AP Exam Review Session I

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AP Exam Review Session I

1. Strategies for Test 2. Math

3. Lab Equipment 4. Uses

5. Atomic Structure/Electron Configuration 6. Nuclear Chem

7. Periodicity

8. Bonding 9. Stoichiometry 10. Gases

11. Solutions/Solubility 12. NIE

13. Phase Change Diagram 14. Organic Chem

I. Strategies for Test

A. Multiple Choice – 3 pass method

B. Identify exactly what the question is asking 1. Especially important in free response! II. Math

A. Calculations without calculators B. Sig Figs

C. Percent Error and Percent Yield III. Lab Equipment

A. Proper piece for proper use B. Sciencegeek.net practice quiz IV. Uses and Properties of :

A. Wavelengths

1. X-ray; Ultraviolet; gamma; visible, infared

B. Chemical Compounds

1. Sulfur, Mercury, ethanol, methanol, hydrogen peroxide, chloroflurocarbons, ozone V. Properties

A. Standard State – what is the standard state for each element 1. Just know the exceptions (Mercury, Bromine, Iodine) B. Elements, families

1. Alkali metals react strongly with water to produce hydrogen gas and basic solutions, shiny, not found alone in nature, quite malleable

2. Alkaline earth, transition metals, metalloids, etc. VI. Atomic Structure-Electron Configuration

A. E-config – Noble Gas Notation

B. Atomic Orbitals – rules (Aufbau, Hunds, Pauli Exclusion)

1. Know what happens when electrons are placed in different orbitals 2. Which atoms or ions are identified with the orbital

C. Quantum Numbers

D. Using above written descriptions 1. Ground State and Excited State 2. Valence electrons

3. Ions


E. Bohr model

F. Diamagnetism vs. paramagnetism

G. Energy and other relationships

1. amplitude

2. frequency

3. momentum

4. photoelectric effect

5. Planck’s constant

6. Rydberg constant

7. Velocity

8. Wavelength

H. Important people, experiments, and theories

1. Rutherford, Millikan, Bohr, Einstein

VII. Nuclear Chemistry

A. Nuclear reactions vs. Chemical reactions B. Emissions and their equations

1. Alpha

2. Beta

3. Positron

4. Electron capture

5. Gamma

C. Balancing a nuclear reaction

D. Nuclear stability and predicting the decay type E. Half Life – dividing or multiplying by two!

F. Binding Energy and mass defect

1. How to write a short paragraph

G. Fission vs. Fusion

1. How to identify

VIII. Periodicity

A. Trends

1. Ionization energy

a. Amount of energy required to remove an electron form an atom to form a positive ion (cation). 1mole of electrons in the gaseous state.

b. Helium has the highest ionization energy

c. Shielding is a factor, reason why decreases as move down the table d. Nuclear Charge another factor, why increases to the right

e. Successive ionization energy tables! 2. Electronegativity

a. The ability to attract electrons - 1mole of electrons in the gaseous state. b. Difference in electronegativities between atoms determines type of bond c. Fluorine highest

d. Shielding and nuclear charge…

3. Electron affinity

a. Amount of energy released when an electron is added to an atom - 1mole of

electrons in the gaseous state.

b. Electron affinity of zero for elements that have full shells

4. Atomic radii

a. Decreases as nuclear charge increases, but more shells increases size. b. Largest radii are left and bottom of table

5. Ionic radii

a. Cations are always smaller than their neutral atom counterpart b. Anions are always larger than their neutral atom counterpart


7. shielding

B. Chemical properties and periodicity

C. Metals vs. nonmetals

IX. Bonding

A. Lewis Dot diagrams 1. General

a. Bi

b. Cl- c.d. CCl4NH3

2. Electron deficient (incomplete)

a. BF3

b. BeH2

3. Expanded octet

a. PF5

b. PO2F2 -B. Ionic Bonding

1. Cation + Anion 2. Naming

3. Coulomb’s Law

a. b.

i. E = Electrostatic energy in J

ii. k = a proportionality constant

iii. (+q) or (Q1) = charge on cation

iv. (-q) or (Q2) = charge on anion

v. d or r = distance between the nuclei of the two atoms in nm

C. Lattice Energy

D. Covalent Bonding 1. Nonmetals 2. Naming 3. Bond length

a. Single bonds are the longest, then double bonds, and triple bonds are the shortest b. Bond Orders

E. Intermolecular Force and Intramolecular Force 1. Covalent

a. Network Covalent Bonds – extensive 3-D structure and no discrete molecular

units. They are very hard and have extremely high melting points.

b. Examples:

i. SiO2 – silicon dioxide, in sand

ii. SiC – silicon carbide, used in grinding and abrasive

iii. C (diamond) – one of the hardest known substances in nature

2. Hydrogen Bonding

a. Stronger than most other intermolecular forces b. Weaker than ionic or covalent bonds

c. H-O, H-N, H-F


a. Variable hardness and melting point b. Most often transition metals (Ni, Fe) 4. dipole-dipole

a. between polar molecules, positive end attracted to negative end (CH2F) b. Measured in units of debye’s. More debye’s, more polar

