AP CHEMISTRY Chapter 1: Matter and

Measurement

The States of Matter

Solid: Particles are highly ordered and can not flow

Liquid: Particles are loosely ordered and are able to flow.

Gas: Atoms are highly disordered and separate from one another.

Plasma: Ionized gas. Bose-Einstein Condensate (BEC): More

about this later

Bose-Einstein Condensate

Physical vs. Chemical Properties

Physical Properties: Can be measured with out changing the composition of the substance.

Chemical Properties: Describe the way a substance may change, or react, to form other substances.

Properties of Matter: Intensive properties: Do not depend on the

amount of the substance being studied. Extensive properties: Do depend on the

amount of the substance being studied.

Physical vs. Chemical Changes

Physical Change: A change that occurs without altering the chemical composition of a substance.

Chemical Change (or reaction): A change that transforms a substance into a chemically different substance.

Making Measurements

SI units: In 1960 an international agreement was made stating which metric units would be used to make scientific measurements.

Mass: Kilogram (kg)

Length: Meter (m)

Temperature: Kelvin (K)

Amount of Substance: Mole (mol)

Electric current: Ampere (A)

Luminous Intensity: Candela (cd)

SI Prefixes

Temperature

Common measurements of temperature: Fahrenheit Celsius Kelvin

Converting between Fahrenheit and Celsius: 0F = 0C (9/5) + 32

Converting between Celsius and Kelvin: K = 0C + 237.15

Derived SI units

Some things can not be directly measured.

Examples: Speed is measured as meters per second. Volume is measured as cubic length. Density is measured as grams per unit

volume.

Uncertainty in Measurment

There are two kinds of numbers in scientific work, exact and inexact.

Exact numbers have a definite value. Examples:

There are exactly 12 eggs in a dozen There are exactly 1000 g in a kg There are exactly 2.54 cm in an inch

Numbers obtained through measurements are always inexact.

This is due to human error as well as error in the measurement tools we use.

Accuracy Vs. Precision

Accuracy is how close a measurement is to the correct or “true” value.

Precision is how close individual measurements are to one another.

To have good precision a measurement device needs to be able to make reproducible measurements. This is why we calibrate instruments.

The precisions of the measurements we make is often expressed in terms of STANDARD DEVIATION.

Significant Figures

All measuring instruments have a certain degree of precision.

Instruments with more subdivisions have greater precision.

We report measurements we take using the smallest subdivision and one guess.

All of the numbers we know for certain and that one guess are called significant figures.

Sig. Fig. Rules

1. Zeros between nonzero digits are always significant. 1005 kg – Has 4 significant figures

2. Zeros at the beginning of a number are never significant. 0.02 g – Has one sig. fig. 0.0025 - Has two sig. figs.

3. Zeros at the end of a number are significant only if there is a decimal in the number. 0.0200 g – Has three sig. figs. 3.0 cm – Has two sig. figs. 100 cm – Has only one sig. fig.

Sig. Figs. in Calculations

When we use measured quantities to do calculations, the least certain measurement limits the certainty of our calculation.

Therefore the number of significant figures in our answer is determined by the number of sig figs in the least certain number.

Rules: For addition and subtraction: The answer has the same

number of decimal places as the number with the least amount of decimal places.

20.42 + 1.322 + 83.1 = 104.842, we round to 104.8 For multiplication and division: The answer has the same

number of sig figs as the number with the smallest number of sig figs.

6.221 x 5.2 = 32.3492, we round to 32

Conversions

The best way to do unit conversions is through dimensional analysis.

Once we know the relationship between two units we can use it to convert.

Example: 1 in. = 2.54 cm Convert 8.50 inches into centimeters.

Conversions with two or more steps

Convert 8.00 m into inches.

CHAPTER 2Atoms, Molecules and Ions

Atomic Theory

The modern atomic theory is credited to and English school teacher by the name of John Dalton.

His theory states that: Each element is composed of extremely small particles

called atoms. All atoms of a given element are identical to one

another in mass and other properties, but the atoms of one element are different from the atoms of all other elements.

The atoms of one element cannot be changed into atoms of another element.

compounds are formed when atoms of more than one element combine.

Subatomic Particles

Expanding on Dalton’s theory we now know that atoms are made of even smaller particles.

Protons Electrons Neutrons

Electrons

The discovery of the electron can be credited to one simple idea:

Particles with the same charge repel one another, whereas particles with unlike charges attract one another.

British scientist J.J. Thomson is credited with the “discovery’ of the electron.

The Cathode Ray Tube

The Mass of an Electron

After making his observations using cathode ray tubes Thomson was able to calculate the value of 1.76 x 108 Coulombs per gram.

He determined that this was the ratio of the electrons charge to its mass.

Using the charge of one electron (1.602 x 10-19 C) Thompson then calculated the mass of an electron.

Radioactivity

Radioactivity is the spontaneous emission of radiation from an atom.

There are three types of radiation α β γ

Each type of radiation reacts differently to an electric field.

