module 5 Flashcards
(106 cards)
1
Q
what is k
A
the rate constant
2
Q
rate of reaction definition
A
the change in concentration of a substance in unit time
units = mol dm3 s-1
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
4
Q
what is initial rate
A
the rate at the start of the reaction where it is fastest
5
Q
how to work out the total order for a reaction
A
add all the individual orders together
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
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
8
Q
effect of temperature on the rate constant
A
increasing the temp increases the value of the rate constant k
9
Q
Arrhenius equation
rearranged version
activation energy version
A
k = Ae^-Ea/RT
lnk = constant - Ea/(RT)
Ea = (lnA-lnK) x RT
10
Q
what does the y intercept show on an arrhenius graph
A
lnA
11
Q
as activation energy gets smaller what happens to rate constant
A
rate constant gets bigger
12
Q
rate determining step definition
A
slowest step of a reaction
13
Q
half life definition
A
time taken for half the reactant to be used up
14
Q
partial pressure definition
A
pressure that the gas would have if it alone occupied the volume occupied by the whole mixture.
15
Q
mole fraction definition
A
The fraction of the total number of moles that each chemical in a reaction is responsible for
16
Q
what is kp
A
the equilibrium constant
only includes gases
unit - atm
17
Q
what affects Kp
A
temperature
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.
19
Q
bronston lowry acid definition
A
a substance that can donate a proton
20
Q
bronston lowry base definition
A
a substance that can accept a proton
21
Q
in reactions with two acids, which one will act as an acid
A
the acid with the bigger ka
22
Q
ionic equation for acid + carbonate = salt + water + carbon dioxide
A
2H+ +CO32- = H20 + CO2
23
Q
ionic equation for acid + alkali = salt + water
A
H+ +OH- = H2O
24
Q
ph =
A
-log(H+)
25
H+
1x10-ph
26
equilibrium for aqueous solutions and pure water
H2O = H+ + OH-
27
what affect does increasing the temperature of water have on pH
- dissociation of water is endothermic
- increasing temp pushes equilibrium to the right, giving a bigger concentration of H+ ions and a lower pH
28
kw =
[H+][OH-]
29
a larger Ka =
stronger acid
30
buffer solution definition
a solution where the ph does not change significantly if small amounts of acid or alkali are added to it
31
how is
- an acidic buffer made
- a basic buffer made
- from a weak acid and a salt of that weak acid
- from a weak base and a salt of that weak base
32
how to make a salt of a
- weak acid
- weak base
- made from reacting the weak acid with a strong base
- made from reacting the weak base with a strong acid
33
how can salt content be added to make a buffer solution
- a salt solution can be added to the acid or some solid salt added
- or by partially neutralising a weak acid with alkali
34
how do buffer solutions work
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
35
explain the buffering action in blood
- give equation
- what are the effects of adding an acid or an alkali
- 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
36
how to construct a pH curve
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
37
choosing an indicator
- 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
38
which indicator is suitable for a weak-acid weak-base titration
No indicator is suitable for a weak acid/weak base titration.
39
enthalpy change of formation definition
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
40
enthalpy of atomisation definition
the enthalpy change when 1 mole of gaseous atoms is formed from the element in its standard state
41
enthalpy of sublimation
enthalpy change for a solid metal turning to gaseous atoms
- numerically the same as the enthalpy of atomisation
42
first ionisation enthalpy definition
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
second ionisation enthalpy definition
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
first electron affinity definition
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
why is first electron affinity exothermic for atoms that form -ve ions
the ion is more stable than the atom
there is an attraction between the nucleus and the electron
46
second electron affinity definition
the enthalpy change when one mole of gaseous 1- ions gains one electron per ion to produce gaseous 2- ions
47
is second electron affinity endo or exothermic
- endothermic because it takes energy to overcome the repulsive force between the negative ion and the electron
48
lattice enthalpy definition
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
enthalpy of hydration definition
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
enthalpy of solution definition
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
what does the strength of enthalpy of lattice formation depend on
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
spontaneous process
proceeds on its own without any external influence
53
entropy definition
measure of disorder within a system
54
which one has a lower entropy
- elements or compounds
- simpler compounds or complex compounds
- pure substances or mixtures
elements
simpler compounds
pure substances
55
when will an increase in entropy occur
- 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
what does an increase in disorder and entropy mean for delta S
- delta s is positive/increases
- positive entropy change
57
entropy equation
- entropy units
s = sum of products - sum of reactants
JK-1mol-1
58
if equilibrium lies to the right hand side what does this mean for the value of k
