unit 8: thermodynamics

10th chemistry

energy

the ability to do work or produce heat (kinetic and potential)

kinetic

- energy of motion
- directly related to the constant random motion of the particles
- temperature \= average kinetic energy (T\= kelvin)

potential

- energy of position
- depends on composition, atoms, bonds, and molecular geometry
- only changes can be measured

chemical potential energy

energy stored in a substance because of its composition

law of conservation of energy

energy can be converted from one form to another, but cannot be created or destroyed (1st law of thermodynamics)

heat

- the total amount of thermal energy that flows (high to low) because of a difference in temperature
- dependent on mass
- symbol \= q
- units \= joules, calorie

thermal equilibrium

same temperature

calculate heat

q\=mcΔT (Joules)
- use when there's a change in temperature

specific heat capacity (c)

- the amount of energy/heat required to raise the temperature of 1 gram by 1°C or K
- always positive
- units \= J/g x °C or J/g x K

low specific heat capacity

- heats up (absorbs energy) and cools down (releases energy) quickly
- ex: metals

high specific heat capacity

- heats up (absorbs energy) and cools down (releases energy) slowly
- absorbs large amounts of energy with relatively little change in temperature
- ex: water (good coolant)

endothermic

- system absorbs heat from surroundings
- +q
- +ΔH

exothermic

* system releases/evolves heat to surroundings
* -q
* -ΔH

change in temperature

ΔT \= Tfinal - Tinitial

calorimetry

the study and measuring of the heat of chemical reactions and physical processes (thermochemistry)

calorimeter

- an insulated device used for measuring the amount of heat absorbed or released during a chemical reaction or physical process
- a known mass of water is placed in an insulated chamber to absorb energy released or provide energy absorbed by the system

system

part of the universe that contains the reaction/metal or process being studied

surroundings

everything in the universe besides the system

universe

system + surroundings

enthalpy

- H
- the heat content of a system at constant pressure
- ΔH: heat of reaction (KJ)
- ΔHrxn \= qrxn/molrxn
- ΔH \= Eproducts - Ereactants

exothermic reactions

products are lower in enthalpy than reactants and are more stable

endothermic reactions

products are higher in enthalpy than reactants and are less stable

bonds

strong bonds take more energy to break

activated complex/transition state

highest point in enthalpy

activation energy

- energy required to reach the complex
- catalysts can lower but won't change ΔH
- high activation \= slow reaction
- low activation - fast reaction
- Ea

thermochemical equation

- includes the physical states of all reactants and products
- energy change expressed as the change in enthalpy (ΔH)

fractional coefficients

represents mole quantities

ΔHfusion

energy/heat required to change 1 mole of solid to liquid at melting point

ΔHvaporization

energy/heat required to change 1 mole of liquid to gas at boiling point

ΔHsolidification

heat released when 1 mole of liquid freezes to a solid at freezing point

ΔHcondensation

heat released when 1 mole of gas condenses to liquid at condensation point

q \= mcΔT

heat is being used to raise the temperature of a solid

q \= ΔHfus n or m

- heat is being used to turn the solid to a liquid
- ΔHfus \= 6.01 KJ/mol
- ΔHfus \= 333 J/g

q \= mcΔT

heat is being used to raise the temperature of the liquid

q \= ΔHvap n or m

- heat is being used to turn liquid to gas
- ΔHvap \= 40.1 KJ/mol

melting

ΔHfus \= 6.01 KJ/mol

freezing

ΔHfus \= -6.01 KJ/mol

vaporization

ΔHvap \= 40.1 KJ/mol

condensation

ΔHvap \= -40.1 KJ/mol

enthalpy of solution

- heat released or absorbed at a solute associates with the solvent during the formation of a solution
- ΔHsoln

enthalpy of combustion

- the enthalpy change for the complete burning of one mole of the substance
- ΔHcomb
- always exothermic \= always -ΔH
1. find the mole ratio of the equation and set it equal to the ΔH (take away the - , just know it releases heat)
2. stoichiometry

hess's law

- flip and multiply given equations to build the equation you want the ΔH for
- flip \= change sign of ΔH
- multiply \= multiply the same \# to the ΔH
- opposite sides \= cross out
- same sides \= add/subtract
- do water equations last

formation

- reactants: elements in their standard state
- products: 1 mol of 1 compound
- ΔHrxn \= ΣΔHproducts - ΣΔHreactants
- diatomic and monatomic elements have a value of 0

standard state

normal physical state at 1 atm at 298K

spontanenous process (TFP)

any physical or chemical process that once begun, occurs with no outside intervention (ex: ball rolling down a ramp, formation of rust)

entropy

* ΔS
* randomness/disorder
* increases as particles become more random: (s) -> (l) - > (g)
* +ΔS: entropy increases (more disordered)
* -ΔS: entropy decreases (more organized)
* side with the most mol of gas has the highest entropy

\+ ΔS

entropy increases (more disordered)

\- ΔS

entropy decreases (more organized)

\- ΔS

- dissolving of a gas in a solvent
- ex: CO2(g) -\> CO2 (aq)

\+ ΔS

- solid or liquid dissolving in a solvent
- NaCl(s) -\> Na+(aq) + Cl-(aq)

\+ ΔS

temperature increasing results in what entropy

gibbs free energy

- energy available to do work and helps determine the TFP of a reaction
- a reaction will proceed spontaneously in the way that lowers its free energy
- ΔG \= ΔH - TΔS
- ΔG \= KJ
- ΔH \= KJ
- T \= Kelvin
- ΔS \= J

\- ΔG

TFP, will happen

\+ ΔG

not TFP, will not happen

\- ΔG

\- ΔH, + ΔS

\+ ΔG

+ΔH, - ΔS