Chem Exam 3

energy

the capacity to do work or transfer heat
units: J

mechanical energy

energy in an object that is attributable to its motion, position, or both

kinetic energy

energy due to motion

thermal energy

energy due do the motion of atoms in a sample. It is proportional to temperature, energy due to the random movement of particles

potential energy

energy due to position

chemical energy

energy due to relative positions of nuclei (protons) and electrons, energy stored in bonds. a type of POTENTIAL energy

Internal energy (E or U)

sum of all kinetic and potential energies stored within a substance.

examples: bond energy, IMFs, ionization energy, energy due to motion (translational, vibrational, rotational)

* when considering any process, we look at a change in internal energy (∆E or ∆U)

units: J

law of conservation of energy

energy can neither be created or destroyed

thermodynamics

study of energy and its transformations

thermochemistry

study of energy related to chemical reactions and processes

First law of thermodynamics

energy cannot be created or destroyed, energy must transfer, the energy of the universe is constant

system

the object of interest, the source of energy, such as a container with chemical reactants and products

surroundings

literally anything else (other than the system)

heat and work

change in internal energy is the sum of heat and work done on/by the system, heat transferred to the system is (+), from system is (-)

heat (q)

the transfer of thermal energy due to a change in temperature, the flow of energy that causes a temperature change in an object or its surroundings

units: J

work (w)

a force acting over a distance, work is the energy resulting from this, work done on the system is (+), by the system is (-)

units:J

state function

a property whose changes depend only on the initial and final states, does not depend on the path taken, describes the current state of a system and is independent of the path taken to achieve its value
ex. distance, temp., pressure, elevation, internal energy, enthalpy
we represent state functions with a delta (∆), energy and enthalpy are state functions

path function

a property whose changes depend on the path taken
ex. heat, work , a function that is dependent on the sequence of steps that move the system from its initial state to its final state
(don't use ∆ here)

pressure-volume work

the work done on or by a system when there is a volume change against an external pressure

Enthalpy (H)

the sum of internal energy of a system and the product of its pressure and volume change, changes in enthalpy for a reaction are often called the "heat of reaction".

Units: Kj/mol

endothermic reaction

(+) ∆H, absorbs heat from surroundings; feels cool to touch

exothermic reaction

(-) ∆H, releases heat into surroundings; feels warm to touch

enthalpy of fusion (melting)

the amount of energy needed to melt one mole of a substance

enthalpy of vaporization

the amount of energy needed to evaporate one mole of a substance

bond enthalpies

the enthalpy change ∆H associated with breaking a specific bond in 1 mole of gaseous molecules

Hess' Law

leverages H being a state function to allow the calculation of ∆H for a process using other ∆H alues

Using ∆Hf values to calculate ∆Hrxn

these are technically a type of Hess' law problem, but made easier because all reactions are formation reactions

standard state

the state in which the substance is most stable at, 25 degrees celsius (298K) and 1 atm

formation reaction

forms one mole of a compound from its elemental components in their natural state

∆Hf

∆Hrxn for a formation reaction at standard state

standard enthalpy of formation

the enthalpy change in the formation of one mole of a compound from its elements in their standard state (∆Hf)

when using ∆Hf we can simply take "products-reactants"

∆Hrxn=[sum of ∆Hf(products)]-[sum of ∆Hf(reactants)]

entropy

S
units: J/K*mol

spontaneous process

proceeds on its own, does not mean that its fast or slow, can be exo or endothermic, all spontaneous processes result in an increase in entropy (S)

second law of thermodynamics

not all of the energy released by a spontaneous process is able to do work. this leads us to define a new function called entropy (S) which is often described as a measure of disorder/randomness

Things that affect entropy (S)

- increase in temp. = increase in entropy
- entropy (s) < entropy (l) << entropy (g)
- increase in number of particles (#moles) = increase in entropy

Third law of thermodynamics

a perfect crystal at 0K (absolute zero) has zero entropy, pure, perfectly ordered, crystalline substance

an application of the 2nd law

- for any isolated process to be spontaneous, the entropy must increase (S>0)
- in a bigger scale, we translate this to say "for a process to be spontaneous, the entropy of the universe must increase"

standard molar entropies

the entropy of 1 mol of substance under standard conditions

gibbs free energy (G)

gibbs proposed a new state function, free energy, that could be used to determine whether or not a reaction is spontaneous, using only properties of the system

units: Kj/mol

Gibbs - spontaneous

∆G < 0, ∆S >0
or
∆G (-), ∆S (+)

gibbs - nonspontaneous

∆G > 0, ∆S< 0
or
∆G (+), ∆S (-)

gibbs at equilibrium

∆G ∆S both = 0

useful observations - spontaneity

∆H - , ∆S +, ∆G - Spontaneous = YES
∆H + , ∆S -, ∆G + Spontaneous = NO
∆H + , ∆S +, ∆G +/- Spontaneous = MAYBE
∆H - , ∆S -, ∆G +/- Spontaneous = MAYBE

rate

how fast something changes per unit in time
ex) velocity - mi/hr, m/s, km/hr

increase in temp. = increase in rate of all rxns
increase [ ] = increased rate

kinetics

study of rates of change in physical and chemical processes

chemical kinetics

study of reaction rates

reaction rate

the change in concentration per unit time
units: M/s, M/min, M/hr, M/yr

reaction rates can be affected by

5 factors: the nature of the reactants, particle size of a solid, concentration of a reactant, temp., and presence of a catalyst

reaction rates can be measured by

the disappearance of a reactant or the appearance of a product

initial rate

we usually define our reaction rates as close to t=0 as possible. This is called the initial rate, instantaneous rates that are measured close to the start of a reaction

rate constant

k
Units: depend on rate order

rate law

a mathematical expression for the relationship between the reaction rate and the concentrations of all reactants
- includes a proportionality constant called the rate constant (k)
- the concentrations of each reactant are raised to the power of their rate order
* the units of the rate constant depend on the overall reaction order
*this must be determined experimentally

