The two types of pure substances are Elements (periodic table)Compounds

MixtureTwo or more pure substances in a container together

Two types of mixtures are Heterogeneous and homogeneous

Heterogeneous mixtures

2 or more phases (separate layers)

Homogeneous mixtures

1 phase (1 layer)

Filtration-use to separate a solid and a liquid

CHAPTER 2: Atoms, Molecules, and Ions

State Dalton’s Atomic Theory 1.

2.

3.

4.

Atoms

Are the smallest particles of an element.

J.J. Thomson

Cathode ray tube-discovered electronBe able to explain his experiment

Robert Milikan

Oil Drop Experiment-discovered charge of electronBe able to explain experiment

Ernest Rutherford

Plum Pudding Model and Gold Foil ExperimentDiscovered nucleus and proton

Discovered the neutron

James Chadwick

Number of protons and electrons

Are equal in an atom (neutral)

Number of neutrons

Mass # - atomic #

Isotopes

Vary in the amount of neutrons, but the number of protons and electrons stay the same

Make sure you can calculate the number of subatomic particles for a set of isotopes.

24

Mg

Shorthand symbols for atoms

12

24 = mass number12 =atomic number

EXTRA PRACTICE page 55

Calculate the atomic weight from isotopes and percent abundances

Ionic compound

Metal and nonmetal

Practice page 57

Molecular compound

Nonmetal and nonmetal

Practice page 57

Write chemical formulas for ionic (including acids and bases) and molecular compounds.

Practice pages 63, 64, 65

CHAPTER 3: Calculations with Chemical Formulas and Equations

Determine the amount of atoms in a compound.

Mg(OH)2

1-Mg2-O2-H

Balance Chemical Equations

Practice page 83

Keep polyatomic ions together

Predict Products of Chemical Reactions

See Reactions Sheet Provided

See page 134 for unstable reactions that create H2O and CO2 are products

Calculate molar mass (aka formula weight)

H2SO4

( 2 x 1.0079 g/mol) + (1 x 32.06 g/mol) + (4 x 15.999) = 98.072 g/mol

Try Mg(OH)2 = 58.319 g/mol

Use Avogadro’s number as a conversion factorAndApply to MOLE METHOD

Use for representative particles only: atoms, ions, molecules, formula unit

rp mol

rpa mol a mol b rpb

Use molar mass as a conversion factorAnd Apply to MOLE METHOD

g mol

ga mol a mol b gb

Use molar volume as a conversion factor(must be a gas and at STOP)And Apply to MOLE METHOD

L mol

La mol a mol b Lb

Determine Empirical Formula from %

See page 95

Determine Molecular Formula

See Page 97

Limiting Reactant

How do you identify that it is a limiting reactant problem?

They give you either grams and/or moles of the reactants!

REMEMBER you must always take your number to moles, compare the amount of moles that you have to the amount of moles that is needed

The limiting reactant is the one that you do not have enough of

See page 104

CHAPTER 6: Electronic Structure of Atoms

Define quantum theory is used to describe the arrangements of electrons in atoms (electronic structure)

The electronic structure of an atom refers to the number of electrons in an atom as well as the distribution of the electrons around the nucleus and their energies

The quantum theory aids in understanding the arrangement of the elements in the periodic table, as well as, periodic trends and formation of bonds between atoms

carries energy through space and is also referred to as radiant energy

Define electromagnetic radiation

The types of electromagnetic radiation

are Gamma rays, X-rays, UV rays, Visible, Infrared, Microwave, and radio waves

You are responsible for knowing the types of electromagnetic radiation for the AP EXAM and the way in which wavelength and frequency increases or decreases!

See page 213

Speed of light (c)

3.00 x 108 m/s

The distance between two crests or two troughs is called the

wavelength

The number of complete wavelengths, or cycles, that pass a given point each second

is the frequency

What is the relationship between wavelength and frequency?

inverse relationship

λ is wavelength

measured in m/s

ν is frequency

measured in Hz, 1/s, or s-1

3.00 x 108 m/s

c is the speed

The equation that connects λ, ν, and c is

c = λν

Why do different forms of electromagnetic radiation have different properties?

Their different wavelengths!

See Sample Exercises 6.1 and 6.2

The emission of light from hot objects (black body radiation)

The object appears black before being heated

What are the three important points used to help us understand how electromagnetic radiation and atoms interact?

The emission of electrons from metal surfaces on which light shines (photoelectric effect)

The emission of light from electronically excited gas atoms (emission spectra)

What is the relationship between E and frequency?

direct relationship!

Define the photoelectric effect

the emission of electrons from a metal surface induced by light

What is happening? If the conditions are right, a photon can

strike a metal surface and be absorbed , causing its energy (work function—as certain amount of energy) to transfer to an electron inside the metal

If the required amount of energy is absorbed the electron will overcome the attractive forces that hold it in the metal, and the electrons are emitted from the metal

If the photons have more energy required to emit the electrons, the excess energy

appears as kinetic energy of the emitted electrons

Calculate the energy of a photon.

See page 217

Write electron configurations

Use the Madalung Rule

Mg- 1s22s22p63s2

When an atom forms an ion it removes electrons from the highest energy level: Mg2+ -- 1s22s22p6

Zn-1s22s22p63s23p63d104s2

Zn2+-1s22s22p63s23p63d10

CHAPTER 7: Periodic Properties of Elements

Valence electrons

Electrons in the highest energy level

Orbitals that hold the valence electrons

Valence orbitals

How is the modern periodic table arranged?

Increasing atomic number

Define effective nuclear charge

the net positive charge experienced by an electron in a many electron atom

(meaning the attraction to the nucleus)

Zeff and Z

Zeff (effective nuclear charge) is smaller than the actual Z used in calculations because the nucleus is shielded by the core electrons

Canceling the attraction of the valence electrons to the nucleus

What does shield mean?

How to calculate Zeff

Zeff = Z -S

Z total electronsS core electrons

The larger the Zeff the greater the attraction of electrons to the nucleus.

What happens to the effective nuclear charge as you go down a group, and sizes of atoms?

Zeff is not affected going down a group (valence electrons stay the same), so when comparing use the amount of energy levels

Example Na and K: K is larger because it has more energy levels, Zeff would not work here because they both have the same amount of valence electrons that are being shielded by the nucleus. So there Zeff is the same (attraction to the nucleus).

What happens to the effective nuclear charge as you go across a period, and sizes of atoms?

Zeff increases (attraction to the nucleus) because there are more valence electrons attracted to the nucleus.

Example Na and Al: Na is larger because it has a smaller Zeff, meaning less attraction to the nucleus creating a larger atomic size.

USE equation to determine Zeff!

DO NOT USE Zeff use the amount of energy levels.

Comparing the size of an atom to its cation.

An atom that lose valence electrons from their highest energy level, creates an ion that has less energy levels, making the cation smaller than the atom from which it came.

Example: Na and Na+, Na is larger because it has more energy levels.

Comparing the size of an atom to its anion.

Energy levels would not work because the atom and the anion will have the same amount.

Anions are always larger than the atom from which they came. Yes, technically they do have a larger Zeff so you would think that they would have a stronger pull to the nucleus; however this is not true. The electron that is gained is not an original electron that the atom had so it actually creates more repulsion between the gained electron(s) and the valence electrons making the anion larger than the atom.

Example: Br and Br-, Br- is larger

Ionization energy

Is the energy required to remove electrons

Ionization energy increases from bottom to top, and from left to right

Which has the higher ionization energy Li or Na?(different period-use energy levels)

Li, it has less energy levels (less shielding) so the electrons are attracted more to the nucleus so it requires more energy to remove its electrons.

