AP Chemistry 2014-2015

CH 1 CHEMICAL FOUNDATIONS

Matter is anything that takes up space and exhibits inertia.Composed of only ~100 types of atoms (ex. water is made of hydrogen and oxygen; running an electric current through it separates it into its constituent elements)

Chemistry is the study of matter and energy—and more importantly, the changes between them

1.1 CHEMISTRY: AN OVERVIEW

Observations Measurement = quantitative observation No number involved = qualitative observation

A hypothesis is a possible explanation for an observation Tested by carrying out an experiment Repetition of experiments is key

A theory comes into existence when hypotheses are assembled in an attempt to explain “why” the “what” happened

We use many models to explain natural phenomena

A scientific law is a summary of observed behavior Law of conservation of mass, law of conservation of

energy, etc.

1.2 THE SCIENTIFIC METHOD

A quantitative observation (measurement) consists of two parts: a number and a unit

UnitsEnglish (U.S., some of Africa)Metric (everyone else, pretty much)SI (Le Système International)—developed in 1960 to improve communication between scientists around the world; based on/derived from metric system

1.3 UNITS OF MEASUREMENT

Volume (derived from length) 1 dm3 = 1 L = 1,000 cm3 = 1,000 mL 1 cm3 = 1 mL = 1 g of water at 4°C

Mass vs. Weight Mass (g or kg)—a measure of the resistance of an object to

a change in its state of motion (a measure of inertia); the quantity of matter in an object

Weight (Newtons)—the response of mass to gravity Gravity varies with altitude; higher at low altitude, lower at high

altitude Every object has a gravitational field proportional to its mass

Precision and accuracyTwo types of errorSee Exercise 1.2

UNITS CONTINUED

RulesNonzero digits are significantA zero is significant if it is

Terminating and right of the decimal (must be BOTH) “Sandwiched” between significant figures

Exact or counting numbers have an infinite amount of significant figures as do fundamental constants

See Exercise 1.3

1.4 SIGNIFICANT FIGURES AND CALCULATIONS

Multiplication and division: the term with the least number of sig figs determines the number of sig figs in the answer Ex. 4.56 x 1.4 = 6.38 6.4

Addition and subtraction: the term with the least number of decimal places determines the number of sig figs in the answer Ex. 12.11 + 18.0 + 1.013 = 31.123 31.1

pH calculations: the number of sig figs in the least accurate measurements determines the number of decimal plces on the reported pH

Round at the end of all calculations

RULES FOR CALCULATING WITH SIG FIGS

Consider a pin measuring 2.85 cm in length. What is its length in inches?

2.54 cm = 1 inch

To convert, multiply your quantity by a conversion factor that “cancels” the undesired unit and puts the desired unit in its place.

2.85 cm x (1 inch/2.54 cm) = 1.12 inches

1.6 DIMENSIONAL ANALYSIS

EXERCISE 1.5 UNIT CONVERSIONS I

EXERCISE 1.6 UNIT CONVERSIONS II

EXERCISE 1.7 UNIT CONVERSIONS III

EXERCISE 1.8 AND 1.9

Three scales Fahrenheit TF = TC x (9°F/5°C) + 32°F

Kelvin TK = TC + 273.15 K

Celsius TC = TK - 273.15 K

See Exercise 1.10 Temperature Conversions

1.7 TEMPERATURE

Density = mass/volume

See Exercise 1.13 Determining Density

1.8 DENSITY

Solid, liquid, gasFluids = liquids and gasesVapor: the gas phase of a substance that is normally a solid or liquid at room temperature; for example, we say “water vapor” but we don’t say “oxygen vapor”

1.9 CLASSIFICATION OF MATTER

Mixtures: can be physically separated Homogeneous Heterogeneous Means of physical separation include filtering,

fractional crystallization, distillation, chromatographyPure substances: compounds and elements

Compounds can be separated into elements by chemical means (ex. electrolysis)

Elements can be broken down into atoms, which can be broken down into the nucleus and electron cloud, which can be broken down into protons, neutrons, and electrons, which can be broken down into quarks and leptons

CLASSIFICATION CONTINUED

AP Chemistry 2014-2015

CH 2 ATOMS, MOLECULES, AND

IONS

Read it if you would like.

