Module 5 Flashcards
(105 cards)
1
Q
Define transition elements
A
D-block elements which form an ion with with an incomplete d-sub-shell
2
Q
Explain the catalytic behaviour of transition elements
A
They can easily change between oxidation states, so give an alternative route with a lower activation energy
3
Q
Define ligand
A
A molecule or ion which donates a pair of electrons to form a coordinate bond
4
Q
Define coordination number
A
The number of coordinate bonds attached to the central metal ion
5
Q
What is the shape of a four-fold coordination complex ion with Pt?
A
Square planer
6
Q
What is the shape of a tetrachloro complex ion?
A
Tetrahedral shape
7
Q
Define a mono-dentate ligand and give some examples
A
A ligand which forms one coordinate bond with the central metal ion, e.g. Water, ammonia
8
Q
Define a bi-dentate ligand and give an example
A
A ligand which forms two coordinate bonds with the central metal atom, e.g. 1,2,diaminoethene (en)
9
Q
Describe the sort of isomerism found in platin
A
Cis-Trans stereoisomerism
10
Q
Describe the use of platin as an anti-cancer drug
A
Cis-Platin forms a platinum complex inside the cancerous cell and binds to the DNA and stops the DNA from replicating
11
Q
Define oxidising agent
A
The species which take electrons from the species being oxidised
12
Q
Define reducing agent
A
The species which gives electrons to the species which is being reduced
13
Q
Define standard electrode potential
A
The e.m.f. of a half cell compared to standard hydrogen half cell under standard conditions
14
Q
Why do some half cells use a platinum electrode instead of an electrode made of the element which the ions are?
A
For when the two ions are not in a solid state, an inert metal is used with the transfer of electrons
15
Q
Describe how to calculate the standard cell potential of two half cells
A
The difference between the two half cells gives the standard cell potential
16
Q
Describe how the feasibility of a reaction can be predicted in terms of kinetics and give the limitations
A
If the standard potential of the cell is positive, then the reaction is kinetically feasible in standard conditions, but different conditions could effect this, e.g. Concentration
17
Q
Explain how a fuel cell works
A
A fuel cell uses the energy from the reaction of a fuel cell with oxygen to create a voltage and the changes that take place at each electrode
18
Q
Discuss the benefits and risks of using a Li-based cells
A
Benefits: High amounts of energy stored easily; rechargeable Risks: Toxic; Fire hazard
19
Q
Define lattice enthalpy
A
Formation of 1 mol of ionic lattice from gaseous ions
20
Q
What does the value for the lattice enthalpy tell us about the ionic lattice?
A
A measure of the strength of ionic bonding in a giant ionic lattice
21
Q
Define enthalpy change of solution
A
The enthalpy change of dissolving one mol of a solute
22
Q
Define enthalpy change of hydration
A
The enthalpy change of dissolving one mol of gaseous ions in water
23
Q
Describe the effect of ionic charge and ionic radius on the exothermic value of lattice enthalpy
A
Increasing the ionic charge increases the attraction between the cations, increasing the melting point, making the value more exothermic Decreasing the ionic radius increases the attraction between the cations, increasing the melting point, making the value more exothermic
24
Q
Describe the effect of ionic charge and ionic radius on the exothermic value of the enthalpy change of hydration
A
Increasing the ionic charge increases the amount of attraction with water molecules, making the value more negative Decreasing the ionic radius increases the amount of attraction with water molecules, making the value more negative
25
Define entropy
A measure of the dispersal of energy in a system which is greater the more disordered a system.
26
Explain the difference in magnitude of entropy of a system for solids, liquids and gases
The solid has the lowest entropy and gas has the highest entropy, for entropy is the measure of dispersal of energy, and the more disordered systems have a higher dispersal of energy
27
Describe what the thermodynamic feasibility of a reaction is depended on
The entropy change and temperature of a system, and the enthalpy change of the system
28
Explain why a reaction might be thermodynamically feasible, but still not happen
A reaction could give a negative value for Gibbs free energy, but not be kinetically feasible, so still won't happen
29
Define rate of reaction
The change of concentration of a reaction over time
30
Define order
The power to which the concentration of a reaction is raised in a rate equation
31
Define overall order
The sum of the individual orders of reactants in the rate equation
32
Define rate constant
The constant which links rate of reaction with the concentrations the reactants raised to the power of their orders in the rate equation
33
Define half-life
The time taken for the concentration of a reaction to decrease to half
34
Define rate-determining step
The slowest step of a reaction mechanism of a multi-step reaction
35
Describe how the order of a reactant can be determined
The concentration of the reactant over time can be recorded and the concentration time graph can be plotted and a linear graph is zero order and exponential graph is first order
36
Describe how the rate of reaction can be determined from a concentration time graph
The gradient = rate of reaction. If graph is non-linear, a tangent can be drawn and gradient of this can be determined
37
Describe how the half-life a first order reaction can be determined
As the half life is always constant, the time taken for the concentration to halve is constant
38
Describe how the rate constant can be determine for a first order reaction graphically
As k = ln(2)/half-life, determine the half life then calculate using the equation
39
Describe how the order of a reaction can be deduced from a rate-concentration graph
Flat = zero order Linear = first order Non-linear = second order
40
Describe how colorimetry can be used to investigate reaction rates
A certain colour wavelength of light is emitted, and the solution absorbs an amount, and this is measured. As the reaction progresses, the change in absorbence, hence colour, is measured, so the rate of reaction can be determined
41
Explain the effect of temperature change on the rate of a reaction and hence the rate constant
* Increasing the temperature shifts the Boltzmann distribution to the right, giving more particles energy above the Ea, allowing more particles to react.