Al, higher effective nuclear charge, so stronger

Which has the higher ionization energy Na or Al?(Same period-Use Zeff)

attraction to nucleus. Requires more energy to remove the electrons.

Lattice Energy

The energy required to break the bond between atoms that have an ionic bond.

Which has the highest lattice energy LiCl or NaCl?(Same group-use energy levels)

They have the same anion so we must look at the cation. Less energy levels create shorter bonds. The nuclei’s of bonded atoms are closer in shorter bonds so the attached atoms are strongly attracted to each other.

LiCl has the higher lattice energy

Define isoelectronic. Is a group of ions all containing the same number of electrons.

See page 264

How does nuclear charge increase in an isoelectronic series? Explain.

The more electrons are removed the stronger the attraction to the nucleus, not much shielding.

Which has the highest lattice energy NaCl or AlCl3?(Same period-use )

Na+ and Al3+ are isoelectronic to Ne. The more electrons being removed the shorter (closer to the nucleus, stronger Zeff) and stronger the bond that is formed between the cation and the anion. Therefore, it has a higher lattice energy.

AlCl3 has a higher lattice energy.

CHAPTER 8: Basic Concepts of Chemical Bonding

CHAPTER 9: Molecular Geometry and Bonding Theories

Ionic compound

Metal and nonmetalElectrons are transferredcontains Ionic bonds

Molecular compound

Nonmetal and nonmetalElectrons are sharedcontains Covalent bonds

Metallic bonding

Attraction between metals and surrounding electrons

Octet rule

Gain, lose, or share electrons to become stable like noble gases (8 valence electrons, except helium)

What do you call the electrons that are involved in chemical bonding?

Valence electrons, the electrons in the highest energy levels

What do you call the simple way of showing valence electrons in an atom?

Lewis symbols

Example: F, would have 7 dots around it

Which sublevels are filled when an atom has 8 valence electrons?

s and p

How many electrons are in a single, double, and triple bond?

Single-2Double-4Triple -6

Distinguish between a polar covalent bond and a nonpolar covalent bond.

Polar covalent- the electrons are shared unequally, creating a overall dipole moment

Nonpolar covalent bond-the electrons are shared equally

REMEMBER, you use the differences in electronegativities to determine the polarity of a substance. Electrons are pulled towards the most

electronegative element.

Electronegativity increases going from the bottom to the top and from left to right on the periodic table.

What must be drawn before determining the polarity of a substance?

LEWIS DOT STRUCTURE

What are the steps for drawing the Lewis dot structure of a compound correctly?

1. sum valence electrons2. place least electronegative in the center,

except hydrogen3. attach all remaining atoms to the central

atom with a single bond4. complete octets for the terminal atoms by

placing the remaining electrons on them (starting with the most electronegative)

5. place all left over electrons on the central atom

6. If the central atom does not have an octet shift from the most electronegative to give it an octet

RULE OF THUMB: Oxygen likes double bonds, nitrogen likes triple bonds

REMEMBER, B tends to have less than an octet (page 322)

REMEMBER, some molecules may have more than an octet (page 323)

Calculate Bond Enthalpies

Page 328, use the chart on page 326

Memorize the Electron Domain Geometries and Molecular Geometries

Pages 347 and 350

Memorize Bond Angles

Page 345

What is the cause of electron domain geometries?

Valence shell electron pair repulsion (VSEPR)

Why do lone pair electrons (nonbonded electrons or unshared electron pairs) exert a greater force on the bonded molecules?

They are not shared.

In order to predict the electron domain geometry and the molecular geometry what must you do?

Draw the Lewis Dot structure!

Electron domain geometry use both bonded and nonbonded domains.

Molecular geometry use only the bonding domains. REMEMBER, although only the bonding domains are considered the shape that is seen is determined by the repulsion exhibited by the nonbonded electrons.

Practice Drawing Lewis Dot Structures of Molecules and Predicting Electron Domain Geometry and Molecular Geometry

Pages 347 and 351

Memorize Electron Domain Geometries and Molecular Geometries

2 Electron Domains

electron domain geometry Linearmolecular geometry is the same Linear

2 bonded, 0 nonbondedDraw an example:

Electron Domain Geometry Trigonal Planar

2 Types of Molecular Geometries:3 bonded and 0 nonbonded Trigonal planarDraw and Example:

3 Electron Domains

2 bonded and 1 nonbondedBentDraw an Example:

4 Electron Domains

Electron Domain Geometry Tetrahedral

3 Types of Molecular Geometries:4 bonded and 0 nonbonded TetrahedralDraw an Example:

3 bonded and 1 nonbonded trigonal pyramidalDraw an Example:

2 bonded and 2 nonbondedBentDraw an Example:

Electron Domain GeometryTrigonal bipyramidal5 bonded and 0 nonbondedTrigonal bipyramidalDraw an Example:

5 electron domains

4 bonded and 1 nonbondedSeesawDraw an Example:

3 bonded and 2 nonbondedT-shapedDraw an Example:

2 bonded and 3 nonbondedLinearDraw and Example:

6 Electron Domains

Electron Domain GeometryOctahedral6 bonded and 0 nonbondedOctahedralDraw and Example:

5 bonded and 1 nonbonded Square pyramidalDraw and Example:

4 bonded and 2 nonbondedSquare planarDraw and Example:

CHAPTER 11: Intermolecular Forces, Liquids, and Solids

London dispersion

What are the four main types of intermolecular forces?

Dipole-dipole Ion-dipole Hydrogen bonding

Increase in strengthLD DD ID HB

What happens to intermolecular forces as the temperature increases?

The intermolecular forces are weakened and the molecules are able to change states (S L G)

Which state of matter has the strongest intermolecular forces?

Solid, therefore it takes more energy to break its intermolecular forces.

The intermolecular forces must be broken in order for the substance to begin to boil. The stronger the intermolecular force the more energy required to break the intermolecular forces. Therefore,

How do intermolecular forces affect boiling points?

substances with hydrogen bonding tend to have boiling points, example (H2O).

Boiling point also increases as the dipole moment increase (increase in polarity). So for substances that have that same types of intermolecular forces you can compare the dipole moments (polarity) to figure out which has the strongest intermolecular forces, then apply it to boiling point.

If the molecules are nonpolar (dispersion forces) then molar mass can be used to determine which substance will have the higher boiling point. Higher the molar mass the higher the boiling point.

What are the two intermolecular forces that are collectively called van der waals forces?

Dipole-dipole and London-dispersion (aka dispersion forces)

Dipole-dipole forces occur between polar molecules. The positive end of one polar molecule attracts to the negative end of another polar molecule (bottom page 440)

Dispersion forces occur between nonpolar molecules. Often referred to as induced dipoles because they are influenced by the motion of electrons (bottom page 441)

Ion-dipole forces

Exists between an ion and the partial charge of a polar molecule

Deals with ionic substances dissolved in a solvent(water surrounding Na+ and Cl- ions in solution)See top of page 440

Hydrogen bonding

Hydrogen must be attracted to the O, N, or F that also has lone pairs on them

SKETCH DIAGRAM HERE

Intermolecular forces-molecule to molecule

Intramolecular force-the bonds that hold a

Distinguish between intermolecular forces and intramolecular forces.

molecule together (ionic bond, covalent bond)

Do intermolecular forces ever disappear?

No, they are always present; however, you may not see them acting on the molecules due to the space between the molecules

Phase changes depend on ____.

Intermoleculer forces

Stronger the intermolecular forces the more energy it takes to move from phase to phase

S L G

Viscosity

The resistance of a fluid to flow

Intermolecular forces

What causes a substance to be very viscous?

