Brainscape's Knowledge GenomeTM


Top Flashcards (Certified)
All Flashcards

Chemistry IB HL (2016+) > Topic 5: Energetics & Thermochemistry > Flashcards

Topic 5: Energetics & Thermochemistry Flashcards

(59 cards)

Start Studying
1
Q

exothermic reaction

A
  • heat energy is released

- energy released in bond formation on product side is greater than energy consumed in bond breaking on reactant side

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
2
Q

enthalpy change of exothermic reactions

A

ΔH = - x

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
3
Q

endothermic rxn

A
  • heat energy is absorbed

- energy released in bond formation on product side is less than energy consumed in bond breaking on reactant side

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
4
Q

enthalpy change of endothermic reactions

A

ΔH = + x

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

system

A
  • a specified part of the universe

- under observation/where a chem rxn is taking place

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

surroundings

A
  • the remaining portion of the universe

- NOT part of the system

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

enthalpy

A

heat content of the system

symbol: H

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

enthalpy change

A
  • heat absorbed/evolved during a rxn
  • measured at a constant temp/pressure
    symbol: ΔH
    unit: kJ mol^-1
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

enthalpy change equation

A 
ΔH = Hp - Hr
ΔH = enthalpy change
Hp = enthalpy of products
Hr = enthalpy of reactants
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

absolute zero

A

-273 degrees Celsius

the temp at which all particle movements cease completely

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

heat

A
  • form of energy
  • measure of total energy in a given amount of substance
  • depends on the amount of substance present
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

temperature

A
  • measures the ‘hotness’ of a substance
  • avg kinetic energy of the substance
  • independent of the amount of substance present
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

specific heat capacity

A 
heat required to raise the temp of 1g of a substance by 1°C/K
unit: J K^-1 g^-1
j = joules
k = kelvin
g = grams
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

heat capacity

A

heat needed to increase the temp of the object by 1°C/K

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

standard enthalpy change

A

enthalpy change when 1 mol of the gaseous bond is broken or formed

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

calorimetry

A

technique of measuring enthalpy change

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
17
Q

what is a calorimeter?

A
  • a well-insulated container (e.g. polystyrene cup)
  • in which temp change of a liquid is measured
  • before and after the change
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
18
Q

assumptions made when making heat calculations

A
  • no heat transfer between the soln, the thermometer, the surrounding air, and the calorimeter itself
  • the solution’s density and specific heat capacity are equivalent to water’s
  • rxn occured fast enough for max temp to be achieved before cooling begins
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
19
Q

problems with calorimetry

A
  • not having the desired rxn occur (e.g. incomplete combustion)
  • loss of heat to surroundings in exothermic reactions
  • absorption of heat from surroundings in endothermic reactions
  • using incorrect specific heat capacity value in enthalpy calculations
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
20
Q

what is a cooling graph?

A
  • for slow rxns (e.g. metal ion displacement), the results will be less accurate
  • due to heat loss over time
  • this can be compensated by plotting a temp-time graph (cooling graph) to extrapolate backwards
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
21
Q

how to draw a cooling graph

A
  • draw a line of extrapolation backwards from cooling curve
  • draw a vertical line where the reactants were mixed (i.e. when the curve begins to rise)
  • the y-coordinate of the point of intersection is the temp that would’ve been reached if no heat was lost to surroundings
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
22
Q

Hess’ Law

A
  • the total enthalpy change for a chem rxn doesn’t depend on the pathway it takes
  • only considers initial and final states
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
23
Q

1st law of thermodynamics

A

AKA: law of conservation of energy

  • energy can neither be created nor destroyed
  • it can only be converted from one form to another
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
24
Q

