radicals are very reactive species

Chemistry Notes

43

(5) Biggest exceptions to the octet rule: Expanded octets (more than 8 electrons around an

atom)

(a) Expanded octets occur when atoms use d orbitals to hold bonding electrons

(b) These orbitals first appear in the third energy level (n=3)

(c) Any atom in Period 3 or beyond (4, 5, etc…) can have an expanded octet

(6) Atoms that must obey the octet rule (unless there is a radical to be formed):

(a) Carbon (unless found in a radical)

(b) Nitrogen (..)

(c) Oxygen (..)

(d) Fluorine (..)

(e) Neon (..)

B) Odd-Electron Species

(1) Molecules with an odd number of electrons in their Lewis structures (used now as LS)

are free radicals

(2) We can’t write good LS for free radicals

(3) For free radicals, we try simply to write the best Lewis structure that we can

(4) This is a LS which minimizes formal charge

C) Incomplete Octets

(1) Another exception to the octet rule involves those elements that tend to form incomplete

octets. The most important of these is boron, which forms compounds with only six

electrons, rather than eight.

(2) With these it is generally better to accept an incomplete octet rather than a formal charge

on more electronegative elements.

D) Expanded octets

(1) Elements in the third row of the periodic table and beyond exhibit expanded octets of up

to 12 and occasionally 14 electrons.

III) Bond Energies and Bond Lengths – 5.6

A) Bond Energy

(1) The bond energy of a chemical bond is the energy required to break 1 mole of the bond in

the gaseous phase.

(2) Generally for atoms bonded by a double bond, those same atoms bonded by a double or

triple bond have a higher bond energy

B) Bond length

(1) Just as we can tabulate average bond energies, which represent the average energy of a

bond between two particular atoms in a large number of compounds, we can tabulate

average bond lengths.

(2) The average bond length represents the average length of a bond between two particular

atoms in a large number of compounds.

(a) Like bond energies, bond lengths depend not only on the kind of atoms involved in

the bond, but also on the type of bond: single, double, or triple.

(b) In general, for a particular pair of atoms, triple bonds are shorter than double bonds,

which are in turn shorter than single bonds.

IV) VSEPR Theory: The Five Basic Shapes – 5.7

A) Electron group – any pair of electrons in a lone pair, any set of electrons found in double,

single, or triple bonds, or any lone electron found in a free radical

(1) Any electron group takes up a single region of space around an atom

B) Electron geometry

(1) Shape generated by the electron groups around an atom

C) Valence shell electron pair repultion (VSEPR) theory is based on the simple idea that

electron groups which we define as: lone pairs, single bonds, multiple bonds, and even single

electrons—repel one another through coulombic forces.

Chemistry Notes

44

(1) The electron groups, of course, are also attracted to the nucleus (otherwise the molecule

would fall apart), but VSEPR theory focuses on the repulsions.

(2) According to VSEPR theory, the repulsions between electron groups on interior atoms of

a molecule determine the geometry of the molecule

D) Linear Geometries:

V) VSEPR Theory: The effect of Lone Pairs – 5.8

A) For every lone pair, there is usually a 2.5° decrease in bond angle between atoms in a

molecule

B) Summarizing VSEPR Theory:

(1) The geometry of a molecule is determined by the number of electron groups on the

central atom (or on each interior atoms, if there is more than).

(2) The number of electron groups is determined from the Lewis structure of the molecule. If

the LS contains resonance structures, we can use any one of the resonance structures to

determine the electron groups

(3) Each of the following counts as a single electron group:

(a) A lone pair

(b) A single bond

(c) A double bond

(d) A triple bond

(e) A single electron

(4) In general, electron group repulsions vary in relative ordering of repulsions as follows:

(a) Lone pair—lone pair > Lone pair—bonding pair > bonding pair—bonding pair

Chemistry Notes

45

(5) Bond angles can vary from the idealized angles because double and triple bonds occupy

more space than single bonds (they are bulkier even though they are shorter), and lone

pairs occupy more space than bonding groups. The presence of lone pairs usually makes

bond angles smaller than the ideal angle for the particular geometry.

VI) VSEPR Theory: Predicting Molecular Geometries – 5.9

VII) Molecular Shape and Polarity – 5.10

A) Polarity in Diatomic Molecules

(1) If a diatomic molecule has a polar bond, the molecule as a whole is polar.

(2) If there is a net dipole moment, the molecule has a polar bond

(3) To determine if there is a net dipole moment in a molecule, find out if one atom is more

electronegative than another. If it is, then there should be a dashed arrow connecting the

two atoms with the head pointing towards the more electronegative atom.

B) Summarizing Determining Molecular Shape and Polarity:

(1) Draw the LS for the molecule and determine its molecular geometry

Chemistry Notes

46

(2) Determine if the molecule contains polar bonds. A bond is polar if the two bonding atoms

have sufficiently different electronegativities. If the molecule contains polar bonds,

superimpose a vector, pointing toward the more electronegative atom, on each bond.

Make the length of the vector proportional to the electronegativity difference between the

bonding atoms.

(3) Determine if the polar bonds add together to form a net dipole moment. Sum the vectors

corresponding to the polar bonds together. If the vectors sum to zero, the molecule is

nonpolar. If the vectors sum to a net vector, the molecule is polar.

