5..., 15.. Flashcards
Heat
q
is form of energy that flows down the temperature gradient from an object of higher temperature to an object of lower temperature
Calorimeter
- an apparatus used to measure the amount of heat being exchanged with surroundings
specific heat capacity
amount of heat required to raise the temperature of one gram of a substance by one degree Celsius
Temperature
is a measure of the average kinetic energy of particles
absolute temperature (in kelvin) = temp (Celsius) + 273
Bond breaking
requires energy
endothermic
heat is absorbed
bond making
releases energy
exothermic
Heat is released
Systems
Open system – matter and energy can be transferred between the system and surroundings.
Closed system – energy can be transferred between the system and surroundings but matter cannot.
Isolated system – neither matter nor energy can be transferred between the system and surroundings.
exothermic reactions
- – ∆H
- ## heat is transferred from the system to the surroundings
Endothermic
- ∆H
- heat is absorbed from the surroundings
Enthalpy change of combustion
(ΔHѳc) is the enthalpy change upon complete combustion of 1 mol of a substance.
always negative as heat is released during combustion
reactants - products
enthalpy change of formation
is the enthalpy change upon formation of 1 mol of a substance from its elements
products - reactants
Limitations of Calculating enthalpy changes
- Heat loss to surroundings
- Incomplete combustion
- Assumptions made about specific heat capacity and density of aqueous solution
Hess’ Law
the total enthalpy change in a chemical reaction is independent of the route by which the chemical reaction takes place, as long as the initial and final conditions are the same.
The enthalpy change for a reaction that is carried out in a series of steps is equal to the sum of the enthalpy changes for the individual steps.
Average bond enthalpy
enthalpy change when one mole of bonds are broken in
the gaseous state, averaged for the same bond in similar compounds
reactants - products
Limitations of average bond enthalpy calculations
Enthalpy changes calculated from average bond enthalpies may be different than enthalpy changes calculated from actual values due to the averaging.
Calculations involving bond enthalpies may be inaccurate because they do not take into account intermolecular forces.
Ozone layer
- The ozone layer in Earth’s atmosphere absorbs harmful UV radiation from the sun. The ozone is in dynamic equilibrium with oxygen and is continually being formed and destroyed.
- The covalent bond in oxygen (O2) is broken by UV radiation. A free radical is a highly reactive species due to the presence of an unpaired electron ().
O2 (g) → O (g) + O* (g) - The highly reactive free radical oxygen atoms react with oxygen (O2) to form ozone (O3).
O2 (g) + O* (g) → O3 (g) - Ozone decomposes to form oxygen (O2) and an oxygen free radical. The bonds in the ozone (O3) molecules are weaker than the double bonds in oxygen (O2), so they can be broken by lower energy UV radiation.
O3 (g) → O2 (g) + O* (g)
Lattice Enthalpy
The standard enthalpy change that occurs on the formation of 1 mol of gaseous ions from the solid lattice. The symbol is ΔHѳlat
- endothermic
- + value
NaCl (s) → Na+ (g) + Cl- (g) ΔHlatꝋ = +790 kJ mol -1
Ionisation energy
standard enthalpy change that occurs on the removal of 1 mole of electrons from 1 mole of gaseous atoms or positively charged ions
- endothermic
- + value
ΔHIE1ꝋAl (g) → Al+ (g) + e– ΔHIE1ꝋ = +577 kJ mol-1
Enthalpy atomisation
standard enthalpy change that occurs on the formation of 1 mole of separate gaseous atoms an element in its standard state
- endothermic
- + value
Na (s) → Na (g) ΔHatꝋ = +108 kJ mol -1
½Cl2 (g) → Cl (g) ΔHatꝋ = +122 kJ mol -1
Electron Affinity
energy change when 1 mole of electrons is gained by 1 mole of gaseous atoms of an element to form 1 mole of gaseous ions under standard conditions
- first: exothermic
Cl (g)+ e– → Cl– (g) ΔHEAꝋ = -364 kJ mol-1
- second: endothermic
O– (g) + e– → O2- (g) ΔHEAꝋ = +844 kJ mol-1
Enthalpy of solution
enthalpy change when 1 mole of an ionic substance dissolves in sufficient water to form an infinitely dilute solution
- can be exothermic or endothermic
LiBr (s) → LiBr (aq) ΔHsolꝋ = -48.8 kJ mol -1
KCl (s) → KCI (aq) ΔHsolꝋ = +17.2 kJ mol -1
Enthaloy of Hydration
nthalpy change when 1 mole of a specified gaseous ion dissolves in sufficient water to form an infinitely dilute solution
Mg2+ (g) → Mg2+ (aq) ΔHhydꝋ = -1963 kJ mol -1
Br- (g) → Br- (aq) ΔHhydꝋ = -328 kJ mol -1
- exothermic
Born-Haber cycles
a specific application of Hess’s Law for ionic compounds and enables us to calculate lattice enthalpy, which cannot be found by experiment
The basic principle of drawing the cycle is to construct a diagram in which energy increases going up the diagram
Born-Haber cycle calculations
ΔHfꝋ = ΔHatꝋ + ΔHatꝋ + IE + EA - ΔHlattꝋ
or
ΔHfꝋ = ΔH1ꝋ - ΔHlattꝋ
where ΔH1ꝋ is the sum of all of the various enthalpy changes necessary to convert the elements in their standard states to gaseous ions
Entropy
A measure of the distribution of available energy among the particles. The more ways energy can be distributed, the higher the entropy.
gases have more entropy than solids
Spontaneous changes regards to entropy
entropy increases as state changes from solid to liquid to gas
Gibbs free energy and spontaneouse reactions
An equation to combine enthalpy, entropy, and temperature of a system
ΔGѳ = ΔHѳ – TΔS
For reaction to be spontaneous, Gibbs Free energy must have negative value
Gibbs free energy
Represents the free energy change when 1 mol of a compound is formed from its elements under standard conditions
If the Gibbs free energies of formation are not known, use enthalpy and entropy data.
ΔGѳ = ΔHѳ – TΔS
Spontaneity (ΔGѳ = ΔHѳ – TΔS)
Spontaneity (ΔGѳ = ΔHѳ – TΔS)
energy transferred
Q is in joules