Strong forces = very viscous

Surface tension

The energy required to increase the surface area of a liquid by a unit of amount

Often referred to as the inward force

What causes a substance to have high surface tension?

Intermolecular forces

What happens to viscosity and surface tension as the temperature increases? WHY?

Decreases, intermolecular forces are being broken

Are intermolecular forces between molecule and

Cohesive forces

molecule (the same type of molecule)

Adhesive forces

Are intermolecular forces between molecule and a surface

What causes capillary action (move up a tube)?

Adhesive forces

CHAPTER 4: Aqueous Reactions and Solution Stoichiometry

A solution in which water is the dissolving medium is called an

aqueous solution

solution

is a homogeneous mixture of two or more substances

The substance present in the greatest quantity is usually the

solvent

Solutes

are said to be dissolved in the solvent

A substance whose aqueous solutions contain ions

Distinguish between an electrolyte and a nonelectrolyte.

is called an electrolyte (ionic substances)

A substance that does not form ions in solution is called a nonelectrolye (molecular, exception of acids, HCl)

What happens to ionic compounds when they are dissolved in water?

When ionic solids dissolve in water, each ion separates from the solid structure and disperses throughout the solution

The ionic solid dissociates into its component ions as it dissolves

The ions are said to be solvated

The solvation is a process in which ions become stabalized in solution and prevents cations and anions from recombining

Water is an electrically neutral molecule, one end of the molecule (O atom) is rich in electrons and has a partial negative charge (delta minus); and the other end (H atom) has a partial positive charge (delta plus)

Draw An Illustration:

Most molecular compounds are

Nonelectrolytes, except acids

STRONG Acids ionize completely when placed in water

2 categories of electrolytes, differ in the extent to which they conduct electricity:

Strong electrolytes-Solutes that exist in solution completely or nearly completely as ions-Essentially all soluble ionic compounds and a few molecular compounds (such as some acids) are strong electrolytes

Weak electrolytes-Solutes that exist in solution mostly in the form of molecules with only a small fraction in the form of ions

We must be careful not to confuse the extent to which an electrolyte dissolves with whether it is strong or weak!

EXAMPLE: CH3COOH is extremely soluble in water, but is a weak electrolyte

Ba(OH)2 is not very soluble, but the amount of the substance that does dissolve dissociate almost completely (Ksp)

The half arrows (double-headed) mean that the reaction is significant in both directions and are used to represent ionization of weak electrolytes

CH3COOH (aq) ⇌ CH3COO- (aq) + H+ (aq)

This balance produces a state of chemical equilibrium in which the relative numbers of each type of ion or molecule in the reaction are constant over time

A single arrow is used to represent ionization of a strong electrolyte, the absence of a reverse arrow indicates that the ions have no tendency to recombine in water

HCl (aq) H+ (aq) + Cl- (aq)

Which solute will cause a lightbulb in an experiment to glow more brightly, CH3OH or MgBr2?

MgBr2 it is an ionic compound, all ionic compounds are strong electrolytes

See Table 4.3 on page 131

Reactions that result in the formation of an insoluble product are called

precipitation reactions

precipitate

is an insoluble solid formed by a reaction in solution

solubility

is the amount of the substance that can be dissolved in a given quantity of solvent at the given temperature

Memorize these solubility rules:

Double-Replacement ReactionsOr

Metathesis ReactionsOr

Exchange Reaction

12 + 34 14 + 32AgNO3 (aq) + KCl (aq) AgCl (s) + KNO3 (aq)

Must use solubility rules to predict the states of matter of the products.

Molecular equations

show the complete chemical formulas of the reactants and the products without indicating their ionic character

ExampleAgNO3 (aq) + KCl (aq) AgCl (s) + KNO3 (aq)

complete ionic equations show all soluble strong electrolytes shown as

Ions

Molecular EquationPb(NO3)2 (aq) + 2KI (aq) PbI2 (s) + 2 KNO3 (aq)

Complete Ionic EquationPb2+ (aq) + 2 No3

-(aq) + 2 K+ (aq) + 2 I- (aq) PbI2 (s) + 2K+ (aq) + 2 NO3- (aq)

Ions that appear in identical forms among both the reactants and products of a complete ionic equation are called

spectator ions

When spectator ions are omitted from the equation, we are left with the

net ionic equation

A net ionic equation includes only the ions and molecules directly involved in the reaction

Are any spectator ions shown in the following chemical equation?Ag+ (aq) + Na+ (aq) + Cl- (aq) AgCl (s) + Na+ (aq)

Yes, Na+ (aq)

Pb2+ (aq) + 2 I- (aq) PbI2 (s)

Write the net ionic equation for the precipitation reaction that occurs when solutions of lead (II) nitrate and potassium iodide are mixed.

Acids

are substances that ionize in aqueous solutions to form hydrogen ions, thereby increasing the concentration of H+ (aq) ions

A hydrogen atom consists of a proton and an electron, H+ is simply a proton; therefore acids are often called

proton donors

H2SO4 is a strong acid, but has 2 H+ to donate, which will occur in two steps; therefore, it will be written as followed:

H2SO4 (aq) H+ (aq) + HSO4- (aq)

HSO4-(aq) ⇌ H+ (aq) + SO4

2- (aq)

Notice that the second equation is written with a double arrow. WHY?

HSO4- is a weak acid

Bases

are the substances that accept H+ ions

Bases produce OH- ions when they dissolve in water

Compounds that do not contain OH- ions may also be bases, for example NH3

When added to water, it accepts an H+ ion from the water molecule and thereby produces an OH- ion

NH3 (aq) + H2O (l) ⇌ NH4+ (aq) + OH- (aq)

Acids and bases that are strong electrolytes are called

strong acids and bases

Acids and bases that are weak electrolytes are called

weak acids and bases

Strong Acids HClHBr

HIHClO3

HClO4

HNO3

H2SO4

Strong Bases

LiOHNaOHKOHRbOHCsOHCa(OH)2

Sr(OH)2

Ba(OH)2

Classify the following as a strong or weak electrolyte, then write the species that would form in a solution:

1) Calcium chloride2) Formic acid (HCOOH)3) HNO3

4) HNO2

1) Strong (bc ionic)2) weak (bc weak acid)3) strong (bc Strong acid)4) weak (bc weak acid)

When a solution of an acid and a solution of a base are mixed, what occurs?

a neutralization reaction occurs

Is a reaction between an acid and a metal hydroxide, it always produces water and a salt

neutralization reaction HCl (aq) + Mg(OH)2 (aq) H2O (l) + MgCl2 (aq)

The salt is MgCl2

What does the term salt mean?

The term salt has come to mean any ionic compound whose cation comes from a base and whose anion comes from an acid

What is the net ionic equation for all neutralization reactions?

The essential feature of the neutralization reaction between any strong acid and any strong base:

Net ionic equation: H+ (aq) + OH- (aq) H2O

Write the net ionic equation for the reaction between aqueous solutions of hydrobromic acid and calcuim hydroxide.

Net ionic equation: H+ (aq) + OH- (aq) H2OAcid-base neutralization reaction.

