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Physical Chemistry Flashcards

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1
Q

What is a permanent dipole?

A

A different in charge across different areas of a molecule, caused by some elements being more electronegative than others, meaning they pull electron pairs closer to themselves

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2
Q

What is an induced dipole?

A

When the electrostatic forces of the permanent dipole in one molecule cause another molecule to also become temporarily polarised

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3
Q

What are London dispersion forces?

A

When electrons randomly move around in an atom, causing some areas to become more charged than others, creating an effect similar to an induced dipole

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4
Q

What is a polar molecule?

A

A molecule with a permanent dipole

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5
Q

What is electronegativity?

A

The tendency of an atom to attract shared electron pairs towards its nucleus

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6
Q

[PPQ] Why do the boiling points of halogens increase down the group?

A

London forces between the molecules increase because the number of electrons in the atoms increases, and more heat energy is required to overcome the stronger intermolecular forces

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7
Q

How do you calculate an atom’s oxidation number?

A

Break all the bonds heterolytically, with the bonding electron pair(s) going to the more electronegative element. The charge on the resulting ion is the oxidation state.

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8
Q

What is the oxidation number of combined oxygen?

A

-2

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9
Q

What is the oxidation number of oxygen in peroxides?

A

-1

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10
Q

What is the oxidation number of combined hydrogen?

A

+1

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11
Q

What is the oxidation number of hydrogen in metal hydrides?

A

-1

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12
Q

What is the oxidation number of a halogen?

A

-1

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13
Q

What is dative covalent bonding?

A

When one of the atoms in the bond supplies both the electrons

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14
Q

What is the octet rule?

A

Elements tend to bond in such a way that each atom has eight electrons in its valence shell

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15
Q

What is a better replacement for the octet rule?

A

Unpaired electrons pair up, and the maximum number of electrons that can pair up is equal to the number of electrons in the outer shell

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16
Q

When does the octet rule not apply?

A

When elements with 2 or 3 electrons on the outer shell form covalent compounds

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17
Q

What determines the shape of a molecule or ion?

A

Number of bonding regions in the outer shell of the central atom. The electron pairs repel each other, meaning that a molecule with 3 atoms bonded to the central atom will have a trigonal planar structure

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18
Q

List the possible shapes of molecules, in order of number of bonding regions on the central atom

A

Linear (planar), linear (planar), trigonal planar, tetrahedral, trigonal bipyramid, octahedral

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19
Q

How do lone electron pairs affect the shape of molecules?

A

Lone electron pairs are slightly more electron dense, meaning they repel more

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20
Q

List the possible shapes of molecules with lone electron pairs, in order of number of lone pairs

A

Pyramidal, non-linear (bent)

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21
Q

What is ionisation energy of an element?

A

How many joules are required to remove an electron from each atom in a mole of gaseous atoms

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22
Q

How do successive ionisation energies change?

A

Each successive ionisation energy is higher than the one before. Ionisation energy increases radically when attempting to remove an electron from a shell closer to the nucleus

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23
Q

What are the factors that affect ionisation energy?

A

Distance from the nucleus of the electron being removed, electron shielding, nuclear charge

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24
Q

How does atomic radius change across the periodic table?

A

Bottom left has the widest nuclei

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25
Q

[PPQ] Explain why ionisation energies show a general increase across Period 3 (Na–Ar)

A

Nuclear charge increases because the number of protons in the nucleus increases

The atomic radius also decreases due to the increase in the number of electrons

The greater nuclear charge and lower atomic radius cause increased attraction between the nucleus and the valence electrons

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26
Q

[PPQ] Why do magnesium and aluminium not follow the trend of increasing ionisation energies in Period 3?

A

The Mg electron is removed from 3s, and the Al electron is removed from 3p. The 3p electron is higher energy than the 3s electron, meaning more energy is required to remove it

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27
Q

[PPQ] What is an ionic bond?

A

Electrostatic attraction between positive and negative ions

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28
Q

What is ΔH?

A

The difference in the energy of the reactants/products before the reaction and after the reaction. If the reaction is exothermic, it’s negative, and vice versa. It’s measured in kJ per mole

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29
Q

What are standard conditions?

