Brainscape's Knowledge GenomeTM

Browse over 1 million classes created by top students, professors, publishers, and experts.


See full index

A-Level Chemistry > 3.2 Physical chemistry > Flashcards

3.2 Physical chemistry Flashcards

You may prefer our related Brainscape-certified flashcards:
Periodic Table flashcards Chemistry 101 flashcards AP® Chemistry flashcards Organic Chemistry 101 flashcards
1
Q

enthalpy definition

A

measure of heat energy stored in a system (H)

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

enthalpy change definition

A

(ΔH) difference between the enthalpy of reactants and enthalpy of products

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

heat definition

A

process whereby thermal energy (J) is transferred from hotter object to a cooler object

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

temperature

A

direction of energy transfer is determined by temperature of objects (measured in K, °C, °F)

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

system definition

A

substances directly involved in a chemical reaction

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

surroundings definition

A

apparatus, laboratory, anything not part of the system

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

universe definition

A

system+surroundings

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

law of conservation of energy

A

energy cannot be created or destroyed

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

exothermic meaning

A

energy released to and gained by surroundings, temperature increased
ΔH is negative
more energy released from making bonds than energy absorbed from breaking bonds

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

endothermic meaning

A

system takes in energy from surrounding, temperature decrease
ΔH is positive
more energy absorbed from breaking bonds than energy released from making bonds

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

enthalpy of reaction definition

A

overall energy change in a reaction

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

activation energy definition

A

minimum energy required for a reaction to take place

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

standard enthalpy of reaction definition

A

overall energy change in a reaction under standard conditions
100 kPa, 298 K / 25°C, conc. of 1.0 mol dm^-3 (for reactions with aqueous solutions)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

standard conditions

A

100kPa,
298K/25°C
1 mol dm^-3 (for solution)

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

standard state definition

A

physical state of substance under standard conditions

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

standard enthalpy change of combustion definition

A

energy change when one mole of substance is combusted in excess oxygen

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

standard enthalpy change of formation definition

A

overall energy change of 1 mol of a compound from its elements

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
18
Q

standard enthalpy change of neutralisation definition

A

overall energy change when 1 mole of water is produced in a reaction with an acid and base

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
19
Q

enthalpy change of reaction formula

A

ΔrH = sum of bond enthalpies of broken bonds - sum of bond enthalpies of bonds made

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
20
Q

average bond enthalpy definition

A

breaking of 1 mol of bonds in gaseous molecules

21
Q

why average bond enthalpy differs from actual bond enthalpy (and less accurate than formation enthalpies)

A

different bond enthalpies are measured in different conditions (no universal, unchanging standard saying which molecules used to determine each bond)

22
Q

Hess law definition

A

if reaction takes place by more than one route from initial to final conditions, total enthalpies change is the same for each route

23
Q

simple collision theory

A

reaction takes place when molecules collide with sufficient energy and correct orientation

24
Q

factors affecting rate of reaction

A

concentration/pressure increases rate of collision (particles closer together, more likely to collide)
temperature increases heat energy (more kinetic energy so faster molecules so higher rate of collision, also more molecules with heat energy equal or more than activation energy)
surface area increases likelihood of collision between molecules

25
Q

Boltzmann distribution curve features

A

area under graph = number of molecules
shaded area = number of molecules able to react
Ea line = activation energy
peak of graph = most likely energy of a molecule

26
Q

how catalyst affects Boltzmann distribution curve

A

moves Ea line to the left (more shaded area)

27
Q

how temperature affects Boltzmann distribution curve

A

lowers peak of the graph then right

28
Q

how dynamic equilibrium exists

A

rate of forward and reverse reactions are equal

concentration of reactants and products do not change

29
Q

Le Chatelier’s principle

A

if a system at equilibrium exposed to a stress (e.g. concentration, pressure, temperature change), system will shift its equilibrium to relieve that stress

30
Q

how concentration affects equilibrium

A

more concentration of reactants, faster rate of forward reaction
equilibrium shifted right
more concentration of products, faster rate of reverse reaction

31
Q

how temperature affects equilibrium

A

when temperature increases
if reaction is exothermic, rate of reverse reaction increases (equilibrium shifts left)
if reaction endothermic, rate of forward reaction increases (equilibrium shifts left)
and vice versa

32
Q

how changing volume/pressure affects equilibrium

A 
lower pressure (more volume) = rate of reaction making more gas molecules increasing
higher pressure (less volume) = rate of reaction making less gas molecules increasing
33
Q

how catalyst affects equilibrium

A

increases rate of forward and reverse reaction
no effect on position of equilibrium
reduces time taken to reach equilibrium

34
Q

how to determine position of equilibrium

A

find out equilibrium constant
if Kc > 1, position of equilibrium closer to the right (products)
if Kc < 1, position of equilibrium closer to the left (reactants)

35
Q

rate of chemical reaction definition

A

how fast a reaction is taking place
how fast a reactant is used up
how fast a product is made

36
Q

what concentration time graphs look like

A 
steepest at start (rate of reaction fastest in the beginning)
gradient lowers the longer the reaction is taking place (reactants concentration lower, less likely to collide and react)
eventually plateaus (no more reaction taking place, one of reactants completely used up)
37
Q

how to measure rate of reaction if gas is produced

A

monitor volume of gas at regular intervals (with gas collection)
monitor loss of mass of reactants with mass balance

38
Q

monitor production of gas with gas collection method

A

place measuring cylinder (held by clamp) submerged under trough of water
measure initial volume in measuring cylinder
place one reactant in conical flask
place bung on top of conical flask
quickly add other reagents to conical flask and replace bung and instantly start time
record volume of gas at regular intervals until reaction is complete (no more gas produced)
plot graph
gradient (found using tangent) = rate of reaction at point of tangent

39
Q

monitor loss of mass of reactants using mass balance method

A

record mass of conical flask and reactants initially and regular intervals
reaction complete when no more gas produced
plot graph

40
Q

catalyst definition

A

substance that changes rate of chemical reaction without undergoing any permanent change itself

41
Q

how catalyst increases rate of reaction

A

provides alternate reaction pathway of lower activation energy

42
Q

catalyst facts

A

not used up in chemical reaction
may react with reactant to form intermediate or provide surface where reaction can take place
catalyst is regenerated at end of reaction

43
Q

homogenous catalyst definition

A

same physical state as reactants

reacts with reactant to form intermediate that breaks down to give products and catalyst

44
Q

heterogenous catalyst definition

A

different physical state from reactants
usually solids in contact with gaseous reactants or reactants in solution
reactant molecules adsorbed onto surface of catalyst, reaction takes place
product molecules leave surface of catalyst via desorption

45
Q

adsorption definition

A

weakly bonding to surface

46
Q

desorption definition

A

molecules leaving surface of another substance

47
Q

autocatalysis definition

A

reaction products acts as catalyst for that reaction

48
Q

importance of catalysts

A

increases rate of industrial chemical reactions by lowering activation energy
also reduces temperature needed for processes and energy requirements (less fossil fuels and electricity used)
cuts cost
more sustainable (higher atom economies, fewer pollutants, cuts CO emissions)

A-Level Chemistry (12 decks)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2024 Bold Learning Solutions. Terms and Conditions