5. dipole-induced dipole

a. ions and a polar molecules (NaCl in water) b. Polarizability

6. London Dispersion – weakest

a. H2,

b. Stronger the larger the molecules (larger the atoms)

F. Radicals

1. NO and NO2 (one unpaired electron) – sources of acid rain

G. Bond Strength


H. Resonance

I. Shapes with VSEPR 1. Dipole Moments 2. Shape of molecules

a. Linear, bent, trigonal planar, tetrahedral, trigonal pyramidal, seesaw, T-shaped, octahedral, square pyramidal, square planar

3. Hybridizaiton

a. Example: d2sp3, or sp3d2 = SF6 4. Angles for each shape

5. Nonbonding vs. bonding pairs, electron repulsion X. Stoichiometry

A. Balancing

1. Regular equations and redox reactions, multiple choice B. Make Solutions

1. From mass, volume, moles and Molarity C. Percent composition

D. Empirical vs. Molecular Formulas

XI. Gases

A. Solving problems with the following gas laws:

1. Avogadro’s Law

2. Boyle’s Law

3. Charles Law

4. Gay-Lussac’s Law

5. Combined Gas Law

6. Dalton’s Law of Partial Pressure - Working with mole fractions a. Grams to moles to total moles, mole fraction and pp 7. Ideal Gas Law

a. Manipulate the equation to find molar mass b. Manipulate the equation to find density

8. Graham’s Law of Diffusion-Effusion

9. Root mean square speed

B. Labs


C. Kinetic Molecular Theory – 5 postulates

D. Properties of Gases

XII. Solubility

A. Solubility Rules

1. Cl, Br, I likes APH;

2. F, watches CBS-PM (at night) 3. Sulfates watch CBS/PBS; B. Solubility of slightly soluble salts

1. Calculate and interpret the value for Ksp

2. calculation of solubility from Ksp and vice versa C. Solve problems with the common ion effect

D. Polar vs. Non-Polar 1. Like dissolves like

2. Polar compounds are immiscible in non-polar compounds and vice versa E. Solubility of a gas

1. Depends upon pressure (why sodas are under pressure, so carbon dioxide is dissolved into the liquid)

2. Collecting gas over water

F. Molarity (moles of solute per liter of solvent) XIII. Net Ionic Equations

A. Combination and Decomposition 1. Redox

a. A + B  AB b. AB  A + B 2. Non-Redox

a. Oxides + Water  Acid (NM) or Base (M) b. Acid (NM) or Base (M)  Oxides + Water c. Oxides 

B. Manganese (IV) oxide is added to a solution of hydrobromic acid.

1. What is the electronic structure of the manganese product? Explain

C. A solution of potassium permanganate is added to a solution of concentrated hydrochloric acid. 1. What substance is oxidized in the reaction?

D. A solution of potassium dichromate is added to an acidified solution of sodium sulfite. 1. What color would you expect to observe.

E. Redox

1. Oxidation Numbers/rules

2. Half reactions in acidic and basic medium (basic just acidic with OH- added at end) 3. iodide ions plus iron (III) ions

4. solid sodium and water

5. tin (II) sulfate and iron (III) sulfate 6. Magnesium in iron (III) chloride

a. Which is oxidized? F. Acids Bases

1. Strong Acid + Strong Base (HCl + NaOH  H2O + NaCl; H+ + OH-  H2O) 2. Strong Acid + Weak Base

a. HCl (aq) + CH3NH2 (aq) CH3NH3Cl (aq) b. H+ (aq) + CH3NH2 (aq) CH3NH3+ (aq) 3. Strong Base + Weak Acid


b. HC6H5CO2 (aq) + OH¯ (aq) C6H5CO2¯ (aq) + H2O (l) G. Solubility/Precipitation

1. calcium hydroxide + magnesium chloride 2. silver nitrate + sodium chromate

3. silver nitrate and sodium chromate

a. What two types of reactions is this?

4. Coordination Complex, Complex Metal Ions (Lewis Acids) a. Common ligands

b. Double metal charge to find number of ligands

5. Disproportionation

6. Combustion, Esterification, Halogenation, Hydrohalogenation

H. Combustion 1. Organic

2. Metals with O2

a. Lithium metal is burned in air.

b. What is the change in oxidation state of the Li? 3. Metals with N2

I. Questions

1. Lewis Dots 2. Thermo

3. Reaction Types

4. Oxidation number changes 5. Reaction orders

XIV. Phase Change Diagram A. Temp vs Time

1. What each line represents

B. Pressure vs. Temp

1. Triple Point

2. Critical Point

3. What each region represents

XV. Organic Chemistry

A. Number of Carbons B. Draw condensed formulas

C. Draw and name various hydrocarbons

1. Alkanes, alkenes, alkynes, cyclic hydrocarbons

2. How to draw and name various ring (aromatic) compounds 3. Functional Groups

a. Alcohols

b. carboxylic acids c. amides

d. amines

e. aldehydes f. ketones g. esters h. ethers

4. Isomers

a. Structural

b. Chain

c. Positional

d. Functional

e. Stereoisomerism

i. Geometric