module 5 Flashcards

1
Q

what is k

A

the rate constant

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2
Q

rate of reaction definition

A

the change in concentration of a substance in unit time
units = mol dm3 s-1

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3
Q

on a graph of concentration of reactant vs time what does the gradient of the curve tell you

A

the rate of reaction

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4
Q

what is initial rate

A

the rate at the start of the reaction where it is fastest

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5
Q

how to work out the total order for a reaction

A

add all the individual orders together

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6
Q

how do orders affect the rate of reaction

A

zero order = no effect on the rate
first order = the rate of reaction is directly proportional to the concentration of A
second order = the rate of reaction is proportional to the
concentration of A squared

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7
Q

how do half lives relate to order

A
  • if the half lives are constant then the order is 1st order
  • if they rapidly increase then the order is 2nd order
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8
Q

effect of temperature on the rate constant

A

increasing the temp increases the value of the rate constant k

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9
Q

Arrhenius equation
rearranged version
activation energy version

A

k = Ae^-Ea/RT
lnk = constant - Ea/(RT)
Ea = (lnA-lnK) x RT

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10
Q

what does the y intercept show on an arrhenius graph

A

lnA

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11
Q

as activation energy gets smaller what happens to rate constant

A

rate constant gets bigger

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12
Q

rate determining step definition

A

slowest step of a reaction

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13
Q

half life definition

A

time taken for half the reactant to be used up

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14
Q

partial pressure definition

A

pressure that the gas would have if it alone occupied the volume occupied by the whole mixture.

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15
Q

mole fraction definition

A

The fraction of the total number of moles that each chemical in a reaction is responsible for

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16
Q

what is kp

A

the equilibrium constant
only includes gases
unit - atm

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17
Q

what affects Kp

A

temperature

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18
Q

Effect of temperature on position of equilibrium and Kc

A

In this equilibrium which is exothermic in the forward direction

If temperature is increased the reaction will shift to oppose the change and move in the backwards endothermic direction. The position of equilibrium shifts left. The value of Kc gets smaller as there are fewer products.

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19
Q

bronston lowry acid definition

A

a substance that can donate a proton

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20
Q

bronston lowry base definition

A

a substance that can accept a proton

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21
Q

in reactions with two acids, which one will act as an acid

A

the acid with the bigger ka

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22
Q

ionic equation for acid + carbonate = salt + water + carbon dioxide

A

2H+ +CO32- = H20 + CO2

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23
Q

ionic equation for acid + alkali = salt + water

A

H+ +OH- = H2O

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24
Q

ph =

A

-log(H+)

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25
Q

H+

A

1x10-ph

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26
Q

equilibrium for aqueous solutions and pure water

A

H2O = H+ + OH-

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27
Q

what affect does increasing the temperature of water have on pH

A
  • dissociation of water is endothermic
  • increasing temp pushes equilibrium to the right, giving a bigger concentration of H+ ions and a lower pH
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28
Q

kw =

A

[H+][OH-]

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29
Q

a larger Ka =

A

stronger acid

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30
Q

buffer solution definition

A

a solution where the ph does not change significantly if small amounts of acid or alkali are added to it

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31
Q

how is
- an acidic buffer made
- a basic buffer made

A
  • from a weak acid and a salt of that weak acid
  • from a weak base and a salt of that weak base
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32
Q

how to make a salt of a
- weak acid
- weak base

A
  • made from reacting the weak acid with a strong base
  • made from reacting the weak base with a strong acid
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33
Q

how can salt content be added to make a buffer solution

A
  • a salt solution can be added to the acid or some solid salt added
  • or by partially neutralising a weak acid with alkali
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34
Q

how do buffer solutions work

A

If small amounts of acid is added to the buffer then the equilibrium will shift in the direction to oppose this removing the H+ ions added

If small amounts of alkali is added to the buffer. The OH ions will react with H+ ions to form water. The Equilibrium will then shift to the right to produce more H+ ions. Overall the concentration of H+ ions and pH remains constant

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35
Q

explain the buffering action in blood
- give equation
- what are the effects of adding an acid or an alkali

