12CHB - Term 1 Flashcards
(86 cards)
1
Q
Gas Law
A
PV = nRT
2
Q
1 atmosphere
A
100kPa
3
Q
0 degrees celcius
A
273 Kelvin
4
Q
Intermolecular Forces
A
Forces between a positive component and a negative component
5
Q
Relative Atomic Mass (Ar)
A
The weighted average of the masses of its isotopes relative to 1/12 of the mass of a carbon-12 atom
6
Q
Relative Molecular mass (Mr)
A
The weighted average of the masses of the molecules relative to 1/12 of the mass of carbon-12 atom
7
Q
Reasons for low experimental yield
A
Loss during Transfer, Equilibrium/Reaction did not complete, side reactions
8
Q
Reasons for high experimental yield
A
Insufficient drying, gain of oxygen (oxidation), side reactions
9
Q
Ideal Gas
A
Particles have negligible volume, No attractive forces between the particles, kinetic energy of the particles is directly proportional to the absolute temperature.
10
Q
Ammonium
A
NH4 +
11
Q
Carbonate
A
CO3 2-
12
Q
Hydrogen Carbonate
A
HCO3 -
13
Q
Hydroxide
A
OH-
14
Q
Nitrate
A
NO3 -
15
Q
Nitrite
A
NO2-
16
Q
Phosphate
A
PO4 3-
17
Q
Phosphite
A
PO3 3-
18
Q
Sulfate
A
SO4 2-
19
Q
Sulfite
A
SO3 2-
20
Q
Ethanedioate (Oxalate)
A
C2O4 2-
21
Q
Peroxide
A
O2 2-
22
Q
Hydrochloric Acid
A
HCl
23
Q
Nitric Acid
A
HNO3
24
Q
Phosphoric Acid
A
H3PO4
25
Sulfuric Acid
H2SO4
26
Ethanoic Acid
CH3COOH
27
Homogenous Mixture
Has the same uniform appearance and composition. Any sample will be exactly the same
28
Heterogenous Mixture
Contains visible different phases, so all samples are not the same
29
Define Temperature
A measure of average kinetic energy
30
Define Heat Energy
The amount of energy exchanged due to a temp difference between 2 substances
31
Define Enthalpy
The heat content of a system, can be transferred to surroundings
32
Define Exothermic reaction
Release heat energy as products form (temp increases), products are more stable
33
Define Endothermic Reaction
Absorbs heat energy as products form (temp decreases, products are less stable
34
Breaking bonds
Requires Energy, Endothermic (S-> L -> G)
35
Forming bonds
Releases Energy, Exothermic (G -> L -> S)
36
∆H Enthalpy for reaction from a potential energy and reaction coordinate graph
Difference between reactants and products
37
Bond Enthalpy equation
∆rH = ΣE products - ΣE reactants
∆H=Enthalpy for reaction
E = Enthalpy
38
Define Bond Enthalpy
Always Positive because breaking bonds requires energy
39
Define Average bond enthalpy
The energy needed to break 1 mole of covalent bond in a gaseous molecule
40
Hess's Law
Energy change in multiple steps, is the sum of enthalpy change of individual steps
41
Enthalpy of enthalpy using heat energy equation
∆H = -q/n
42
Define General specific heat capacity
The amount of energy, a substance absorbs to raise its temp by 1 degree
43
Define Enthalpy of Combustion
Energy released when 1 mole of substance burns with eccess oxygen under standard conditions
44
Causes of Incomplete combustion
low O2, rx is too quick, depletes O2, Low combustion temp, large fuel particles
45
Complete combustion qualities
Blue flame, higher temp, no smoke
46
Incomplete combustion
Yellow flame, Less energy per mole. Harmful byproducts, Unreacted fuel is a hazard
47
Carbon Footprint
total greenhouse gas
47
Photosynthesis
6CO2 + 6H2O + λ -> C6H12O6 + 6O2
Chlorophyll containing pigments absorb wavelength, energy used to make glucose
48
Biofuels
Plant material converted to energy. Breakdown of larger organic molecules into short chain hydrocarbons
49
Enthalpy of Formation
Change in enthalpy when 1 mole of a substance is formed from its elements in their standard states
50
Enthalpy of Formation equation
∆rH = Σ∆fH products - Σ∆fH reactants
51
Bond enthalpy vs Formation Enthalpy
Bond Enthalpy result uses average values, doesn't take into account different environments so its inaccurate
52
Enthalpy of the reaction using standard combustion enthalpies
∆Hr = Σ∆Hc reactants - Σ∆Hc products
53
Ionisation Energy