C) Molecular polarity

(1) Polar: charge separation

(2) Molecular polarity comes from EN of atoms and the arrangement of atoms (molecular

geometry) in a molecule

(3) Polar molecules or non-polar molecules

(4) To determine if a molecule is polar without thinking about vectors, draw a 3D

representation of the molecule:

(a) Determine the physical centers of the most and the least electronegative atoms in the

molecule

(b) If the centers are in different physical spots, then the molecule is polar; label the

polarity with a dipole arrow

(c) If the centers of the atoms are in the same location, then the molecule is nonpolar;

thus, draw no polarity arrow (a dipole arrow points in the direction of the

electronegativity and the other side has a plus sign attached to it by the least

electronegative atom)

(5) Example determining polarity: BH3 is a non-polar molecule, CH2O is a polar molecule,

) Chapter 6: Chemical Bonding II—Valence Bond Theory and Molecular Orbital Theory

I) Oxygen: A Magnetic Liquid – 6.1

A) Oxygen is magnetic. If something is magnetic, we know that it contains unpaired electrons

which orient towards a magnetic field when applied to a molecule or atom. The Lewis

Structure for oxygen gas fails to predict the correct magnetic properties of oxygen but the

molecular orbital theory—the most sophisticated and accurate bonding theory—correctly

predicts the magnetic properties of oxygen. When we apply molecular orbital theory to the O2

molecule, it shows that O2 contains two unpaired electrons. These unpaired electrons are

responsible for the magnetic behavior of oxygen. We will also examine bonding models that

apply to metals and semiconductors in this chapter.

II) Valence Bond Theory: Orbital Overlap as a Chemical Bond

A) In valence bond theory, a chemical bond is the overlap between two half-filled atomic

orbitals (AOs).

B) In some cases, these are s, p, d, and f orbitals. In other cases, these are hybridized atomic

orbitals, a kind of blend or combination of two or more standard atomic orbitals.

C) Summarizing Valence Bond Theory:

(1) The valence electrons of the atoms in a molecule reside in quantum-mechanical atomic

orbitals. The orbitals can be s, p, d, and f orbitals, or they may be hybrid combinations of

these.

(2) A chemical bond results from the overlap of two half-filled orbitals and spin-pairing of

the two valence electrons (or less commonly the overlap of a completely filled orbital

with an empty orbital)

(3) The geometry of the overlapping orbitals determines the shape of the molecule

III) Valence Bond Theory: Hybridization of Atomic Orbitals

A) Valence Bond theory accounts for the bonding in many polyatomic molecules by

incorporating orbital hybridization.

Chemistry Notes

47

(1) Hybridization is a mathematical procedure in which the standard atomic orbitals are

combined to form new atomic orbitals called hybrid orbitals that correspond more closely

to the actual distribution of electrons in chemically bonded atoms. Hybrid orbitals are still

localized on individual atoms, but they have different shapes and energies from those of

standard atomic orbitals.

B) Here are some general statements regarding hybridization:

(1) The number of standard atomic orbitals added together always equals the number of

hybrid orbitals formed The total number of orbitals is conserved.

(2) The particular combinations of standard atomic orbitals added together determine the

shapes and energies of the hybrid orbitals formed

(3) The particular type of hybridization that occurs is the one that yields the lowest overall

energy for the molecule. Since actual energy calculations are beyond the scope of this

book we use electron geometries as determined by VSEPR theory to predict the type of

hybridization.

C) Valence bond theory – bonds are created when orbitals holding one electron overlap

(1) When atoms bond to one another, the orbitals used to form bonds are called “hybrid

orbitals.” These are no longer called “atomic orbitals.”

(2) Hybrid orbitals are directly linked to electron geometry

D) Explanation of different hybrids:

(1) 2 electron groups – sp hybrid orbitals

(a) Combination (of shape and energy) of an s and one p orbital

(2) 3 electron groups – sp2 hybrid orbitals

(a) Combination of an s and two p orbitals

(3) 4 electron groups – sp3 hybrid orbitals (probably the most common set of hybrid orbitals

in nature)

(a) Combination of one s and three p orbitals

(4) 5 electron groups – sp3d hybrid orbitals

(a) Combination of an s three p orbitals, and one d orbital

(5) 6 electron groups – sp3d2 hybrid orbitals

(a) Combination of an s three p orbitals, and two d orbitals

E) Note:

(1) CH4 can only be made using sp3 orbitals on the C atom

(2) Hybridization is the mathematical averaging of orbitals. When hybridization happens,

orbitals are always conserved (not lost or gained). There simply amount a certain number

of hybrid orbitals and the remaining orbitals exist as they were before.

(3) Any lone pair in a Lewis structure must occur in a hybrid orbital

(4) When there are multiple bonds, then one pair of shared electrons must occur in a hybrid

orbital. The remaining pairs will be in unhybridized atomic orbitals

F) Hybridization Chart:

Chemistry Notes

48

Hybridization Scheme Hybrid orbitals Unhybridized orbitals Total Orbitals

Sp 2 2 (p orbital) 4

Sp2 3 1 (p orbital) 4

Sp3 4 0 4

Sp3d 5 4 (d orbitals) 9

Sp3d2 6 3 (d orbitals) 9

Chemistry Notes

49

G) Sigma (σ) bond

(1) A sigma bond is a bond that occurs when electrons are found directly in between two

nuclei of atoms (at the midpoint) – theoretically the sigma bond cannot be “seen”

(2) Single, unpaired electrons must occur in hybrid orbitals. This is the first way that atoms

link together covalently – sigma bonds are very strong bonds

H) Pi (π) bonds

(1) Electrons are found outside the internuclear distance – theoretically, pi bonds can be

“seen”

(2) Pi bonds only happen when there are single, unpaired electrons in unhybridized orbitals

