1
Q
products of combustion
A
CO2, H2O, energy
2
Q
products of incomplete combustion
A
CO or C (s), H2O, energy
3
Q
thermal energy
A
a form of kinetic energy that results form the motion of molecules
4
Q
law of conservation of energy
A
the total energy of the universe is constant, it can be transferred from one substance to another
5
Q
system vs. surrounding
A
part of universe being studied vs. everything else
6
Q
when is energy needed/release in a chemical reaction (bonding making and breaking)
A
- needed to break chemical bonds
- release when new bonds are made
7
Q
flow of energy in exothermic processes
A
- heat flows out of system into surroundings, increasing temp (feels hot to us, but system is losing energy)
8
Q
flow of energy in endothermic processes
A
- heat absorbs into system from surroundings, decreasing temp (feels cold to us, but system is gaining energy)
9
Q
open vs closed vs isolated systems
A
matter and E transferred vs. E transferred but not matter vs. matter and E cannot escape/enter
10
Q
specific heat capacity
A
the heat needed to increase the temperature of unit mass of material by 1K
- reflects ability to store energy
11
Q
types of motion of molecules
A
translational (linear), rotational (spins on axis), vibrational (bond shortening/lengthening)
12
Q
motion + SHC of monatomic gases
A
- translational
- low SHC
13
Q
motion + SHC of diatomic gases
A
- all 3 motions (in 1 bond)
- medium SHC
14
Q
motion + SHC of triatomic gases
A
- all 3 motions (in 2 bonds)
- high SHC
15
Q
relationship bw motion of molecules and SHC
A
the more ways a molecule can move, the higher the SHC
16
Q
freedom of motion of molecules in gas
A
all free
17
Q
freedom of motion of molecules in liquids
A
trans: some
rot: some
vib: free
18
Q
freedom of motion of molecules in solids
A
trans: none
rot: very little
vib: some/free
19
Q
enthalpy, where is it stored?
A
heat content of a chemical system
- in the chemical bonds/intermolecular forces as potential energy
20
Q
enthalpy change
A
amount of heat released/absorbed when a chemical reaction occurs at constant pressure
21
Q
where is energy (product/reactant) in an exothermic and endothermic reaction?
A
exo: product side
endo: reactant side
22
Q
types of enthalpy change
A
physical change, chemical change, nuclear reaction
23
Q
molar enthalpy
A
ΔH per mole of a substance undergoing a change
24
Q
how to calculate the amount of energy involved in a reaction
A
ΔH = nΔHx
25
potential energy direction for endo and exo reactions
endo: arrow points upwards
exo: arrow points downwards
26
standard molar enthalpy change of combustion
enthalpy change for the complete combustion of one mole of a substance in its standard state in excess oxygen under standard conditions
27
how can enthalpy changes of reactions in aqueous solutions be found experimentally
- carry out rxn in insulated system
- heat absorbed/released by the reaction can be measured from the temp change of water (q system = -q surroundings)
28
bond enthalpy
the energy needed to break one mole of bonds in *gaseous covalent molecules* under standard conditions
29
define average bond enthalpy
the average energy required to break one mole of the bond in similar compounds, in gaseous state
30
limitations of using average bond enthalpy in calculations
- they will all be slightly wrong so overall enthalpy change will not be precise (estimation)
31
how to calculate enthalpy change
ΔH = Hproducts - Hreactants
32
how to calculate enthalpy change using bond enthalpies
ΔH^or= sum(bonds broken, reactants) - sum(bonds formed, products)
33
reasons why calculated bond enthalpies is different from experimental values
- calculations: assumes rxn takes place in gas state w not intermolecular forces
- experiments: combustion involves octane and water in liquid states
- bond enthalpy values are averaged across wide range of compounds, approximation of true value
34
Hess's law
no matter route of chemical reaction, the enthalpy change will always be the same, as long as the initial and final states of system are the same
35
what is needed to draw an enthalpy diagram
- 2 equations must have same products
- 2/3 equations must have known energy change
36
standard enthalpy of formation
amount of energy associated with the formation of one mole of a substance from its elements in standard state
37
standard state
most stable form of an element
38
standard enthalpy of formation of an element at standard state
zero
39
born haber cycle
- application of Hess's law
- show energy changes in the formation of an ionic compound
40
lattice enthalpy/lattice dissociation enthalpy
enthalpy change when 1 moles of a solid ionic compound breaks down to form gaseous ions under standard conditions
41
what does lattice enthalpy represent
a measure of the strength of the ionic bonds bw the ions
42
what does a larger lattice enthalpy value mean
that more energy is required to overcome the attraction bw the ions and the stronger bond
43
what 2 factors does the magnitude of lattice enthalpy depend on
- ionic charge (charge of ions)
- ionic radii (size of atom/ions)
44
relationship between lattice enthalpy and ionic charge
- directly proportional
- as ionic charge increases, the ionic attraction (lattice enthalpy) bw the ions increase
45
relationship between lattice enthalpy and ionic radii
- inversely proportional
- as ionic radius increases, the ionic attraction (lattice enthalpy) decreases
46
can lattice enthalpy value be determined experimentally
no, must find using born-haber cycle
47
diff steps of born haber cycle
- enthalpy of formation
- enthalpy of atomization
- ionisation energy
- electron affinity
48
enthalpy of atomization
enthalpy change when one mole of gaseous atoms is formed from an element in its standard state
49
ionisation energy vs. electron affinity
energy required to remove one mole of electrons from one mole of gaseous atoms vs. energy release when one mole of electrons are added to one mole of gaseous atoms
50
entropy
a measure of the molecular disorder, randomness, or energy dispersal within a system (measures number of ways you can rearrange systme and have it look the same)
- entropy of universe increases for all spontaneous events
51
relationship bw entropy and disorder
- the greater the disorder of the particles, the greater the entropy
- particles naturally adopt more disordered state
52
large entropy vs small entropy vs entropy =0
random, less random, perfectly ordered
53
ΔS=neg vs. ΔS=pos.