Acid-Base Reactions with Gas Formation Carbonates and bicarbonates react with acids to form CO2 gas

Carbonic acid is unstable and will decompose to form H2O (l) and CO2 (g)

HCl (aq) + NaHCO3 (aq) NaCl (aq) + H2CO3 (aq)

H2CO3 is unstable, so the following occurs

H2CO3 (aq) H2O (l) + CO2 (g) THEREFORE,

Net ionic equation: H+ (aq) + HCO3- (aq) H2O

(l) + CO2 (g)

A reaction in which electrons are transferred between reactants are called

oxidation-reduction, or redox reactions

LEO

Lose electrons oxidation

GER

Gain electrons reduced

Assigning Oxidation Numbers

Elemental form always 0

Examples: N2 or Mg

Monatomic ions, charge of the ion

Assigning Oxidation Numbers

Assigning Oxidation Numbers

Nonmetals1. oxygen is -2, except with peroxides (O2

2-) it is -1

2. hydrogen is +1 with nonmetals, and -1 with metals

3. fluorine always -1, other halogens -1 in BINARY COMPOUNDS, otherwise you must use summation process

Assigning Oxidation Numbers

Use summation SOLVE FOR X!

For a compound set equal to 0 (neutral)

For an ion set equal to the charge of the ion

What happens in a single displacement reaction?

A + BX AX + B

The ion in solution is displaced or replaced through oxidation of an element

What is oxidized and what is reduced in the following equation:

Mg (s) + 2H+ (aq) Mg2+ (aq) + H2 (g)

Mg is oxidized

H+ is reduced

Reduced

Whenever one substance is oxidized, some other substance must be

Give the net ionic equation for the following reaction:

Fe(s) + Ni(NO3)2 (aq) Fe(NO3)2 (aq) + Ni (s)

Fe(s) + Ni2+ (aq) Fe2+ (aq) + Ni(s)

What must be used to predict whether or not single replacement reactions will take place?

Activity Series-Metals can displace anything that is beneath it (metals displace metals)

Nonmetals displace nonmetals, must use the periodic table, halogen displace halogens-can displace anything beneath it

When can hydrogen be displaced?

Metals from Li to Na will replace H from acids and water

Mg to Pb will replace H from acids only

Yes!

Mg (s) + FeCl2 (aq) MgCl2 (aq) + Fe (s)

Will an aqueous solution of iron (III) chloride oxidize magnesium metal? If so, write the balanced molecular and net ionic equations for the reaction.

Mg (s) + Fe2+ (aq) Mg2+ (aq) + Fe (s)

What term is used to designate the amount of solute dissolved in a given quantity of solvent or quantity of solution?

Concentration

The __________ the amount of solute dissolved in a certain amount of solvent, the more concentrated the resulting solution.

greater

Molarity (M) expresses the concentration of a solution as the number of moles of solute in a liter of solution

M = moles solute Volume of solution in liters

Mol/L

Which is more concentrated, a 1.00 x 10-2 M solution of sucrose or a 1.00 x 10-4 M solution of sucrose?

1.00 x 10-2 because it has a larger value of molarity (more solute dissolved in a liter)

1.32 M sodium sulfate

Calculate the molarity of a solution made by dissolving 23.4 g of sodium sulfate in enough water to form 125 mL of solution.

What are the molar concentrations of each of the ions present in a 0.025 M aqueous solution of calcium nitrate?

0.025 M Ca2+

0.050 M NO3-

How many grams of sodium sulfate are required to make 0.350 L of 0.500 M sodium sulfate?

24.9 g sodium sulfate

Solutions of lower concentrations can be obtained by adding water, this process is called

Dilution

Mconc x Vconc = Mdil x Vdil

Moles of solute before dilution = moles of solute after dilution

When diluting a concentrated acid or base, the acid or base should be added to the water and then further diluted by adding water.

Adding water directly to a concentrated acid or base will cause spattering caused by the intense heat generated.

How many milliliters of 3.0 M sulfuric acid are needed to make 450 mL of 0.10 M sulfuric acid?

15 mL sulfuric acid

How many grams of calcium hydroxide are needed to neutralize 25.0 mL of 0.100 M nitric acid?

.0926 g calcium hydroxide

Remember anytime you are going from one substance to a new substance you must use a balanced chemical equation!

MOLE METHOD!

To determine the concentration of a particular solute in a solution, chemists often carry out a

Titration

which involves combining a sample of the solution with a reagent solution of known concentration, called a standard solution

The point at which stoichiometrically equivalent quantities are brought together is known as the equivalence point of the titration

CHAPTER 10: Gases

Typical characteristics of gases:

Indefinite shape (takes shape of container)Indefinite volume (takes the volume of the container)

VERY, VERY weak intermolecular forces, so weak that we call our gases ideal (IDEALLY NO INTERMOLECULAR ATTRACTIONS)

But when you look closely they are present, but they cause very little attraction between the molecules because of the distance between the gas particles.

Boyle’s Law

Practice Problems WHITE TEXTBOOK CHAPTER 5!

states that the volume of a fixed quantity of gas maintained at constant temperature is inversely proportional to the pressure

T and n are held constant

P1V1 = P2V2

Indirect relationship (increase P, decrease V)

Charle’s law

Practice Problems WHITE TEXTBOOK CHAPTER 5!

states that the volume of a fixed amount of gas maintained at constant pressure is directly proportional to its absolute temperature

P and n are held constant

V1 = V2

T1 T2

Direct relationship (Increase T, Increase the V)

Gay-Lussac’s Law

Practice Problems WHITE TEXTBOOK CHAPTER 5!

(Avogadro’s Law) is that with a constant amount of gas and volume, the pressure of the gas is proportional to temperature

V and n are held constant

P1 = P2

T1 T2

Direct relationship (increase T, increase P)

Avogadro’s Hypothesis

states that equal volumes of gases at the same temperature and pressure contain equal numbers of molecules

ideal gas equation

(Gray Textbook Chapter 10)

combines Boyle’s, Charle’s, and Avogadro’s Laws

PV = nRT

(use this equation when making qualitative predictions on the exam!)

Ideal gas

(Gray Textbook Chapter 10)

is a hypothetical gas whose pressure, volume, and temperature behavior are described completely by the ideal-gas equation

R

is the gas constant, its value depends on the units of P, V, n, and T

R is typically 0.08206 L atm/mol K, but other units may be used

The conditions 0oC and 1atm are referred to as the standard temperature and pressure (STP)

molar volume

The volume occupied by one mole of ideal gas at STP, 22.4 mol/L,

Must be a gasMust be at STP

combined gas law

(Gray Textbook Chapter 10)

P1V1 = P2V2

T1 T2

The ideal-gas equation allows the ability to calculate gas density from the molar mass, pressure, and temperature of the gas

(Gray Textbook Chapter 10)

d = P M RT

The pressure exerted by a particular component of a mixture of gases is called

(Gray Textbook Chapter 10)

the partial pressure of that gas

Dalton’s law of partial pressuresPt = P1 + P2 + P3 + …

mole fraction, X,

(Gray Textbook Chapter 10)

is a dimensionless number that expresses the ratio of the number of moles of one component to the total number of moles in the mixture, so

P1 = X1Pt

Collecting Gases over Water

(Gray Textbook Chapter 10)

P total = P gas + PH2O

Kinetic molecular theory

1. gas are in continuous random motion2. no matter how many different gas

molecules that are in the same container they all have the same volume

3. we assume there are no intermolecular forces between gas molecules (ideal gas)

4. energy is transferred when the gas molecules collide, the collision are perfectly elastic

-Although the molecules in a sample of gas have an average kinetic energy and hence an average

speed, the individual molecules move at varying speeds, colliding-Momentum is conserved in each collision, but one of the colliding molecules might be deflected off at high speed while the other is nearly stopped (balancing the overall average speed)

5. temperature and kinetic energy have a direct relationship (increase T, increase Ek)

What causes gas pressure?

Gas particles colliding with the container

If Ne and N2 are in the same container at a temperature of 263 K, what do you know about the kinetic energy of both gas molecules?

They have the same kinetic energy.