(a) Pi bonds are always found in double and triple covalent bonds

I) Pi and Sigma bonds

(1) Pi bonds have higher energy than sigma bonds and are the first to be changed when

reactions occur

J) Sigma and pi bonds

(1) Single bond = 1 σ

(2) Double bond = 1 σ + 1 π

(3) Triple bond = 1 σ + 2 π

IV) Chapter 7 – Solution Chemistry

A) Chemical reactions are a series of symbols that convey a chemical change

(1) Reactants → (the arrow is a “yields” sign) Products

(2) Show the reactants and products in a chemical equation using correct chemical formulas

(3) Show the physical states of reactants and products

(a) Solid (s)

(b) Liquid (l)

(c) Gas (g)

(d) Aqueous (aq) – the substance that is dissolved in water (H2O)

(4) Coefficients are shown in chemical reactions. Coefficients indicate how many molecules

are needed to satisfy the law of conservation of matter to be met

B) Example: phosphoric acid (H3PO4) and barium hydroxide solutions react to form solid barium

phosphate and water. Write a balanced chemical equation.

C) Example: ethanol (C2H6O) combusts in the presence of oxygen to form gaseous carbon

dioxide and water vapor

D) Balancing chemical equations:

(1) Use pencil and eraser for the coefficients

(2) Begin with any atom which is NOT Hydrogen or Oxygen

(3) Look for the easiest type of atom to balance

(4) Find the simplest whole-number ratio of the chemical equation

(5) “Ones” are understood to exist when there is no coefficient

E) Coefficients in a balanced equation

(1) Coefficients indicate the ratio of reactant and product molecules involved in a reaction

(2) Coefficients also indicate the ratio of moles on each side of the equation

(3) Mole ratio is a conversion factor that relates any two compounds or elements that are

involved in a chemical reaction. The coefficients in front of a mole ratio are the number

entities from the balanced chemical reaction

F) Combustion reactions – complete combustion

(1) Fuel (compound with carbon and hydrogen)

(2) Oxygen

(3) Example combustion reaction = Fuel + O2 (g) → CO2 (g) + H2O (g)

G) Limiting Reagents – Any reactant that is completely consumed during a chemical process

(1) Limiting reagents are responsible for creating the smallest amount of product in a

chemical reaction (in any unit of measure)

Chemistry Notes

50

(2) To find the limiting reagent:

(a) Calculate the amount of A product which could be generated from each staring

material (from each reactant)

(3) The limiting reagent makes the smallest amount of product

H) Theoretical Yield – the amount of product (smallest amount of product) that is possible from

any chemical process. The theoretical yield comes from a limiting reagent calculation

I) Percent yield – percentage relationship between the actual yield in a reaction and the

theoretical yield in a reaction times 100%

(1) %Yield=actual yield x 100%/theoretical yield

(2) The actual yield in a real life reaction is less than the theoretical yield and the percent

yield is always below 100%

V) Excess Reagents – reactants which are completely consumed during a chemical process

A) Calculating the amount of excess reagent:

(1) Calculate the amount of excess reagent used in a reaction

(2) Subtract: (initial amount) – (amount used in reaction) = amount of excess reagent

) Lecture Notes 11/7/17

I) Balancing Chemical Equations (problem in Tro Textbook)

A) Fe3O3(s) + CO(g) → Fe(s) + CO2(g) – unbalanced chemical equation

22.55g 14.78g

B) Fe3O3(s) + 3CO(g) → 2Fe(s) + 3CO2(g) – balanced chemical equation

(1) How much excess reactant?

(2) Find the theoretical yield of Fe(s) in grams.

(a) 2Fe + 3O = (2*55.8452 g/mol)+(3*15.999 g/mol)→ 159.687 g/mol Fe3O3

(b) 22.55g Fe3O3 x (1mol Fe3O3/158.687 g) = 0.1412 mol Fe3O3

(c) 14.78g CO x (1mol CO/28.010g) = 0.5277 mol CO

(d) 0.1412 mol Fe3O3 x (3mol CO/1 mol Fe3O3) = 0.4236 mol CO needed

(e) 0.4236 mol CO are needed, 0.5277mol CO are available

(f) Therefore, Fe3O3 is the limiting reagent

(g) Excess CO = (0.5277 mol available – 0.4236 mol needed) = 0.1041 mol excess CO

(h) 0.1041 mol CO x (28.010g/1mol) = 2.916g CO in excess

(i) 0.1412 mol Fe3O3 x (2mol Fe/1mol Fe3O3) = 0.2824 mol Fe

(j) 0.2824 mol Fe x (55.8452g Fe/1mol) = 15.77g Fe

II) Chapter 8: Solution Chemistry

A) Solution: solute dissolved in a solvent

B) Concentration of solute in a solution is quantified usually with molarity (M)

(1) 𝑀𝑜𝑙𝑎𝑟𝑖𝑡𝑦 ≡𝑚𝑜𝑙

𝐿

(2) Molarity is an intensive property of a solution

C) Practice example: 0.25g NaCl dissolve to form a 500.0 milliliter solution. What is the

concentration in moles/L of that solution?

