AP Chemistry Exam

The same as the number of protons in the nucleus of an element; it is also the same as the number of electrons surrounding the nucleus of an element when it is neutrally charged.

The sum of an atom's neutrons and protons

Atoms of an element with different numbers of neutrons

6.022×10²³ particles per one mole

grams/molar mass

Pressure = 1 atm Temperature = 273 K

I mole of gas = 22.4 L

Moles = (molarity)(liters of solution)

The percent by mass of each element that makes up a compound. It is calculated by dividing the mass of each element or component in a compound by the total molar mass for the substance.

• Represents the simplest ratio of one element to another in a compound

• Start by assuming a 100 g sample

• Convert percentages to grams

• Convert grams into moles

• Divide each mole value by the lowest of the values

• These values become the subscripts

• Determine the molar mass of the empirical formula

• Divide that mass into the molar mass x = m/e x = molar mass/ empirical mass

• Multiply all subscripts in the empirical formula by the value of x

States that when building up the electron configuration of an atom, electrons are placed in orbitals, subshells, and shells in order of increasing energy.

States that the two electrons which share an orbital cannot have the same spin. One electron must spin clockwise, and the other must spin counterclockwise.

States that when an electron is added to a subshell, it will always occupy an empty orbital if no one is available. Electrons always occupy orbitals singly if possible and pair up only if no empty orbitals are available.

The amount of energy that an electron has depends on its distance from nucleus of an atom. While on the exam, you will not be required to mathematically calculate the amount of energy a given electron has, you should be able to qualitatively apply Coulomb's Law.

Essentially, the greater the charge of the nucleus, the more energy an electron will have.

Max Planck figured out that electromagnetic energy is quantized. That is, for a given frequency of radiation (or light), all possible energies are multiples of a certain unit of energy, called a quantum (mathematically, that's E = hv). So, energy changes do not occur smoothly but rather in small but specific steps.

ΔE = hv = hc/λ

ΔE = energy change h = Planck's constant, 6.626×10⁻³⁴ J∙s v = frequency of the radiation λ = wavelength of the radiation c = the speed of light, 3.00×10⁸ m/s

c = λv

Inversely proportional

c = speed of light in a vacuum (2.998×10⁸ m/s) λ = wavelength of the radiation v = frequency of the radiation

The amount of energy necessary to remove an electron from an atom.

A chart of the amount of ionization energy for all electrons ejected from a nucleus.

The y-axis describes the relative number of electrons that are ejected from a given energy level.

The x-axis shows the binding energy of those electrons.

High temperature and low pressure

s-subshell holds two electrons

p-subshell holds six electrons

d-subshell holds 10 electrons

f-subshell holds 14 electrons

The complete description of the energy level and subshell that each electron on an element inhabits

States that it is impossible to know both the momentum of an electron at a particular instant.

Atomic radius decreases across a period

Atomic radius increases down a group

Cations are smaller than their atoms

Ions are larger than their atoms

The energy required to remove an electron from an atom.

Refers to how strongly the nucleus of an atom attracts the electrons of other atoms in a bond.

Across the periods

• atomic radius decreases

• ionization energy increases

• electronegativity increases

Down the periods

• atomic radius increases

• ionization energy decreases

• electronegativity decreases

An ionic solid is held together by the electrostatic attractions between ions that are next to one another in a lattice structure.

Occurs between a metal and a nonmetal; electrons are not shared, they are given up by one atom and accepted by another.

Substances with ionic bonds are usually solids at room temperature and have very high melting and boiling points.

Ex. NaCl

1. Charge on ions - a greater charge leads to a greater bond energy

Ex. MgO will have a higher melting point than NaCl

2. Size of ions - smaller ions will have greater attraction

Ex. LiF will have a greater melting KBr

Metal atoms with two different radii combine

Ex. In steel, much smaller carbon atoms occupy the interstices of the iron atoms

Forms between atoms of similar radii

Ex. Atoms of zinc are substituted with copper atoms to create an alloy

Bonding in which two atoms share electrons. Each atom counts the shared electrons as part of its valence shell to achieve complete outer shells.

The first covalent bond formed between two atoms is called a sigma bond.

Bond designation: One sigma

Bond order: One

Bond length: Longest

Bond energy: Least

Bond designation: One sigma and one pi

Bond order: Two

Bond length: Intermediate

Bond energy: Intermediate

Bond designation: One sigma and two pi

Bond order: Three

Bond length: Shortest

Bond energy: Greatest

In a network solid, atoms are held together in a lattice of covalent bonds.

They are very hard and have very high melting and boiling points.

Ex. The most commonly seen network solids are compounds of carbon (such as diamond or graphite) and silicon (SiO₂ quartz)

Much stronger than dipole-dipole forces.

