The final AP chemistry exam will consist of multiple ...



AP Chemistry Final Exam Study Guide

The final AP chemistry exam will consist of 50 multiple choice questions that will assess your mastery of the topics and key concepts outlined below. All equations and any data will be provided in the reference tables or the question itself. You are responsible, however, for being able to determine oxidation states and number of valence electrons for an element using the Periodic Table.

Phase Diagram Interpretation

• Be able to identify triple point ( point where solid, liquid, gas coexist), critical point (temp above which no gas can be liquefied, regardless of the pressure applied)

sublimation/deposition line, melting/freezing line, evaporation/condensation line.

Thermochemistry/Energy and Spontaneity

• Be able to calculate ΔS° given the standard molar entropies or ΔH° given the standard molar enthalpies for the products and reactants. Remember that S, H, G are state functions meaning that they depend on the amount of the substance; this requires that you multiply each enthalpy or entropy value by the coefficient in the balanced equation.

ΔS° = ∑ ΔS°products - ∑ ΔS°reactants remember ° means standard conditions (25°C, 1 atm)

ΔH° = ∑ ΔH°products - ∑ ΔH°reactants

Example: Given the following standard molar entropies measured at 25°C and 1 atm

pressure, calculate ΔS° in (J/K) for the reaction

2 Al(s) + 3 MgO (s) 3 Mg(s) + Al2O3(s)

given Al(s) = 28.0 J/K, MgO(s) = 27.0 J/K, Mg(s) = 33.0 J/K, Al2O3(s) = 51.0 J/K.

ΔS° = ∑ ΔS°products - ∑ ΔS°reactants

= [3(33.0) + 51.0] – [2(28.0) + 3(27.0)]

= 13.0 J/K

• The change in the free energy of a system is given by the Gibbs-Helmholtz equation,

ΔG° = ΔH° - T ΔS°. It also predicts temperature’s effect on spontaneity.

The sign of ΔG can be used to predict the spontaneity of a reaction at constant

temperature and pressure. If ΔG is negative, the reaction probably spontaneous; if ΔG is

positive , the reaction is improbable; if ΔG is 0, the system is at equilibrium and there is

no net reaction. Will it

ΔH ΔS ΔG happen? Comment

• The relationship of the variables in line one of the chart above explains the drive in nature to achieve a minimum of free energy and can be interpreted as the driving force of a spontaneous chemical reaction. Remember when predicting the spontaneity of a reaction, the tendencies in nature are towards lowest energy (-ΔH) and greater entropy (+ΔS).

• Hess’s Law of Heat Summation Calculation

* *

Example:

Be sure to reverse the sign for an equation that is reversed!!

Le Chatelier’s Principle: If a stress is placed on a system in equilibrium, the system will react in the direction that will relieve the stress. Be able to predict the shift in response to various conditions (stress):

Decrease volume (increases the pressure) shifts to side with fewer moles

Increase in [Reactant] favors forward reaction (to remove [R])

Increase in [Product] favors reverse reaction (to remove [P])

Decrease in [Reactant] favors reverse reaction (to restore [R])

Decrease in [Product] favors forward reaction (to restore [P])

Be familiar with the balanced equation for the Haber process for producing ammonia and

apply Le Chatlier’s principle under various conditions as stated above.

3H2 + N2 2 NH3

• Potential Energy Diagrams

Exothermic negative ΔH

AC = energy of products is lower than that of reactants.

PE b Reactants d

c

a Products

e

a = potential energy of the reactants

b = activation energy of the forward reaction

c = ΔH = H products – H reactants

d = activation energy of the reverse reaction

e = potential energy of the products

AC = activated complex

• A catalyst increases the rate of a reaction by providing an alternative reaction pathway

with a lower activation energy but is not permanently changed or used up in the reaction.

- A catalyst does not change the energy of the reactants or products (ΔH).

Bonding

• Recall that bond order is inversely related to bond length: single bonds (bond order of 1) are

longer than double bonds (bond order of 2) and triple bonds (bond order of 3) are shortest.

• Ionic bonds (transfer of e-) electronegativity differences greater than 1.7 (metal-nonmetal)

• Covalent bonds (sharing of e-) electronegativity differences less than 1.7 (nonmetal-nonmetal).

Unequal electronegativity differences between sharing electrons results in a polar covalent bond.

Equal sharing results in a nonpolar bond (ex. Cl2)

• Electronegativity differences determine polarity, but the structure is important. In the case of

CO2 , even though oxygen is electronegative and the molecule has 2 polar bonds, since there

is a line of symmetry, the charges cancel leaving the molecule nonpolar.

_

Water is polar O CO2 is nonpolar O = C = O

H H 2 polar bonds line of symmetry

+ +

• Hydrogen bonding (H attracted to electronegative elements O, F, and N) is responsible for

water’s unusually high boiling point.

