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Gr. 12 Energy Changes and Rates of Reaction Cheat Sheet by


Thermo­che­mistry: the study of the energy changes that accompany physical or chemical changes in matter
Types of energy:
Ep (potential energy)
the energy of an object due to its positi­on/­com­pos­ition
Ek (kinetic energy)
the energy of an object due to its motion
Thermal energy (Eth): the total quantity of Ek and Ep in a substance; depends on how fast the particles are moving: more energy = more speed = more Eth
Heat: the transfer of Eth from a warm object to a cool object
Temper­ature: measure of the average Ek of the particles in a substance
Law of Conser­vation of Energy: energy cannot be created or destroyed, only converted from one form to another
Note: Temper­ature ≠ Eth! A cup of water at 90°C has a higher temper­ature than a bathtub of water at 40°C, but the water has more Eth since it has more molecules

System­/Su­rro­undings and Reactions

System: the group of reactants and products being studied
Surrou­ndings: all the matter that is not a part of the system
Types of systems:
Open system
both energy and matter are allowed to enter and leave freely
Closed system
energy can enter and leave the system, but matter cannot
Isolated system
neither matter are allowed to leave the system (complete isolation is impossible)
Types of reactions:
energy from the surrou­ndings is absorbed by the system
energy from the system is released into the surrou­ndings

Specific Heat Capacity and Calori­metry

Specific heat capacity: the amount of energy required to raise the temper­ature of 1 g of a substance by 1°C (measured in J/g°C); depends on type and form of substance
Calori­metry: the experi­mental process of measuring the ΔEth in a chemical or physical change
Calori­meter: device used to measure ΔEth
Types of calori­meters:
Polyst­yrene (styro­foam)
Reasonably accurate and inexpe­nsive
More precise, used for reactions that involve gases
Used for combustion reactions

Calori­metry Calcul­ations

4 assump­tions when performing calori­metry calcul­ations:
1. Any thermal energy transf­erred from the calori­meter to the outside enviro­nment is negligible
3. All dilute, aqueous solutions have the same density as water (D = 1.00 g/mL)
2. Any thermal energy absorbed by the calori­meter itself is negligible
4. All dilute, aqueous solutions have the same specific heat capacity as water (c = 4.18 J/g°C)
Calori­metry formula:
Q = mcΔT
m = mass of the substance (g)
c = specific heat capacity of the substance ( J/g°C)
ΔT = temper­ature change experi­enced by the system; ΔT = Tfinal - Tinitial (°C)
Q = total amount of Eth absorb­ed/­rel­eased by a chemical system ( J )
Value of Q has two parts:
The number: how much energy is involved
The sign: the direction of the energy transfer (important to show, even if it is positive!)

Because of the law of conser­vation of energy, the total thermal energy of the system and the surrou­ndings remain constant:

Qsystem + Qsurrou­ndings = 0
Qsystem = - Qsurrou­ndings

Enthalpy Change (ΔH)

Enthalpy (H): the total amount of Eth in a system; not directly measurable
Must measure enthalpy change (ΔH) by measuring the ΔT in the surrou­ndings
Enthalpy change (ΔH): the energy released to/abs­orbed from the surrou­ndings during a chemic­al/­phy­sical change; can be measured using calori­metry data
As long as pressure is constant, the enthalpy change of a chemical system is equal to the flow of thermal energy in or out of the system
Enthalpy change formula:
ΔH = |Qsystem|
ΔH = ±|Qsurrou­ndings|
If ΔH > 0, the reaction is endoth­ermic
If ΔH < 0, the reaction is exothermic
If there is more than one substance making up the surrou­ndings (i.e. bomb/flame calori­meters), then
Qsurrou­ndings = Σ Qsubstances

Molar Enthalpy Change (ΔHx)

Molar enthalpy change (ΔHx): the enthalpy change associated with a physic­al/­che­mical change involving 1 mol of a substance (J/mol)
x = type of change (vapor­iza­tion, neutra­liz­ation, combus­tion, etc.)
Molar enthalpy change formula:
ΔH = nΔHx

Repres­enting Enthalpy Change

4 ways to represent ΔH:
1. Thermo­che­mical equations with energy terms
CH4 + 2 O2 CO2 + 2 H2O + 890.8 kJ
2. Thermo­che­mical equations with ΔH terms
CH4 + 2 O2 CO2 + 2 H2O ΔH = -890.8 kJ
3. Molar enthalpies (ΔHx)
ΔHcomb = -890.8 kJ/mol
4. Potential energy (Ep) diagrams

Hess' Law

Enthalpy change (ΔH) is determined by initial and final conditions of a system; it is indepe­ndent of the pathway
The total ΔH of a multi-step reaction is the sum of the ΔH of its individual steps
Hess's Law formula:
ΔHreaction = Σ ΔHsteps
This formula can be used in cases where the overall reaction is not feasible to be done in a calori­meter (i.e. reaction is too slow/too fast/too violent)
1. If a reaction is flipped, flip the ΔH value's sign
2. If a reaction is multiplied, multiply the ΔH value

