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chemistry > year 13 physical > Flashcards

year 13 physical Flashcards

(106 cards)

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1
Q

total pressure

A

sum of all partial pressures (in pascals/kPa)

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2
Q

mole fraction

A

no of moles of one gas ÷ total no of moles

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3
Q

partial pressure

A

mole fraction x partial pressure

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4
Q

key point (ratios when doing partial pressure)

A

ratios can be compared when on the same side of the equation but not to the other side

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5
Q

kp
(+ units)

A

products ÷ reactants
to the power of moles
curly brackets
little pp
units = kpa (same method as Kc)

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6
Q

what effects Kp

A

temperature
not pressure!!!

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7
Q

what does movement of equilibrium position do to Kp?

A

shift to right = larger
shift to left = smaller

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8
Q

which direction does equilibrium move when temp increases?

A

endothermic direction
(counteract the change)

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9
Q

which direction does equilibrium move when temp decreases?

A

exothermic direction
(counteract the change)

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10
Q

effect of pressure on Kp?

A

none

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11
Q

effect of catalyst on Kp?

A

none

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12
Q

bond dissociation enthalpy

A

1 mole
covalent bonds broken
gaseous

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13
Q

enthalpy of lattice formation

A

1 mole
solid ionic compound
formed from ions in gas state

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14
Q

enthalpy of lattice dissociation

A

1 mole
solid ionic compound
dissociates into ions in gas state

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15
Q

why is 1st electron affinity exothermic

A

attraction between nucleus and electrons

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16
Q

why is 2nd electron affinity endothermic

A

electron repulsion between nucleus and outer shell
both already negatively charged

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17
Q

effect of ionic charge on lattice enthalpy

A

larger causes higher enthalpy change

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18
Q

effect of size on lattice enthalpy

A

smaller causes higher enthalpy change

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19
Q

difference between theoretical/experimental value

A

perfect ionic model/covalent character

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20
Q

1st/2nd electron affinity

A

enthalpy change
1 mole
gaseous atoms
gain one election

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21
Q

enthalpy change of solution

A

1 mole of ions
minimum amount of solvent
dissolved
to ensure no further enthalpy change

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22
Q

resultant of enthalpy change of solution

A

very saturated solvent
ions become hydrated

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23
Q

enthalpy of solution =

A

enthalpy of dissosiation + enthalpy of hydration
LiCl (s) -> Li+(g) + Cl-(g) -> Li+(aq) + Cl-(aq)

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24
Q

entropy

A

measure of disorder (more disorder = higher entropy)