decrease in randomness vs. increase in randomness
54
2 factors to predict change in entropy
1. physical state (solid=most ordered, gas=most random)
2. # of particles (more particles=more random)
55
standard entropy value
entropy change from heating substance from absolute zero to thermodynamic standard temp of 298K
- perfect crystal has standard enthalpy of 0, everything else has positive value
56
how to calculate standard enthalpy change
ΔS^o= sum(ΔSproducts) - sum(ΔSreactants)
57
what is a spontaneous process
a process that occurs without the addition of energy other than that required to overcome the initial energy barrier (activation energy)
58
what is a non-spontaneous reaction
one that needs a constant input of energy to occur
59
how to know is something is a spontaneous process
- total entropy of system and surroundings must increase
- ΔStotal = ΔSsystem+ΔSsurroundings >= 0
60
increase in entropy of the surroundings
- usually in form of heat release from system
- depending on temp of surroundings, heat can have small or large effect (cool =large effect, warm=small efect)
61
what reaction is favoured at high temps, why
- greater kinetic energy --> more motion of particles --> high randomness --> endothermic reaction is preferred
- to achieve maximum disorder with minimum energy
62
what reaction is favoured at low temps, why
- exothermic reaction
- to achieve maximum disorder with minimum energy
63
ΔG<0 vs ΔG>0 vs ΔG=0
spontaneous vs non spontaneous vs at equilibrium
64
gibbs energy
the combination of enthalpy, entropy, temp of a system which helps determine if a reaction is spontaneous under varying temperature conditions
65
spontaneity of reaction if ΔH= - and ΔS= +
yes, at any temperature
66
spontaneity of reaction if ΔH= + and ΔS= -
no, at any temperature
67
spontaneity of reaction if ΔH= - and ΔS= -
probable at low temperatures
68
spontaneity of reaction if ΔH= + and ΔS= +
probably at high temperatures
69
how to find gibbs energy at 298K
ΔG^or= sum(ΔGproducts) - sum(ΔGreactants)
- otherwise use databooklet
70
fossil fuels
- non-renewable and non-sustainable energy sources as it takes millions of years to form, used faster than they can be replaced
71
what is crude petroleum
a complex mixture of gas, liquid, solid hydrocarbons
- formed by pressure and temp on marine life sediments
72
what is natural gas
hydrocarbon mixture made mostly of methane with some higher alkanes and small amounts of CO2
73
alkanes
organic compounds composed entirely of single-bonded carbon and hydrogen atoms
74
trend of molecular mass of alkanes
- increases by a fixed amount going up the series as neighboring members differ by CH2
75
importance of relationship bw boiling point and # of carbon atoms of alkane in oil industry
more carbon atoms = higher boiling point
- allows for fractional distillation (ex. crude oil is heated, smaller hydrocarbons boil off first)
- different fractions are used for diff purposes
76
homologous series and chemical properties + examples
- members of the same homologous group have the same functional group --> show similar chemical properties/reactivity
- alcohols (-OH) can all oxidize
- corboxylic acids (-COOH) are acids
77
relationship bw number of carbon atoms of alcohol and enthalpy of combustion
- more carbon atoms = more energy released (value is more negative)
78
consequences of using fossil fuels
- release a lot of CO2 (product of combustion) --> GHG
- only makes up 0.04% of atmosphere but effects can be felt
79
how does CO2 work as a GHG
- CO2 absorbs infrared radiation --> vibrates bonds within molecule
- after, molecule will emit infrared radiation back into atmosphere --> some will be directed to Earth's surface, inc. global temp
80
specific energy of fuel + example
amount of heat energy released per mass of fuel
- wood has lowest se, natural gas highest
81
how to calculate specific energy
SE= (energy released from fuel) / (mass of fuel consumed) KJ/Kg
82
why are longer hydrocarbons more likely to undergo incomplete combustion
- have reduced volatility due to stronger LDF
- affects how hydrocarbons interact with oxygen molecules --> affects type of combustion
83
what are biofuels
- renewable resources produced from organic compounds which are generated from carbon dioxide during biological processes
84
biological carbon fixation + example
- production of organic compounds from CO2
- ex. green plants use photosynthesis to absorb CO2--> glucose --> fermentation into ethanol (a biofuel)
85
advantages of using biofuels
- renewable + sustainable (more CO2 absorbed than released)
- reduced GHG emissions
- wide range of plant materials can be used
- less depedence on oil supplies
86
disadvantages of using biofuels
- use of agricultural land/water/resources --> potential deforestation --> less food production
- high production cost
- monocultures reduce biodiversity
