ch10 - thermochemistry

Law of Conservation of Mass

mass cannot be created nor destroyed

Law of Conservation of Energy

Energy (ex: heat) is conserved; it can change form but cannot be created or destroyed.

Energy

The capacity to do work.

Radiant Energy

Energy from the sun; Earth’s primary energy source.

Kinetic Energy

Energy of motion; KE = 1/2 mv².

Thermal Energy

Energy associated with the random motion of atoms and molecules.

Chemical Energy

Energy stored in the bonds of chemical substances.

Nuclear Energy

Energy stored within the protons and neutrons of an atom.

Gravitational Energy

stored potential energy due to height of object

Heat

Transfer of thermal energy between bodies at different temperatures.

Temperature

Measure of a substance’s thermal energy.

what is the difference between heat and temperature?

Heat refers to the transfer of thermal energy between bodies, while temperature is a measure of a substance’s thermal energy.

Thermochemistry

Study of heat changes in chemical reactions focusing on initial and final states (intermediate steps don’t matter).

System

Specific part of the universe under study.

Open System

Exchanges both mass and energy with surroundings.

Closed System

mass is contained, but the system exchanges energy with the surroundings.

Isolated System

Exchanges neither mass nor energy with surroundings.

Exothermic Process

Releases thermal energy from system to surroundings (ΔH < 0).

In a reaction equation, energy is a product.

Endothermic Process

Absorbs thermal energy from surroundings into system (ΔH > 0).

In a reaction equation, energy is a reactant.

define the First Law of Thermodynamics and its corresponding equations.

Energy can change forms but cannot be created or destroyed; the total energy of an isolated system remains constant.

• ΔEsystem + ΔEsurroundings = 0

• ΔEsystem = -ΔEsurroundings

• ΔE = q + w

  • w = -PΔV

define the variables in the formula:

ΔE = q + w

ΔE = the change in internal energy

q = heat exchanged between system and surroundings

w = work done on/by the system

• w = -PΔV

what does positive work mean? (+w)

work on system by surroundings

what does negative work mean? (-w)

work by system on surroundings

what does positive heat mean? (+h)

endothermic

heat absorbed by system from surroundings

what does negative heat mean? (-h)

exothermic

heat released by system to surroundings

define enthalpy and its corresponding variable and equations.

used to quantify the heat flow into or out of a system at constant pressure

• Enthalpy = H

• q = ΔH

  • q = msΔt

• ΔE = ΔH - PΔV (same as ΔE = q+w)

• ΔH = Hproducts − Hreactants

how do you know if ΔH is endothermic or exothermic?

ΔH = Hproducts − Hreactants

• ΔH > 0 = endothermic

• ΔH < 0 = exothermic

define the variables in q = msΔt

q = amount of heat absorbed/released

m = mass

s = specific heat

Δt = Tfinal - Tinitial = change in temperature

Bond Dissociation Energy

Energy required to break a specific chemical bond.

Lattice Dissociation Energy

Energy needed to break an ionic solid into its gaseous ions.

Heat of Formation

Heat required to form 1 mole of a compound from its elements in standard states.

compare ΔH and ΔE

ΔH refers to the enthalpy change, or the change in heat flowing in or out of a system at constant pressure.

ΔE represents the change of internal energy in a system.

relationship: ΔE = ΔH - PΔV

Specific Heat (s)

Heat required to raise 1 g of a substance by 1 °C.

Heat Capacity (c)

Heat required to raise the temperature of a given quantity by 1 °C.

Hess’s Law

Total enthalpy change for a reaction is the same regardless of the pathway taken.

• only final and initial states of the system are compared, the intermediate states don’t matter (given that the conditions are the same)