Map of AP Chemistry

1.1 Moles and Molar Mass

1 mole = 6.022140857×10^23

1.1

Mass Spectroscopy of Elements

1.2

Elemental Composition of Pure Substances

1.3

1.4

1.5

1.6

1.7

1.8

Types of Chemical Bonds

2.1

Ionic Bonds: Attraction between cation (+) and anion (-).

Covalent Bonds:

• Non-Polar (Electrons shared equally)
• Polar (Electrons shared unequally)

• Metallic Bonds: Positive metal ions surrounded by sea of electrons

Understand Partial Charges!

Intramolecular force and potential energy

Intramolecular = forces are the forces that hold atoms together within a molecule

Examples of Potential Energy Graphs

2.2

Justifications using Coulomb's Law will be needed.

Correct

Structure of Ionic Solids

A 3D array that maximizes the attractive forces among cations and anions while minimizing the repulsive forces.

2.3

Incorrect

Structure of Metals and Alloys

2.4

Substitutional Alloy

• Atoms of similiar radii (size)
• One atom substitutes for another

Interstitial Alloy

• Atoms of different radii
• Smaller atoms fill space between larger atoms

Lewis Diagrams

Steps:

Electrons represented by dots, and by bonds.

2.5

Resonance and Formal Charge

A group of two or more Lewis structures that collectively represent a single polyatomic species

Forms with the lowest charge

is the most dominant/common

2.6

VSEPR and Bond Hybridization

2.7

Memorize structures and angles for each number of electron domains

Hybridization: Domains when atoms bond.

-->

SP^(number of electron domains) -1

Intermolecular Forces

3.1

Forces between molecules

Types of intermolecular forces by strength:

• Ionic - Polar
• Hydrogen Bond
• Polar - Polar
• Ion-induced polar
• London Dispersion Forces

3.2

3.3

3.4

3.5

3.6

3.7

3.8

3.9

3.10

3.11

3.12

3.13

Introduction for Reactions

4.1

Physical

No Composition change

Changes in state of matter (solid, liquid, gas, etc)

Chemical

Compoisition change

Change in : Energy (Heat and/or light) (or) Color (rust etc)

Formation of: precipitant (or)

Gas

Net Ionic Equations

4.2

Equations in chemistry must be balanced since you cannot destroy or add molecules from nothing (Law of Conservation of Mass).

There are different ways to represent equations (you can cancel certain things)

--->

Watch videos online for help on this

Representations of Reactions

4.3

There are different ways to represent reactions (in the exam):

• Chemical Symbols (with states of matter)

• Visual Representations

(Try out Questions on AP classroom)

Physical and Chemical Changes

4.4

Physical Changes:

Involve changes (breaking and forming) in Intermolecular Forces

Chemical Changes:

Involve breaking/forming chemical bonds

Dissolution can be reffered as either change. (Salt's ionic bond is broken and ion dipole interactions are made).

Stoichiometry

4.5

Used to predict amount of products produced or reactants needed in a reaction etc.

In reactions there are limiting reactants:

Introduction to Titration

4.6

A laboratory technique used to determine the concentration of a substance in a solution by reacting it with a standardized solution of another substance.

Slowly add the titrant to the analyte, and measure the pH of the solution (until solution turns pink).

At this equivilance point (pH of 7), there is equal number of base and acid, so using the chemical equation you can find the number of moles of the sunstance there is. Then find the molarity.

Look at AP classroom for redox titrations.

Types of Chemical Reactions

There are 3 important types of chemical reactions:

• Acid-Base
• Redox
• Precipitation

Will go into more depth in later sub-units.

4.7

Introduction to Acid-Base Reactions

When an acid and base react, a neutralization reaction occurs

Bronsted Lowry Acids and Bases:

A bronsted lowry acid is a proton (H+) donor, and a base is a proton (hH+) acceptor

4.8

Oxidation - Reduction Reactions (Redox)

4.9

In a redox reaction electrons are transferred from one speicies to another

Sunstance that

Loses electon = Oxidized

Gains electron = Reduced

(OILRIG)

Look how to do half reactions online.

5.1

5.2

5.3

5.4

5.5

5.6

5.7

5.8

5.9

5.10

5.11

9.1

9.2

9.3

9.4

9.5

9.6

9.7

9.8

9.9

9.10

8.1

8.2

8.3

8.4

8.5

8.6

8.7

8.8

8.9

8.10

Introduction to Equilibrium

7.1

• In reversible reactions, the forward and reverse rate will proceed at the same rate.

--> Equilibrium

• No observable change.

• Concentration stays the same

Direction of Reversible Reactions

Rate of Forward --> More Products

Rate of Reverse --> More Reactants

7.2

Reaction Quotient and Equilibrium Constant

7.3

Reaction Quotient Q = The ratio (amount) of products and reactants present during a reaction at a particular point in time

Equilibrium Constant K =The ratio (amount) of products and reactants present during equilibrium

Solids and Liquids are not included

Calculating the Equilibrium Constant

K(eq) values change with temperatrue (will not be on test)

7.4

Magnitude of the Equilibrium Constant

7.5

If K>1 there is a higher conc. of products

If K<1 there is a higher conc. of reactants

If K is extremely large --> Goes to completion

If K is extremely small --> Forward reaction doesn't occur

Properties of the Equilibrium Constant

7.6

• Reverse a reaction = Inverse of K value

• For more than one elementary reaction = K of the overall process is the products of the K values for each step

• If stoichiometric coefficients are multiplied by a factor = exponents in K are also raised to that factor

• Applies to Q and K(p) etc.

Calculating Equilibrium Concentrations

Use ICE (Or RICEK'S) to solve for initial concentrations/pressures

7.7

Representations of Equilibrium

Graphic models representing

equilibrium

7.8

Introduction to Le Châtelier’s Principle

7.9

Changing conditions can knock system out of equilibrium (stressing)

Molarity, partial pressure, temperature, volume/pressure

Keq does not change.

When there is an increase in pressure, the equilibrium will shift towards the side of the reaction with fewer moles of gas.

Temperature

Reaction Quotient and Le Châtelier’s Principle

When equilibrium is shifted Q is no longer equal to K

Use Q to justify response of shift

7.10

Introduction to Solubility Equilibria

K(sp) represents the level at which a solute dissolves in solution.

The more soluble a substance is, the higher the K(sp) value it has.

7.11

Common-Ion Effect

7.12

Common Ions

Example: NaCl, and KaCl

Q(sp) > K(sp) precipatants form.

pH and Solubility

Some solubility equilibria are sensitive to changes in pH.

(addition of OH- ion etc)

7.13

Free Energy of Dissolution

7.14

Exothermic

Entropy

Temperature

Negative Gibbs Free energy = Spontaneous reaction

(constant temperature and constant pressure.

6.1

6.2

6.3

6.4

6.5

6.6

6.7

6.8

6.9