If two different gases are at the same temperature, their molecules have the same average kinetic energy

Consider three samples of gas HCl at 298 K, H2 at 298 K, and O2 at 350 K. Compare the average kinetic energies of the molecules in the three samples.

The average kinetic energies depend only on temperature and not on the identity of the gas; therefore, HCl = H2 < O2

Root-mean-square (rms) speed, u

is the speed of a molecule possessing average kinetic energy

The rms speed is not quite the same as the average (mean) speed; however, the difference between the two is relatively small

The rms speed is important because the average kinetic energy of the gas molecules in a sample, E, is related directly to u2

E k= ½ mu2

What is the relationship between root-mean-square speed and temperature?

Direct relationship

The increase in the average kinetic energy as the temperature increases implies that the rms speed of molecules likewise increase as temperature increases

A) Will larger particles or smaller particles at the same average kinetic energy at the same temperature?

B) What about rms?

A) YesB) No

A gas composed of lightweight particles will have the same average kinetic energy as one composed of much heavier particles, provided that they are at the same temperature

The mass of the particles in the lighter gas is smaller than that in the heavier gas, but the particles of the lighter gas has a higher rms speed than the particles of the higher one, as shown in the following equation:u = √3RT/M

Because the molar mass is in the denominator, the less massive the gas molecules, the higher the rms speed

5.15 x 102 m/s

Calculate the rms speed of an N2 molecule at 25oC.

Effusion

which is the escape of gas molecules through a tiny hole into an evacuated space

Diffusion

which is the spread of one substance throughout a space or throughout a second substance

What is the relationship between effusion and rms?

the rate of effusion is directly proportional to the rms speed of the molecules

this is also true for diffusion

Diffusion is faster for lower mass molecules than for higher mass ones?

The have a higher rms

The average distance traveled by a molecule between collisions is called the

mean free path of the molecule

Will the following changes increase, decrease, or have no effect on the mean free path of the gas molecules in a sample of gas?

a) Increasing pressureb) Increasing temperature

a) mean free path decreases because the molecules are crowded closer together

b) -NO effect. The molecules are moving faster at the higher temperature, but they are not crowded any closer together

When do real gases behave NONIDEALLY?

High pressure and low temperature The attractive forces between molecules

come into play at short distances, as when molecules are crowded together at high pressures

As a gas is cooled, the average kinetic energy of the molecules decreases, but intermolecular attractions remain constant (may began to turn into liquids)

Would you expect helium gas to deviate from ideal

B- lowest temperature and highest pressure

behavior more at a) 100 K and 1 atmb) 100 K and 5 atmc) 300 K and 2 atm

CHAPTER 14: Chemical Kinetics

4 factors allow us to change the rates at which particular reactions occur:

The physical state of the reactantsThe concentration of the reactantsThe temperature at which the reaction occursThe presence of a catalyst

The speed of a chemical reaction- its reaction rate- is the change in the concentration of reactants or products per unit of time

Units for reaction rate are usually molarity per second (M/s)

Average rate of appearanceB = ∆ [B] t

A = - ∆ [A]

Average rate of disappearance t

Calculate the Average Rate of Reaction Sample Exercise 14.1, 14.2, 14.3

An equation that shows how the rate depends on the concentrations of reactants is called a

rate law

The constant k in the rate law is called the rate constant

What do the exponents m and n represent?The exponents m and n in a rate law are called reaction orders; these exponents indicate how the rate is affected

by the concentration of each reactant

m and n are experimentally determined

overall reaction order is the sum of the orders with respect to each reactant in the rate law

What do the units of the rate constant depend on? The units of rate constant depend on the overall reaction order of the rate law.

Determining Reaction Orders and Units for Rate Constants

Sample Exercise 14.5, 14.6 (especially)

is one whose rate depends on the concentration of a single reactant raised to the first power

Equationln[A]t – ln[A]0 = -kt

first order-reactionThe above equations can be used with any concentration units, as long as the units are the same for both [A]t and ln[A]0

Using the Integrated First-Order Rate Law Sample Exercise 14.7

For a first-order reaction, therefore, a graph of ln[A]t versus time gives a

straight line with a slope of –k and a y-intercept of ln[A]o

second –order reactionis one whose rate depends on the reactant concentration raised to the second power or on the concentrations of two different reactants, each raised to the first power

1 = kt + 1 [A]t [A]0

If the reaction is second order, a plot of 1/[A]t versus t will yield a

straight line with a slope equal to k and a y-intercept equal to 1/[A]0

Determining Reaction Order from the Integrated Rate Law Sample Exercise 14.8

Half-life of a reaction, t1/2

is the time required for the concentration of a reactant to reach one-half of its initial value

half-life of a first order reaction t1/2 = 0.693 k

Determining the Half-life of a First-Order Reaction Sample Exercise 14.9

What affects the reaction rate? Reaction rates are affected both by the concentrations of reactants and by temperature

Collision Model

The central idea of the collision model is that molecules must collide to react

the greater the number of collisions occurring per second, the greater is the reaction rate

As the concentration of reactant molecules increases, the number of collisions increases, leading to an increase in reaction rate

Orientation Factor

In most reactions, molecules must be oriented in a certain way during collisions for a reaction to occur

The relative orientations of the molecules during their collisions determine whether the atoms are suitably positioned to form new bonds

activation energy, Ea

The minimum energy required to initiate a chemical reaction

Activated complex

The particular arrangement of atoms at the top of the barrier

What do you know about the speed of a reaction that has a low Ea ?

fast reaction

Why isn’t collision frequency the only factor affecting a reaction rate?

Molecules must not just collide, but they must collide in the proper orientation and with an energy greater than the activation energy for the reaction.

Relating Energy to Activation Energies and Speeds of Reaction

Sample Exercise 14.10

Know the areas of the Energy Profile (Page 595)

reaction mechanism The process by which a reaction occurs

elementary reactions

Product ion of products occur in a single event or step

MolecularityThe number of molecules that participate as reactants in an elementary reaction

Unimolecular If a single molecule is involved, the reaction is

Bimolecular Elementary reactions involving the collision of two reactant molecules

Termolecular(rare)

Elementary reactions involving the simultaneous collision of three molecules are

What is the molecularity of this elementary reaction? Explain.NO (g) + Cl2 (g) NOCl (g) + Cl (g)

Because the elementary reaction involves two molecules, it is bimolecular.

Some compounds are neither an original reactant nor a final product in the overall reaction- it is formed in one elementary reaction and consumed in the next- it is called an

Intermediate

Determining Molecularity and Identifying Intermediates

Sample Exercise 14.12

Elementary reactions are significant in a very important way: if a reaction is an elementary reaction, then its rate law is based directly on its

molecularity

Predicting the Rate Law for an Elementary Reaction

Sample Exercise 14.13

The overall rate of a reaction cannot exceed the rate of the slowest elementary step of its mechanism; the slow step limits the overall reaction rate, it is called the

rate-determining step (or rate limiting step)

Determining the Rate Law for a Multistep Mechanism

Sample Exercise 14.14

Deriving the Rate Law for a Mechanism with a Fast Initial Step

Remember the fast step is at equilibrium so you must solve and then substitute

Sample Exercises14.15

catalyst is a substance that changes the speed of a chemical reaction without undergoing a permanent chemical change itself

Distinguish between a catalyst and an intermediate

A catalyst is there at the start of the reaction, whereas the intermediate is formed during the course of the reaction

Catalyst written over the arrow

Why does a catalyst speed up the reaction rate?

A catalyst lowers the activation energy by providing a different, lower-energy pathway.

Remember activation energy and the rate of a reaction has an inverse relationship.

