Upload
alfred-singleton
View
213
Download
0
Tags:
Embed Size (px)
Citation preview
Chemistry Unit 3
Atomic Structure (Ch.3)
3-1 Early Models of AtomsDemocritus (450
BC)Proposed that all
matter was made of tiny indivisible particles.
He called these particles atomos (meaning indivisible).
We call them atoms.Good looking guy!
Atom
An atom is the smallest particle of an element that retains the identity of that element.
If we repeatedly cut a piece of Al, the smallest possible piece is an atom of Al.
Classic model of an atom
Aristotle
Didn’t agree with Democritus.
Believed matter was continuous and made up of only one substance called “hyle”
It wasn’t until the 1700’s when his ideas were reexamined.
Newton and Boyle (1600s)
Published articles stating their belief in the atomic nature of elements
They had no proof
Antoine Lavoisier (1770’s)
French, The “Father of Modern Chemistry”
Discovered the law of conservation of matter.
Matter is neither created nor destroyed.
Joseph Proust (1799)
French ChemistDeveloped The
Law of Definite Proportions
Compounds always contain elements in the same proportion by mass.
Law of Definite Proportions
H20 (by mass is always)
88.9% Oxygen, 11.1% Hydrogen
If we had an 80g sample of H20 how much is O?
.889 x 80 = 71gHow much is H?.111 x 80 = 9g
John Dalton (1803)
Proposed the atomic theory of matter, which is the basis for present atomic theory
John Dalton,
English schoolteacher
Atomic Theory of Matter
Each element is composed of extremely small particles called atoms.
All atoms of a given element are identical, but differ from those of any other element.
Which element is this?
Atomic theory of matter
When elements unite to form compounds, they do so in a ratio of small whole numbers. This is called the Law of Multiple Proportions.
Ex: C and O can combine to form CO or CO2, but not CO1.8.
Dalton’s Model of an Atom
All matter is composed of tiny particles
J.L. Gay-Lussac (early 1800s)
Observed that working with gas reactions at constant volume, temperature and pressure are directly related.
He named the discovery of this relationship Charles Law, which is represented by P1/T1=P2/T2.
Amadeo Avogadro (early 1800s) – Italian Physicist
Explained Gay-Lussac’s work using Dalton’s theory: Equal volumes of gases at the same temp/pres have the same number of gas molecules.
Michael Faraday (1839)Suggested that
atomic structure was related to electricity.
Atoms contain particles that have electrical charges.
Positive (+)Negative (-)Opposite charges
attractLike charges repel
William Crookes (1870’s)
English Physicist
Developed the cathode ray tube to find evidence for the existence of particles within the atom.
J.J. Thomson (1896)Used a cathode
ray tube (CRT) to identify negatively charged particles, called electrons.
Determined the ratio of an electron’s charge to its mass.
Developed the “plum pudding” model of an atom.
Cathode ray bending toward a positive charge
+-
+
+
++
++
-
-
--
-
Plum Pudding Model
Atoms are composed of randomly arranged charged particles
Robert Millikan (early 1900s)
US PhysicistUsed the oil drop
experiment to prove the electron has a negative charge
Was able to determine the charge of the electron
Millikan’s Oil Drop Experiment
Bothe/Chadwick (early 1930s)
English
Found high energy particles with no charge with the same mass as the proton called neutrons.
Ernest Rutherford (1909)Used the gold foil
experiment to prove the atom is mostly empty space.
Rutherford concluded that all of an atom’s positive charge, and most of its mass is located in the center, called the nucleus. Analogy: thumb nail and the 50 yard line.
98% of the particles passes straight through 2% of the particles deflected off at varying angles 0.01% of the particles bounced straight back
+
-++++
+ +
-
-
--
-
Rutherford’s PlanetaryModel of an atom
Positive charge and majority of mass located in the nucleus.
Negatively charged electrons orbit the nucleus like planets. Most of an atom is empty space!
Problem
He thought a moving electrical charge (-) in a curved path should lose energy (give off light).
If it did, it would fall into the (+) nucleus.
Why don’t the (-) electrons fall into the (+) nucleus?
Atom:The smallest particle of an element that has the properties of that element. Make up of nucleus
consists of protons and neutrons
Surrounded by an electron cloud
Electron cloud
Sub-Atomic Particles
Protons Positively (+) charged The number of protons
in an atom refers to the atomic number (Z)
Composed of 3 quarks (2 up, 1 down)
Mass= 1.6726 x 10-
27kg Atomic mass 1
amu (µ)
Sub-Atomic Particles
Neutrons Found in nucleus Neutral (no) charge composed of 3
quarks (1 up, 2 down)
Atomic mass 1 amu (µ)
Isotopes- atoms of the same element that have a different number of neutrons.
