IB Chemistry SL > R1 > Flashcards

R1 Flashcards

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1
Q

Entropy

A

(S) A measure of the amount of disorder of the particles in a system.

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2
Q

Particle behavior in a system with high entropy

A

The particles are more spread out and moving more.

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3
Q

Particle behavior in a system with low entropy

A

The particles are not as spread apart and are moving less

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4
Q

Order of states of matter for increasing entropy

A

Solid, liquid, gas

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5
Q

Combination of what substances lead to increased entropy

A

The combination of pure substances to form homogeneous or heterogeneous mixtures

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6
Q

An increase in the number of moles of gas effect on entropy

A

Increase in entropy

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7
Q

An decrease in the number of moles of gas effect on entropy

A

Decrease in entropy

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8
Q

The same number of moles of gas effect on entropy

A

Entropy with stay the same

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9
Q

Sign for standard entropy values

A

S⦵

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10
Q

Sign for standard entropy change

A

ΔS⦵

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11
Q

The standard entropy of a substance

A

Entropy change from heating the substance from absolute zero (0 K) to the thermodynamic standard temperature of 298 K.

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12
Q

Standard entropy of a perfect crystal at absolute zero (0 K)

A

Zero

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13
Q

Standard entropy value of all substances

A

positive

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14
Q

Equation for the standard entropy change (ΔS⦵) for a reaction

A

ΔS⦵ = ΣS⦵ (products) − ΣS⦵ (reactants)

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15
Q

When spontaneous reactions will happen

A

Once the energy needed to start the reaction is provided, activation energy

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16
Q

Spontaneous reaction

A

A chemical reaction that favors product formation once the activation energy has been provided to start the reaction

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17
Q

Spontaneous reaction ability to occur without energy other than activation energy

A

Yes

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18
Q

Non-spontaneous reaction ability to occur without energy other than activation energy