D) Solutions – mixtures of matter

(1) Can be multiple Solute(s) – elements or compounds found in smaller quantities (than

solvent)

(2) Usually only one solvent – elements or compounds found in a solution that is present in

the largest quantity

(3) Aqueous solution – the solvent is water

(a) Example NaCl (aq)

E) Dilution – changing the amount of solvent in a solution to reduce the overall concentration of

the solution. In a solution, the number of moles of a substance remain the same

(1) Equation for dilution: 𝑀𝑐𝑜𝑛𝑐 ∗ 𝐿𝑐𝑜𝑛𝑐 = 𝑀𝑑𝑖𝑙 ∗ 𝐿𝑑𝑖𝑙 F) MW(NaCl)= 22.98976 g/mol + 35.4532 g/mol = 58.4430 g/mol

Chemistry Notes

51

G) 0.25g NaCl x (1mol NaCl/58.4430g NaCl) = 4.2 x 10-3 mol NaCl

H) Concentration of NaCl solution = [NaCl(aq)] = 4.2 x 10-3 mol NaCl/5000 mL

I) 4.28 x

J) Performs experiment in class with trying to turn on a light bulb with using water to connect a

circuit, a water NaCl solution, and a water sugar solution. The light bulb only turns on with

the water NaCl solution

(1) Water did not turn on the lightbulb

(2) NaCl (aq) turned on the lightbulb

(3) Sugar (aq) did not turn on the lightbulb

(4) Sugar in H2O

(a) Discrete, neutral

(b) Molecules surrounded by water molecules

(5) NaCl (aq)

(a) Dissolves to form solvated ions

(b) Na+

(c) Solvated ions are much better at transporting charge

(6) Species that dissolve in water to form solvated ions are called electrolytes

(7) Species that are present as neutral molecules in a solution are called non-electrolytes

K) Types of solutes

(1) Electrolytes – solutes that are ionic compounds which are soluble in water and conduct

electricity.

(2) Non-Electrolytes – solutes which do not produce conductive solutions when dissolved in

water

(a) Any non-electrolyte could be a molecular compound

(b) It could also be any compound which is insoluble in water

(3) Conductivity of electrolytes comes from a dissociation reaction

(a) Cations and anions separate into their constituent parts and dissociate in water

(i) For example, in water, NaCl separates into Na+ and Cl-

(b)

L) Reactions:

(1) Combustion:

(2) Double-displacement:

(3) Precipitation:

(4) Acid-base reactions:

(5) Redox reactions:

M) Spectator ions: anything that stays the same on both sides of a chemical equation (must

cancel what is soluble and keep what maintains a reaction or is insoluble)

N) Net ionic equation – anything that remains in a chemical equation

O) Tro Textbook examples for chemical equations

(1) NaNO3 (aq) + KCl (aq) → KNO3 (aq) + NaCl (aq) – no reaction

(a) Double displacement (swap partners)

(2) NaCl (aq) + Hg2(C2H3O2)2 (aq) → NaC2H3O2 (aq) + Hg2Cl2 (s) – unbalanced chem. Eqn.

(3) 2NaCl (aq) + Hg2(C2H3O2)2 (aq) → 2NaC2H3O2 (aq) + Hg2Cl2 (s) – balanced chem. Eqn.

(4) Net ionic eqn. Hg2+(aq) + 2Cl- (aq) → Hg2Cl2 (s)

P) Reactions in Solutions:

(1) Double displacement (replacement) reaction:

(a) Cations and anions replace one another to form new products

(b) These can lead to solids (precipitates) – a precipitate is an insoluble solid

(c) These can also lead to H2O (shown as H2O(l) )

(d) General formula for a DD reaction: AX + BY → AY + BX where A,B-cations:X,Y-

anions

Chemistry Notes

52

(e) In double displacement reaction, if a change occurs in physical states (solubility),

then a reaction occurs. If no state change occurs then there is no reaction

(f) All DD reactions can be shown as either:

(i) Molecular equations—all reactants and products are shown as neutral molecules

1. Write the chemical formulas of all entities correctly based on charges

2. Look up physical states of the reactants and products

3. Balance the equations using coefficients

(ii) Ionic equations—also called a complete ionic equation, all electrolytes are shown

as dissociated ions

1. Include physical states

2. Balance the equations using coefficients

(iii) Net ionic equation—only shows cations, anions, or molecules which change

(physical states) during a reaction

1. In a net ionic equation leave out spectator ions—spectator ions are ions

which do not change physical states in a chemical process

2. Electrolytes break up—show these

(iv) Example of reactions(performed in class): CuSO4(aq) + 2NaOH(aq) →

Cu(OH)2(s) + Na2SO4(aq) ---this solution is an electrolyte

(v) Example 2 of reactions(ionic equation): Cu2+(aq) + SO42-(aq) + 2Na+(aq) + 2OH-

(aq) → Cu(OH)2(s) + Na+(aq) + SO42-(aq)

(vi) Example 3 of reactions(net ionic equation): Cu2+(aq) + 2OH-(aq) → Cu(OH)2(s)

(g) To form H2O, we must have hydrogen cations and hydroxide anions in our reactants.

These are also called “Acid-Base” reactions or “neutralization” reactions

(i) Shown as: H+ + OH- → H2O not HOH

(h) An acid is any compound that contains H+ ions

(i) A base is any compound that contains OH- anions

(2) Redox reaction (stands for “oxidation-reduction” reaction)

(a)

) Lecture Notes 11/9/17

I) Double displacement reaction

A) AB (aq) + CD (aq) → AD (aq) + BC (aq)

B) Need a change from reactants to products for reaction to occur

(1) Formation of a precipitate

(a) Ex: AFNO3 (aq) + NaCl (aq) → AgCl (s) + NaNO3 (aq)

(2) Neutralization (acid/base reaction)

(a) HX (aq) + MOH (aq) → MX (aq) + H2O (l)

(b) Nitric acid is neutralized by potassium hydroxide solution. Show this chemical

equation.

(i) HNO3 (aq) + KOH (aq) → KNO3 + H2O (l)

II) Naming Acids

A) Naming binary acids

(1) Binary acids – composed of only 2 types of atoms. One of these is H+

(2) Begin with hydro- prefix. Include the name of the anion

(3) Change the -ide ending to -ic then add the word “acid.”

