Exam 1 Flashcards by Tres Schuler

Molality

moles of solute/mass of solvent in kg

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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

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Finding 🔼Hsolution from 🔼Hlattice

🔼Hlattice>0: 🔼Hsol=🔼Hhydration+🔼Hlattice

🔼Hlattice<0: 🔼Hsol=🔼Hhydration-🔼Hlattice

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Finding 🔼Hlattice from 🔼Hf

🔼Hlattice=🔼Hf(products-reactants)

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🔼Hlattice trend

Bigger charge different and higher up on periodic table=higher 🔼Hlattice

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Lattice enthalpy

Energy change that occurs when 2 gaseous atoms come together to form a solid compound

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Lattice enthalpy trend

Bigger difference in charges=bigger lattice enthalpy

if theyre the same then

Smaller radius=large lattice enthalpyp

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Under which conditions are gases more soluble in water?

High pressure, low temperature

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Ion-dipole interaction

Interaction of ions with polar solvents

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dipole dipole interactions

attractive force between polar molecules

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dipole-induced dipole

attraction caused by proximity of polar molecule

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Equation associated with 🔼Hhydration(+1 and +2 charges)

M+(g) + X-(g)▶️M+(aq) + X-(aq)

M2+(g) + 2X-(g)➡️M2+(aq) + 2X-(aq)

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🔼Hhydration trends

Smaller ions and ions with higher charges have greater 🔼Hhydration

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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

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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

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What are ideal solutions formed by?

Substances with similar intermolecular forces and similar structures. They are formed as a result of increased entropy

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Ideal solution

Solution that obeys Raoultms law, 🔼Hsol=0

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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*

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

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

1. Physical state of reactants: liquids and gases are faster than solids 2. Concentration of reactants: Interactions increase as concentration increases, increasing likelihood of a reaction 3. Temperature 4. 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)