Sub-Atomic Particles
Electrons Found in electron
clouds surrounding the nucleus.
Negative (-) charge Mass = .00091
x 10-27 kg 1800 times smaller
than protons & neutrons
Mass 0 amu (µ)
Sub-atomic particles
Electrons Orbit the nucleus at
very high speed in energy levels (electron clouds).
Negatively (-) charged
Have no mass (when compared to protons and neutrons)
Atomic Number = Protons
The atomic number of an element is the number of protons an element has.
Located above the symbol of the element
The number of protons determines the identity of the element.
Each element has a different atomic number
Neils Bohr (1913)
Improved Rutherford’s work by saying electrons do not lose energy in the atoms so they will stay in orbit
Stated there are definite levels in which the electrons follow set paths without gaining or losing energy (Planetary Model)
Each level has a certain amount of energy associated with it and the electrons can only jump levels if they gain or lose energy
Could not explain why (-) electrons don’t fall into the (+) nucleus.
Energy Levels In the ground state
for an atom, electrons are at their lowest, most stable energy levels.
In the excited state, atoms require energy and electrons move to a higher energy level.
How many electrons are in an atom?
For an atom to have an overall neutral charge the number of electrons must equal the number of protons.
#Protons=#electrons
What element is this?
Mass number
The Mass number of an atom is the sum of the mass of protons and neutrons
Located below the symbol of the element
Atomic mass is measured in amu’s, (atomic mass units)
Based on Carbon having a mass of 12
Mass = Protons + Neutrons
How many neutrons are in an atom?
Mass=Protons+Neutrons 195= 78 + Neutrons 195-78= Neutrons Platinum has 117
Neutrons Find the number of
neutrons in: Hydrogen Carbon Helium
Potassium Boron Gold
Mass =Protons + Neutrons Hydrogen (H) 1 =1 + Neutrons
Hydrogen has 0 neutrons Helium (He) 4 = 2 + Neutrons
Helium has 2 neutrons Boron (B) 11 = 5 + Neutrons
Boron has 6 neutrons Carbon (C) 12 = 6 + Neutrons
Carbon has 6 neutrons Potassium (K) 39 = 19 + N
Potassium has 20 neutrons Gold (Au) 197 = 79 + N
Gold has 118 neutrons
Atomic MassThe average
mass of all of the isotopes of an element.
Aka: average atomic mass number, or atomic weight.
Isotopes:Atoms of the same element with different masses.
Average Atomic Mass
Ne-20 has a mass of 19.992 amu (u), and Ne-22 has a mass of 21.991 amu (u). In any sample of 100 Ne atoms, 90 will be Ne-20. Calculate the average atomic mass of Ne.
.90 x 19.992 = 17.9928.10 x 21.991 = 2.1991avg mass = 20.1919 amu
IonsAn atom that has
gained or lost an electron.
It acquires a net electrical charge.
If an atom loses an electron (oxidation) it has more protons than electrons and has a net positive charge. (cation)
11 P
11 e-
11 P
10 e-Na+
Na
Ions
If an atom gains an electron (reduction) it has more electrons than protons and has a net negative charge.(anion)
Full octet7 valence e-
Ionic ChargesCharge of ion = # protons - #
electronsWhat is the charge of a magnesium
atom that loses 2 electrons?Number of protons 12-Number of electrons 10charge of ion +2Mg2+ or Mg+2
Charge is written to the upper right of the symbol.
Representations of atoms
General form: (Elemental Notation)
X = Element SymbolA = Atomic Mass
(P + N)Z = Atomic Number
(P)Ionic Charge
XA
Z
Charge
What is the atomic structure?
Determine the number of:
P =N =e =
Na23
11
+
What is the atomic structure?
Determine the number of:
P = 11N = 12e-= 10
Na23
11
+
What is the atomic structure?
Determine the number of:
P =N = e- =
I127
53
-
What is the atomic structure?
Determine the number of:
P = 53N = 74e- = 54
I127
53
-
Put into elemental notation
Atomic # = 29Atomic Mass = 64Ionic charge = +2 ?
How many electrons?
Atomic # = 29Atomic Mass = 64Ionic charge = +2# of electrons =
Cu64
29
2+
Put into elemental notation
37 Protons48 Neutrons36 Electrons
?