Units of k depending on the order

0: m/s
1: 1/s
2: 1/mxs

integrated rate law expressions

used to determine the concentration of reactant at anytime needed to reach a certain concentration of reactant

rate law shows the relationship between

rate and the concentrations of each reactant and must be determined experimentally

integrated rate law shows the relationship between

concentration and time

Half life (t 1/2)

the time it takes for a reactant to end up half its original concentration

activation energy (Ea)

an energy barrier that exists for many reactions that must be overcome in order for the reactants to make products, minimum amount of energy needed for the reactant to reach the transition state; also called barrier energy

*there is an inverse relationship between the values of Ea and k
when Ea increases, reaction rate decreases

units: J/mol

transition state

a high-energy state in the rearrangement of bonds that occurs during chemical reactions; also called activated complex

Arrhenius equation

describes the relationship between the rate constant and the activation energy of a chemical reaction

reaction mechanisms

a series of elementary steps that make up a chemical reaction
- a series of reactions that make up an overall reaction
- these reactions are called elementary steps and they represent the actual collisions happening throughout the reaction
- the elementary steps must add up to form the overall reaction (like last step of Hess' law problem)
- the rate law for the mechanism must equal the overall reaction

reaction mechanisms must satisfy the following criteria

- elementary steps must sum up to the overall equation
- the rate law for the rate-determining step must be consistent with the observed rate law, which cannot contain intermediates

elementary steps

an individual molecular event with a transition state and rate law, typically part of a series that makes up a reaction mechanism

rate-determining step

the slowest elementary step in a reaction mechanism
* determines the overall reaction rate

intermediate

a high-energy, unstable species formed by one elementary step and consumed by the next elementary step in a reaction mechanism
- forms in middle of mechanism and is consumed later
- not present in overall reaction

higher the activation energy (Ea) the

slower the reaction

catalysts

a substance added to a reaction that speeds it up by providing a second pathway/mechanism that has a lower activation energy
- we observe that catalysts are not present in overall reaction. They are present at beginning of each mech., are then consumed, and then produced again

homogeneous catalysts

a catalyst in the same phase as the reactants

heterogeneous catalyst

a catalyst in a different phase than the reactants

Chloroflurocarbons (CFCs)

a type of substance that can act as a catalyst

chemical equilibrium

when the rate of the forward reaction equals the rate of the reverse reaction, we represent this by using double arrows

equilibrium

the situation in which the forward and reverse reactions are occurring at the same rate. It's a dynamic state

dynamic state

a state in which two opposite processes occur at equal rates

extent of a reaction

the balance between the amount of reactants and products at equilibrium

K (equilibrium constant)

a constant that tells how far a reaction will proceed until it reaches equilibrium

observations for chemical equilibrium

- k >> 1, prod >> react @ equilibrium
- k << 1, prod << react @ equilibrium
- k ≈ 1, prod ≈ react @ equilibrium

if using gases, what is used to represent concentration?

partial pressures

heterogenous equilibria

an equilibrium in which the species are in different phases

Kc

do NOT include solids or liquids in equilibrium expression , include aqueous solutions and gases

Kp

Only include gases in equilibrium expression

Reaction quotient (Q)

the ratio of the products to the reactants, where each species is raised to its stochiometric coefficient , at any time during a reverse reaction, used to predict how the concentrations or partial pressures of products and reactants must change to reach equilibrium
- this is calculated the exact same way as k, but it may not be at equilibrium
- a reaction will always adjust in a manner to make Q equal to J
- by comparing Q to k, we can deduce how the reaction will adjust in order to reach equilibrium

Q<K

rxn must produce more products than reactants to reach equilibrium
"shifts right/towards products"
- it will form more products

Q>K

rxn must produce more reactants to reach equilibrium
"shifts left toward reactants"
- it will form more reactants

Q=K

reaction is already at equilibrium
"no shift"

Le Chatelier's Principle

describes what a reaction will do to re-establish equilibrium if its equilibrium is disturbed , when a stress is applied to a system at equilibrium, the reaction shifts in a direction to relieve that stress
- it does this by changing concentration, P and V, or temperature

Le Chatelier's Principle - changing concentration

- adding reactant/product will result in a shift away to consume the excess added
- removing reactant/product will result in a shift toward to re-create what was lost

Le Chatelier's Principle - Changing P and V (only affects gases)

- compressing a gas mixture (decrease V and increase P) causes the rxn to shift in the direction of fewer moles of gas to relieve the excess pressure

- expanding gas mixture (increase V, decrease P) causes the rxn to shift in the direction of more moles of gas to recoup the "lost" pressure

*if there are equal moles of gas on both sides of reaction, expanding and compressing has no effect on equilibrium

Le Chatelier's Principle - Changing Temp.

endothermic (+∆H)
- increase T shifts right
- decrease T shifts left

exothermic (-∆H)
- increase T shifts left
- decrease T shifts right

*k is only altered by temp change and not any other of the conditions

frequency factor

A

Universal gas constant (R)

8.3145
units: J/mol*K