2.1 CONTAINS HISTORICAL INFORMATION.  

Antoine Lavoisier was the fi rst chemist to insist on quantitative experimentation. He got guillotined, but not for that reason.

The law of conservation of mass : matter is neither created nor destroyed.

The law of defi nite proportions : a given compound always contains exactly the same proportions of elements by mass.

The law of multiple proportions : when two elements combine to form a series of compounds, the ratios of the masses of the second element that combine with one gram of the fi rst element can always be reduced to small whole numbers. You can see an example of this on the right. The likely formulas for these compounds would be CO and CO 2 .

2.2 FUNDAMENTAL CHEMICAL LAWS

Mass of Oxygen that combines with 1

gram of CCompoun

d I1.33 g

Compound II

2.66 g

The following data were collected for several compounds of nitrogen and oxygen: Mass of Nitrogen That Combines With 1 g of OxygenCompound A 1.750 g Compound B 0.8750 g Compound C 0.4375 g Show how these data illustrate the law of multiple proportions.

EXERCISE 2.1 ILLUSTRATING THE LAW OF MULTIPLE

PROPORTIONS

Dalton’s Theory (partially correct, partially not) All matter is made of atoms. These indivisible and

indestructible objects are the ultimate chemical particles. All the atoms of a given element are identical, in both

weight and chemical properties. However, atoms of different elements have different weights and different chemical properties.

Compounds are formed by the combination of different atoms in the ratio of small whole numbers.

A chemical reaction involves only the combination, separation, or rearrangement of atoms; atoms are neither created nor destroyed in the course of ordinary chemical reactions

Two modifications were made when subatomic particles and isotopes were discovered.

2.3 DALTON’S ATOMIC THEORY

Avogadro’s HypothesisAt the same temperature and pressure, equal volumes of different gases contain the same number of particles .

The electron J.J Thomson found that when high voltage was applied to

an evacuated type, a “ray” he called a cathode ray was produced. The ray was produced at the electrode (also called the cathode) and was repelled by the negative pole of an applied electric field. He postulated that the ray was a stream of negative particles (now called electrons). He then measured the deflection of beams of electrons to determine the charge-to-mass ratio. Thomson discovered that he could repeat this deflection and calculations using different metal electrodes, showing that all metals contain electrons and all atoms contain electrons. He also deduced that since atoms were neutral, there must be a positive charge within the atom, giving rise to the “plum pudding” model.

2.4 EARLY EXPERIMENTS TO CHARACTERIZE THE ATOM

Next up, Robert Millikan sprayed charged oil drops into a chamber. He halted their fall (due to gravity) by adjusting the voltage across two charged plates. He used the stop-drop voltage and Thomson’s charge-mass ratio to determine the charge on one drop of oil, which was a whole number multiple of the electron charge.

The mass of an electron is 9.11 x 10 -31 kg.

MILLIKAN’S OIL DROP EXPERIMENT

Henry Becquerel famously (and accidentally) discovered radiation when he left a uranium ore in a closed drawer with a photographic plate. When he realized that the plate had been exposed, he realized that a form of radiation other than light had penetrated it. The uranium, of course, was the culprit.

RADIOACTIVITY

Three types of radioactive emission Alpha (particles): helium nuclei, relatively massive and

slow, poorly penetrating, somewhat dangerous Beta (particles): electrons, relatively light and fast,

moderately penetrating, a little more dangerous Gamma (rays): just energy, most penetrating, most

dangerous These are not the only kinds of radioactive emission. We

will discuss more in the spring.