it must be greater than one
59
gibbs free energy equation
△G = △H - T△S
g = kjmol-1
h = kjmol-1
s = J K-1 mol-1
T = kelvin
60
negative △G =
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
calculating temperature at which a reaction is feasible
T = △H/△S
62
effect of temp on feasibility
- decrease in entropy
- increase in entropy
- △S close to zero
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
applying straight line equation to free gibbs
-△S =
△H =
positive gradient =
△G = O
gradient
c
means deltaS is negative
above this the reaction is spontaneous
64
△solH =
-△LEH + sum of △hydH
65
why are hydration enthalpies exothermic
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
what increases hydration enthalpy
higher charge densities eg smaller ions or ions with larger charges
67
what does it mean if △solH is exothermic or endothermic
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
what does it mean if
△H solution exothermic
△H solution endothermic
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
which conditions will mean a reaction is never feasible
△H = positive
△S = negative
temp = high
70
reducing agent definition
oxidising agent definition
electron donors, it is itself oxidised
electron acceptors, it is itself reduced
71
thiosulfate redox reaction equation
starch indicator
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
manganate redox titration equation
MnO4- + 8H+ + 5Fe2+ → Mn2+ + 4H2O + 5Fe3+
manganate goes from purple to colourless
73
why use a high resistance voltmeter
- to stop the current flowing in the circuit
74
which acid to use for manganate titration
dilute sulfuric acid
75
what is a salt bridge made of
what does the salt bridge do
why is a wire not used
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
hydrogen electrode equilibrium
H2 (g) = 2H+ + 2e- (reversible)
77
components of a standard hydrogen electrode
1. hydrogen gas at pressure of 100Kpa
2. solution containing the hydrogen ion at 1 moldm-3
3. temp at 298K
78
standard conditions when measuring standard electrode potentials
all ion solutions at 1moldm-3
temp 298k
gases at 100Kpa pressure
no current flowing
79
more negative half cell...
more positive half cell...
oxidise (go backwards) and are reducing agents
reduce (go fowards) and are oxidising agents
80
Ecell =
Ereduced - Eoxidised
81
what is e.m.f.
a measure of how far from equilibrium the cell reaction lies
- a more positive emf the more likely a reaction will occur
82
effect of concentration on cell emf
increasing conc of reactants increases EMF and vice versa
83
fuel cell definition
uses the energy from the reaction of a fuel with oxygen to create a voltage
84
why do fuel cells maintain a constant voltage
they are continuously fed with fresh O2 and H2 so maintain a constant concentration of reactants
85
advantages and limitations of fuel cells
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
why is Zn not a transition metal
why is Sc not a transition metal
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
typical properties of transition elements
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
complex definition
central metal ion surrounded by ligands
89
ligand definition
an atom, ion or molecule which can donate a lone electron pair
90
co-ordinate bonding definition
when the shared pair of electrons in the covalent bond comes from only one of the bonding atoms
91
co-ordination number definition
the number of co-ordinate bonds formed to a central metal ion
92
uni vs bi vs multidentate ligands
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
two types of isomerism shown by complexes
cis-trans and optical isomerism
- optical isomerism shown by complexes with 3 bidentate ligands
94
cis-platin
- what is it used for
- how does it do this
- give equation
- 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
how does Fe(II) enable oxygen to be transported in the blood
- 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
why is CO toxic to humans
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
Which apparatus could be used to determine the effect of the concentration of CuSO4(aq) on the
rate of reaction?
colorimeter
98
why would an excess amount of solution ensure that that solution would be first order
concentration would remain constant
99
Explain how a student could determine the activation energy, Ea, for the reaction
graphically using values of k and T.
1. plot graph using lnk and 1/T
2. measure gradient
3. Ea = -R x gradient
100
Explain why a small proportion of molecules in water have a relative molecular mass of 20.
different O/H isotopes are present
101
Explain the significance of the expression: Kp > 1
equilibrium position far to the right
102
How would the lattice enthalpies of magnesium chloride and calcium chloride differ?
lattice enthalpy of mgcl2 is more exothermic than cacl2
because the magnesium ion is smaller
therefore the attraction between mg2+ and cl- is greater
103
is the bond enthalpy of the C-H bond endo or exothermic
endothermic as it is bond breaking
104
When the student prepares the buffer solution, the volume of solution increases slightly.
Suggest whether the pH of the buffer solution would be the same, greater than, or less
pH is the same
ratio of [HA]/[A–] is the same
105
percentage dissociation equation
H+/acid
then times by 100
106
Predict how the enthalpy changes of hydration of F– and Cl – would differ
enthalpy of hydration of F- is more exothermic as F- has a smaller size
smaller sized ion has a greater attraction to H20