(4) Example: hydrosulfuric acid = H2S

(5) When naming acids, start by naming the anion. A monatomic anion, such as Cl−, is

simple named by replacing the ending of the element name with ide. For example, Cl− is

named chloride and O2− is named oxide.

Chemistry Notes

53

(6) Once the anion is named, name the acid based on this anion. Start by adding the prefix

hydro, and then change the ide suffix to ic. For example, HCl would require the anion

chloride to be changed to chloric. Adding the prefix hydro creates hydrochloric acid.

B) Naming oxoacids

(1) Oxyacids – contain H+ and a polyatomic ion

(a) Begin with anion name

(i) If anion has an -ate suffix, change -ate to -ic

(ii) If anion has an -ite suffix, change -ite to -ous

(b) Add the word “acid”

(c) Example sulfuric acid = H2SO4

(d) Example: sulfurous acid = H2SO3

(2) Oxoacids contain an oxygen atom in the anion. To name them, again, start with the name

of the anion when naming acids. Polyatomic anions are named according to the following

rules:

(a) If an element can only form two oxyanions then the one with the least amount of

oxygens is given the ending ite. The one containing the most oxygens is given the

ending ate.

(b) Elements that can form four polyatomic oxyanions use prefixes in addition to

the ite and ate endings.

(c) The oxyanion with the least amount of oxygens is given the prefix hypo and the

ending ite (i.e., ClO− is hypochlorite).

(d) The oxyanion with the second least amount of oxygens is simply given the ite ending

(i.e., ClO2− is chlorite).

(e) The oxyanion with the second most amount of oxygens is simply given

the ate ending (i.e., ClO3− is chlorate).

(f) Finally, the oxyanion with the most amount of oxygens is given the prefix per and

the ending ate (i.e., ClO4− is perchlorate).

(3) When naming the acid of an oxoanion the endings of the anion are modified as follows:

(4) ate changes to ic (retain prefixes) and

(5) ite changes to ous (retain prefixes).

(6) The word acid is added following the new name of the anion to complete the name of the

acid.

C) Acids of monoatomic anions, HX, hydro___ic acid

(1) Ex: HCl – hydrochloric acid

(2) HBr – hydrobromic acid

(3) HI – hydroiodic acid

(4) HCN – hydrocyanic acid

(5) HCl – hydrogen chloride gas

(6) HCl (aq) – hydrochloric acid

D) Naming acids of polyatomic oxoanions HxMOy

Anion Acid

Per_____ate Per_____ic acid

____ate -ic acid

-ite -ous acid

Hypo____ite Hypo___ous acid

E) Examples of naming acids:

Chemistry Notes

54

(1) HClO3 – chloric acid

(2) HClO – hypochlorous acid

(3) HNO2 – nitrous acid

(4) H2SO3 – sulfurous acid

F) Examples of writing chemical formulas from acid names:

(1) Phosphoric acid – H3PO4

G) A 15.3 mL solution of NaOH (aq) is titrated/neutralized with a stock solution of HCl (aq). It

takes 27.3 mL of the 1.5 M HCl (aq) to neutralize the NaOH (aq) solution. What was the

concentration of the NaOH (aq) solution?

(1) HCl (aq) + NaOH (aq) → NaCl (aq) + H2O (l)

III) Displacement reaction

A) M + AB → A + MB

(1) Cu (s) + HCl (aq) → H2 + CuCl2 (aq)

IV) Redox reactions: involve transfer of one or more electrons from one species to another with a

change in oxidation states

V) Titration – experiment carried out to figure how much solute is in any solution. Titration is the

experimentation designed to determine the amount of solute in any solution

A) Equivalence point – there is no limiting reagent present (all of the things in the solution

which we are worried about completely react with another compound). Solute in solution

completely reacts with another molecule or ion.

B) Add an indicator solution (a compound which changes color at the equivalence point) to let

you know when something has completely reacted. One of these indicator solutions is called

phenolphthalein

) Lecture Notes 11/14/17

I) Redox reactions – redox reactions occur when electrons are transferred from one atom to another

A) Oxidation—any process which involves the loss of electrons

(1) When this occurs, atoms increase in charge

B) Reduction—any process which involves the gain of electrons

(1) When this occurs, atoms decrease in charge

C) When electron transfer occurs, one atom is oxidized and another atom, by necessity, is

reduced

D) Acronym: OIL RIG: Oxidation involves loss (of electrons), Reduction involves gain (of

electrons)

E) Acronym: LEO the lion says GER: Loss of Electrons is Oxidation, Gain of Electrons is

Reduction

II) Oxidation states—hypothetical charge; this is a value assigned to atoms based on the assumption

that all bonds are ionic (whether they are ionic or covalent does not matter)

III) In the cases of an oxidation state where you arrive at a fractional oxidation state, report the

fractional oxidation state as it is

IV) Rules for assigning oxidation states (four rules which can be summed up in three):

A) Free elements have an oxidation state of 0

B) Monatomic ions have an oxidation state equal to their charge.

