Thermochemistry

1st law of thermodynamics

Energy cannot be created nor destroyed, only converted from one form to another

1st law of thermodynamics equation

E_universe = constant

Potential energy

Energy due to position or composition that can be converted to work

Kinetic energy

Energy due to the motion of an object

Kinetic energy equation

KE = (1/2)mv²

Temperature

Reflects random motion of particles and is related to the kinetic energy of the system, measure of motion

Heat

Actual transfer of energy between 2 objects due to a temperature difference, flow of energy

System

Main subject of focus

Surroundings

Everything else in the universe

Universe

System and surroundings

When heat exchange occurs

1. Heating and cooling of a substance

2. Phase changes

3. Dissolving solutes

4. Chemical reactions

Exothermic processes

Heat flows out of the system to the surroundings

Endothermic processes

Heat flows into the system from the surroundings

Exothermic phase changes

Freezing, condensation, deposition

Endothermic phase changes

Melting, vaporization, sublimation

Exothermic processes at the molecular level

Some kind of bond or attractive force is forming

Value of ΔH⁰ in exothermic processes

Negative

Endothermic processes at the molecular level

Some kind of bond or attractive force is breaking

Value of ΔH⁰ in endothermic processes

Positive

Energy diagram in endothermic reaction

Products’ potential energy is higher

Energy diagram in exothermic processes

Reactants’ potential energy is higher

Change in heat in energy diagram

Difference in potential energy from reactants to products

Thermal equilibrium

Kinetic energy from the warmer object is transferred to the cooler object until both objects have the same average kinetic energy or temperature

Specific heat capacity (c)

Amount of energy required to raise the temperature of 1 gram of a substance by 1 degree

Specific heat units

J(g * °C)

Molar heat capacity

Amount of energy required to raise the temperature of 1 mole of a substance by 1 degree

Molar heat capacity units

J/(mol * °C)

Heat capacity (C_p)

Heat absorbed per degree of a substance

Law of conservation of energy equation

-q_A = q_B

Specific heat equation

q = mcΔt

q in specific heat equation

Heat or energy in joules

m in specific heat equation

Mass of object in grams

ΔT in specific heat equation

Change in temperature (K or °C)

Calorimeters

Measure enthalpy or heat changes

Substance inside calorimeter

Releases or absorbs heat

Water in calorimeter

Surrounds substance inside

Flow of heat

Results in a change in the water’s temperature until thermal equilibrium is reached

What is needed to figure out specific heat of a substance

Starting temperatures, ending temperatures, masses, and specific heat of water

Increasing temperature on a heating curve

Substance is in one phase and temperature keeps increasing as energy is added

Flat temperature on a heating curve

Substance is changing phases which requires energy

Melting and boiling points on a heating curve

Temperatures where phase changes occur (flat areas)

Equation for increasing temperature on a heating curve

q = mcΔt

Equation for flat temperature on a heating curve

q = amount * ΔH_fus or q = amount * ΔH_vap

Correct signage

In endothermic processes total heat should be positive, in exothermic processes total heat should be negative

Enthalpy

Same as heat or energy

ΔH vs ΔH_rxn

ΔH is the energy flow as heat at constant pressure, ΔH_rxn is the heat of the reaction or the change in enthalpy of a reaction

Signage of ΔH_rxn

Positive in endothermic processes and negative in exothermic processes

ΔH_rxn equation

ΔH_rxn = H_products - H_reactants

Representing enthalpy in equations

Denote the ΔH at the end of the equation based on the moles of reaction

Reversible reactions

Enthalpy of a reversed reaction is the negative of the enthalpy of the normal reaction

Bonds

Forces that hold groups of atoms together and make them function as a unit

Bond energy or enthalpy

Energy stored in a chemical bond

Single, double, and triple bonds ranked in energy

Triple bonds > double bonds > single bonds

Bond breaking heat transfer process

Endothermic

Bond formation heat transfer process

Exothermic

Enthalpy change with bond energies equation

ΔH = ΣD(bonds broken) - ΣD(bonds formed)

Bond energy cases

Bonds formed are being subtracted because bond formation releases energy, so only use the magnitudes of the energies

Standard enthalpy (ΔH°)

Enthalpy change at standard conditions

Standard conditions

1. For a gas pressure is 1 atm

2. For a solution concentration is 1 M

3. Temperature is 25 °C

Element types

Elements used to create these compounds are naturally-occurring or standard state elements found in nature

Calorimetry

Doing the reaction and finding out how much heat was produced or absorbed

Enthalpy of formation

Heat required to form

Formation enthalpies to determine ΔH° equation

ΔH_rxn° = Σn_pΔH_f°(products) - ΣnpΔH_f°(reactants)

Meaning of ΔH_rxn°

Standard enthalpy change for reaction

Meaning of Σn_pΔH_f°

Sum of all standard enthalpies of formation of products/reactants multiplied by their respective amounts

Why enthalpy is a state function

Change in enthalpy is the same whether the reaction takes place in one step or a series of steps

Hess’s law

Elementary steps can be combined to find the overall reaction, adding the enthalpy changes of the steps finds the enthalpy change for the overall reaction

Reversal in Hess’s law

If a reaction is reversed, the enthalpy change is also reversed

Multiplication in Hess’s law

If the coefficients of a reaction are multiplied by an integer, the enthalpy change is multiplied by the same integer