Substances that have hydrogen bonds have higher melting and boiling points. Ex. water, H₂O and ammonia, NH₃

*Water is less dense as a solid than as a liquid because its hydrogen bonds force the molecules in ice to form a crystal structure, which keeps them apart than they are in the liquid form

Forces that occur when the positive end of one polar molecule is attracted to the negative end of another polar molecule.

Molecules with greater polarity have higher melting and boiling points.

Dipole-dipole attractions, however, are relatively weak, and these substances melt and boil at very low temperatures.

Forces that occur between all molecules.

These very weak attractions occur because of the random motions of electrons on atoms within molecules. At a given moment, a nonpolar molecule might have more electrons on one side than on the other, given it an instantaneous polarity.

Molecules with more electrons will experience will experience greater London dispersion forces, and therefore have generally higher melting and boiling points.

London forces are even weaker than dipole-dipole forces, so substances that have only London dispersion forces melt and boil at extremely low temperatures and tend to be gases at room temperature.

1. Network Covalent Bonds

2. Ionic Bonds (based on Coulombic attraction) a. Greater Ion Charge b. Smaller Atom Size

3. Covalent Bonds (based on molecular polarity) a. Hydrogen Bonds b. Non-Hydrogen Bond Dipoles c. London Dispersion Forces (temporary dipoles) i. Large molecules are more polarizable because they have more electrons.

1. Substances with weak intermolecular forces (LD), tend to be gases at room temperature.

2. Substances with strong intermolecular forces (HB) tend to be liquids at room temperature.

3. Because ionic bonds are generally significantly stronger tan intermolecular forces in covalent molecules, ionic substances are usually solid at room temperatures. Ionic substances do not experience intermolecular forces.

Arises from the fact that molecules inside a liquid are in constant motion. If those molecules hit the surface of the liquid with enough kinetic energy, they can escape the intermolecular forces holding them to the other molecules and transition them into the gas phase.

This is not to be confused with a liquid boiling. In order for vaporization to occur, no outside energy needs to be added.

Some atoms are stable with less than eight electrons in their outer shell.

- Hydrogen only requires two electrons

- Boron is considered to be stable with only six electrons, as in BF₃

In molecules that have d subshells available, the central atom can have more than eight valence electrons, but never more than twelve.

• PCl₅

• SF₄

• XeF₄

If more than one valid Lewis structure exists for a molecule, formal charge can be used to determine the more likely structure.

The amount of heat lost or gained when one mole of a compound is formed from its constituent elements.

The enthalpy change for a reaction is equal to the sum of the enthalpy of formation of all the products minus the sum of the enthalpy of formation of all the reactants.

Electrons repel each other, so when atoms come together to form a molecule, the molecule will assume the shape that keeps its different electron pairs as far apart as possible.

• Central atom with 2 electron pairs

• Zero lone pairs

• sp hybridization

• Ex. BeCl₂ and CO₂

B - A - B

• Central atom with three electron pairs

• Zero lone pairs

• sp² hybridization

• 120° bond angles

• Ex. BF₃, SO₃, NO₃⁻, CO₃²⁻

• Central atom with three electron pairs

• One lone pair

• sp² hybridization

• 120° bond angle

• Ex. SO₂

The amount of heat required to raise the temperature of one mass unit of a substance by 1.00°C

• Central atom has four electron pairs

• Zero lone pairs

• sp³ hybridization

• 109.5° bond angle

• Ex. CH₄, NH₄⁺, ClO₄⁻, SO₄²⁻, PO₄³⁻

• Central atom has four electron pairs

• One lone pair

• sp³ hybridization

• 109.5° bond angle

• Ex. NH₃, PCl₃, AsH₃, SO₃²⁻

• Central atom has four electron pairs

• two lone pair

• sp³ hybridization

• 109.5° bond angle

• Ex. H₂O, OF₂, NH₂⁻

• Central atom has 5 electron pairs

• Zero lone pairs

• dsp³ hybridization

• Ex. PCl₅, PF₅

• Central atom has 5 electron pairs

• One lone pair

• dsp³ hybridization

• Ex. SF₄, IF₄⁺

• Central atom has 5 electron pairs

• Two lone pairs

• dsp³ hybridization

• Ex. ClF₃, ICl₃

• Central atom has 5 electron pairs

• Three lone pairs

• dsp³ hybridization

• Ex. XeF₂, I₃⁻

• Central atom has 6 electron pairs

• Zero lone pairs

• d²sp³ hybridization

• Ex. SF₆

• Central atom has 6 electron pairs

• One lone pair

• d²sp³ hybridization

• Ex. BrF₅, IF₅

• Central atom has 6 electron pairs

• Two lone pairs

• d²sp³ hybridization

• Ex. XeF₄, ICl₄⁻

A measure of molecular randomness, or disorder

• The kinetic energy of an ideal gas is directly proportional to its absolute temperature: The greater the temperature, the greater the average kinetic energy of the gas molecules.