(1) The sum of the oxidation states of all the atoms in a compound is 0

(2) The sum of the oxidation states of all the atoms in a polyatomic ion equals the charge on

the ion

V) Oxidation – Reduction reactions – reactions in which electrons are transferred between reactants

A) Fe(s)+O2(g) → Fe2O3(s) – unbalanced reaction

B) 4Fe(s)+3O2(g) → Fe2O3(s) – balanced reaction

(1) Electrons are being transferred in this reaction

(2) Fe is being oxidized in this reaction (loss of electrons). The O atoms are being reduced

(gain of electrons)

Chemistry Notes

55

VI) Assigning Oxidation Numbers (or oxidation states) – General Rules

A) All atoms in their elemental state have an oxidation state of zero

(1) For example: O2 oxidation number = 0

(2) And: S8(s) oxidation number = 0

B) If something is a monoatomic ion, the oxidation state is equal to the charge of the ion

(1) For example: Cl1- oxidation number = -1

(2) And: Fe3+ oxidation number = +3

C) Nonmetals usually have a negative oxidation state

(1) Ex: O has an o.s. of -2 (except peroxides which are usually “-1”)

(2) H has an o.s. of “+1” when bonded to nonmetals and “-1” when bonded to metals

(i) HCl = +1; KH = -1

(3) F has an o.s. of -1 generally

D) Sum of oxidation states must be equal to the overall charge of the molecule

VII) Examples of assigning oxidation numbers

Molecule Atom Oxidation Numbers Overall Charge

H2O H → +1 0

O → -2

NaClO4 Na → +1 0

Cl → +7

O → -2

NH4Cl N → -3 0

H → +1

Cl → -1

Molecule Chlorine oxidation number

Cl-1 -1

ClO- +1

ClO2- +3

ClO3- +5

ClO4- +7

GENERAL OXIDATION

RULES:

ATOMS IN THEIR

ELEMENTAL STATE:

Oxidation state of 0

MONOATOMIC IONS: Oxidation state equal to the

charge of the ion

NONMETALS:

Oxygen (O) has an oxidation

state of -2

Peroxides have oxidation states

of -1

When bonded to nonmetals,

hydrogen (H) has an oxidation

state of +1

When bonded to metals,

hydrogen (H) has an oxidation

state of -1

Fluorine has an oxidation state

of -1

FOR ALL MOLECULES: The sum of the oxidation states

of all the atoms must be equal to

the overall charge of the

molecule

Chemistry Notes

56

VIII) Performed experiment in which this chemical reaction was made to happen:

A) Cu (s) + 2AgNO3 (aq) → 2Ag (s) + Cu(NO)2 (aq)

(1) Ag+1 is being reduced to Ag

(2) Cu is being oxidized to Cu2+

B) M + HX → MX + H2

(1) 2Na (s) + 2HCl (aq) → 2NaCl (aq) + H2 (g)

(a) Na is being oxidized to Na+1

(b) H+ is being reduced to H2 (g)

Are these reactions Redox reactions? Label for below Answer

2 Li (s) + Cl2 (g) → 2LiCl (s) A YES

2 Al (s) + 3 Sn2+ (aq) → 2 Al3+ (aq) + 3 Sn (s) B YES

Pb(NO3)2 (aq) + 2LiCl (aq) → PbCl2 (s) + 2LiNO3 (aq) C NO (DD reaction)

C (s) + O2 (g) → CO2 (g) D YES

For reaction A

Li is the reducing agent

Cl2 is the oxidizing agent

Li is oxidized to Li+1

Cl is reduced to Cl-1

) Thermochemistry – Spring 2018

I) Energy

A) Definition – The ability to do work and or transfer heat

B) Energy units: Joules (J), Kilojoules (kJ), scientific calorie = 4.184 Joules, nutritional calorie

(Cal); relationship: 1 Cal = 1000 cal = 4184 Joules

C) Forms of Energy

(1) Kinetic Energy – The energy due to motion

(a) Examples: motion of electrons around the nucleus

(b) The seemingly random motion of atoms in a sample

(2) Potential Energy – Energy due to position. This can be referred to as stored energy.

(a) Examples: Bonds between atoms (whether ionic, covalent, etc. there is still energy in

bonds)

(b) Relative positions between atoms

(c) Attraction between electrons and the nucleus

(d) Attraction or repulsion that occurs between electrons and protons

II) First Law of Thermodynamics

A) The total energy in the universe is constant. Energy is conserved. Energy cannot be created

nor destroyed.

B) The universe consists of a system and its surroundings: definition next

(1) System – reactants and products involved in a process

(2) Surroundings – container, solvent, anything that comes in contact with the system

C) Mathematically: ΔEuniverse=0; Universe = system + surroundings; ΔEuniverse=

ΔEsystem+ΔEsurroundings ; and so ΔEsystem=-ΔEsurroundings

(1) Example: burning an eraser would be an exothermic reaction because energy is being

released into the surroundings (through fire in this case)

III) Internal Energy – Internal energy is the sum of all kinetic and potential energy possible in a

system

A) This value cannot be measured very accurately at all

B) Changes in internal energy will have an impact on surroundings and can be measured

C) ΔEsystem = heat + work; will get simplified to q + w

Chemistry Notes

57

(1) Sign conventions – signs are reported in terms of the system in chemistry (they are not

displayed in formulas but are implied)

(2) So, if q is negative, then heat is released by the system

(a) This is called an exothermic process (releasing heat; consider the prefix “exo”)

(3) If, w is negative, then the system does work on the surroundings

(a) Expansion will occur when a system has a gas involved

) Lecture Notes 11/16/17 – Chapter 9: Thermodynamics

I) Thermodynamics – the study of energy and its transformations

II) Thermochemistry – the study of chemical reactions and the energy change associated with a

chemical reaction

III) Energy – capacity to do work or transfer heat

(1) We usually use the unit of the Joule (J)

(2) 1 J = 1 kg m2/s2

(3) We also use the caloric unit

(4) 1 calorie (cal) = 4.184 J

(a) This is different from the nutritional calorie

(5) 1 nutritional calorie (Cal) = 1000 cal = 4184 J

IV) Calorimetry

A) Type of experiment designed to measure the heat that’s absorbed or released by a system.