A

100 kPa, 273K (0 degrees)

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30
Q

What is enthalpy change of reaction?

A

Energy change in the reactants/products

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31
Q

What is enthalpy change of formation (ΔfH)?

A

Energy change when 1 mole of a compound is formed from its elements. You might need to use fractions in the equation

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32
Q

What is enthalpy change of combustion (ΔcH)?

A

Energy change that takes place when 1 mole of a substance is combusted (usually negative)

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33
Q

What is enthalpy change of neutralisation (ΔneutH)?

A

Energy change when 1 mole of water is formed in a neutralisation reaction (always negative)

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34
Q

What is the formula for enthalpy change?

A 
q = mcT
q = Energy exchanged with the surroundings (joules)
m = mass
c = SHC
T = change in temperature (kelvin)
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35
Q

How do you calculate ΔH?

A

Bond enthalpies of reactants minus bond enthalpies of products

The enthalpy of a bond changes depending on where it is found, so the average enthalpy of the bond is used in calculations

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36
Q

What is Hess’ law?

A

The enthalpy change of a chemical reaction is independent of the route it takes

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37
Q

How is the rate of a reaction measured?

A

change in concentration over time (moldm^3 s^1)

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38
Q

How do catalysts work?

A

They increase the rate of a chemical reaction by decreasing the activation energy by providing an alternative route for the reaction to follow

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39
Q

What is a homogeneous catalyst?

A

A catalyst which is in the same phase as the reactants (e.g. both liquid)

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40
Q

What are heterogeneous catalysts?

A

A catalyst which isn’t in the same phase as the reactants (e.g. the reactant is a gas passed over a solid catalyst)

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41
Q

What is the Boltzmann distribution?

A

The mathematical distribution of molecules’ energy

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42
Q

When is a chemical system said to be in equilibrium?

A

The concentrations of the reactants and the products remain constant

The rate of the forwards reaction is the same as the rate of the backwards reaction

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43
Q

What is Le Chatelier’s principle?

A

When a system in equilibrium is subjected to a change, the position of equilibrium will shift to minimise the change

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44
Q

What is the effect of concentration on equilibrium?

A

Increasing a side’s concentration causes the system to produce more of the opposite side

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45
Q

What is the effect of temperature on equilibrium?

A

Increasing the temperature produces more of the chemicals on the higher-enthalpy side

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46
Q

What are the theoretical optimum conditions for the Haber process?

A

High pressure, low temperatures

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47
Q

What conditions are used for the Haber process?

A

~450 degrees, 200 atmospheres, iron catalyst

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48
Q

What is the equilibrium constant?

A

The ratio of products to reactants (a measure of where the equilibrium point is) (solids are not counted in the equation)

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49
Q

How is Kc calculated?

A

aA + bB => cC + dD

Kc = ([C]^c x [D]^d)/([A]^a x [B]^b)

[A] = concentration of A

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50
Q

If ΔH is positive, the reaction is…?

A

Endothermic

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51
Q

If ΔH is negative, the reaction is…?

A

Exothermic

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52
Q

[PPQ] Advantages/disadvantages of chlorine in water treatment

A

Kills bacteria, toxic (forms carcinogenic chlorinated hydrocarbons)

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53
Q

[PPQ] Why does phosphorus have a higher boiling point than chlorine?

A

More electrons, therefore stronger London forces between molecules which require more energy to overcome

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54
Q

[PPQ] What are the differences between sigma bonds and pi bonds?

A

Sigma bonds are between bonding atoms, while pi bonds are above and below
Sigma bonds have head-on overlap of orbitals while pi bonds have sideways overlap
Pi bonds are weaker than sigma bonds

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55
Q

[PPQ] What is a disproportionation reaction?

A

A reaction in which the same element is both oxidised and reduced.

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56
Q

Name the test for chlorine

A

Insert red or blue litmus paper into the gas and see if it loses its colour

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57
Q

Name the colours produced by burning different cations

A

Lithium = Crimson red, Sodium = Orange/yellow, Potassium = Lilac, Calcium = Red/orange, Copper = green

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58
Q

How would you use sodium hydroxide to identify cations?

