AP Chem Unit 6 Test

Lattice energy

The energy required to break apart a crystal lattice of an ionic compound to form gaseous ions or the energy released when an ionic compound crystal lattice is formed form gaseous ions

How does Coulomb’s law work with lattice energy?

higher attractive force means greater lattice energy, higher charge and lower distance (r² in denominator) means higher attractive force

Bonding in ionic solids

Ionic (attraction between cations and anions)

examples of ionic solids

NaCl, Cs2O

Bonding in noble gases

london dispersion forces

Properties of noble gases

• Very low boiling point (gases at room temp)

• Don’t conduct in any phase

• Insoluble in water but soluble in nonpolar solvents

Bonding in molecular solids

dispersion, dipole-dipole, or H-bonding (IMFs)

Properties of molecular solids

• Generally low boiling point

  • Gas, liquid, or solid at room temperature

• Don’t conduct in any phase

• Some are insoluble in water but soluble in nonpolar solvents; depends on polarity

Examples of molecular solids

Sugar, H2O, N2

Bonding in metallic solids

Delocalized electrons

Properties of metallic solids

• Generally high melting point (usually solids at room temp)

• Conduct in all states

• Insoluble in water

• Malleable

• Ductile

Bonding in network covalent solids

Directional covalent (like one giant molecule)

Properties of network covalent solids

• High melting point

• Does not conduct in any state

• Insoluble

• Hard

• Brittle

examples of network covalent solids

Diamond, graphite

Alloy

A mixture of two or more elements, where at least one is a metal

Substitutional alloy

An alloy where the two elements are similar in size/radii and one type of atom replaces the other 

e.g. for brass made of Cu and Zn, Zn replaces a couple Cu atoms)

Interstitial alloy

An alloy where some the smaller atom fits into interstices (holes)

e.g. for carbon steel, carbon fits into holes between Fe, disrupting sea of electrons and making it harder for metal atoms to move past each other

P vs V

inverse, exponential

P vs 1/v

direct

P vs n

direct

V vs T

direct

P vs T

direct

Difference between barometer and manometer

A barometer measures the pressure of the atmosphere in mmHg, and a manometer measures differences in pressure

How to read a manometer

Pgas = Patm ± height (depending on if gas or atm is pushing harder)

Combined Gas Law

P1V1/T1 = P2V2/T2

Ideal Gas Law

PV=nRT

Avogadro’s principle

V=kn where k is a constant

Graham’s law

rateA/rateB = root(MMB/MMA)

Dalton’s Law

The partial pressure of a gas in a mixture is proportional to its mole fraction

Useful equation for gas density/MM calculations

MM=dRT/P

formula for finding pressure of gas when its collection over water

Pgas = Ptotal - PH2O

Relationship between temp and KE

Two gases at the same temp have the same average KE

Ideal gases

• Gas molecules are in constant random motion

• Gas molecules do not attract or repel each other

• Gas molecules are far apart, and so small compared to the distance between them that the individual particles’ volume is negligible

• Gases at the same temperature have the same average kinetic energy of molecules

When do ideal gases start behaving like real gases?

• Low temperature, high pressure

  • When particles are moving slowly enough that they become attracted to each other

  • The container is small so the size of the gas particles isn’t actually that small in comparison to the distance between the particles (particle volume not negligible)

What are non-ideal conditions for gases?

high pressure (particles close together so their volume isn’t negligible) and low temperature (particles get attracted to each other)

Boiling point

The temperature at which the vapor pressure equals the external pressure

miscible vs immiscible

solutions mix vs do not mix

heat of solution

the energy gained or lost when one mole of compound is dissolved in water

energy of hydration

energy gained or lost when gaseous ions form IMFs with water (ion-dipole interaction)

How to calculate delta H of solution

Lattice energy of compound + Hydration energy of ion A + Hydration energy of ion B

How does solubility of solids vs gases vary with temperature

For solids: higher temp, higher solubility

For gases: higher temp, lower solubility

unsaturated solution

a solution that can dissolve more solute

saturated solution

a solution that can’t dissolve more solute (at a given T, solution can dissolve only a fixed amount of solute)

supersaturated solution

a solution where more solute is dissolved in it than should be at that temp

how are supersaturated solutions created?

heat a saturated solution, dissolve more solute until saturated, and then slowly cool

dilution rule 

M1V1=M2V2

common mistake when using dilution rule

if its asking for how much water added, subtract the initial volume from V2

properties of ionic solids

• High melting point (always solids at room T)

• Solids don’t conduct but molten and aqueous phases do

• Many soluble in water

• Brittle

factors that affect solubility (AMOUNT that can dissolve)

just temperature

factors that affect rate of dissolution (how FAST it dissolves)

stirring (more stirring = faster), surface area (less SA = faster), temp (higher temp=faster)

colloid vs suspension

particles stay dispersed vs larger particles sink to the bottom

avogadro’s hypothesis in gas stoichiometry

under the same T and P before and after reaction, all gases have volume in proportion to their moles, so coefficients in balanced equations represent volume ratios as well as mole ratios