B) Calculating quantities of heat (heat absorbed or released)

(1) Heat transfer depends on a change in temperature. We show this as ΔT

(a) ΔT=Tfinal-Tinitial

(2) This produces the correct sign for any q we would like to calculate

(3) We need to know the mass of the matter that is undergoing change

(4) We also need to know the heat capacity of a substance, otherwise known as Specific Heat

(or called Specific Heat Capacity), which is denoted by Cs.

(a) The amount of heat that one gram of a sample can absorb and show a temperature

change of 1°C

(b) Units are 𝐽

𝑔°𝐶

(c) Q=mCsΔT → relationship between heat, mass, specific heat, and temperature change

C) Heat transfer: heat flows from areas of hotter temperature to areas of lower temperatures.

This is an intensive property of heat flow and does not depend on the amount of matter

present.

(1) Heat flows until two samples of matter reach a point of thermal equilibrium, where the

temperatures are the same.

V) System and Surroundings

(1) A system is any well-defined limited region of study

(2) The surroundings are anything not part of the system (the rest of the universe)

B) Types of Systems

(1) Open – a system in which matter and energy can be transferred from a system to the

surroundings or from the surroundings to the system

(2) Closed – a system in which only energy can transfer back and forth from the

surroundings

(3) Isolated – a system in which neither matter nor energy can be transferred to the

surroundings

(a) Example: Thermos

C) Work: energy used to move an object against a force

(1) Ex: PV work

D) Heat: energy transferred from a hotter object to a colder object

VI) Laws of Thermodynamics:

Chemistry Notes

58

A) First Law of Thermodynamics

(1) Energy cannot be created nor destroyed (energy is conserved)

(2) 𝐸 ≡ 𝑖𝑛𝑡𝑒𝑟𝑛𝑎𝑙 𝑒𝑛𝑒𝑟𝑔𝑦 𝑜𝑓 𝑎 𝑠𝑦𝑠𝑡𝑒𝑚

(a) Sum of all kinetic and potential energy

(3) ∆𝐸 = 𝐸𝑓𝑖𝑛𝑎𝑙 − 𝐸𝑖𝑛𝑖𝑡

(a) If ΔE>0 total internal energy of a system has increased

(b) If ΔE<0 system has transferred energy to the surroundings

(c) ∆𝐸 = 𝑞 + 𝑤

(i) Change in internal energy of a system = heat added to the system + work done on

the system

B) Qualification of heat and work

(1) If (in the system):

(a) q<0→ heat is released

(b) w<0→ work is done on the surroundings (surroundings expand)

(c) q>0→ heat is absorbed

(d) w>0→ the surroundings do work on the system (surroundings contract)

(2) Practice Problem: When the H2 and O2 are ignited the system loses 1150J of heat to the

surroundings. Reaction also raises piston exerting 480J of work on surroundings. What is

ΔE?

(a) ΔE = q + w

(b) ΔE = -1150J + (-480J)

(c) ΔE = -1630J

(3) Practice Problem: imagine a balloon filled with hot air (the system). The air absorbs 387 J

of heat and the internal energy of the system decreases by 731 J. Calculate the amount of

work associated with this process. Is the system expanding or compressing?

(a) Heat, q = -381J

(b) Energy, delta E = -731J

(c) Work=-1120 (with 3 sig. figs) J

(d) The system is expanding

(e) This is an endothermic reaction

Second Law of Thermodynamics

) Thermochemistry review 12/5/17

I) E=KE+PE (KE=molecular motion; PE=position of nuclei and electrons (chemical bonds))

II) Chemists generally more interested in change in internal energy ΔEsys

III) ΔEsys= - ΔEsurroundings = q + w

A) If system transfers heat to surroundings qsystem<0 (exothermic; if qsystem>0, it is endothermic)

B) If system does work on surroundings w<0

IV) First law of T.D. tells us that energy is conserved

A) When heat flows into a system the temperature increases by an amount proportional to the

heat capacity

B) Qsys = Csys * ΔTsys

C) Qsys = msys * Ssys * ΔTsys

D) **ΔH = qf

V) If we define reactants and products as system then heat of reaction qrxn flows into solution

(surroundings)

A) Qrxn = - msoln * Ssoln * ΔTsoln (Ssoln = 4.18 J/g °C)

VI) Enthalpy= H = E + PV

VII) ΔH= Ef – Ei + PfVf - PiVi

A) At constant pressure Pf=Pi

VIII) ΔH = ΔE + PΔV

IX) zazz

Chemistry Notes

59

X) ΔH = q – PΔV + PΔV

XI) ΔH = qf

) Problem:

I) In a coffee cup calorimeter (pressure remains constant), the heat from the reaction qrxn is

transferred to the solution causing a temperature change (ΔTsolution)

II) Qrxn = - qsoln = -msoln Ssoln ΔTsoln

III) Hess’s Law:

A) Allows us to calculate the change in enthalpy ΔH°rxn of an unknown reaction by combining

ΔH°rxn of known reactions.

(1) Special type of Hess’s Law ΔH°rxn= [summation from chemistry information sheet]

) General Chemistry I, Spring 2018

I) Chemistry is..

A) The study of matter and its changes

II) Succeeding in this course

A) Go to SI sessions and arrange a study group**

) Thermochemistry in Class notes from 4/16/2018

I) Types of Calorimetry

A) Constant volume (Bomb Calorimetry) – experiment designed to measure heat transfer

(1) Heat transfer occurs between the reaction (the system) and the surroundings (the

calorimeter and everything else).