CHAPTER 15: Chemical Equilibrium

Chemical equilibrium

occurs when opposing reactions are proceeding at equal rates: the rate at which the products are formed from the reactants equals the rate at which the reactants are formed from the products

energy and concentrations are conserved

The equilibrium mixture results because the reaction is

Reversible

the rate laws for elementary reactions can be written from their

chemical equations

Forward reaction: N2O4 (g) 2 NO2 (g) rate f = kf [N2O4]

Reverse reaction: 2 NO2 (g) N2O4 (g) rate r = kr [NO2]2

Important points about equilibrium

At equilibrium, the concentrations of reactants and products no longer change with time

For equilibrium to occur, neither reactants nor products can escape from the system

At equilibrium a particular ratio of concentration terms equals a constant

From which direction can equilibrium be reached?

The equilibrium condition can be reached from either direction

How do we know when equilibrium has been reached in a chemical reaction?

Concentrations no longer change with time indicates that equilibrium has been reached.

equilibrium-constant expression

Kc = [D] d [E] e (products) [A] a [B] b (reactants)

What does the equilibrium constant depend?

the equilibrium-constant expression depends only on the stoichiometry of the reaction, not on its mechanism

What does the equilibrium constant depend on?

The value of the equilibrium constant depends only on the particular reaction and on the temperature

What is not included in equilibrium expressions?

Solids and liquids

Whenever a pure solid or a pure liquid is involved in a heterogeneous equilibrium, its concentration is not included in the equilibrium-constant expression for the reaction; therefore, the equilibrium –constant expression for the above equation is

Give the equilibrium expression for Ag+ (aq) + 2NH3 (aq) ↔ Ag(NH3)2

+ (aq)

Kc = [Ag(NH3)2+ ]

[Ag+] [NH3]2

Are units given for the values of Kc?

no units

What do the symbols Kc and Kp represent?

Kc- obtained when equilibrium concentrations expressed in molarity are substituted in the equilibrium-constant expression

Kp- obtained when equilibrium concentrations expressed in atmospheres are substituted into the expression

Kp = Kc(RT) ∆n

Relationship between Kp and Kc

The change in n is the number of moles of gas in the chemical equation for the reaction

Converting Between Kc and Kp

Sample Exercises 15.2

K>1

If K>1 (that is, large K): Equilibrium lies to the right; products predominate

*REMEMBER*Opposing rates, not concentrations, are equal at equilibrium; thus, equilibrium constants for different reactions can span a very wide range

K<1

If K<<1 (that is, small K): Equilibrium lies to the left; reactants predominate

*REMEMBER*Opposing rates, not concentrations, are equal at equilibrium; thus, equilibrium constants for different reactions can span a very wide range

Evaluating an Equilibrium Constant When an

For Equilibrium Reactions (Kc and Kp):1. When you flip the equation you take the

reciprocal of the K value.2. When you multiply the equation by a

coefficient you raise the K to that power, so if you multiply the equation by 2 you square the K

3. To get the K for the overall reaction you

Equation is Reversed

Combining Equilibrium Expressions

multiply the K values together

Sample Exercises 15.4, 15.5

Remember for thermochemistry:1. When you flip the equation you flip the

sign on H2. When you multiply the equation by a

coefficient you raise the H to that power, so if you multiply the equation by 2 you multiply the H by 2

3. You add the H values for each reaction to get the H for the overall reaction.

Write the equilibrium-constant expression for SnO2 (s) + 2 CO (g) ↔ Sn (s) + 2 CO2 (g)

Kp = PCO22

PCO2

Calculating K When All Equilibrium Concentrations Are Known

Sample Exercises 15.8

Sample Exercise 15.9The following steps outline the procedure used to do this (PG 642):-Tabulate all known initial and equilibrium concentrations of the species that appear in the equilibrium-constant expression

Calculating K from Initial and Equilibrium Concentrations

-For those species for which both the initial and equilibrium concentrations are known, calculate the change in concentration that occurs as the system reaches equilibrium

-Use the stoichiometry of the reaction to calculate the changes in concentration for all the other species in the equilibrium

-From the initial concentrations and the changes in concentration, calculate the equilibrium concentrations. These are then used to evaluate the equilibrium constant.

The equilibrium constant also allows us to

Predict the direction in which a reaction mixture will proceed to achieve equilibrium

Calculate the concentrations of reactants and products when equilibrium has been reached

reaction quotient

is a number obtained by substituting reactant and product concentrations or partial pressures at any point during a reaction into an equilibrium-constant expression; therefore, the general reaction

Use to determine whether the reaction will produce reactants or products to reach equilibrium

The reaction quotient will equal the equilibrium constant only if the system is already at equilibrium

Q = K

Q > K

The concentration of the products is too large and that of reactants too small. Thus, substances on the right side of the chemical equation will react to form substances on the left; the reaction moves from right to left in approaching equilibrium

Q < K

The concentration of products is too small and that of reactants too large. Thus, the reaction will achieve equilibrium by forming more products; it moves from left to right

Predicting the Direction of Approach to Equilibrium

Sample Exercise 15.10

Calculate Equilibrium Concentration

Sample Exercise 15.11

Calculate Equilibrium concentrations from Initial Concentrations

Sample Exercise 15.12

Note the initial concentrations

Construct a table in which we tabulate the initial concentrations

Use stoichiometry of the reaction to determine the changes in concentration that occur as the reaction proceeds to equilibrium. Represent the change in concentration by the variable x (The balanced chemical equation tells us the relationship between the changes in the concentrations of reactant/product)

Use the initial concentrations and the changes in concentrations, as dictated by the stoichiometry, to express the equilibrium concentrations

Substitute the equilibrium concentrations into the equilibrium-constant expression and solve for the unknown, x -Expand this expression to obtain a quadratic equation in x, this will lead to two solutions

-Remember, a negative concentration is not chemically meaningful, we reject this solution

Le Chatelier’s principle

If a system at equilibrium is disturbed by a change in temperature, pressure, or the concentration of one of the components, the system will shift its equilibrium position so as to counteract the effect of the disturbance

-Adding or removing a reactant or productIf a chemical system is at equilibrium and we

Three ways that a chemical equilibrium can be disturbed:

increase the concentration of a substance (either a reactant or product), the system reacts to consume some of the substance

-Changing the pressure by changing the volumeIf a system is at equilibrium and its volume is decreased, thereby increasing its total pressure, Le Chatelier’s principle indicates that the system will respond by shifting its equilibrium position to reduce the pressure (creating less gas particles)

-Changing the temperatureThe rules for the temperature dependence of the equilibrium constant by applying Le Chatelier’s principle; we do this by treating heat as if it were a chemical reagent

Endothermic: Reactants + heat ↔ products

Exothermic : Reactants ↔ products + heat

The equilibrium shifts in the direction that consumes the excess reactant (or product), namely heat

What happens if we add a catalyst to a chemical system that is at equilibrium?

A catalyst lowers the activation barrier between reactants and products

The catalyst thereby increases the rates of both the forward and reverse reactions

As a result, a catalyst increases the rate at which equilibrium is achieved, but it does not change the composition of the equilibrium mixture

NO, catalysts have no effect on the position of an

Does the addition of a catalyst have any effect on the position of an equilibrium?

equilibrium, although they do affect how quickly equilibrium is reached.

