Chem 105 Midterm 1

Sublimation

When a solid goes straight to a gas without the liquid stage

Physical Property of Gases

Structureless

No fixed volume

no definite shape/ fills the shape of the container

Kinetic Energy of Gases

Due to the movement of molecules

Ideal Gases

Perfectly follow kinetic molecular theory

No intermolecular attractions and molecules take up no space

Key principles of Ideal Gases

Gas molecule or atom size is negligible

Gases kinetic energy based completely on temperature

Intermolecular forces between gases negligible

Gas collisions are elastic/no energy is lost upon collision

Boltzmann’s Distribution

Gas molecules constantly colliding

Collisions result in a distribution of possible velocities

Plot of velocity of molecules vs number of molecules

Changing temperature shifts the Boltzmann’s diagram

Root Mean Square Velocity

Urms = root(3RT/Mm)

kinetic Energy of Gases

E kinetic = 3/2RT

Mean Free Path

The distance gas molecules travel before colliding with other molecules

Urms/collision frequency

Gas Compressibility

Gases can be compressed due to the distance between gas molecules

Gas Pressure

Gas molecules move randomly

They will collide with each other and the walls of container

Exert an outward force against walls

P = F/A

Boyles Law

As pressure increases, volume decreases (inversely proportional)

Moles of gas and temperature are held constant

P1V1=P2V2

Ex. Mercury adds pressure and decrease volume

Charles Law

As temperature increases, volume increases (directly proportional)

Moles of gas and pressure are held constant

V1/T1=V2/T2

Increased temperature increases collisions with the surface, causing the surface to expand

The Combined Gas Law

Combination of Boyles and Charles gas laws

P1V1/T1=P2V2/T2

Avogadro’s Law

As moles increases, the volume increases

Temperature and pressure are held constant

V1/n1=V2/n2

The relationship is true regardless of the type of gas

The Ideal Gas Law

PV=nRT

STP

Standard temperature and pressure

0 C

1 atm

Total Gas Pressure

Sum of individual gas pressures (partial pressures)

Diffusion

The process where gases intermingle with eachother

Effusion

The process where gases pass through a porous solid or a pin hole

Speeds up the Rates of Diffusion and Effusion

A greater root mean square velocity

R1/R2=ROOT(M1/M2)

Kinetic Molecular Theory of Gases: Real Gases

Gases are small compared to distances between them

No intermolecular forces between gas molecules

Assumptions hold true at standard temperature and pressure

High Pressure Real Gases

Distances between molecules small enough that molecule size is important

Low Pressure Real Gases

Decreases kinetic energy

Intermolecular forces no longer overcome by gas movement

Van Der Waals

Version of the ideal gas law

Accounts for molecule size and intermolecular forces

\[p + an^2/v^2\](V-nb) = nRT

Supercritical Fluids

A phase of matter with no distinction between the liquid and gas phase

Unique Properties of Supercritical Fluids

Allow some molecules, like caffein, to be dissolved without affecting other molecules

Critical Temperature

Minimum temperature necessary to completely overcome intermolecular forces

Critical Pressure

Pressure necessary to condense into the supercritical phase

Reaction Rate

The speed at which a chemical reaction occurs

Chemical Kinetics

The rate at which a reaction progresses can be influenced by several factors

Factors that Influence Chemical Kinetics

Temperature

Pressure

Reactant concentration

The addition of a catalyst

Mechanical force

Average Reaction Rates

\-Reactant/time = Product/time

The rate is proportional to its stoichiometric coefficient

Instantaneous Rate

The rate at a particular time during the reaction

Determined from the slope of curve at a certain time

Is the slope of the tangent

Reaction Order

The relationship between the concentrations of species and the rate of the reaction

Zero Order

Indicates that the concentration of that species does not affect the rate of the reaction

m+n=0

Rate = k

At=-kt + Ao

Positive Integer

Indicates that the concentration of that species directly effects the rate of the reaction

Negative Integer

Indicates that the concentration of that species inversely affects the rate of the reaction