37 Protons48 Neutrons36 Electrons Rb
85
37
+
Put into elemental notation
Max Planck (early 1900s)
Proposed Planck’s Theory which says that energy is given off in little packets or particles called quanta which is based on the particle nature of light
Each quantum of energy corresponds to the different energy levels for electrons.
Proposed the equation: E=hf, where E is energy, f is frequency, and h is Planck’s constant (6.63 x 10^-34 J/Hz)
De Broglie (1923)
Suggested that if waves can have a particle nature, particles can have a wave nature, known as the “wave-particle duality” principle
Wondered why the positive nucleus and negative electrons do not attract. Proposed that electrons moved so fast (speed of light) that they had properties of waves instead of particles.
The Study of WavesWave: a progressive disturbance
propagated from point to point in a medium or space without progress or advance by the points themselves
Types of Waves
Mechanical: a wave that requires an energy source and an elastic material medium to travel.
Electromagnetic: a wave that does not require a material medium to travel; it propagates by electric and magnetic fields
Wave Travel
Transverse: displacement of the medium is perpendicular to the direction of propagation of the wave.
Longitudinal: displacement of the medium is parallel to the direction of propagation of the wave
Properties of Waves
Wavelength (ג): The distance between any part of the wave (peak) and the nearest part that is in phase with it (another peak). Standard unit is meters (m).
Frequency (f ): The number of peaks which pass a given point each second. Standard unit is cycles per second which is a hertz (Hz).
Amplitude (A): The maximum displacement of a vibrating particle from its equilibrium position. Standard unit is meters (m).
Velocity (v): the distance a wave (peak) travels in a given time. Standard unit is meters per second (m/s).
Energy (E): The energy of a single photon of radiation of a given frequency. Standard unit is the joule (J).
Some relationships between the properties of waves are represented by the equations:
V=f*ג and E=h*f , where h=6.63x10^-34 J/Hz
Werner Heisenberg (1927)
Proposed his “Heisenberg uncertainty principle”, which says that the position and momentum of an electron cannot simultaneously be measured and known exactly.
The arrangement of electrons is discussed in terms of the probability of finding an electron in a certain location.
Erwin Schrodinger (1926)
Studied the wave nature of the electron and developed mathematical equations to describe their wave-like behavior.
The most probable location of the electrons can be found and the plot of this probability is called the charge cloud model.
The four quantum numbers
Principal Quantum Number (n) Refers to the energy
levels in the atom which is the distance from the nucleus and designated with a positive whole number (n=1,2,3,etc)
Wavelength of emitted photon is determined by the “energy jump” between energy levels
Energy levels (or shells) means electrons are contained in an area where the probability of finding the electron is 90%
Angular Momentum Quantum Number (l ) Refers to the sublevel
(within an energy level) which is one or more “partitions” each with a slightly different energy.
The number of sublevels in a particular energy level is equal to the principal quantum number (n).
The types of sublevels include: s, p, d, f, etc.
The four quantum numbers (continued)
Magnetic Quantum Number (m)
Refers to the orientation in space of the electrons in a sublevel
Can have any whole number value from -1 to +1 which will tell how many orbitals are in a sublevel.
A maximum of 2 electrons per orbital.
Sublevel # of Orbitals Total # of electronsspdf
Four Quantum Numbers (continued)
Spin Quantum Number (s) Refers to the spin of the electron. Pauli Exclusion Principle : if two
electrons occupy the same orbital, they must have opposite spin. Half-filled orbital: _____Filled orbital: _____
Permissible Values of Quantum Numbers for Atomic Orbitals
n l m Orbital # of Subshells #of Orbitals Max # of Electrons1 0 0 1s 1 1 2
2 0 0 2s 2 1 2 1 -1,0,1 2p 3 6
3 0 0 3s 3 1 2 1 -1,0,1 3p 3 6 2 -2,-1,0,1,2 3d 5 10
4 0 0 4s 4 1 2 1 -1,0,1 4p 3 6 2 -2,-1,0,1,2 4d 5 10
3 -3,-2,-1,0,1,2,3 4f 7 14
Distribution of Electrons for Different Elements (Electron Configuration)
Electrons will occupy the lowest energy levels and sublevels first.
Notation:
2py
2
Type of Orbital (sublevel)
Principal Quantum Number, n (energy level)
Number of electrons
Orientation of Orbital
Long Notation: Pyramid Filling
“Rule of thumb” for filling electrons at the lowest energy level possible.
Give the long notation electron configuration for:
O
Ca
Ag
Give the short notation
O
Ca
Ag
Orbital Diagrams
Usually only done for the outer shell electrons, which always includes the s and p orbitals.
Electron Dot DiagramsShows the outer shell electrons
for elements.