CHAPTER 16: Acid-Base Equilibria

Arrhenius Acid and base definitions

Acid-a substance that when dissolved in water increases the H+ ion in solution

Base- a substance that when dissolved in water increases the OH- ion in solution

Acid- is a substance that donates a H+ (proton) to another substance

Bronsted-Lowry acid and base definitions

Base- is a substance that accepts a H+ (proton) from another substance

Conjugate acid-base pairs

An acid and a base that differs only in the presence or absence of the proton (H+)

Conjugate acid(always on the product side of the equation)

Formed by the adding of a proton to the base

Conjugate base(always on the product side of the equation)

Formed by removing a proton (H+) from the acid

See pages 670 and 671

Label the conjugate acid-base pairs

What another abbreviation for a H+ ion?

H3O+

Amphiprotic or amphoteric

A substance that can act as an acid or a base

H2O

The stronger an acid

the weaker its conjugate base

the weaker is its conjugate acid

the stronger the base

Weak acid produce

Weak conjugate base

The auto ionzation of water leads to

Ion-product constant (Kw) for H2O

Kw = [H3O+][OH-] = 1.0 x 10-14

In acidic solutions [H+] exceeds [OH-]In basic solutions [OH-] exceeds [H+]

The [H+] and [OH-] in a neutral solution is

[H+] = 1.0 x 10-7

[OH-] = 1.0 x 10-7

[H+] > 1.0 x 10-7

In an acidic solution the [H+]

In a basic solution the [H+]

[H+] < 1.0 x 10-7

Calculating [H+] from [OH-]

Sample Exercise 16.5

Kw = [H3O+][OH-] = 1.0 x 10-14

Calculating pH from [H+]

Sample Exercises 16.6 and 16.7

pH = -log[H+]ORpH = -log [H3O+]

pOH and Other “p” Scales

pOH = - log [OH-]ThereforepH + pOH = 14.00 (at 25oC)

Memorize the 7 STRONG ACIDS

HCl, HBr, HI, HNO3, HClO3, HClO4; and H2SO4

How do strong acids and bases ionize? Completely; therefore, when calculating the pH or pOH of strong acids/bases you can use

pH = -log[H+] for acids

pOH = - log [OH-] for bases

pH + pOH = 14.00 (at 25oC)

Kw = [H3O+][OH-] = 1.0 x 10-14

BUT NOT SO FOR WEAK ACIDS AND BASES. They ionize partially which means they form equilibrium reactions; therefore I-C-E tables must be used. Remember products/reactants and coefficients as powers for writting the equilibrium expression.

Calculate the pH and pOH for Strong acids and bases.

Sample Exercises 16.10, 16.11, 16.12, 16.13, 16.15, and 16.16.

Relationship between Ka and Kb.

Ka x Kb = [H+] [OH-] = Kw

Ka x Kb = Kw

REMEMBER Ka and Kb are for weak acids and bases only!

The larger the value of Ka and Kb the stronger the acid (weak) and base (weak).

In terms of ionization, what does a large Ka or Kb indicate about the amount of H+ or OH- ions in the solution?

Large amount of H+ ions for weak acid in solution, and the same for the OH- ions concerning bases.

pKw Is the potential ion-product constant

pKa is the potential acid-dissociation constant

pKb is the potential base-dissociation constant

pKa + pKb = pKw =14.00

The equation above connects pH to Ka and Kb

Calculate Ka or Kb for a Conjugate acid-base pair

Sample Exercise 16.17

How do anions of a strong acid affect the pH of a solution?

An anion that is the conjugate base of a strong acid will not affect the pH (it will be a spectator)

How do anions of a weak acid affect the pH of a solution?

An anion that is the conjugate base of a weak acid will cause an increase in pH

How do cations of a weak base affect the pH of a solution?

A cation that is the conjugate acid of a weak base will cause a decrease in pH

Other than the cations of group 1A and heavier members of group 2A, how will the other ions have an effect on the pH of a solution?

The cations of group 1A and heavier members of group 2A (Ca2+, Sr2+, and Ba2+) will not affect pH ( will be spectator ions)

will cause a decrease in pH

When a solution contains both the conjugate base of a weak acid and the conjugate acid of a weak base how will the pH be affected?

the ion with the larger equilibrium constant, Ka or Kb, will have the greater influence on the pH

What effect will each of the following ions have on the pH of a solution: NO3

- and CO32-?

NO3- is the anion of a strong acid and will not affect

the pH

CO32- is the anion of a weak acid and it will increase

the pH of the solution

Which of the following cations has no effect on the pH of a solution: K+, Fe2+, or Al3+?

K+ ion, an alkali metal cation, does not affect pHMost transition metal ions with a 2+ charge or higher form acidic solutions

Determine whether aqueous solutions of each of the following salts will be acidic, basic, or neutral:

(a) Ba(CH3COO)2

(b) NH4Cl(c) CH3NH3Br(d) KNO3

(e) Al(ClO4)3

Ba2+, heavy alkaline earth metal and will not affect the pH; CH3COO-, is the conjugate base of the weak acid and will produce OH- ions increasing the pH, so the solution is basic

NH4+, is the conjugate acid of a weak base and

decreases the pH (acidic), Cl- is the conjugate base of a strong acid and would not affect the pH, so the solution will be acidic

CH3NH3+ is the conjugate acid of a weak base and

is acidic, Br- is the conjugate base of a strong acid and has no affect on the pH, so the solution will be acidic

K+ is a group 1 A metal and has no affect on the pH, NO3

- is the conjugate base of a strong acid and will not affect the pH, so the solution is neutral

Al3+ is not in group 1A or 2A and will cause the solution to become acidic, ClO4

- is the anion of a strong acid and will not affect the pH, so the solution will be acidic

Ka > Kb

the ion will cause the solution to be acidic

Kb > Ka

the solution is basic

CHAPTER 17: Additional Aspects of Aqueous Equilibria

When a weak electrolyte and a strong electrolyte contain a common ion, the weak electrolyte ionizes less than it would if it were alone in

Common ion effect

solution

Calculate the pH when a common ion is involved.

Use tabulation

Sample Exercise17.1 and 17.2

Buffered solutions (buffers)

Solutions that contain a weak conjugate acid-base pair and can resist drastic changes in pH upon the addition of small amounts of strong acid or strong base

Will H2SO4 and HSO4- be a good buffer? Explain.

No, because buffers consists of weak acids and their conjugate bases.

Henderson-Hasselbalch Equation

pH = pKa + log [conj base] [acid]

The best equation to use for calculating the pH of a solution that contains a buffer.

If the [conj base] is equal to the [acid] what do we notice about the pH?

pH = pKa

Calculate the pH of a buffer

Sample Exercise 17.3

What happens to the pH when there are large amounts of the conjugate acid-base pairs in the solution?

The pH is more resistant to change. (can place more acid or base in the solution before the pH starts to change)

Acids and bases

What type of substances are involved in an acid-base titration?

What is the purpose of acid-base indicators?

They help identify the equivalence point of acid-base reactions.

What do you call the substance inside of the buret?

Tirant

If HCl is in the buret, and NaOH is in a beaker, What happens to the pH of the solution in the beaker when the titrant is added?

Decreases the OH- is being removed from solution! Trying to reach neutralization!

What are the three types of titrations that have characteristic shapes?

Strong acid-Strong baseWeak acid –Strong basePolyprotic acids

1. initial pH (low on the pH scale)2. b/w initial pH and the equivalence point

What are the four regions of a Strong acid-Strong base titration curve?

3. equivalence point (where the H+ ion is removed)

4. after the equivalence point

What are the four regions of a Weak acid-Strong base titration curve?

1. initial pH (high on the pH scale, but lower than 7)

2. between the initial pH and the equivalence point

3. the equivalence point (above 7), because the anion that is formed is a weak base (where the H+ ion is removed)

4. after the equivalence point

Distinguish between the equivalence points of a strong acid- strong base titration and a weak acid-strong base titration.

