General Chemistry Final Exam

Pi bonds are formed from

unhybridized p orbitals

Pi bonds occur due to a ________ interaction

side on

Valence Bond Theory vs. Molecular Orbital Theory

Valence Bond Theory:

• Both depend on geometry and shape

• Electrons are localized in a given bond/lone pair

• Bond is overlap of half filled atomic orbitals

Molecular Orbital Theory

• Both depend on geometry and shape

• Electrons are delocalized across the whole molecule (Resonance Structures)

• Bonding is a mix of atomic orbitals on all atoms, and molecular orbital is spread over multiple atoms

Lewis Structures fail to explain which concept?

Diamagnetic vs Paramagnetic (example of O2 which appears to be diamagnetic, but it is actually paramagnetic)

Diamagnetic

All electrons are paired

Paramagnetic

At least one electron is unpaired

Constructive Interference

occurs when two atomic orbitals of the same phase meet each other. the electron cloud (molecular orbital) interacts with both nuclei and is more stable.

What bonding character occurs with constructive interference?

Bonding Molecular Orbital

Destructive Interference

occurs when two atomic orbitals of opposite phase meet each other. electron cloud (molecular orbital) interacts less with each nuclei, a node is seen, and it is less stable.

What bonding character occurs with destructive interference?

Antibonding molecular orbital

Bonding orbitals are ____ stable than antibonding orbitals

more stable, lower in energy

Correlation Diagrams

Bond Order

(Number of bonding e - number of antibonding e) / 2

As bond order increase, bond length…

decreases

As bond order increases, bond strength…

increases

What does a bond order of zero mean?

Means that the molecule cannot form. Noble gases are example of this because they have bond order zero, and prefer to exist with an octet of lone pairs instead of forming bonds

Why can core electrons be left out of correlation diagrams?

1. They do not contribute to bonding

2. Core electrons do not overlap well

σ_1s

σ^*_1s

σ_2pz

σ^*_2pz

π_{2px or y}

π^*_{2px or y}

Molecular Orbital Diagram for O₂

Homonuclear diatomics ordering

For diatomics N and earlier, σ_2pz is higher in energy than π_{2px and y}

For diatomics O and later, σ_2pz is lower in energy than π_{2px and y}

This is due to the downward trend in Ionization Energy 1 as your move across a period (since zeff increases and electrons are pulled closer to nucleus)

Polar bonds in heteronuclear diatomics are formed from…

electrons that do not cancel each other out (do not have a bonding/antibonding cancellation) SHOWN IN GREEN

Lone pairs in heteronuclear diatomics are formed from…

electrons that get cancelled between bonding and antibonding orbitals SHOWN IN ORANGE

What attracts a molecule to another molecule?

1. Ion-Ion

2. Intermolecular Forces

What to consider when determining molecular polarity?

1. Lewis Structure

2. Shape (Molecular Geometry)

3. Bond Dipoles (Electronegativity differences)

4. Do Bond Dipoles Cancel? (only if symmetric shape and all bonded atoms on outside are the same)

   1. If cancel, non-polar

   2. If do not cancel, molecule is polar

Types of Intermolecular Forces

1. Dipole Dipole

2. Hydrogen Bonding

3. Dispersion Forces

Dipole Dipole

Interactions between partial charges on polar molecules. Can be attractive (opposite poles) repulsive (same pole) or no impact (one molecule upright and one sideways)

Hydrogen Bonding

H bonding is a particularly strong dipole-dipole interaction between electron poor H (bound to O, N, or F) with a particularly small electron rich element (O, N, or F) with a lone pair

Dispersion Forces

weak attractive forces between atoms and molecules caused by temporary induced dipoles. Always attractive regardless of orientation

Stronger dispersion forces are found in _______

Molecules with bigger electron clouds

Why is ice less dense than liquid water?