RADIOACTIVITY CONTINUED

Rutherford’s famous gold foil experiment proved that a positively-charged and somewhat bulky nucleus could be found in the center of an atom. He also found that atoms are mostly empty space.

THE NUCLEAR ATOM

Elements All matter composed of only one type of atom is an

element. 92 elements are naturally-occurring; the rest are manmade.

Atoms The atom is the smallest particle of an element that

retains the chemical properties of that element. It consists of a bulky, dense nucleus (protons and neutrons) and electrons shells/clouds (which of course contain electrons).

2.5 THE MODERN VIEW OF ATOMIC STRUCTURE (AN INTRODUCTION)

We can find a few pieces of information about each element using isotope notation. Mass number = #protons + #neutrons for specific

isotopes of an element Actual mass is not an integral number! mass defect--

causes this and is related to the energy binding the particles of the nucleus together

Atomic number = #protons = #electrons in a neutral atom = identity of the element

ATOMS AND ISOTOPES

Write the symbol for the atom that has an atomic number of 9 and a mass number of 19. How many electrons and how many neutrons does this atom have?

EXERCISE 2.2 WRITING THE SYMBOLS FOR ATOMS

Isotopes are atoms that have the same number of protons (and therefore are the same element) but different numbers of neutrons (and therefore different masses). Most elements have at least two stable isotopes.

Exceptions include Al, F, P. Hydrogen isotopes are important because they have

special names. 0 neutrons = hydrogen 1 neutron = deuterium 2 neutrons = tritium

ISOTOPES

Electrons are responsible for bonding and chemical reactivity. Chemical bonds—forces that hold atoms together Covalent bonds—atoms share electrons and make

molecules [independent units]; H2, CO2, H2O, NH3, O2, CH4 to name a few.

Molecule--smallest unit of a compound that retains the chem. characteristics of the compound; characteristics of the constituent elements are lost.

Molecular formula--uses symbols and subscripts to represent the composition of the molecule. (Strictest sense--covalently bonded)

2.6 MOLECULES AND IONS

Structural formula—bonds are shown by lines [representing shared e- pairs]; do not always indicate shape

Ions--formed when electrons are lost or gained in ordinary chem. reactions; dramatically affect size of atom

Cations--(+) ions; often metals since metals lose electrons to become + charged

Anions--(-) ions; often nonmetals since nonmetals gain electrons to become - charged

Polyatomic ions--units of atoms behaving as one entity--MEMORIZE formula and charge!

Ionic solids—Electrostatic forces hold ions together. Strong ions held close together solids.

MOLECULES AND IONS CONTINUED

Metals—malleable, ductile & have luster; most of the elements are metals—exist as cations in a “sea of electrons” which accounts for their excellent conductive properties; form oxides [tarnish] readily and form POSITIVE ions [cations]. Why must some have such goofy symbols?

2.7 AN INTRODUCTION TO THE PERIODIC TABLE

Groups or families--vertical columns; have similar physical and chemical properties (based on similar electron configurations!!) Group A—Representative elements Group B--transition elements; all metals; have

numerous oxidation/valence states Periods --horizonal rows; progress from

metals to metalloids [either side of the black “stair step” line that separates metals from nonmetals] to nonmetals

PERIODIC TABLE CONTINUED

ALKALI METALS—1A

ALKALINE EARTH METALS—2A

HALOGENS—7A

NOBLE (RARE) GASES—8A

MEMORIZE

Binary Ionic Compounds (Type I and Type II) In general—consist of a metal cation and a nonmetal anion. The cation is written first. The charges from the

cation and anion must cancel; we use subscripts to make this happen.

The names of ionic compounds do not contain prefixes such as mono- or di- unless that is part of the name of a polyatomic ion in the compound.

Monatomic ions end in –ide. Ex. NaF is sodium fluoride.