• There are no forces of attraction between the gas molecules in an ideal gas.

• Gas molecules are in constant motion, colliding with one another and with the walls of their container without losing any energy.

PV = nRT

R = .08206 Latm/molK

As pressure increases, temperature increases

As pressure increases, volume decreases, and vice versa

As temperature increases, volume increases

Ptotal = Pa + Pb + Pc + ...

Pa = (Ptotal)(Xa)

Xa = moles of gas A/total moles of gas

At low temperature and/or high pressure, gases behave in a less-than-ideal manner. This is because the assumptions made in the kinetic molecular theory become invalid.

D = m/v

Expresses the concentration of a solution in terms of volume.

M = moles of solute / liters of solution

Mole fraction gives the fraction of moles of a given substance (S) out of the total moles present in a sample.

Mole Fraction (Xa) = moles of substance S / total number of moles in solution

"Like dissolves like"

A basic rule to remember which solutes will dissolve in which solvents.

Polar or ionic solutes (such as salt) will dissolve in polar solvents (such as water).

Nonpolar solutes (such as oils) are best dissolved in nonpolar solvents.

When an ionic substance dissolves, it breaks up into ions in a process called dissociation. Free ions in a solution are called electrolytes because they can conduct electricity.

Electron loss

Electron gain

Reactant that is reduced (gains electrons)

Reactant that is oxidized (loses electrons)

Not requiring an outside source of energy to proceed

The rate of diffusion of a gas molecule is inversely proportional to the square root of that molecule's mass.

The amount of energy in a system that is available to do useful work.

-ΔG is spontaneous

+ΔG is nonspontaneous

When ΔG = 0 the reaction is at equilibrium

A solution consisting of a weak acid plus its conjugate base or a weak base plus its conjugate acid.

This solution resists changes to its pH

Solid that has its atoms arranged in an orderly way.

Solids whose particles have no orderly pattern.

1. Compounds with an alkali metal cation (Na⁺, Li⁺, K⁺, etc) or an ammonium cation (NH₄⁺) are always soluble.

2. Compounds with a nitrate (NO₃⁻) anion are always soluble.

When bonds are broken, energy is released.

When bonds are formed, energy is absorbed.

If the products have stronger bonds than reactants, then the products have lower enthalpy than the reactants and are more stable; in this case, energy is released by the reaction, or the reaction is exothermic.

If the products have weaker bonds than the reactants, then the products have higher enthalpy than the reactants and are less stable; in this case, energy is absorbed by the reaction, or the reaction is exothermic.

The amount of energy required to reach the transition state, the highest point on the graph. At this point, all reactant bonds have been broken, but no product bonds have been formed, so this is the point in the reaction with the highest energy and lowest stability.

Speed up a reaction by providing the reactants with an alternate pathway that has a lower activation energy.

A catalyst lowers the activation energy, but it has no effect on the energy of the reactants, the energy of the products, or the ΔH of the reaction.

In a galvanic cell, a favored redox reaction is used to generate a flow of current.

Two half-reactions take place in separate chambers, and the electrons that are released by the oxidation reaction pass through a wire to the chamber where they are consumed in the reduction reaction. That's how the current is created.

If the concentration of the products in a voltaic cell increases, the voltage decreases. If the concentration of the reactants increases, the voltage increases.

Oxidation takes place at the anode

Reduction takes place at the cathode

Where reduction occurs and the solution is becoming less positively charged, the positive cations from the salt bridge solution flow into the half-cell.

Where oxidation occurs and the solution is becoming more positively charged, the negative anions from the salt bridge solution flow into the half-cell.

1. If you know the current and time, you can calculate the charge in coulombs.

I = q/t

I = current (amperes, A) q = charge (coulombs, C) t = time (seconds, s)

2. Once you know the charge in coulombs, you know how many electrons were involved in the reaction.

moles of electrons = (coulombs) / (96,500 coulombs/mol)

3. When you know the number of moles of electrons and you know the half-reaction for the metal, you can find out how many moles of metal plated out.

4. Once you know the number of moles of the metal, you can convert this to the number of grams of the metal.

Rate = k[A]

The rate law for a first order reaction uses natural logarithms.

The use of natural logarithms in the rate law creates a linear graph comparing concentration and time.

The slope of the line is given by -k and the y-intercept is given by ln[A]₀

Describes the amount of time it takes for half of a sample to react.