A
  • a carbonic acid-hydrogencarbonate equilibrium acts as a buffer in the control of blood pH
  • the H2CO3/HCO3- buffer is present in blood plasma, maintaining a ph between 7.35 and 7.45
  • H2CO3 (aq) ⇌ H+(aq) + HCO3–(aq) (reversible)
  • Adding alkali reacts with H+ with the equation
    H+ + OH- → H2O
    so the above equilibrium would shift right forming new H+ and more HCO3-
  • Adding acid shifts the above equilibrium left.
    The reaction is
    H+ + HCO3- → H2CO3
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36
Q

how to construct a pH curve

A
  1. transfer 25cm3 of acid to a conical flask with a volumetric pipette
  2. measure initial pH of the acid with a pH meter
  3. add alkali in small amounts (2cm3) noting the volume added
  4. stir mixture to equalise the pH
  5. measure and record the pH to 1 dp
  6. repeat steps 3-5 but when approaching endpoint add in smaller volume of alkali
  7. add until alkali in excess

calibrate the pH meter first by measuring known pH of a buffer solution

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37
Q

choosing an indicator

A
  • choose an indicator whose end point coincides with the equivalence point for the titration
  • phenolphthalein: strong bases, not weak acids
    colour change: colourless acid to pink alkali
  • methyl orange: strong acids, not weak acids
    colour change : red acid to yellow alkali
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38
Q

which indicator is suitable for a weak-acid weak-base titration

A

No indicator is suitable for a weak acid/weak base titration.

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39
Q

enthalpy change of formation definition

A

the energy transferred when 1 mole of the compound is formed from its elements under standard conditions (298K and 100kpa), all reactants and products being in their standard states

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40
Q

enthalpy of atomisation definition

A

the enthalpy change when 1 mole of gaseous atoms is formed from the element in its standard state

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41
Q

enthalpy of sublimation

A

enthalpy change for a solid metal turning to gaseous atoms
- numerically the same as the enthalpy of atomisation

42
Q

first ionisation enthalpy definition

A

the enthalpy change required to remove 1 mole of electrons from 1 mole of gaseous atoms to form 1 mole of gaseous ions with a +1 charge

43
Q

second ionisation enthalpy definition

A

the enthalpy change to remove 1 mole of electrons from one mole of gaseous 1+ ions to produce one mole of gaseous 2+ ions.

44
Q

first electron affinity definition

A

enthalpy change when one mole of gaseous atoms gain 1 mole of gaseous electrons to form 1 mole of gaseous ions with a 1- charge

45
Q

why is first electron affinity exothermic for atoms that form -ve ions

A

the ion is more stable than the atom

there is an attraction between the nucleus and the electron

46
Q

second electron affinity definition

A

the enthalpy change when one mole of gaseous 1- ions gains one electron per ion to produce gaseous 2- ions

47
Q

is second electron affinity endo or exothermic

A
  • endothermic because it takes energy to overcome the repulsive force between the negative ion and the electron
48
Q

lattice enthalpy definition

A

the standard enthalpy change when 1 mole of an ionic crystal lattice is formed from its constituent ions in gaseous form.
- can be used as a measure of ionic bond strength

49
Q

enthalpy of hydration definition

A

Enthalpy change when one mole of gaseous ions become aqueous ions
- this is always exothermic as bonds are made between the ions and water molecules

50
Q

enthalpy of solution definition

A

the standard enthalpy change when one mole of an ionic solid dissolves in
a large enough amount of water to ensure that the dissolved ions are well separated and do not interact with one another.