enthalpy of formation

Study These Flashcards
A
  • denoted by ΔHf
  • enthalpy change when 1 mole of a substance is formed from its elements
  • all substances being in standard states
  • enthalpy of formation of every element in its standard state is assumed to be 0!
25
what happens to enthalpy change when you reverse a rxn
signs are reversed | e.g. negative value turns positive and vice versa
26
Energy
measure of the ability to do work
27
conditions of measuring standard enthalpy change
- pressure: 100kPa - conc.: 1 mol/dm3 for all solutions - temp.: 298K (usually) - all substances in standard state
28
how to calculate reaction enthalpies from temp changes
ΔH (reaction) = - ΔH (water) = - m (H2O) x c(H2O) x ΔT(H2O) - it's negative bc water gains heat (positive enthalpy change) from the reaction - so if water's ΔH is positive then the reaction's ΔH is negative
29
how to calculate enthalpy change using Hess' Law
ΔH1 = ΔH2 + ΔH3 ΔHf (products) = ΣΔH (reaction) + ΣΔHf (reactants) ``` ΔH1 = enthalpy change of formation (of products) ΔH2 = enthalpy change of formation (of reactants) ΔH3 = enthalpy change of reaction ```
30
average bond enthalpy
- energy needed to break one mole of a bond in a gaseous molecule - can be used to calculate enthalpy change
31
limitations of using bond enthalpies
- can only be used if all reactants and products are in gaseous state; in other states, more heat would be evolved - bond enthalpies are averaged values that are obtained considering a number of compounds containing the bond, but in reality the bond energy varies between compounds
32
formula for specific heat capacity
q = mcΔT q: heat change m: mass c: specific heat capacity ΔT: change in temperature
33
formula for heat capacity
C = q/ΔT C: heat capacity q: heat change ΔT: change in temperature
34
assumptions made in enthalpy change calculations for solutions
``` ΔH(system) = ΔH(water) + ΔH(reaction) = 0 ΔH(reaction) = - ΔH(water) ```
35
enthalpy change: does change in temperature differ depending on volume of reactants?
volume halved = enthalpy halved = heat required also halved = change in temperature will be constant!
36
lattice enthalpy
AKA ΔH(lat) - endothermic: enthalpy change when 1 mol of gaseous ions form from 1 mol of a solid crystal - exothermic: enthalpy change when 1 mol of solid crystal forms from 1 mol of gaseous ions - sign (+/-) of ΔH(lat) denotes whether it's endothermic/exothermic
37
factors affecting lattice enthalpy/enthalpy of hydration
↓ size + ↑ charge = ↑ ΔH(lat) ↓ size + ↑ positivity of charge = ↓ ΔH(hyd) -- because the enthalpy of hydration is always negative
38
enthalpy change of hydration
- enthalpy change occurring when 1 mol of gaseous ions is dissolved to form an infinitely dilute solution of 1 mol of aqueous ions - always exothermic so ΔH(hyd) is always negative
39
entropy
- i.e. randomness/disorder - denoted by S - defined as the distribution of available energy among particles - higher energy distribution (i.e. diffused state) = higher entropy - all entropy values are positive except for perfectly ordered solids at absolute zero (S = 0)
40
how to predict ΔS with relation to state
- conversions from s --> l/g have positive ΔS | - vice versa for g --> l/s
41
formula to calculate ΔS
ΣΔS(system) = ΣΔS(products) - ΣΔS(reactants)
42
why are exothermic reactions more common than endothermic reactions?
- because ΔS(surroundings) ∝ -ΔH(system) | - exothermic reactions cause an increase in the entropy of the surroundings
43
relationship between ΔS and temperature
ΔS(surroundings) ∝ 1/T | - i.e. entropy of surroundings is inversely proportional to absolute temperature
44
formulae for ΔS
ΔS(total)* = ΔS(system) - ΔS(surroundings) *total denotes total entropy of the universe. ΔS(surroundings) = - ΔH(system) / T
45
how to deduce whether a reaction is feasible
as long as ΔS(total) > 0, a reaction will occur spontaneously
46
Gibbs free energy
- criterion of feasibility of a reaction occurring - feasible reactions = spontaneous - ΔG is always negative for spontaneous processes and positive in the reverse case - ΔG is NOT the same as ΔS
47
formula for Gibbs free energy
ΔG(system) = -TΔS(total) ΔG(system) = ΔH(system) - TΔS(system) T must be in Kelvin (in standard conditions, T is 298K)
48
how can ΔG(system) predict feasibility?
- as ΔS is inversely proportional to T, the T value adjusts the importance of ΔS(system) in the gibbs free energy equation - at low temps, ΔS is high but T is low so overall TΔS(system) ≈ 0. Thus ΔG(system) = ΔH(system) - at high temps ΔH is negligible so ΔG(system) = TΔS(system)
49
formula for ΔG(reaction)
ΣΔG(reaction) = ΣΔG(products) - ΣΔG(reactants)
50
Relationship between ΔG(reaction) and equilibrium
as ΔG(reaction) becomes more negative (i.e. as reaction feasibility increases), the equilibrium position shifts more to the right
51
linking ΔS to ΔG
ΔS(total) = ΔS(system) - ΔS(surroundings) Knowing ΔS(surroundings) = - ΔH(system) / T, we can substitute that: ΔS(total) = ΔS(system) - ΔH(system) / T -TΔS(total) = ΔH - TΔS(system) Knowing ΔG(system) = -TΔS(total), we can substitute that: ΔG(system) = ΔH(system) - TΔS(system)
52
enthalpy of atomization
ΔH(at)* - standard enthalpy change when 1 mol of gaseous atoms are formed from the element in its standard state under standard conditions - for diatomic molecules the ΔH(at) value is half that of the bond dissociation enthalpy *standard enthalpies like ΔH(atom) are always written with a theta in superscript
53
difference between bond dissociation enthalpy and bond enthalpy
- bond dissociation enthalpy refers to the enthalpy of a specific bond in a specific compound - bond enthalpy refers to the average enthalpy of bond, taking into account different compounds - a specific bond's energy differs somewhat between compounds, so bond dissociation enthalpy is more specific/accurate
54
enthalpy of solution
ΔH(sol) | - enthalpy change when 1 mol of ionic substance dissolves in water to form a solution of infinite dilution
55
Hess' Law for the solubility of salts
ΔH(sol) = ΔH(lat) + sum of ΔH(hyd) of ALL constituents | NOTE: ΔH(sol) and ΔH(lat) are values of compounds while ΔH(hyd) is individual values of ions
56
why can't H2O be added to SO3?
- taking SO3's non-ionic nature into account, its ΔH(lat) value will be small - considering ΔH(sol) = ΔH(lat) + ΔH(hyd of ALL constituents), remember the fact that ΔH(hyd) is always negative (exothermic) - as ΔH(lat) will be small, the overall ΔH(sol) value for H2O + SO3 will be highly exothermic - the evolved heat could be high enough to cause the product H2SO4 to boil
57
factors affecting ΔH(hyd)
- ionic radius | - charge of ion
58
factors affecting ΔH(hyd): ionic radius
smaller ionic radius = higher ΔH(hyd)
59
factors affecting ΔH(hyd): charge of ion
more positive charge = higher ΔH(hyd) | - due to increased attraction between the positive ion and the partially negative O atoms in water

Chemistry IB HL (2016+) (13 decks)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2025 Bold Learning Solutions. Terms and Conditions