Kinetics and Equilibrium - Honors Chemistry

kinetics

study of the rate of chemical reactions

reaction rate

speed at which reactants become products

chemical reactions are balanced in

mass charge and energy

chemical reactions have what type of energy

potential energy stored in chemical bonds

chemical reactions start with

certain amount of energy and products end with different amount of energy (endo vs. exo)

heat of reaction (∆H)

net energy absorbed or released in a chemical reaction; change in enthalpy of a system (∆H = H products - H reactants)

nature and chemistry favors

low energy - stability

entropy

measure of randomness or disorder of a system

nature and chemistry favors

high entropy - chaos (phase change S→ L → aq → G)

spontaneous reactions

favor lower enthalpy and higher entropy; if both factors present in reaction then it’s spontaneous

Gibbs Free Energy formula

∆G = ∆H - T∆S

∆G is

kJ/mol - Gibbs free energy

∆H is

kJ/mol - enthalpy

T is

k - temperature

∆S is

kJ/mol * k - entropy

natural processes proceed in

direction that lowers free energy

describe how (-) exothermic and (+) disorder would look like in Gibbs Free Energy equation

∆G = ∆H - T∆S; if ∆H is negative and subtracting positive T (kelvin can’t have negative numbers) multiplied by positive ∆S then it’s ∆H(-) - T∆S (+) so ∆G is always negative

collision theory

in order for reactions to occur molecules and atoms must collide with each other

for reactions to occur they need

enough energy and proper orientation

do you want factors to increase or decrease to make reactions occur faster

factors that increase energy and orientation

effective collisions

reactants become products

ineffective collisions

reactants remain unchanged

sufficient kinetic energy

brings reactants together but potential energy brings them apart

proper orientation

orientation which mot likely breaks the bond is most effective or if given what product is supposed to look like then orientation that will give product

how do you increase reaction rate

increasing number of collisions

factors that affect reaction rate

temperature, concentration, surface area, pressure, nature of reactants, catalyst

how does temp affect reaction rate

higher temp = more collisions

how does concentration affect reaction rate

higher concentration = more collisions

how does surface area affect reaction rate

greater surface area = more collisions

how does nature of reactants affect reaction rate

ionic substances generally react faster than molecules and dissolving speeds up reactions

how does a catalyst affect reaction rate

substances that speeds up reaction by giving new path with lower activation energy - facilitate reaction but not part of it so always in reactant and product without changing

reaction mechanism

steps or pathway by which reacting particles rearrange themselves to form productsne

net reaction

chemical equation that shows the reactants and products onlyac

activated complex

temporary arrangement of particles of chemical reactions

activation energy

minimal amount of energy needed to form an activated complex and have chemical reaction occur

kinetic energy diagram

higher temp increases number of particles that can have effective collisions by increasing their kinetic energy

thermodynamics

study of energy in chemical reactions

enthalpy

heat of content in a system

endothermic reaction has heat of reaction that’s

positive; ex. dissolving potassium nitrate

heat is on which side for it to be endothermic

on reactant side being absorbed

exothermic reaction has heat of reaction that’s

negative; ex. combustion of propane

Table I

net energy of equation is heat of reaction at 101.3 kJ and 298K; endo ∆H = (+) and exo ∆H = (-)

in Table I energy is proportional to

moles

if reaction reverses in Table I then

sign reverses; ex. opposite of endo (+) is exo (-)

why is most of Table I exothermic

nature wants lower energy so there’s more reactions where heat is released

How many kJ of energy are absorbed for 1 mole of nitrogen dioxide?

33.2 kJ