exam one

ch12, ch13, ch14.1-14.3

gas’s shape and volume?

variable shape and volume

liquid’s shape and volume?

variable shape, fixed volume

solid’s shape and volume?

fixed shape and volume

intramolecular forces?

• strong

• within molecules

• covalent and ionic bonds

• remains after a physical process

intermolecular forces

• weaker

• between molecules

• affected from physical processes

polar molecule

• uneven distribution of charges

• permanent dipole moment on molecule

nonpolar molecule

• even distribution of charges

• no permanent dipole moment on molecule

London dispersion forces

• short-lived attractions between instantaneous and induced dipoles

• occurs in all molecules because electrons are always moving

• larger atoms = stronger LDFs = higher boiling points

• strongest in heavier molecules and molecules with greater surface area (therefore linear molecules have stronger LDFs than rings)

instantaneous dipoles

occur due to asymmetric distribution of electrons

induced dipoles

asymmetric charge in one molecule causes asymmetry in an adjacent molecule

dipole-dipole attractions

electrostatic attractions between oppositely charged ends of polar molecules

hydrogen bonds (H-bonds)

attraction between partially positive H on one molecules and partially negative F,O,N on another molecule

ion-dipole attractions

electrostatic attraction between an ion and polar molecule 

strongest IMF

what does △H stand for in phase changes?

• enthalpy change for one mol of the substance

• enthalpy = energy absorbed/released during the phase change

+△H

endothermic

energy must be absorbed from surroundings

-△H

exothermic

energy is released to surroundings

name and enthalpy of phase change:

solid → gas

sublimation, +△H

name and enthalpy of phase change:

gas → solid

deposition, -△H

name and enthalpy of phase change:

liquid → gas

vaporization, +△H

name and enthalpy of phase change:

gas → liquid

condensation, -△H

name and enthalpy of phase change:

solid → liquid

melting/fusion, +△H

name and enthalpy of phase change:

liquid → solid

freezing, -△H

heating/cooling curve

shows how temperature changes as pure substance is heated

• flat regions = phrase changes

how to calculate energy during a constant phase vs during a phase change

• at a constant phase: q = mc△T

• during a phase change: q = n△H_vap

(both are not on formula sheet)

vapor pressure definition

pressure of vapor on container walls when at equilibrium

occurs when evaporation + condensation are balanced

what affects vapor pressure?

• increasing temperature = molecules move faster = more molecules become vapor = increased vapor pressure

• stronger IMFs = harder for molecules to escape phase = lower vapor pressure

• weak IMFs = high vapor pressure

boiling point

temperature when vapor pressure of liquid = external atmospheric pressure (equilibrium)

what affects boiling point?

stronger IMFs = higher boiling point because more energy is required to break the bonds

relationship between IMFs, vapor pressure, and boiling point

IMFs ∝ 1/VP ∝ boiling point

• NOTE: not a direct linear relationship

surface tension

tendency of a liquid to bead up rather than spread out

viscosity

liquid’s resistance to flow (honey vs water)

what affects surface tension?

• strong IMFs = high surface tension

what affects viscosity?

• strong IMFs = high viscosity

enthalpy of vaporization formula

• P_{vap, T} = vapor pressure

• △H_vap = enthalpy of vaporization

• R = universal gas constant = 8.314 J/mol*K

• T = temperature (in K)

NOTE: this formula is on equation sheet, but is not given for ACS exam

phrase diagram

phase of a substance under all pressure-temperature combinations

• x-axis: temperature

• y-axis: pressure

critical point on phase diagram

above this point, the Temperature and Pressure conditions allow liquid and gas phase properties to merge and become a supercritical fluid

supercritical fluid

fluid that exhibits has properties/behaves as both a liquid and a gas

triple point

the pressure and temperature where all three phases (solid, liquid, gas) coexist in equilibrium

normal melting point

temperature that solid ↔ liquid equilibrium occurs when pressure = 1 atm

normal boiling point

temperature that liquid ↔ gas equilibrium occurs when pressure = 1 atm

sublimation

solid → gas

deposition

gas → solid

vaporization

liquid → gas

condensation

gas → liquid

melting

solid → liquid

freezing

liquid → solid

solution

• any homogeneous mixture

• made of almost any 2 phases of matter

solute

the dissolved substance

solvent

the substance that the solute dissolved in

how do you know if a solute will dissolve in the solvent?