51
Q

what does the strength of enthalpy of lattice formation depend on

A
  1. the sizes of the ions: the larger the ions, the less negative the enthalpies of lattice formation. As the ions get larger the charges become further apart and so have a weaker attractive force between them
  2. the charges on the ion: the bigger the charge of the ion, the greater the attraction between the ions so the stronger the lattice enthalpy (more -ve value)
52
Q

spontaneous process

A

proceeds on its own without any external influence

53
Q

entropy definition

A

measure of disorder within a system

54
Q

which one has a lower entropy
- elements or compounds
- simpler compounds or complex compounds
- pure substances or mixtures

A

elements
simpler compounds
pure substances

55
Q

when will an increase in entropy occur

A
  • there is a change of state from solid or liquid to gas
  • when a solid increases in temperature
  • there is an increase in the number of molecules between product and reactants
56
Q

what does an increase in disorder and entropy mean for delta S

A
  • delta s is positive/increases
  • positive entropy change
57
Q

entropy equation
- entropy units

A

s = sum of products - sum of reactants
JK-1mol-1

58
Q

if equilibrium lies to the right hand side what does this mean for the value of k

A

it must be greater than one

59
Q

gibbs free energy equation

A

△G = △H - T△S
g = kjmol-1
h = kjmol-1
s = J K-1 mol-1
T = kelvin

60
Q

negative △G =

A

reaction is feasible

if the reaction does not occur and delta g is negative, this may be due to a high activation energy, meaning the reaction occurs too slowly

61
Q

calculating temperature at which a reaction is feasible

A

T = △H/△S

62
Q

effect of temp on feasibility
- decrease in entropy
- increase in entropy
- △S close to zero

A
  1. decrease in entropy means △S is negative. increasing temp will make it more less likely that △G is negative and less likely for the reaction to occur
  2. increase in entropy means △S is negative. increasing temp means its more likely △G is negative and more likely that the reaction occurs
  3. temp will not have a large effect on the feasibility as -T△S will be small and △G wont change much
63
Q

applying straight line equation to free gibbs
-△S =
△H =
positive gradient =
△G = O

A

gradient
c
means deltaS is negative
above this the reaction is spontaneous

64
Q

△solH =

A

-△LEH + sum of △hydH

65
Q

why are hydration enthalpies exothermic

A

as energy is given out as water molecules bond to the metal ones

the negative ions are attracted to the partially positive hydrogens on the polar water molecules and the positive ions are attracted to the partially negative oxygen on the polar water molecules

66
Q

what increases hydration enthalpy

A

higher charge densities eg smaller ions or ions with larger charges

67
Q

what does it mean if △solH is exothermic or endothermic

A

exothermic = soluble.
endothermic = insoluble. this is because the lattice enthalpy is much larger than the hydration enthalpy and it is not energetically favourable to break up the lattice

68
Q

△H solution exothermic =
△H solution endothermic

A

salt will always dissolve at all temps
salt may dissolve depending on whether the -T△S value is more negative than △H is positive

-increasing temp will make it more likely that △G will become negative, making the reaction more feasible

69
Q

which conditions will mean a reaction is never feasible

A

△H = positive
△S = negative
temp = high

70
Q

reducing agent definition
oxidising agent definition

A

electron donors, it is itself oxidised
electron acceptors, it is itself reduced

71
Q

thiosulfate redox reaction equation

starch indicator

A

2S2O3 2- + I2 → 2I- + S4O6 2-
iodine goes from a yellow brown to colourless

a starch indicator is added near the end point when the iodine fades a pale yellow to emphasise it
with starch added, the colour change is from blue/black to colourless

72
Q

manganate redox titration equation

A

MnO4- + 8H+ + 5Fe2+ → Mn2+ + 4H2O + 5Fe3+

manganate goes from purple to colourless

73
Q

why use a high resistance voltmeter

A
  • to stop the current flowing in the circuit
74
Q

which acid to use for manganate titration

A

dilute sulfuric acid

75
Q

what is a salt bridge made of
what does the salt bridge do
why is a wire not used

A

made from a piece of filter paper soaked in a salt solution, usually potassium nitrate
- the salt should be unreactive with the electrodes and electrode solutions

used to connect up the circuit - the free moving ions conduct the charge

the metal wire would set up its own electrode system with the solutions

76
Q

hydrogen electrode equilibrium

A

H2 (g) = 2H+ + 2e- (reversible)

77
Q

components of a standard hydrogen electrode

A
  1. hydrogen gas at pressure of 100Kpa
  2. solution containing the hydrogen ion at 1 moldm-3
  3. temp at 298K
78
Q

standard conditions when measuring standard electrode potentials

A

all ion solutions at 1moldm-3
temp 298k
gases at 100Kpa pressure
no current flowing

79
Q

more negative half cell…
more positive half cell…

A

oxidise (go backwards) and are reducing agents
reduce (go fowards) and are oxidising agents

80
Q

Ecell =

A

Ereduced - Eoxidised

81
Q

what is e.m.f.