* Increasing the temperature increases the particles speed, so collisions become more frequent
42
Describe how the activation energy of reaction can be determined
* A graph of ln k against 1/T is plotted
* Gradient can be determined by plotting a line of best fit through the points, and working out the rate of change
* Gradient = -Ea/R, so -gradient\*8.314 = Ea.
43
Define mole fraction of a gas
It's proportion by volume of all of the gases present
44
Define partial pressure
The contribution which the gas makes to the overall pressure
45
Describe the effect on equilibrium constants of changing temperature for an exothermic reaction
* The equilibrium constant decreases with increasing temperature
* Raising the temperature decreases the equilibrium yield of products
46
Describe the effect on equilibrium constants of changing temperature for an endothermic reaction
* The equilibrium constant increases with increasing temperature
* Raising the temperature increases the equilibrium yield of products
47
Describe the effect on equilibrium constants of changing pressure
The equilibrium constant remains at the same value
48
Describe the effect on equilibrium constants of changing concentration
The equilibrium constant remains unchanged
49
Describe why a catalyst does not affect equilibrium constants
* Catalysts affect the rate of a chemical reaction, but not the position of equilibrium
* Catalysts speed up both the forward and reverse reactions in the equilibrium by the same factor
* So equilibrium is reached quicker, but the position is not changed
50
Define Brønsted-Lowry acid
A species that donates a proton
51
Define Brønsted-Lowry base
A species which accepts a proton
52
Define conjugate acid-base pair
Two species that can be interconcerted by transfer of a proton, e.g. H2O and H3O+ (water and hydronium ion)
53
What is a monobasic acid?
An acid which dissociates once, releasing 1 H+ (Proton)
54
What is a dibasic acid?
An acid which dissociates twice, releasing 2 H+ (Protons)
55
What is a tribasic acid?
An acid which dissociates thrice, releasing 2 H+ (Protons)
56
Describe the role of H+ in reactions with acids
When showing ionic equations, as the rest of the acid is a spectator ion, it can cancel, so just H+ can be written.
57
What is the general formula for Ka?
Ka=[H+(aq)]\*[A-(aq)]/[HA(aq)]
58
Describe the relationship between the strength of an acid and its Ka value
The greater the Ka value, the stronger the acid
59
Describe the relationship between the strength of an acid and its pKa value?
The lower the pKa value, the stronger the acid
60
Describe the relationship between Ka and pKa?
pKa = -log(Ka);
Ka = 10-pKa
61
What are the units for Ka?
Mol dm-3
62
Define pH
pH = -log([H+])
63
Describe how pH can be used to determine the concentration of protons?
[H+] = 10-pH
64
Write an equation for Kw and explain how this can be used to work out the pH of base
Kw = [H+(aq)]\*[OH-(aq)]
[OH-(aq)] is worked out and -log(Kw/[OH-(aq)]) is used, as Kw/[OH-(aq)] = [H+(aq)]
65
State and describe the approximations in calculations using weak acids
Assumption 1:
* The dissociation of water is negligible
* So [H+(aq)] = [A-(aq)]
Assumption 2:
* The concentration of acid is much greater than the concentration of H+
* [HA]eqm = [HA]start - [H+]start ≈ [HA]eqm = [HA]start
66
Describe when the assumption which assumes the concentration of acid is much greater than the H+ concentration at equilibrium breaks down
* Stronger weak acids, where Ka \> 10-2 mol dm-3
* Very dilute solutions
67
Define a buffer solution
A system that minimises pH changes on addition of small amounts of an acid or a base
68
Describe the formation of a buffer solution from a weak acid and a salt of a weak acid
HA(aq) ⇌ H+(aq) + A-(aq)
MA + (aq) → A-(aq) + M+ aq)