First Order Reaction

Proportional to the concentration of the reactant

m+n=1

Rate = kA

lnAt= -kt + lnAo

Second Order

Proportional to the square of the concentration of the reactant

m+n=2

Rate = KA^2

1/At=kt+1/Ao

K Unit for 0 Order

m/s^-1

K Unit for 1st Order

S^-1

K unit for 2nd Order

M^-1/S^-1

Half Life

The time required for a reaction to go to half its completion

Half Life 0 Order

A/2k

Half Life 1st Order

ln2/k

Half Life 2nd Order

1/kAo

Collision Theory

Must collide

Must have the correct spatial orientation

Must have sufficient energy

Exothermic

Release energy

Reactants Ea higher

Endothermic

Requires energy

Products Ea higher

Intermediate

Any molecule that does not appear in the overall reaction

Rate Limiting Step

The slowest most energy needing step in the reaction

Arrhenius Equation

K=Ae^-Ea/RT

lnK1/K2= Ea/R (1/T2-1/T1)

Catalyst

Lowers the activation energy of a reaction by assisting reactant molecules to be in proper orientation

Homogenous Catalysts

Are the same chemical state as the compounds they will react with (substrates)

Ex. Enzymes

Heterogenous Catalysts

Are in different chemical states then their substrates

Ex. Precious metals in petroleum refining

Features of Catalysts

Don’t appear in the overall balanced equation

Can be reused many times

Are used very early in the mechanism

Don’t change enthalpy

Increase the rate of forward and reverse reactions equally

Chemical Equilibrium

The simultaneous occurrence of a forward and reverse reaction at the same rate

How to Calculate Kp

Kp = Kc(RT)^n

Product Favoured at Equilibrium

K>1

Reactant Favoured at Equilibrium

0

Reactants and Products are Equally favored at Equilibrium

K=1

Reaction Quotient

When concentrations or pressures are not at equilibrium for the k value

Reaction Forms More Products

K>Q

Reaction forms more reactants

K

More Reactant is Added

Shifts right

Reactant Taken Away

Shift left

More Product Added

Shifts left

Product is Taken Away

Shifts right

Reactant Partial Pressure Increases

Shifts right

Reactant Partial Pressure Decreases

Shifts left

Product Partial Pressure Increases

Shifts left

Product Partial Pressure Decreases

Shifts right

Volume Goes Down

Pressure spikes

Shifts right

Volume Goes Up

Pressure drops

Shifts left

Temperature Increases Exothermic

Shift left

Temperature Decreases Exothermic

Shifts right

Temperature Increases Endothermic

Shifts right

Temperature Decreases Endothermic

Shifts left

Increase Concentration of Substance

Away from substance

Decrease Concentration of Substance

Towards substance

Increase Pressure of a System

Towards fewer moles of gas

Decrease Pressure of a System

Towards more moles of gas

Increase Temperature of a System

Away from heat/energy

Decrease Temperature of a System

Towards heat/energy

Add a Catalyst

NO SHIFT

Method of Successive Approximations

If x < 5% the initial concentration/pressure then x can be ignored in addition or subtraction steps

Can also be used when K

Bronsted-Lowry Acid

Donates a proton to water

HCl + H2O → H3O + Cl

Bronsted-Lowry Base

Accepts a proton from water

NH3 + H20 → NH4 + OH

Amphiprotic

Molecule that can either gain or lose a proton

Can act as an acid or a base

Ex. Water, all period 4 elements

Strong Acids

Donate a proton to water and fully dissociate into a proton and an anion

Ex. HCl, HBr, HI, HNO3 ,H2SO4, HClO4, HClO3

Strong Bases

Accept a proton from water and fully dissociate into a cation and hydroxide

Ex. LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH)2, Ba(OH)2, Sr(OH)2

Alkali metal + OH or alkali earth metal + OH

Weak Acids

Do not completely dissociate

Weak electrolytes

E. HF, H3PO4, Carboxylic acids

Weak Bases

Do not completely dissociate

Reverse reactions are relevant

Weak electrolytes

Ex. Amines

How to Calclate PH

\-log \[H3O+\] or 10^-pH

Larger Ka Value/ Lower pKa Value

The stronger the acid

Ka> 1

Larger Kb Value/ Lower pKb Value

The stronger the base

Kb>1

Reactant Favoured

0

Neither Side Favoured

K=1

Product Side Favored

K>1