A

a measure of how far from equilibrium the cell reaction lies
- a more positive emf the more likely a reaction will occur

82
Q

effect of concentration on cell emf

A

increasing conc of reactants increases EMF and vice versa

83
Q

fuel cell definition

A

uses the energy from the reaction of a fuel with oxygen to create a voltage

84
Q

why do fuel cells maintain a constant voltage

A

they are continuously fed with fresh O2 and H2 so maintain a constant concentration of reactants

85
Q

advantages and limitations of fuel cells

A

advantages:
1. less pollution and less co2
2. greater efficiency
limitations
1. storing and transporting hydrogen in terms of safety
2. limited lifetime
3. use of toxic chemicals in their production

86
Q

why is Zn not a transition metal

why is Sc not a transition metal

A

Zn can only form a 2+ ion. in this ion, the Zn2+ has a complete d orbital and so does not meet the criteria of having an incomplete d orbital

Sc can only form a 3+ ion. In this ion, the Sc3+ has an empty d orbital and so does not meet the criteria of having an incomplete d orbital

87
Q

typical properties of transition elements

A
  1. the existence of more than one oxidation state for each element in its compounds
  2. the formation of coloured ions
  3. the catalytic behaviour of the elements and their compounds and their importance in the manufacture of chemicals by industry
    - eg Iron is used as a catalyst in the haber process to produce ammonia
88
Q

complex definition

A

central metal ion surrounded by ligands

89
Q

ligand definition

A

an atom, ion or molecule which can donate a lone electron pair

90
Q

co-ordinate bonding definition

A

when the shared pair of electrons in the covalent bond comes from only one of the bonding atoms

91
Q

co-ordination number definition

A

the number of co-ordinate bonds formed to a central metal ion

92
Q

uni vs bi vs multidentate ligands

A

unidentate: can only form one coordinate bond per ligand eg H20
bidentate: have two atoms with lone pairs and can form two coordinate bonds per ligand
multidentate: more than two

93
Q

two types of isomerism shown by complexes

A

cis-trans and optical isomerism
- optical isomerism shown by complexes with 3 bidentate ligands

94
Q

cis-platin
- what is it used for
- how does it do this
- give equation

A
  • used as an anticancer drug
  • binds to DNA of cancer cells and stops replication
  • also binds to healthy DNA, stopping the replication of healthy cells which may lead to side effects such as hair loss
  • Pt(NH3)2Cl2 + H20 = [Pt(NH3)2Cl(H2O)]+ + Cl-
95
Q

how does Fe(II) enable oxygen to be transported in the blood

A
  • found in haemoglobin
  • haem in an Iron(II) complex with a multidentate ligand
  • O2 bonds to Fe2+ ions in the haemoglobin and when required O2 is released
96
Q

why is CO toxic to humans

A

can form a strong coordinate bond with haemoglobin
- this is a stronger bond than that made with oxygen and so it prevents the oxygen attaching to the haemoglobin
- with CO, the stability constant is greater than with the complex in O2

97
Q

Which apparatus could be used to determine the effect of the concentration of CuSO4(aq) on the
rate of reaction?

A

colorimeter

98
Q

why would an excess amount of solution ensure that that solution would be first order

A

concentration would remain constant

99
Q

Explain how a student could determine the activation energy, Ea, for the reaction
graphically using values of k and T.

A
  1. plot graph using lnk and 1/T
  2. measure gradient
  3. Ea = -R x gradient
100
Q

Explain why a small proportion of molecules in water have a relative molecular mass of 20.

A

different O/H isotopes are present

101
Q

Explain the significance of the expression: Kp > 1

A

equilibrium position far to the right