like dissolves like → substances dissolve if solute and solvent can form intermolecular attractions

energetics of solution formation

1. break solute-solute interactions (endothermic, △H₁ >0)

2. break solvent-solvent interaction (endothermic, △H₂ >0)

3. form solute-solvent interactions (exothermic, △H₃ <0)

sum of all 3 steps = △H_sol = enthalpy of solution

saturated solution

holds max amount of solute at a given temperature

supersaturated solution

holds more solid than is stable

unsaturated solution

holds less than the maximum amount of solute

relationship between temperature and solubility of gases and solids

as temperature increases, solubility of gases decreases and solubility of solids increase

• this is bc KE increases, so more molecules escape into gas phase

Henry’s Law

C_g = kP_g

• c = solubility of dissolved gas in solution (M)

• P = partial pressure of the gas above the liquid (atmosphere)

• k = Henry’s law constant (M/atm)

• increasing gas solubility = increasing gas partial pressure (direct relationship)

• if the temperature is constant, gas solubility is directly proportional to its pressure (so k isn’t needed)

*NOT on equation sheet

percent by mass formula

mass of solute / total mass of solution 100%

*NOT on equation sheet

mole fraction

X_A= moles of A / moles of A + moles of B + …

*NOT on equation sheet

molality

m = moles of solute / kg of solvent

*NOT on equation sheet

strong electrolytes

dissociate into ions when dissolved in water

• the subscripts in chemical formula show mole ratio

• ex: 1.0M AlCl₃ contains 1.0M Al³⁺ and 3.0M Cl^-

dissociation of strong electrolytes in water

strong electrolytes fully dissociate in water

dissociation of weak electrolytes in water

weak electrolytes incompletely dissociate in water

dissociation of non-electrolytes in water

non-electrolytes do not dissociate in water

colligative properties

physical properties that change depending on concentration of solute particles

• includes boiling point elevation, freezing point depression, vapor pressure depression, osmotic pressure

ex: 1M Na2S(aq) dissolves to Na+, Na+, and S2-, which is 3M ions.

• colligative properties don't just depend on molarity, it depends on the number of particles in solution. having more particles = stronger effect on colligative properties

vapor pressure depression equation

Raoult’s Law: P_solution = X_solvent * P°_solvent

• P_solution = mole fraction of the solvent

• P°_solvent = vapor pressure of pure solvent

boiling point elevation

△T_b=K_bm

freezing point depression

△T_f=K_fm

how to find final boiling/freezing point of solution

• T_b=T_b° + △T_b

• T_f=T_f° - △T_f

as concentration increases, what happens to final boiling and freezing points of solution?

increase in concentration = increase BP, decrease FP

osmosis

flow of solvent into solution through a semipermeable membrane

semipermeable membrane

solvents (ex: water) can cross, no solute molecules can pass through

• water molecules move from high → low concentration until concentrations are equal

• high solute concentration = low water concentration

osmotic pressure

result of increased hydrostatic pressure on the solution than on the pure solvent

how to calculate osmotic pressure

pi = iMRT

• if solute is non electrolyte, assume i = 1

• on formula sheet!

how to solve for vapor pressure

1. calculate for mole fraction (X_A)

   1. if electrolyte: X_H2O= moles H2O/ moles H2O + moles of ions

2. plug into Raoult’s Law: P_z = X_ZP_Z°

van’t hoff factor (i)

number of particles in solution when 1 mole of substance dissolves in water

• accounts for increase in concentration when electrolyte dissociates

i = moles of particles in solution / moles of solute dissolved

• if solute is strong electrolyte, I is included as multiplier (ex: iK_bm = △T_b)