Exam 1 Flashcards
Molality
moles of solute/mass of solvent in kg
Properties of ionic compounds
- High melting points
- Hard and brittle
- Do not conduct electricity in solid form, do conduct it in molten form
- Solutions are soluble in polar solvents and conduct electricity
Finding 🔼Hsolution from 🔼Hlattice
🔼Hlattice>0: 🔼Hsol=🔼Hhydration+🔼Hlattice
🔼Hlattice<0: 🔼Hsol=🔼Hhydration-🔼Hlattice
Finding 🔼Hlattice from 🔼Hf
🔼Hlattice=🔼Hf(products-reactants)
🔼Hlattice trend
Bigger charge different and higher up on periodic table=higher 🔼Hlattice
Lattice enthalpy
Energy change that occurs when 2 gaseous atoms come together to form a solid compound
Lattice enthalpy trend
Bigger difference in charges=bigger lattice enthalpy
if theyre the same then
Smaller radius=large lattice enthalpyp
Under which conditions are gases more soluble in water?
High pressure, low temperature
Ion-dipole interaction
Interaction of ions with polar solvents
dipole dipole interactions
attractive force between polar molecules
dipole-induced dipole
attraction caused by proximity of polar molecule
Equation associated with 🔼Hhydration(+1 and +2 charges)
M+(g) + X-(g)▶️M+(aq) + X-(aq)
M2+(g) + 2X-(g)➡️M2+(aq) + 2X-(aq)
🔼Hhydration trends
Smaller ions and ions with higher charges have greater 🔼Hhydration
Factors affecting vapor pressure and their trends
- Temperature: As temp increases, KE and therefore vapor pressure increases
- Intermolecular forces: Stronger forces=higher KE needed to enter gas phase, decreasing vapor pressure
- Presence of nonvolatile solute: Affects rate of evaporation, decreases vapor pressure of solution compared to pure solvent
Raoult’s law
Psolvent=xsolvent,l•p*solvent
Psolvent=partial pressure exerted by solvent vapor above a solution
xsolvent,l= mole fraction of solvent in a liquid phase
P*solvent=vapor pressure of pure solvent
What are ideal solutions formed by?
Substances with similar intermolecular forces and similar structures. They are formed as a result of increased entropy
Ideal solution
Solution that obeys Raoultms law, 🔼Hsol=0
Volatile
measure of how easily something evaporates. more volatile components evaporate more easily
How to read graphs of boiling points of solute-solvents?
- Contant temp: Area above both lines is liquid, area in between is vapor and liquid, and area below is a vapor
- Constant pressure: Area below both lines is liquid, area in between is vapor and liquid, and area above both lines is a vapor
When do vapor pressure deviate negatively and positively from ideal behavior?
- Negative deviation: solute-solvent interactions and stronger than solvent-solvent interactions and prevent solvent from escaping solution
- Positive deviation: solute-solvent interactions are weaker than solvent-solvent interactions. Interactions of solvent molecules are disrupted and makes it easier to escape into vapor state
Colligative properties
properties that depend on the total concentration of solute particles in a solution and not their identity or other factors
Vapor pressure lowering
the difference between vapor pressure with and without a solute, proportional to the mole fraction of a solute
How do nonvolatile solutes effect vapor pressure?
They decrease it
🔼Hlattice enthalpy equation
M+(g) + X-(g)▶️MX(s) when 🔼Hlattice is negative and bonds are forming
MX(s)▶️M+(g)+X-(g) when
🔼Hlattice is positive and bonds are breaking
Boiling point elevation
A colligative property in which the boiling point of the solution is higher than the boiling point of pure solvent, proportional to molality
🔼Tb=Tb-Tb*
Tb=boiling point of solution
Tb*=normal boiling point of pure solvent
🔼Tb is always positive so Tb>Tb*
Kb
proportionality constant for molality and boiling point elevation, unique to each solvent, does not depend on solute
🔼Tb=Kb•m
Freezing point depression
A colligative property in which the freezing point of a solution is lower than the freezing point of the pure solvent, proportional to molality
🔼Tf=Tf-Tf
Tf=normal freezing point of pure solvent
Tf=freezing point of solution
🔼Tf is a positive quantity so Tf
Kf
proportionality constant for freezing point depression and molality, unique for each solvent, does not depend on solute
🔼Tf=Kf•m
Most common nonelectrolytes
sugar and alcohol
Van’t hoff factor
Number of ions in a formula unit(i), i corrects for the number of ions
🔼Tb=i•Kb•m
or
🔼Tf=i•Kf•m
Why are theoretical i values typically higher than experimentally determined i values?
Once compounds break apart in solution, some of the ions reattract to form other compounds, and i assumes all ions are available, while some are actually attracting to each other
Osmotic pressure formula
|-|=iMRT
R=gas constant
T=temp in kelvin
Dispersion force trend
Larger atoms/molecules have more dispersion forces
Factors affecting reaction rates
- Physical state of reactants: liquids and gases are faster than solids
- Concentration of reactants: Interactions increase as concentration increases, increasing likelihood of a reaction
- Temperature
- Presence of a catalyst
Average rates of consumption
Denoted by 🔼{A}\🔼t
Reactants: -🔼{A}\🔼t=-{Afinal}-{Ainitial}\tfinal-tinitial
Products: same formula, removed negative signs before equation
Rate of consumption and formation formulas given aA▶️cC
Consumption: -🔼{A}/🔼t=a/c • 🔼{C}/🔼t=rate
Formationa: 🔼{C}/🔼t=c/a • (-🔼{A}/🔼t)=rate
Formulas for how much something reacted in a given time given aA▶️cC
{A}0={A}reacted+{A}t for reactants
{C}formed=(c/a)•{A}reacted
Instantaneous reaction rate
The rate when 🔼t approaches 0
Instantaneous rate=-lim(🔼t▶️0)🔼{A}/🔼t=-d{A}/dt
Integrated rate law for a first order reaction
ln(At)=-kt+ln(Ao)
Half life equation for first order reactions
t(1/2)= ln(2)/k
Integrated rate law for second order reactions
1/(At)=kt+1/(Ao)
half life equation for second order reactions
t(1/2)=1/(k•Ao)
integrated rate law for 0 order reactions
At=(Ao)-kt
🔼Hsolution equation
(s)▶️(g)+(g)