When adding alkali, the weak acid removes added alkali
When adding acid, the conjugate base removes added acid
69
Explain how a buffer solution works when a small amount of acid is added
1. H+ concentration increases
2. These ions react with the conjugate base
3. The equilibrium poisition shifts to the left, removing most of the H+
70
Explain how a buffer solution works when a small amount of alkali is added
1. OH- concentration increases
2. The small concentration of H+ reacts with these ions, forming water
3. HA dissociates shifting the equilibrium position to the right to restore most of the H+ ions.
71
Why is there a buffer solution in blood plasma?
For pH of blood needs to be kept between 7.35-7.45
72
Which buffer solution is present in blood plasma?
Carbonic acid-hydrogencarbonate buffer system
73
What is the equation for Gibbs' free energy?
ΔG = ΔH - TΔS
74
Describe the thermodynamic feasibility of an exothermic reaction of increasing entropy
As ΔG will always be positive, this reaction is always thermodynamically feasible
75
Describe the thermodynamic feasibility of an endothermic reaction with decreasing entropy
As ΔG is always positive, this reaction is never thermodynamically feasible
76
Describe the thermodynamic feasibility of an exothermic reaction with decreasing entropy
ΔG will be negative for low temperature only, so is only thermodynamically feasible at low temperatures
77
Describe the thermodynamic feasibility of an endothermic reaction with increasing entropy
ΔG will be negative for high temperatures only, so will only be thermodynamically feasible at high temperatures
78
Describe the pH curve for a strong acid with strong base
Acid regtion is around 1
Base region is around 13
Equivalnce point is around 7
79
Describe the pH curve for a strong acid and weak base
Acid region around 1
Base region around 10
Equivalence point around 4
80
Describe the pH curve for a weak acid with a strong base
Acid region is around 4
Base region is around 13
Equivalence point is around 9
81
Describe the pH curve for a weak acid with a weak base
Acid region is around 4
Base region is around 10
No clear equivalence point
82
Suggest a suitable indicator for the titration of a strong acid with a strong base
Almost any indicator will work, e.g. methyl orange, phenolphthalein, litmus etc.
83
Suggest a suitable indicator for the titration of a weak acid with a strong base
An indicator with a pH range roughly between 7 and 11, e.g. phenelphthalein, metacresol purple etc.
84
Suggest a suitable indicator for the titration of a strong acid with a weak base
An indicator with a pH range roughly between 3 and 7, e.g. Methyl orange, methyl red, bromophenol blue etc.
85
Suggest a suitable indicator for the titration of a weak acid with a strong base
As there is no clear equivalence point, no indicator is suitable
86
Explain how an acid-base indicator works
* The indicator is a weak acid (HA)
* The HA and A- have clearly distinct colours
* At the end point of a titration, the indicator contains equal concentrations of the HA and A-, so the colour will be between the two colours
87
Explain the difference between the end point and equivalence point
* The end point is where the indicator changes colour
* The equivalence point is where the volume of one solution has reacted exactly with the volume of the second solution
88
Why is the pKa value of a weak acid being used as an indicator useful?
For this value is the same as the pH of the end point
89
Describe the procedure of measuring pH with a pH meter/probe
* Place the electrode of the pH meter in the conical flask, containing either the acid or the base
* Add the other chemical from the burette 1cm3 at a time, and swirl
* Let the pH meter settle, and then add record the pH value
* Repeat until the pH starts to change more rapidly, then add the other chemical drop wise and record the pH and amount of chemical added
* Continue with 1cm3 once the pH starts to change less rapidly
90
Describe the procedure of a redox titration with potassium manganate(VII)
* The KMnO4 is added to the burrette
* The solution being analysed is added to a conical flask with an excess of H2SO4 to provide the acid required for the reduction of the mangante(VII) ions
* As the KMnO4 is added, it decolourises as it is being added
* The end point is decided as the place where the first permanent colour (of pink) is seen
* The titration is repeated until concordent titres are obtained
91
Describe the problems which arise from titrating with potassium manganate(VII) and explain how these problems can be overcome