The lattice-like structure that forms as a result of hydrogen bonding in solid water forces molecules to be further apart, which means it is less dense than liquid water

Ice: 2 h-bonds per water molecule

Liquid water: 1.8 H-bonds per water moleule

As strength of IMFs increases, boiling point ______.

increases

As strength of IMFs increases, surface tension ______.

increases

As strength of IMFs increases, viscosity ______.

increases. Inversely proportional to temp

As strength of IMFs increases, heat of vaporization ____.

increases

Comparing strength of IMFs

• Always compare molecules of similar size (similar number of large, non H atoms)

• Then, H bonds > Dipole-Dipole > Dispersion (since force is the same in both molecules)

Evaporation

molecules on the surface with sufficient energy break IMFs. Needs no input of outside energy (always happening in a sample at some rate)

Rate of Evaporation

• Constant at a given temperature

• Increases with increasing temperature (avg. kinetic energy increases)

Condensation

If a gas particle hits the surface with low enough energy, it will not have enough energy to move away due to IMFS

• rate increases with increased SA

• rate increases with increased IMFs

Vapor Pressure

Gas pressure when the rate of evaporation = rate of condensation; where P_H20,vap=P_{H20, liquid} Increases as temp increases

Lower vapor pressure means _____ IMFs

stronger (because rate of condensation is higher with stronger IMFs meaning rate evap = rate condensation faster

What occurs in a sealed container

At Start

• Total pressure = 0 (no gas)

• Rate of evaporation > Rate of Condensation (Because no gas currently)

As Time Passes:

• H20 liquid decrease

• H20 gas increase

• Rate of evaporation stays the same

• Rate of condensation increases as P_H20 increases over time

Eventually

• Reach the vapor pressure of water for whatever temp this happens at

• Rate evaporation = rate condensation

Boiling Point

Temperature at which vapor pressure=external pressure

• If pressure increases, the boiling point increases

If given temperature and vapor pressure for a sealed container, can determine how much water will evaporate….

Vapor pressure = Pressure H20 (g)

1. Set up PV=nRT using vapor pressure for P and volume of container for V to find moles of H20 gas at the end

2. Use moles H20 gas to find grams H20 gas (then if its water, density of 1 g/mL means this amt can apply for mL or grams), then subtract this from initial water mass and add to initial mass of gas

Boiling

occurs throughout when P_atm= P_vap

Boiling point is higher when pressure is higher

Use m_objC_sΔT when…

finding q for warming/cooling, ΔKE

Use n_rxnΔH_rxn when…

finding q for a phase change

ΔH_fusion compared to ΔH_vaporization

ΔH_fusion < ΔH_vaporization because for fusion to occur, only need to break some IMFs vs vaporization where all IMFs must be broken

Solid to liquid

ΔH_fusion Endothermic. Can also be known as melting

Liquid to gas

ΔH_vaporization Endothermic. Can also be known as boiling.

Gas to liquid

-ΔH_vaporization Exothermic. Can also be known as condensation

Liquid to solid

-ΔH_fusion Exothermic. Can also be known as freezing

Solid to gas

Sublimation

Gas to Solid

Deposition

Showing change in temperature with phase change (heating curve)

Must get substance to freezing/boiling point using m_objC_sΔT, then use n_rxnΔH_rxn for the phase change at that temperature. Slopes represent heating/cooling and horizontal lines represent phase changes. Shown at constant temp

Phase diagrams

Show phases at different temperatures and pressures (on left).