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25
two factors which effect entropy
no. of moles (more moles = more disorder) state of molecules
26
calculation entropy
S(products)- S(reactants)
27
positive entropy value (delta S)
entropically favourable feasible
28
Gibbs free energy equation units
G-TS J/mol
29
why might a reaction not occur even if Gibbs works?
activation energy too high rate too slow
30
value of k relation with rate
larger = faster rate
31
initial rate (calculate from graph)
gradient of tangent at 0 mins
32
finding rate equation
repeat one experiment several times but changing the conc of one reactant at a time find initial rate using graph (measure product formed/change in pH/colour change/reactant used) place in table and calculate orders
33
horizontal line on rate time graph
from straight line down on conc time graph zero order
34
where does a rate time graph come from
derivative of conc time graph
35
positive gradient straight line time graph
from curving down line on conc time graph first order
36
positive gradient curved line time graph
from steep curving down line on concrete time graph second order
37
rate equation and rate determining step
reactants in rate determining step are always in the rate equation
38
intermediate in rate determining step (with respect to rate equation)
reactants which made intermediate must be in rate equation
39
x and y axis of Arrhenius plots
x = 1/T y = ln K
40
m and c in Arrhenius plot
m = activation energy c = Ln A
41
y = mx + c in Arrhenius
Ln k = Ln A - Ea/RT
42
explain why increasing temp is better for increasing reaction than increasing concentration of one species
reaction occurs when all molecules have energy above activation energy doubling temp affects this of all molecules whereas doubling conc only affects one species
43
change to rate equation when one species largely in excess
species removed from rate equation conc virtually constant
44
units of k
(mol dm-3)n s-1
45
what should be done to samples of reactions when finding orders
stop the reaction quench dilution/cooling/species to react with a reactant/removing catalyst
46
explaining order from graph
describe gradient link to order e.g. used at constant rate
47
how to find Ea from Arrhenius graph
gradient = -Ea/R multiply by 8.31, divide by 1000
48
units for A (Arrhenius constant)
s-1
49
bronsted lowry acids
proton donors H+ ions released when aqueous forms H3O+ in water HA + H2O <-> H3O+ + A- (equill) water behaves as a base
50
bronsted lowry bases
proton acceptors react with water to form OH- B + H2O <-> BH+ + OH - (equill)
51
ethanoic acid equilibrium equation
CH3COOH <-> CH3OO- + H+ lies heavily to left backwards reaction favoured
52
hydrochloric acid equillibrium equation
HCl <-> H+ + Cl- forwards reaction favoured lies heavily to right
53
acid base reaction equillibrium
HA + B <-> BH+ + A- HA = acid B = base BH+ = salt
54
Kw
ionic product of water
55
constant of water
1 x 10 (-14) mol2dm-6
56
pH of strong acid
conc of H+ = conc of strong acid
57
monoprotic strong acid
[H+] = [acid] pH = -log 10 [acid]
58
diprotic strong acid
2[H+] = [acid] pH = -log [2acid]
59
pH of strong base
dissociate fully Kw = [OH-][H+] [OH-] = [base] find [H=] pH = -log10[H+]
60
acid dissociation constant
HA <-> H+ + A- (lies to left) Ka = {[H+][A-}}/[HA] or Ka = [H+}2/[HA] moldm-3
61
pH of weak acids
Ka = [H+]2/[HA] [HA] = conc of acid
62
assumption with Ka
[acid] (start) = [acid] (equillibrium) dissociation of acid is greater than dissociation of water (all H+ ions from acid)
63
calculate Ka/concenctration of a weak acid
find [H+] (from pH) use Ka equation
64
pKa
-log10 Ka
65
equivalence point
fully neutralised acid / base
66
half neutralisation point
half way between 0 and equivalence point [HA] = [A-] cancels out in Ka Ka = [H+] pKa = pH
67
diprotic curves
two equivalence points H+ ions released separately
68
acidic buffer
keeps pH below 7 from weak acid and its salt CH3COOH <=> CH3COO- + H+ (acid, lies to left) CH3COO-Na <-> CH3COO- + Na (salt, lies to right)
69
add H+ to acidic buffer
react with negatively charged ions (from salt) forms undissociated acid (equillibrium lies to left)
70
add OH - to acidic buffer
react with H+ ions equillibrium counteracts change
71
basic buffer
weak base and its salt pH above 7 NH3 + H2O <-> NH3+ + OH- (base, lies to left) NH4+Cl- <-> NH4+ + Cl- (salt, lies to right)
72
add H+ to basic buffer
reacts with OH- makes water equillibrium re-established
73
add OH - to basic buffer
reacts with positive ions in solution (high conc from salt) shifts base equillibrium to the left
74
calculating pH of a buffer
find Ka expression use equillibrium concentrations not initial [salt] = A- find concs input to find [H+] then pH
75
why is water with pH of near 7 not acidic
[H]+ = [OH]-
76
why do buffer solutions have a constant pH even when diluted
ratio [HX]/[X-] remains constant
77
difference in enthalpy of hydration
size to charge ratio weaker attraction between polar water molecule and ion
78
difference in enthalpy of lattice formation reasons
smaller ion - greater size:charge ratio ratio of metal:non metal
79
high pKa
weaker acid
80
postive electrode
anode
81
negative electrode
cathode
82
positive ion
cation
83
negative ion
anion
84
molten electrolysis
only one anion and cation
85
which cations are more easily reduced
more positive cations
86
which anions are more easily oxidised
more negative anions
87
standard conditions of reference half cell (S.H.E.)
1 mole H+ ions hydrogen gas platinum metal 298 K 1 atm/ 100kPa
88
requirements of voltmeter
high resistance prevent current flowing
89
use of salt bridge
allows ions to pass from solutions completes circuit
90
why can't metal be used as salt bridge
create its own equillibrium prevent ions flowing
91
suitable salt bridges
filter paper soaked in ions agar jelly soaked in ions KNO3 used commonly generally inert highly soluble ions
92
where is p.d. measured by electrochemical cells
between metal and solutions
93
change to charge when equillibrium of electrode shifts to left
more electrons build on electrode negative charge
94
change to charge when equillibrium of electrode shifts to right
electrons used up positive charge
95
large positive value of electrode
indicates strong oxidising agent
96
large negative value of electrode
indicates strong reducing agent
97
positive lithium cell equation
Li+ + CoO2 + e- -> Li+[CoO2]-
98
negative lithium cell equation
Li -> Li+ + e-
99
negatives to lithium cells
very short supply of lithium lithium fires begin to loose charge more quickly over time
100
use of lithium cell
rechargeable batteries
101
what is a fuel cell
provides electrical current does not need to be recharged
102
benefits of hydrogen alkaline fuel cell
water = only product room temp and no combustion no energy loss
103
cons to hydrogen alkaline fuel cell
hydrogen is highly flammable finite and non renewable resources need to carry fuel in large containers not carbon neutral - CO2 used to make H2
104
overall equation hydrogen alkaline fuel cell
2h2 + O2 -> 2H2O
105
why not aqueous lithium hydrogen fuel cells
lithium reacts with water electrode potential of lithium more negative than water
106
how does a hydrogen fuel cell produce a current
electrons from hydrogen

chemistry (9 decks)

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