SA-SB titrations the equivalence point is 7WA-SB titration the equivalence point is above 7

Can you identify the equivalence point on titration curves?

See page 731 and 734

How can you identify the end point?

It is the point in the titration where the indicator changes color

7, Strong acid and a strong base

What is the pH of the equivalence point when 0.10 M H2SO4 is added to a solution of NaOH?

Calculate the pH for a strong acid-strong base titration

Sample Exercise 17.6

Calculate the pH for a weak acid-strong base titration

Sample Exercise 17.7

Why is the choice of indicator more crucial for a weak acid-strong base titration than for a strong acid-strong base titration?

It’s a weak acid so it takes less base to make it reach its equivalence point so an indicator that can reach its color change in such a short time needs to be used.

The indicator that is chosen should be one that will provide signals (change color for a split second then disappears) before it actual change color (end point).

How do you identify a titration curve for a polyprotic acid?

It has multiple equivalence points, which is where the H+ ion is removed.

One in which the solution is in contact with undissolved solute.

Saturated solutionSketch a diagram of a saturated solution:

Solubility-product constant (Ksp)

Dissolution of a solid, the equilibrium of a solid with its aqueous solution

Ksp indicates how soluble a solid is in water

The larger the Ksp the more soluble the substance in water (a lot can be dissolved in water)

Write Solubility-product Expressions

Calculate Ksp from Solubility

Calculate Solubility from Ksp

Sample Exercises 17.9, 17.10, 17.11

How can equilibrium be achieved?

From either side (starting with reactants or products)

What is used to determine the direction in which a reaction must proceed to reach equilibrium?

Q (reaction quotient)

Q > Ksp, precipitation occurs (falls out of solution)

The 3 relationships between Q and Ksp

until Q = Ksp (more products than reactants)

Q< Ksp, solid dissolves until Q = K (more reactants than products)

Q = Ksp, equilibrium (saturated solution)

Saturated solution

Equilibrium between the ionic solid and its aqueous solution

Predicting Whether a Precipitate Will Form

Calculating Concentrations for Precipitation

Sample Exercises 17.15, 17.16

CHAPTER 5: Thermochemistry

System What we are concerned with

Surroundings Everything else

Open systemMatter and energy can be exchanged with surroundings

Closed system Can exchange energy, but not matter with surrounding

Isolated system Neither matter nor energy can be exchanged with surroundings

Ef > Ei System gained E from surroundings; endo

Ef < Ei

System released energy to the surroundings; exo

Relate heat and work to changes of internal energy.

Sample Exercise 5.2

Distinguish between endothermic and exothermic.

Endothermic energy absorb by the system from the surroundings; surroundings doing work on the system

Exothermic energy leaves the system and goes to the surroundings; system doing work on the surroundings

Know how to interpret Energy Diagrams Page 177 and 178

What does enthalpy measure? It measure heat flow to and from the system.

SO, q = H

Relate H to Quantities of Reactants and Products Sample Exercise 5.4

Know how to use the specific heat equation to calculate grams, mole, or H?

q = C x m x T

Remember T is equal to K, so no conversion to K is needed when the temperature is given in oC, simply change the unit to K!

Practice Problems Sample Exercises 5.5, 5.6, 5.7 is BOMB calorimetry so q = -cal x T

Hess’s Law (flipping or multiplying through by coefficients to obtain an overall reaction)

Sample Exercises 5.8 and 5.9

Remember for thermochemistry:4. When you flip the equation you flip the

sign on H5. When you multiply the equation by a

coefficient you raise the H to that power, so if you multiply the equation by 2 you multiply the H by 2

6. You add the H values for each reaction to get the H for the overall reaction.

For Equilibrium Reactions (Kc and Kp):4. When you flip the equation you take the

reciprocal of the K value.5. When you multiply the equation by a

coefficient you raise the K to that power, so if you multiply the equation by 2 you square the K

6. To get the K for the overall reaction you multiply the K values together

Look at equations and determine if the enthalpy changes represent a standard enthalpy of formation.

Must be in the proper states at room temperature.

Sample Exercise 5.10

Calculate Ho Ho = nHo(products) - mHo (reactants)

CHAPTER 19: Chemical Thermodynamics

Spontaneous To proceed on its own

Spontaneous reactions can be ___. Fast or slow

Spontaneous processes are ___.

Irreversible

Entropy (S)Randomness in a systemState function

What happens to entropy in a spontaneous process?

It increases

S > 0, spontaneous (irreversible)

2nd law of thermodynamics Irreversible processes result in an overall increase in entropy, whereas a reversible process results in no overall change in entropy

What happens to the movement of molecules at higher temperatures?

They move faster

What are the three kinds of motion a molecule may have?

Translation

Vibrational

Rotational

Microstate

Is a single possible arrangement of the positions and kinetic energies of the gas molecules when the gas is in a specific thermodynamic state

Each thermodynamic state has a certain number of microstates associated with it

W = microstates

S = k ln W

What happens to entropy as the number of microstates increase?

Entropy increasesDirect relationship

SO we can say an increase in entropy represents an increase in randomness of the system

What are the three things that have a direct relationship with the number of microstates?

Temperature, volume, number of independently moving particles

When can we expect an increase in entropy?

Gases are formed from solids and liquids

Liquids or solutions are formed from solids

The number of gas molecules increase during a reaction

What happens to the entropy of a system as the temperature is decreased?

It is decreased, the particles move slower and have less degrees of freedom, stronger intermolecular forces!

Use the summation equation to calculate So.

So = nSo(products) - mSo (reactants)

Given on AP EXAM

Are spontaneous processes that result in a decrease in the system’s entropy exothermic or endothermic?

Exothermic, energy is leaving the system

What connects enthalpy and entropy? Go

If Go is negative then…Spontaneous in the forward direction(exothermic)

If Go is positive then… Nonspontaneous in the forward direction, work must be done on the system by the surroundings(endothermic)

How to calculate Go

Go = nGo(products) - mGo (reactants)

or

Go = Ho - TSo

BOTH EQUATIONS PROVIDED ON AP EXAM

How to calculate G (nonstandard conditions)

G = Go + RT ln Q

How to calculate Q… Products /reactants

Relationship between Go and K

Go < 0 then K > 1, equilibrium lies to the right (more products than reactants) is spontaneous in the forward direction

AND

VICE VERSA

Equation that connects Go and K Go = - RT ln K

CHAPTER 20: Electrochemistry

OxidationLEOLoss of electrons

ReductionGERGain of electrons

Half-cells  One half of the galvanic cell (e.g. the oxidation half-cell or the reduction half-cell).

Anode  Electrode at which oxidation takes place.

Cathode Electrode at which reduction takes place.

Salt Bridge 

Connecting the two half-cells, and will contain an aqueous solutions (ionic compound dissolved in it)

The cations, positive ions, flow toward the cathode to replace the cations that are being picked up at the electrode.

The anions, negative ions, flow toward the anode to balance the positive charge of the cations that are released from the electrode.

Voltmeter is a device to detect the amount of current flowing through the cell (in volts).

What is happening to the mass of Zn electrode?

What is happening to the mass of Cu electrode?

Explain what is happening in the salt bridge.

Eocell > 0

Spontaneous

Calculate Eored(cathode), Eo

red(anode), or Eo

cell

Equation provided on AP EXAM

Sample Exercises 20.5, 20.6, 20.7

The more positive the Eored…

The less reactive, meaning it is reduced

Reduction occurs at the more positive half-reaction

Calculate Eocell Sample Exercises 20.9

Determining ∆Go and K Sample Exercises 20.10