The Nucleus

Thomson reasoned that since electrons contribute only a very small fraction of the mass of the atom then they were probably only responsible for a small fraction of the atoms size….he was wrong

The Rutherford Experiment

The modern view of the atom

Since the time of Rutherford scientists have learned much about the structure of the atom.

Atoms are composed of protons, electrons and neutrons.

We have already seen that the atoms of one element are all the same, and are different than atoms of other elements.

Atoms of different elements have a characteristic number of protons.

Atomic number:

Mass number:

Isotopes:

Average Atomic Masses

Most elements occur in nature as a mixture of isotopes.

We can determine the average atomic mass of an element by using the masses of its various isotopes and their relative abundances.

The Periodic Table

Periodic Table

Molecules and Molecular Compounds

Even though the atom is the smallest possible sample of an element only the noble gasses are normally found in nature as isolated atoms.

A molecule is an assembly of two or more atoms.

Molecular Formulas:

Empirical Formulas:

Ions and Ionic Compounds

The nucleus of an atom can not be changed by chemical processes, but some atoms can gain or lose electrons.

If electrons are removed or added to an atom we are left with a charged particle called an ion.

Cation Anion When two or more ions combine they

form an ionic compound.

Polyatomic Ions

Naming Ions

Cations (+): Cations formed from metal atoms have the same name as the metal

Examples: Na+ - Sodium Ion Zn 2+ - Zinc Ion Al3+ - Aluminum Ion

If a metal can form different ions the charge is indicated by a roman numeral

Examples: Fe2+ - Iron (II) ion Fe3+ - Iron (III) ion

Cations formed from non-metal atoms have names that end in –ium.

Examples: NH4

+ - Ammonium ion H3O+ - Hydronium ion

Anions that are formed from a single atom end in –ide H- - Hydride O2- - Oxide N3- - Nitride

A few simple polyatomic anions also end in –ide OH- - Hydroxide CN- - Cyanide O2

2- - Peroxide

Polyatomic anions containing oxygen end in -ate or –ite

NO3- - Nitrate

NO2- - Nitrite

SO42- - Sulfate

SO32- - Sulfite

Naming Ionic Compounds

Names of ionic compounds consist of the cation name followed the anion name.

CaCl2 – Calcium Chloride Al(NO3)3 – Aluminum Nitrate Cu(ClO4)2 – Copper (II) Perchlorate

Names and Formulas of Acids

Acids containing anions whose names end in –ide are named by changing the –ide to –ic, and adding the prefix hydro.

HCl – Hydrochloric acid H2S – hydrosulfuric acid Acids containing anions whose names

end in – ate or – ite are named by changing –ate into –ic, and - ite into –ous.

HClO3 – Chloric acid HClO2 – Chlorous acid

Names and Formulas of Binary Molecular Compounds

The name of the element farther to the left in the periodic table is usually written first.

If both elements are in the same group in the periodic table, the one having the higher atomic number is named first.

The name of the second element is given and –ide ending.

Greek prefixes are used to indicate the number of atoms of each element.

Some Simple Organic Compounds.

Hydrocarbons

Alkanes:

Alcohols:

CHAPTER 3Stoichiometry

Chemical Equations

H2 + O2 H2O

Types of chemical reactions: Combination (or synthesis):

Mg(s) + O2(g) MgO(s)

Decomposition: CaCO3(s) CaO(s) + CO2(g)

Combustion: C3H8(g) + O2(g) CO2(g) + H2O(g)

Formula Weights (aka molar mass)

The formula weight of a substance is the sum of the atomic masses of each atom in it’s chemical formula.

Percent composition is simply the percent that each element in a compound contributes to the total formula weight.

Calculate the percent composition of C12H22O11

42.1% - C 6.4% - H 51.5% - O

Avogadro’s Number and the Mole

The mole is just a counting number. Just like there are 12 things in a dozen there

are….

6.02 x 1023 things in a mole.

1 mole of Carbon atoms = 6.02 x 1023 carbon atoms

1 mole of H2O molecules = 6.02 x 1023 H2O molecules

1 mole of NO3- ions = 6.02 x 1023 NO3

- ions

Empirical Formula

To find the empirical formula of a compound we need to know its percent composition.

Mercury and chlorine combine to for a compound that is 73.9% mercury, and 26.1% chlorine.

Molecular Formula

If we know a substances molar mass we can determine if it’s empirical formula is the same as its molecular formula.

The empirical formula of a hydrocarbon was found to be C3H4. The molar mass of this compound was found to be 121 g/mol. Is the empirical formula also the molecular formula? If not what is the molecular formula.

Stoichiometry

2H2(g) + O2(g) 2H2O(l)

How many moles of H2O can be produced from 1.57 moles of O2

2 C4H10(l) + 13 O2(g) 8 CO2(g) + 10 H2O(l)

Calculate the mass of CO2 produced when 1.00g of C4H10 is burned.

Limiting Reactants

Sometimes one reactant runs out before the reaction is complete.

2 H2(g) + O2(g) 2H2O(l)

If 10 moles of H2 and 7 moles of O2 react

Percent Yield

Percent Yield = actual yield/theoretical yield x 100