A

No

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19
Q

Ball pushed down a hill type of reaction

A

Spontaneous

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20
Q

Ball pushed up a hill type of reaction

A

Non-spontaneous

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21
Q

The second law of thermodynamics

A

The total entropy of the universe tends to increase

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22
Q

The total entropy of the system and surroundings in spontaneous reaction

A

Must increase

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22
Q

Energy form in spontaneous reaction

A

A concentrated form to a dispersed form

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23
Q

Spontaneous proccess entropy equation

A

ΔS total = ΔS system + ΔS surroundings ≥ 0

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24
Entropy of the system and suroundings in a spontaneous process
The system can have a decrease in entropy by the surrounding's entropy will always be so large that the total increases
25
Gobbs energy change equation
ΔGo = ΔHo - (T)(ΔSo)
26
ΔGo units
kilojoules per mole (kJ mol–1)
27
ΔHo
kilojoules per mole (kJ mol–1)
27
T
Temperatur Kelvin
28
Gibbs Energy
Energy associated with a chemical reaction that can be used to do useful work, ΔG
29
If the Gibbs energy change, ΔG, is negative
The reaction is spontaneous
30
If the Gibbs energy change, ΔG, is positive
The reaction is non-spontaneous
31
If the Gibbs energy change, ΔG, is zero
The reaction is at equilibrium.
32
Exothermic reactions that result in an increase in entropy
The value of the ΔG will be negative at any temperature
33
Endothermic reactions with a decrease in entropy
The value of the ΔG is positive at any temperature
34
Exothermic reactions with a decrease in entropy
Only spontaneous at low temperatures
35
Endothermic reactions with an increase in entropy
Only spontaneous at high temperatures
36
equilibrium
Ratio between the concentration of products and reactants for a reversible reaction at equilibrium.
36
reversible
Reactions that can go in both directions, i.e. reactants to products (forward reaction) then back from products to reactants (reverse reaction). These reactions are represented by the symbol ⇌.
37
forward reaction
Transformation of reactants into products in a reversible reaction.
38
backward reaction
Transformation of products into reactants in a reversible reaction.
39
equilibrium constant
Ratio between the concentration of products and reactants for a reversible reaction at equilibrium, K
39
reaction quotient
Ratio between the concentration of products and reactants for a reversible reaction not at equilibrium, Q
40
equilibrium constant for the general reaction
if aA + bB ⇌ cC dD then K= (C)^c(D)^d/(A)^a(B)^b
41
for the equilibrium constant, what do a, b, c, and d represent
the stoichiometric coefficients in the balanced equation
42
Compared to the equilibrium constant formula, how is the reaction quotient one
written the same way, but the concentrations of products and reactants used should be the ones when the system is not at equilibrium.
43
Q=K
Reaction in equilibrium
43
Q>K
Backward reaction favored
44
Q
Forward reacttion favored
45
equation used to calculate the Gibbs energy change for a reaction not at equilibrium
ΔG = ΔG° + RT ln Q
46
R
Universal gas constant (8.314 J/mol·K)
47
equation used to calculate the Gibbs energy change for a reaction at equilibrium
ΔG° = -RTlnK
48
how to take J to kJ
divide by thousand
49
when the forward reaction is favored
spontaneous
50
When the position of the equilibrium lies to the right on a reversible reaction
he forward reaction is spontaneous
51
For a spontaneous reaction, the position of equilibrium
lies towards the products side,K>1
52
For a non-spontaneous reaction, the position of equilibrium
lies towards the reactants side, K<1
53
redox reaction steps
1) split reaction into reduction and oxidation half reaction 2) balance all non-oxygen and non-hydrogens atoms first 3) balance oxygens by adding H2O 4) baland hygrogens by adding H+ 5) balance charge by adding e- 6) multiply reaction if needed to make sure e- lost gained are equal 7) add reactions togther and cancel out the same things on each side 8*) Add OH- to nuetrilize existing H+ 9*) cancle out H2O if needed
54
Complete combustion
Sufficient oxygen supply, Carbon Dioxide, Water, Clean, blue, and hot flame, Highest energy release, most efficient, no smoke, production fully oxidized
54
Incomplete combustion
Insufficient oxygen, Carbon Monoxide, Soot (unburnt Carbon), and Water, Yellow/orange, sooty flame, with glowing carbon particles, Lower energy release, less efficient, dangerous, O2 supply limited
55
Superincomplete combustion
extreme lack of O2, slow flameless, Primarily soot and CO, very little CO2 or heat.
56
bond breaking
energy is absorbed endothermic +ΔH
57
bond forming
energy is released exothermic -ΔH
57
equation to solve for thermal energy
Q = mcΔT
58
types of exothermic reaction
combustion neutralization
58
equation to calculate the change in enthalpy
ΔH = -Q/n
59
types of endothermic reaction
photosynthesis thermal decomposition
60
commonaltiy of exothermix and endothermic reactions
exothermic, because the activation energy is lower and productions are more stable
61
energy profile axis
y - potential energy x - reaction coordinate
62
energy profile stability
reactant/product has a smaller potential energy reactant - endothermic product - exothermic
63
activation energy on energy profiles
bottom of reactant to maximum potential energy
63
ΔH on energy profiles
bottom of reactant to bottom of productc
64
ΔG⦵ always negative
ΔH⦵ - negative ΔS⦵ - positive -TΔS⦵ - negative Spontaneous at T - all T
65
ΔG⦵ always positive
ΔH⦵ - positive ΔS⦵ - negative -TΔS⦵ - positive Spontaneous at T - non-spontaneous at all T
66
ΔG⦵ negative at low T
ΔH⦵ - negative ΔS⦵ - negative -TΔS⦵ - positive Spontaneous at T - only at lox T
67
ΔG⦵ negative at high T
ΔH⦵ - positive ΔS⦵ - positive -TΔS⦵ - negative Spontaneous at T - only at high T
68
Coal: State at room temp and Composition
Solid Mainly carbon, traces of hydrogen, oxygen, nitrogen, sulfur and metals
69
Crude Oil: State at room temp and Composition
Liquid Mixture of medium to long chain hydrocarbons, traces of sulfur, nitrogen and oxygen and metals
70
Natural Gas: State at room temp and Composition
Gas Mostly methane, traces of other small hydrocarbons, water and carbon dioxide
71
Greenhouse Gas Effect
a phenomenon that occurs when certain gases, such as carbon dioxide, methane or nitrogen oxides, absorb infrared (IR) radiation that has been emitted from the Earth’s surface after it has been heated by the Sun’s rays
72
largest tendency to undergo incomplete combustion
coal
73
smallest tendency to undergo incomplete combustion
natural gas
74
biofuels
A fuel derived directly from living matter.
75
fuel cell
A device that uses an external source of fuel and oxygen (or air) to transfer energy from a chemical store by electricity
76
biofuels examples
ethanol and biodesiel
77
carbon chains of clean combustions
shorter
78
biofuel advantage
Release less CO2 compared to fossil fuels Inexpensive Can be used directly after synthesis – no mining or processing at a refinery as with fossil fuels
79
biofuel disadvantage
Uses land suitable for agriculture Combustion still releases CO2 into the atmosphere Growth of crops requires the use of fossil fuels
80
bond enthalpy equation
reactants - products
81
combustion enthalpy equation
reactants - products
82
formation enthalpy equation
products - reactants
IB Chemistry SL
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