A

Add sodium hydroxide solution to a solution of the substance being analysed, and some ions will produce a precipitate

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59
Q

Name the colours produced by adding NaOH to different cations

A

Blue = copper, Green = Iron(II), Brown = Iron(III)

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60
Q

Tetrahedral bond angle

A

109.5

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61
Q

Trigonal pyramidal bond angle

A

107

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62
Q

Bent bond angle

A

104.5

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63
Q

Octahedral bond angle

A

90

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64
Q

What is the ideal gas equation?

A 
pV = nRT
p = pressure (Pascals, RTP = 100000 or 101325 Pa)
V = volume (m^3 = 1000dm^3)
n = moles of gas
R = constant (8.31)
T = temperature (RTP = 298K/25C)
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65
Q

What is the effect of pressure on the point of equilibrium?

A

When there is an increase in pressure, the equilibrium will shift towards the side of the reaction with fewer moles of gas (more of that chemical will be produced)

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66
Q

Test for carbonate ions

A

Add a dilute acid, and there will be fizzing due to formation of CO₂ if CO₃(2-) ions are present. This can be confirmed by bubbling the gas through limewater and watching it turn cloudy
CO₃(2-) + 2H(+) => CO₂ + H₂O

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67
Q

Test for sulfate ions

A

Add barium chloride or barium nitrate. If a white precipitate forms, the solution contains sulfate ions
Ba(2+) + SO₄(2-) => BaSO₄

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68
Q

Test for halide ions

A

Add dilute nitric acid (HNO₃) followed by silver nitrate solution (AgNO₃) and record the colour of the precipitate formed
Ag(+) + X(-) => AgX
Add aqueous ammonia and shake the solution. If it does not dissolve, it’s iodide.

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69
Q

Test for ammonia

A

Ammonia is alkaline so it will turn red litmus paper blue

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70
Q

Test for ammonium ions

A

Add some sodium hydroxide to the substance and warm the mixture. If ammonium ions are present, ammonia will be produced. You can test for ammonia by holding wet litmus paper over the test tube and watching it turn blue

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71
Q

Balancing combustion reactions

A

C₆H₁₄ + O₂ -> CO₂ + H₂O
Balance carbon (6 carbons on the left so 6CO₂)
Balance hydrogen (14 hydrogens on the left so 7H₂O)
Balance oxygen (there are now 19 oxygens on the right, so 9.5O₂)
Remove fractions
2C₆H₁₄ + 19O₂ -> 12CO₂ + 14H₂O

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72
Q

What is oxidation?

A

Loss of electrons, gain of oxygen or loss of hydrogen

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73
Q

How do electrons fill up orbitals?

A

They fill orbitals with the same energy singly before they start sharing

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74
Q

[PPQ] Why does the reactivity of the Group 2 elements Mg–Ba increase down the group?

A

More electron shells, increasing the effect of electron shielding. Distance to nucleus increases, decreasing nuclear attraction. This means less energy is required to remove electrons.

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75
Q

[PPQ] Explain why the boiling point of PH3 is lower than the boiling point of NH3

A

NH3 has hydrogen bonding while PH3 does not, and more energy is required to overcome the hydrogen bonds in NH3

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76
Q

[PPQ] Explain why the boiling point of PH3 is lower than the boiling point of AsH3

A

AsH3 has more electrons, so it has more induced dipole-dipole interactions. More energy is required to overcome these dipoles.

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77
Q

How does NMR spectroscopy work?

A

A sample of compound is placed in a strong magnetic field and exposed to a range of different frequencies of radio waves. The nuclei of certain atoms in the molecule absorb energy from the waves. The amount of energy a nucleus absorbs tells you what environment it’s in

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78
Q

What are the two types of NMR spectroscopy?

A

Carbon-13 NMR (gives information about the number of carbon atoms and their environments), and proton NMR (number of hydrogen atoms)

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79
Q

Why do different carbon atoms absorb different amounts of energy?

A

They are shielded from the effects of external magnetic fields by surrounding electrons. The amount of shielding is also affected by which atoms the C or H is bonded to

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80
Q

How is tetramethylsilane used in NMR?

A

TMS produces a single peak at a lower frequency than almost everything else. This peak is given the value of 0, and peaks are measured as chemical shifts relative to this

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81
Q

What is chemical shift?