Triple Point

Temperature and pressure (shown by all lines intersecting) where all phases are present

Critical Point

Above this point, you have a supercritical fluid

Supercritical fluid

• density similar to liquid

• viscosity similar to gas

• fills container like a gas

If solvent-solute interactions > solvent-solvent and solute-solute interactions

ΔH_soln< 0 (exothermic)

Solution forms

If solvent-solute interactions = solvent-solvent and solute-solute interactions

ΔH_soln = 0

Solution forms

If solvent-solute interactions < solvent-solvent and solute-solute interactions

ΔH_soln > 0 (endothermic)

Solution may form

ΔH_soln=

break solvent-solvent interactions (endo) + break solute-solute interactions (endo) + form solute-solvent interactions (exo)

DRIVING FORCE IS ENTROPY

Entropy

measurement of matter or energy dispersal in a system

• more arrangements, more entropy

solubility

maximum amount of solute able to be dissolved in a solvent under given conditions

solubility = max amt of solute / amt solvent

unsaturated solution

concentration < solubility

rate dissolution > rate of recrystallization

saturated solution

concentration = solubility

rate dissolution = rate of recrystallization

oversaturated solution

concentration > solubility

rate dissolution < rate recrystallization

miscible

mix in any ratio (no solubility limit)

solubility: gases with gases

always miscible

solubility: liquids with liquids

like dissolves like (in terms of polarity)

concentration

amt solute / amt solution

• molarity: mol/L

• mass %: mass/mass x 100%

• mole fraction: moles/total moles

Energy diagrams for solution formation

Dissolution of Ionic and Covalent Compounds

Ionic: dissociate into individual ions which are surrounded by solvent

Covalent: dissolve as whole molecules (not breaking covalent bond)

electrolyte

ions in solution conduct electricity (produce ions when put in water)

ionic species and dissolves —> strong electrolyte

strong electrolyte

100% of products form

weak electrolyte

< 100% of products form

non-electrolyte

no products form (aka pathetic)

some molecular solutes form ions by reacting

Solution equilibrium

solubility curves for solids

solubility of a substance v. temperature

With Increased Temperature:

• increased rate of dissolution (by a lot)

• increased rate of recrystallization (by a smaller factor)

• so, solubility typically increases for solids at higher temperatures

temperature dependence of solubility of gases in water

increase of temperature, decrease of solubility

With Increased Temperature:

• increased rate of dissolution (by a smaller factor)

• increased rate of gas bubbling out (by a lot!)

• so, solubility typically decreases for gases at higher temperatures

impact of temperature on solubility of solids v. gas

pressure dependence of solubility of gases in water

(refers to the partial pressure of that gas above water: not dependent on pressures of other gases)

as pressure of that gas increases, solubility increases (more will end up dissolves in the water)

Henry’s Law

Concentration: molarity of gas in aqueous solution

P_gas: pressure of gas above liquid

molarity (M)

moles solute / L solution

molality (m)

moles solute / kg solvent

mole fraction (x)

moles solute / moles solution

(then multiply this by 100% for mole percent)

colligative properties

Nothing to do with the identity of the solute, only dependent on the amount

• Boiling Point T_b

• Freezing Point T_f

• Vapor Pressure (P_vap)

• Osmotic Pressure

Freezing Point Depression: ΔT_f=

i x m x k_f

i: Van’t Hoff Coefficient (for electrolytes), moles of particles formed in solution / moles solute added

m: molality of particles

k_f : freezing point constant (unique to the solvent)

Boiling Point Elevation: ΔT_b=

i x m x k_b

i: Van’t Hoff Coefficient (for electrolytes), moles of particles formed in solution / moles solute added

m: molality of particles

k_b : boiling point constant (unique to the solvent)

Equation for vapor pressure after dissolution of a compound (when solutes are nonvolatile)

P_{vap, soln} = x_solvP°_vap,solv

x_solv: mole fraction (aka how much of the surface is still solvent particles)

P°_vap,solv : vapor pressure of pure solvent

Equation for vapor pressure after dissolution of a compound (when solutes are volatile)

This means that evaporation of solute must be accounted for!

P_{vap, soln} = x_AP°_vap,A + x_BP°_vap,B

same as P_tot = P_A + P_B

Osmotic Pressure

Important Equations to Consider with Ch. 11 Questions

Equation for average rate of consumption or production of a substance, A

Δ[A] / Δt