It can be hard to see the bottom of the meniscus through the deep purple colour, so the top of the meniscus is used. As it is the difference between two readings, the difference is the same, so the total volume used is still accurate
92
Describe the procedure for the iodine/thiosuflate titration
* Add a standard solution of Na2S2O3 to the burette
* Prepare the solution to be analysed by adding the solution to a conical flask and adding excess potassium iodide
* The solution oxidises the iodide ions to produce iodine
* Titrate this solution with the thiosulfate and when the brown colour sartes to fade, add starch as an indicator
* The starch turns the solution blue-black, and fades as the iodine is reduced to iodide, so the end point is when the solution turns colourless
93
Describe and explain the ligand substitution of [Cu(H2O)6]2+ from [Cu(NH3)4(H2O)2]2+
* In the first stage, the light blue solution turns to a pale blue precipitate of Cu(OH)2, which then dissolves in excess ammonia to form a dark blue solution
* [Cu(H2O)]2+(aq) + 4NH3(aq) → [Cu(NH3)4(H2O)2]2+(aq) + 4H2O
94
Describe and explain the ligand substitution of [Cu(H2O)6]2+ to [CuCl4]2-
* Concentrated HCl is added for a source of the chloride ions and the pale blue solution turns into a yellow solution. As this is a reversable reaction, adding water turns the solution back to blue, and green colour is formed when both of the solutions are present
* [Cu(H2O)6]2+(aq) + 4Cl-(aq) ⇌ [CuCl4]2-(aq) + 6H2O(l)
95
Describe and explain the ligand substitution reaction of [Cr(H2O)6]3+ with ammonia
* Initially, a grey-green precipitate of Cr(OH)3 is formed from the original pale purple solution
* The precipitate then dissolves in excess ammonia, forming the complex ion [Cr(NH3)6]3+
* [Cr(H2O)6]3+(aq) + 6NH3(aq) → [Cr(NH3)6]3+(aq) + 6H2O(l)
96
Explain the role of iron in haemogloblin and hence how carbon monoxide poisoning can occur
* O2 gas binds to the Fe2+ iron as a ligand and is carried around the body, forming oxyhaemoglobin, and this releases the oxygen when required into body cells.
* The carbon dioxide then forms a coordinate bond with the Fe2+ and is carried out of the lungs.
* CO can also bine to the Fe2+ in a ligand substitution reaction, replacing the oxygen, forming carboxyhaemoglobin
* As the CO binds more strongly than oxygen, a small concentration of carbon monoxide can prevent a large proportion of the haemaglobing molecuels for carrying oxygen
* As the bond is so strong, the process is irreversible, so if the concentration of carboxyhaemoglobin becomes too high, oxygen trasprot is prevented, leading to death
97
Describe the reaction between aqueous Cu2+ ions with aqueous sodium hydroxide
The pale blue solution forms a pale blue precipitate, Cu(OH)2
Cu2+(aq) + 2OH-(aq) → Cu(OH)2(s)
98
Describe the reaction with aqueous Fe2+ ions and aqueous sodium hydroxide
Pale green solution reacts to form a green precipitate Fe(OH)2
Fe2+(aq) + 2OH-(aq) → Fe(OH)2(s)
As the air can oxidise the Fe2+ to Fe3+, the green precipitate can turn brow at the surface as it forms Fe(OH)3
99
Describe the reaction between aqueous Fe3+ ions and aqueous sodium hydroxide
The pale yellow solution reacts to form an orange-brown precipitate Fe(OH)3.
Fe3+ +3OH- → Fe(OH)3(s)
100
Describe the reaction between aqueous Mn2+ ions and aqueous sodium hydroxide
The pale pink solution reacts to form a light brown precipitate Mn(OH)2
Mn2+(aq) + 2OH-(aq) → Mn(OH)2(s)
101
Describe the reaction between aqueous Cr3+ ions and aqueous sodium hydroxide
The violet solution reacts to from a grey-green precipitate Cr(OH)3
Cr3+(aq) + 3OH-(aq) → Cr(OH)3(s)
The precipitate then dissolves in excess sodium hydroxide to form a dark green solution
Cr(OH)3 + 3OH-(aq) → [Cr(OH)6]3-(aq)
102
Describethe interconcersions between Fe2+ and Fe3+
* Fe2+ is oxidised by H+/MnO4- and the iron(II) ions decolourise the purple manganate ions
* Fe3+ is oxidised by iodide, but the colour change of the iron ions is obscured due to the iodide being oxidised into iodine, which is yellow/brown
103
Describe the interconversions between Cr3+ and Cr2O72- ions
* Cr3+ ions are oxidised to Cr2O72- ions by hot alkaline peroxide, turning from
104
Describe the reduction of Cu2+ to Cu+
Pale blue Cu2+ is reduced to Cu+ with iodide ions, forming a white precipitate and brown iodine
105
Describe the disproportionation of Cu+
* Cu+ readily disproportionates in aqueous conditions
* Copper(II) oxide also disproportionates when reacted with sulfuric acid forming copper metal and blue copper sulfate