A

Difference in the radio frequency absorbed by TMS and that absorbed by the atoms in the molecules being analysed. It’s measured in ppm

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82
Q

How do you interpret a Carbon-13 NMR diagram?

A

Count the number of peaks that aren’t 0. This tells you how many different carbon environments there are. Look up chemical shifts in the data sheet to determine which bonds are present, and then try out possible structures that match the chemical formula given.

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83
Q

What does the relative area under each peak tell you in a proton NMR diagram?

A

The number of H atoms in each environment. If two peaks have areas in the ratio 2:1, the number of H atoms in the second peak is 2x the number of H atoms in the first

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84
Q

What is spin-spin coupling?

A

The number of mini-peaks in a peak is one more than the number of hydrogens on all the surrounding carbons not in the environment definition

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85
Q

What solvents are used for proton NMR?

A

Deuterated solvents, because deuterium doesn’t absorb radio wave energy like hydrogen does. This solvent is often CDCl3

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86
Q

How can OH and NH protons be identified? (since their peaks are so wide)

A

Run two spectra, one with D2O added. The deuterium should turn the OH and NH into OD and ND, causing the peak to disappear

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87
Q

What is meant by the order of reaction with respect to a specific reactant?

A

A number that tells you how that reactant’s concentration affects the rate of reaction. The rate is proportional to C^r (C = concentration, r = order).

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88
Q

How do you measure how each reactant affects the rate?

A

Add all the other reactants in excess, and then continuously measure the concentration of that reactant. The concentration-time graph if zero-order should be a straight line, and like e^-x if first-order

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89
Q

How do you work out the rate from a concentration-time graph?

A

Take the derivative. The initial rate is calculated from the tangent at t = 0

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90
Q

What is a rate equation?

A

An equation that tells you how the rate is affected by the concentrations of reactants

Rate = k[A]^m[B]^n
k = rate constant (bigger it is, faster the reaction)
m and n are orders of the reaction

The overall order of the reaction is m+n, and the order with respect to A is m

If the overall reaction is first-order, the rate constant is equal to the gradient of the rate-concentration graph

91
Q

How do you find the overall order of a reaction?

A

Adding up the individual orders

92
Q

How can the half life be calculated for a first-order reaction?

A

Reading the concentration-time graph

93
Q

What equation links the half life of a first-order reaction with the rate constant?

A

k = ln(2)/h

94
Q

What is the official definition of relative atomic mass?

A

The average mass of all isotopes of a certain element

95
Q

How can you reduce the percentage uncertainty?

A

Using larger masses/quantities

96
Q

What is the difference between a qualitative and a quantative test?

A

Qualitative tests measure physical qualities such as colour, while quantative tests measure quantities such as mass

97
Q

How do catalysts speed up reactions?

A

They lower the activation energy by providing a different way for the bonds to be broken and remade. This means that more particles will have enough energy to react.

98
Q

How do you work out the empirical formula of a compound using the percentages by mass of its elements?

A

Divide each element’s fraction of the total mass by the its relative atomic mass:

H: 1/20 (/1) = 1/20
N: 7/20 (/14) = 0.7/28
O: 3/5 (/16) = 3/80
(1/20):(0.7/28):(3/80) = (2/40):(1/40):(1.5/40)
= H4N2O3
99
Q

What are p-block elements?

A

Elements whose highest energy electrons are in a p orbital (B, C, N, O, F, Ne and all elements below them)

100
Q

What is a standard solution?

A

A solution of known concentration

101
Q

Formula for percentage uncertainty

A

200u/r (u = +/- uncertainty, r = reading)

102
Q

What is the recommended way of making a standard solution?

A

Using a volumetric flask instead of a beaker

103
Q

Electronic structure of radon

A

1s²2s²2p⁶3s²3p⁶3d¹⁰4s²4p⁶4d¹⁰5s²5p⁶4f¹⁴5d¹⁰6s²6p⁶

104
Q

[PPQ] Explain why iodine is less reactive than bromine

A

Iodine has a larger atomic radius
Iodine has greater shielding / more shells
Iodine has weaker / less nuclear attraction (on electron
gained than bromine)

105
Q

When you increase the concentration of a reactant, what happens to the rate of the reaction?