(2) Qcal (surroundings) = Ccal ΔT → units: J/°C (there are no units of mass because a calorimeter

has a precise mass of liquid to measure the heat transfer of the reaction)

(3) qcal=-qrxn B) Constant Pressure Calorimetry – coffee cup calorimetry

(1) Another experiment designed to measure the heat transfer between substances

(2) Suitable for reactions that take place in solution. Any reaction that occurs in solution is a

good match for coffee cup calorimetry (combustion reactions would set fire to the

Styrofoam cups)

(3) Heat transfer occurs between reaction and the solution in which the reaction occurs

(4) Qsolution (surroundings) = mCsΔTsolution (the solution is usually water – 4.184J/g°C)

(5) Qreaction (calorimeter) = -qsolution

II) Enthalpy – enthalpy is the heat that is absorbed or released by a specific amount of reactant or

product involved in a system. Enthalpy is a thermodynamic quantity equivalent to the total heat

content of a system. It is equal to the internal energy of the system plus the product of pressure

and volume.

A) Symbol for enthalpy: ΔH → units: J or kJ (or J/mol or kJ/mol)

B) Relationship between q and ΔH

(1) Sign conventions are the same

(2) When q<0, heat is released, and the reaction is exothermic

(3) When q>0, heat is absorbed, and the reaction is endothermic

(4) ΔH is linked to a specific sample size (similar to a unit price such as for gas)

(5) Q is calculated for a random sample size (similar to the actual cost like at the gas pump)

C) Thermochemical equations – a balanced chemical equation that includes states and an

entropy value for the reaction or a ΔH value.

(1) Example: 4Fe(s) + 3O2(g)→2Fe2O3(s) ΔH=-1644.4 kJ

(a) The entropy value shown in the equation shows that for every 4 mol Fe, 3 mol O2,

etc., 1644.4 kJ heat is produced.

(b) Example conversion factor: 4 mol Fe/-1644.4 kJ

(2) If a hand warmer consists of 150 grams Fe, use the above thermochemical equation to

calculate the amount of heat that is absorbed or released by this reaction

Chemistry Notes

60

(3) Calculation: 150g Fe (1 mol Fe/55.84g Fe)(-1644.4 kJ/4 mol Fe)=-1.1*10^3 kJ

(a) Notice that this final quantity is a q value

D) Equation: nΔH=q

III) Properties of ΔH

A) ΔH is an extrinsic property which depends on the amount of reactant and product

B) ΔH values depend heavily on physical states.

C) ΔH values can be multiplied or divided

D) ΔH values can be reversed in terms of their sign when a reaction is reversed

IV) Energy diagrams for reactions

A) Done in class – showing that in a reaction, a ΔH<0 implies that heat energy is released from

the reaction and that a ΔH>0 shows that heat energy is absorbed by a reaction

V) Hess’s Law – a way of finding out the change in enthalpy for a reaction. Using Hess’s Law we

can calculate the change in enthalpy for a reaction without conducting a calorimetry experiment

A) If a reaction is broken down into a series of steps, the enthalpy change for the reaction is

equal to the sum of all the steps of the reaction. Definition: the sum of the enthalpy changes

for a series of steps involved in a reaction is the same as the enthalpy change for the overall

reaction.

VI) Bond enthalpies – to find the enthalpy of a generic reaction

A) Energy required to break a bond

B) Bond enthalpies are always positive quantities measured in kJ/mol

C) When we use bond enthalpies to calculate the change in enthalpy for reaction (ΔHrxn), we

take steps

(1) Imagine all the chemical bonds as being broken to generate free, unstable atoms that can

rearrange themselves to form new products

(2) Imagine all the free atoms forming new bonds in products

D) Equation: ΔH°rxn = Σ nΔH bonds broken – Σ nΔH bonds formed. Where n represents

moles of a molecule shown in a chemical equation

E) When figuring out the enthalpy of a general reaction using a table, draw the Lewis structures

and use these to figure out the enthalpy for the reactants and products. Substitute these into

the equation and figure out ΔH.

VII) Standard enthalpies of formation – to find the ΔH of a reaction

A) The amount of heat that it takes to form one mole of a compound in their standard states

B) Definition: the amount of heat energy absorbed or released when 1 mol of a compound is

formed from elements in their standard states

(1) Symbol: ΔH°f where the “f” indicates formation and the “°” indicates standard states of

formation

(2) A standard state refers to the state of a compound found at 25 °C and 1 atm of pressure

(a) For example, the standard state of water is liquid, of gold is solid, of oxygen is gas, of

CO2 is gas, etc.

C) Formation Reactions: thus far we’ve talked about redox, acid-base, double displacement, etc.,

reactions. A formation reaction takes place when elements are in their standard states and the

product (only one product) is 1 mol of compound also in its standard state.

(1) Write the formation reaction for NO2 (g)…. ½ N2 (g) + O2 (g) → NO2 (g)

(2) Write the formation reaction for N2O4 (g)… N2 (g) + 2O2 (g) → N2O4 (g)

D) ΔH°f values

(1) Enthalpy changes that apply strictly to formation reactions

(2) Many ΔH°f values are known and tabulated

(3) In appendix 2 of the Tro textbook there is a table of ΔH°f values at 25 °C and 1 atm

pressure

(4) Elements in their standard states have a ΔH°f of 0 by definition

(5) ΔH°rxn = Σ n(ΔH°f products) – Σ n(ΔH°f reactants)

Chemistry Notes

61

(a) Example: 2 CaO (s) + 2 SO2 (g) + O2 (g) → 2 CaSO4 (s)

(i) ΔH°rxn = -1006 kJ