• Know these molecules as examples of the basic geometries:

linear (180°) CO2 and BeCl2

trigonal planar (120°) BF3

tetrahedral (109.5°) CH4 (Remember CCl4 is non-polar because of symmetry)

trigonal pyramidal (107°) NH3 (1 lone pair e-)

bent (105°) H2O (2 lone pairs of e-)

Note: the progressive decrease in the bond angles in the series of molecules CH4, NH3,

and H2O results because of the increased number of unshared pair of e-,

CH4 (0 e- pairs), NH3 (1 e- pair), and H2O (2 e- pairs)

Gas Laws and Kinetic Molecular Theory

• Be able to solve gas law problems using the combined gas law P1V1 = P2V2

T1 T2

Boyle’s Law, Charles’ Law, Gay-Lussac’s Law can all be derived from this equation

Remember to convert temperature to Kelvin by adding 273 to ºC

• Be able to use the Ideal Gas Law, PV = nRT when given grams or moles of a gas; be sure to atm units for pressure and L for volume since the gas constant is derived using these units: The gas constant, R, will be given (0.082 L-atm/mol.K). n= mass/MW

• Graham’s Law of Gas Diffusion states that the rate of the diffusion (or effusion) of a gas is inversely proportional to the square root of its molecular mass.

Example: Compare the rate of diffusion of hydrogen to that of oxygen.

Rate H2 = √ 32 = √16 = 4 Hydrogen diffuses 4 times as fast as oxygen.

Rate O2 √ 2 √1 1

• Dalton’s Law of Partial Pressures

• Kinetic Molecular Theory

- Helps to define ideal gases which are assumed to have negligible volume compared to the volume of the container and have no attractive forces.

- A real gas is most like an ideal gas under conditions of high temperature and low pressure.

- As temperature increases, the KE of the gas increases.

Equilibrium

Be able to set up and solve a simple Keq expression from a balanced equation given the concentration of reactants and products.

moles/liter

[P]

[R]

Ignore solids!

Example Problem:

H2 + I2 2HI

I 3 3 0

C -x -x +2x

E 3-x 3-x 2x

Acids, Bases, and Buffers

• Arrhenius acid generates H+ as the only positive ion in solution

Arrhenius base generates OH- as the only positive ion in solution

• The strength of an acid/base is determined by the degree of ionization.

Strong acids/bases ionize almost completely. Example: HCl, H2SO4 strong acids

NaOH, LiOH strong bases

• A buffer solution resists changes in H+ concentration

• H3O+ is the hydronium ion; OH- is the hydroxide ion

• pH = -log [H+] and pOH = -log [OH-]

Example: What is the pH of a 0.1M HCl solution?

pH = -log [10-1]

pH = - (-1)

pH = 1

• pH + pOH = 14; in the above example, the pOH is 13.

• A titration is a neutralization reaction in which a base of known concentration is slowly added to an acid of unknown concentration

• The equivalence point, or endpoint, is the point in the titration when exactly enough base has been added to neutralize the acid originally present. For a strong acid and a strong base, the equivalence point is exactly 7.

Periodic Table and Trends

• Remember that electronegativity (EN) and ionization energy follow the same trends decreasing down a group and increasing across a period.

• Similarly, atomic radius and metallic character follow the same trends decreasing from left to right across a period and increasing from top to bottom in a group.

• The number of valence electrons remains the same as you go down a column on the

Periodic Table. As a result, elements in the same group or family (columns) react similarly.

• The most electronegative element (strongest attraction for e-) is the nonmetal, fluorine with

an EN of 4.0.

• The least electronegative element (lowest ionization energy) is francium (group 1 bottom).

• Know the basic groups alkali metals (Group 1), alkaline earth metals (Group 2), halogens (Group 17), noble gases (Group 18)…

• Know the 7 diatomic elements !!! Some of you are still writing these as monatomic:

Br, I, N, Cl, H, O, F

General Knowledge

• Be able to determine the oxidation number of a particular atom in a compound.

Example: Cr in K2CrO4 is +6 because 2(+1) + 1(x) + 4(-2) = 0, x = +6

K+1 Crx O-2

• Be able to balance a chemical equation.

• Be able to calculate the molar mass (also called formula mass, gram formula mass).

• Be able to determine the number of electrons from a given electron configuration and identify the element

Example: 1s22s22p63s1 total electrons are the sum of all the exponents (11 e- is Na)

valence electrons are the number of electrons in the outermost

principal energy level (1 e- in principal energy level 3)

• Remember that a dilute solution can be saturated and a concentrated solution can be unsaturated.

• Be able to name ionic compounds using the stock numbering system.

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