Standard Enthalpy of Formation (ΔH°f)

The standa­rdized ΔH when 1 mol of a substance is formed (synthe­sized) directly from its elements to its standard state at SATP
The elements themselves have a ΔH°f of 0 (elements cannot be synthe­sized)

Bond Energies (D) and Bond Enthalpy

Bond Energies
Stability of a molecule is related to the strength of its covalent bonds
The strength is determined by the energy required to break that bond
Bond Enthalpy:
ΔH for breaking a particular bond in 1 mol of a gaseous substance
Always positive because energy is always required to break bonds
Used for predicting reaction types before the reaction is performed (not entirely accurate)
Formula for predicting reaction type using D and bond H:
ΔH = Σ (nDbonds broken) - Σ (nDbonds formed)

Reaction Rates

The speed at which a reaction occurs
Can be fast (10-15s) or slow (years)
Measured by the change in the amount of reactants consumed or products formed at a given time interv­al(s)
Can be measured by volume, mass, colour, pH, and electrical conduc­tivity
Often expressed as a positive value for conven­ience, regardless of what is being measured
Average rate of reaction: rate of a chemical reaction between two points in time (one time interval); calculated from the slope of the secant of the time interval on a concen­tra­tio­n-time graph
Average rate of reaction formulas:
How fast a reactant disappears
- Δ[A]/Δt
How fast a product appears
Δ[A], Δ[B], Δt = [A]2 - [A]1, [B]2 - [B]1, t2 - t1
Instan­taneous rate of reaction: rate of a chemical reaction at a single point int time; calculated from the slope of the tangent of the time position on a concen­tra­tio­n-time graph

Collision Theory

States that chemical reactions can only occur if the reactants have the right kinetic energy (speed) and orient­ation to break reactant bonds and form product bonds
Effective collision: a collision that has sufficient energy and correct orient­ation of colliding particles to start a reaction
Ineffe­ctive collision: a collision where the particles rebound, unchanged in nature
Activation energy (Ea): the minimum energy required for reactants to have for a collision to be effective
Activated comple­x/t­ran­sition state: unstable arrang­ement of atoms containing partially formed and partially broken bonds; maximum Ep point in the reaction
Rate of a reaction depends on the frequency of collisions and the fraction of those collisions that are effective.
Rate = frequency of collisions x fraction of collisions that are effective

Increasing Reaction Rates

5 factors that can increase a reaction rate: chemical nature of reactants, concen­tration, surface area, temper­ature, and catalysts
Chemical nature of reactants
For any reactant, the activation energy required depends on the bond type (single vs double vs triple), the bond strength (D value), the number of bonds, and the size and shape of the molecu­le(s)
Concen­tration of reactants
Concen­tration = amount of substance per unit volume (mol/L); applies only to solutions
[reactant] = collisions = rate
Rate α [reactant] - as the concen­tration increases, the rate increases, and vice versa
Surface area
Surface area = total area of all the surfaces of a solid figure
SA = collisions = rate
Rate α SA - as the surface area increa­ses, the rate increa­ses, and vice versa
Temper­ature of system
T = collisions + fraction of effective collisions = rate
Rate α T - as the temper­ature increa­ses, the rate increa­ses, and vice versa
Catalyst = substance that increases the rate of a reaction without itself being consumed in the reaction; provide an alternate pathway for the reaction with a lower Ea
Ea = fraction of effective collisions = rate
Rate α 1/Ea - as the catalyzed activation energy decreases, the rate increa­ses, and vice versa

Rate Law

Mathem­atical relati­onship between the reaction rate and the concen­tration of reactants; needs experi­mental data
Formula: Rate = k[A]a[B]b[C]c
[A]/[B­]/[C] = concen­tration of reactants (only reactants are relevant); k = rate constant
Orders of Reaction
Order of reaction: the exponent used to describe the relati­onship between the [ i ] of a reactant and the rate of reaction; tells us how quickly the rate will increase when [conc] increases
Zero order
Rate = k[A]0; slope is flat; rate is not affected by [A]
First order
Rate = k[A]1; slope is an increasing straight line; rate α [A]
Second order
Rate = k[A]2; slope is an increasing curve; rate α [A]2
Total order of reaction = the sum of the exponents in the rate law equation
The only accurate data for concen­tration and rate is the initial rate, because as soon as the reaction starts, products are formed and the reverse reaction starts, making any rate measured after t = 0 affected by the products.

Reaction Mechanisms

Chemical reactions usually occur as a sequence of elementary steps that, when added, result in the overall reaction
Mechanism is dependent on the slowest elementary step - the rate-d­ete­rmining step
Elementary step = a single molecular event in the reaction mechanism
3 criteria for a proposed reaction mechanism:
The elementary steps must add up to the overall reaction
The elementary steps must be physically reasonable - there should not be more than 2 reactants
The rate-d­ete­rmining step must be consistent with the rate law equation


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