2.8 NAMING SIMPLE COMPOUNDS

Type I contain non-transition metals, which have only one charge when they are cations. Group 1A = +1, Group 2A = +2, Aluminum = +3 Zinc, silver, and cadmium also fit into this category;

silver ions always have a +1 charge, while zinc and cadmium ions always have a +2 charge.

Writing the name of a Type I Binary Ionic compound is simple. Ex. MgCl2 is magnesium chloride. The formulas are also simple, but you have to swap-and-drop to get the correct formula. Ex. sodium oxide is Na2O, and calcium nitride is Ca3N2.

TYPE I BINARY IONIC COMPOUNDS

Type II contain transition metals, as well as a few others such as lead, tin, and mercury. These ions have variable charges which are

reflected in the formula using roman numerals. For example, FeCl3 would be iron (III) chloride and SnO2 would be tin (IV) oxide. Conversely, lead (II) chloride would be PbCl2.

Some of them are real weirdoes. For example, the mercury (II) ion is Hg2+ which makes sense, but the mercury (I) ion is Hg2

2+.

TYPE II BINARY IONIC COMPOUNDS

Name each binary compound: a.CsF b.AlCl3

c.LiH

EXERCISE 2.3 NAMING TYPE I BINARY COMPOUNDS

Give the systematic name of each of the following compounds. a. CuCl b. HgO c. Fe2O3

d. MnO2

e. PbCl2

EXERCISE 2.4 NAMING TYPE II BINARY COMPOUNDS

Give the systematic name of each of the following compounds. a.CoBr2

b.CaCl2

c.Al2O3

d.CrCl3

 

EXERCISE 2.5 NAMING BINARY COMPOUNDS

Same as the other ionic names/formulas we’ve seen, but you need to look out for polyatomic ions. I’ve given you a sheet of them, but you will not be given a list for the AP Exam.

IONIC COMPOUNDS WITH POLYATOMIC IONS

Give the systematic name of each of the following compounds. a. Na2SO4

b. KH2PO4

c. Fe(NO3)3

d. Mn(OH)2

e. Na2SO3

EXERCISE 2.6 NAMING COMPOUNDS CONTAINING

POLYATOMIC IONS

a.Na2CO3

b.NaHCO3

c.CsClO4

d.NaOCl e.Na2SeO4

f. KBrO3

EXERCISE 2.6 CONTINUED

Consist of two nonmetals bonded together

Use prefixes: mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca

Don’t forget the –ide ending

BINARY COVALENT COMPOUNDS

Name each of the following compounds. a. PCl5

b. PCl3

c. SF6

d. SO3

e. SO2

f. CO2

EXERCISE 2.7 NAMING TYPE III BINARY COMPOUNDS

Hydrogen is listed first in the formula; the anion is listed second

-ide →hydro [negative ion root]ic ACID

-ate →-ic ACID -ite → -ous ACID

ACIDS

Give the systematic name for each of the following acids.a. H2SO4

b. HClO3

c. HNO3

d. H3PO4

e. HClf. H2CO3

g. H2SeO3

h. HBrO2

EXERCISE 2.8 NAMING ACIDS

Give the formula for each of the following acids.a. Hydrobromic acidb. Perchloric acidc. Sulfurous acidd. Acetic acide. Iodic acidf. Dichromic acid  

EXERCISE 2.9 WRITING ACID FORMULAS

Water (easy)Ammonia NH3

Hydrazine N2H4

Phosphine PH3

Nitric oxide NONitrous oxide (laughing gas) N2O

ANNOYING THINGS THAT PEOPLE CAN’T LET GO OF

Give the systematic name for each of the following compounds. a. P4O10

b. Nb2O5

c. Li2O2

d. Ti(NO3)4

EXERCISE 2.10 NAMING VARIOUS TYPES OF COMPOUNDS

Given the following systematic names, write the formula for each compound. a. Vanadium(V) fluorideb. Dioxygen difluoridec. Rubidium peroxided. Gallium oxide

EXERCISE 2.11 WRITING COMPOUND FORMULAS FROM

NAMES