The energy required to remove one mole of the most loosely held electrons from one mole of gaseous atoms/ions to produce 1 mole of gaseous ions
54
Electron affinity
the enthalpy changed when one mole of electrons is added to one mole of gaseous atoms/ions. Mostly exothermic
55
Atomisation Enthalpy
Enthalpy changes when one mole of gaseous atoms form from elements in their standard state
56
Hydration enthalpy
Exothermic. The amount of energy released when one mole of gaseous ions undergoes hydration
57
Enthalpy of Solution
The enthalpy change of solution is the enthalpy change when 1 mole of ionic substance dissolves in water to give a solution of infinite dilution
58
Define Lattice Enthalpy
A measure of the strength of the forces between the ions in an ionic solid. The greater the lattice enthalpy, the stronger the forces. Depends on size of the ion and the amount of charge
59
Gibbs free energy equation
ΔG = H - TΔS (Negative = Spontaneous)
60
Define Entropy
A Measure of the dispersal of energy
61
Entropy: Ordered state
Low energy distribution = Low Entropy
62
Entropy: Higher Disordered state
Higher energy distribution = High Entropy
63
Amount of Entropy in system Equation
ΔS = ΣS products - ΣS reactions
64
How to identify changes in Entropy
- Change in num. of moles (↑mol = ↑entropy)
- Phase Change (Most prominent) (↑phase (solid→liquid) = ↑entropy)
- Dissolving/Mixing (Mixing of things = ↑entropy)
65
Exothermic reaction coordinate
Gives off heat, surrounding entropy increases
66
Endothermic reaction coordinate
Absorbs heat, cold, surroundings will decrease in entropy
67
ΔS (total) equation
ΔS total = ΔS system + ΔS surroundings (total increase = spontaneous)
68
Define Gibbs Free Energy
The amount of energy avaliable to do work. If positive = useful energy of the universe increasing, negative decreasing
69
Calculate change in Gibbs free energy
ΔGf = ΣGf products - ΣGf reactants
70
How to measure rate of reaction (5)
- Change in Colour
- Volume change
- Conductivity
- Change in pH
- Change in mass
71
Collision Theory
- Particles need to collide
- Need to collide with correct orientation
- Collide with enough energy > Activation energy
72
Factors changing rate of reaction
- Concentration
- Pressure
- Catalyst
- Temperature
Surface area
73
Effect of Concentration on ror
Higher concentration means more particles in a given volume can react
74
Effect of Pressure on ror
Increase in pressure, more particles in a given volume and will result in more frequent collisions
75
Effect of Catalyst on ror
Provides an alternate pathway, that is lower in activation energy, so more particles have successful collisions
76
Effect of Temperature on ror
2 effects
- Particles have higher average kinetic energy, so more likely that energy > activation energy
- Particles are moving faster, so there are more frequent collisions
77
Effect of Surface area on ror
More exposed particles, so greater proportion of particles available to collide, so more frequent collisions
78
Define Activation energy
The difference between reactants and the top of the peak, energy required for the reaction to occur
79
Problems with measuring ror
Reactions that are very fast, timing is more difficult to be accurate, and will have a higher percentage error.
80
Define Spontaneous
Happens without an external force being applied
81
Define Non-Renewable Fuels
Not replaceable in a reasonable amount of time
82
Define Renewable Fuels
Easily replaceable in a short amount of time
83
Examples of Non-Renewable Fuels
Crude oil, Natural Gas, Coal
84
Example of Renewable Fuels
Ethanol (From Sugars), Biodiesel (Vegetable oil)
85
Differences between fuels (5)
- Higher proportion of H = more energy per gram
- More hydrogen = less CO2 Produces for the same amount of energy, reduced climate impact
- Short chains clean burn
- Long chains more stable, less likely to explode
- Long chains easier to trasport