A

If the reactant is order zero, the rate stays the same
If the reactant is order one, the rate increases linearly
If the reactant is order two, the rate increases quadratically

106
Q

How does the order of a reactant affect its concentration/time graph

A

Order zero: Straight line like y=-x
Order one: Curve like 1/x
Order two: Curve like 1/x^4

107
Q

How can you use Kc to tell if a reaction is endothermic or exothermic?

A

If Kc gets bigger as the temp increases, it’s endothermic. Temperature is the only thing that can change Kc

108
Q

What is the rate of a reaction?

A

Change in the amount of reactants or products per unit time. It can be calculated by measuring the volume of gas evolved, change in mass, or pH change. You can also draw a tangent to a concentration/time graph

109
Q

Describe the iodine clock reaction

A

A small amount of sodium thiosulfate and starch are added to an excess of hydrogen peroxide and iodide ions in acid solution. The sodium thiosulfate reacts immediately with any iodine that forms. When the sodium thiosulfate runs out, the iodine will immediately turn the starch indicator blue-black. Varying the iodide or hydrogen peroxide concentrations will give different times for the colour change

110
Q

Zero order, first order, second order

A

Order 0: Double reactant, rate stays the same
Order 1: Double reactant, rate doubles
Order 2: Double reactant, rate quadruples

111
Q

What is the rate-determining step?

A

The slowest step in a multi-step reaction. A reactant’s order is how many molecules of it show up in the reactants of this step. It does not necessary have to occur first or last. When writing step equations, you’re allowed to make unstable intermediates and strange equations (e.g. Fe+ + Fe3+ => 2Fe2+)

112
Q

What is the Arrhenius equation?

A 
k = Ae^(-Ea/RT)
A = pre-exponential factor (constant)
R = gas constant
k = rate constant
Ea = activation energy (J mol^-1)
T = temperature (K)
113
Q

What is the Arrhenius equation in logarithmic form?

A

ln(k) = -Ea/RT + ln(A)

114
Q

What is an Arrhenius plot?

A

A plot with 1/T on the x-axis and ln(k) on the y-axis

The y-intercept is ln(A), and the gradient is -Ea/R

115
Q

What is a conjugate pair?

A

Two species linked by the gain of a proton (e.g. a Cl- ion and HCl)

116
Q

History of acid-base theory

A

At first it was thought that oxygen caused a compound to be acidic, but then HCl was discovered
Arrhenius suggested that acids always release H+ in solution and bases always release OH-. This is often true, but not in cases such as NH3, which is basic.

117
Q

What is the Bronsted-Lowry theory of acids?

A

Acids donate protons, and bases accept them. Acids and bases react to form 2 conjugate pairs, rather than just a salt and water

118
Q

pH =

A

-log10( [H+] ), [H+] = 10^-pH

For strong monoprotic acids, [H+] = [HA]

119
Q

What is Kw?

A

The equilibrium constant for the dissociation of water. This is given by Kw = [H+][OH-]. The only thing that can change Kw is temperature, not concentration. It is 10^-14 at 298K.

120
Q

What is Ka?

A

The equilibrium constant for the dissociation of a weak acid (HA <=> H+ + A-). It’s given by [H+][A-]/[HA]

121
Q

How can Ka be approximated?

A

[H+] is roughly [A-], so Ka is roughly [H+]^2/[HA]. This formula can be used in calculations, but not in “give the expression” questions

122
Q

What is pKa and how is it calculated?

A

pKa is pH but for weak acids. pKa = -log10(Ka)

123
Q

What is a buffer?

A

A solution that resists changes in pH when small amounts of acid or base are added.

124
Q

How are acidic buffers created?

A

They have a pH less than 7, and are made by setting up an equilibrium of a weak acid and the salt of its conjugate base. This can also be done by mixing an excess of weak acid with a strong alkali

125
Q

What does a buffer solution contain?

A

HA <=> H+ + A-

Lots of undissociated acid, lots of the conjugate base, and enough H+ to cause acidity

126
Q

What happens when you add H+ to a buffer?

A

The extra H+ ions combine with the A- ions to form HA, shifting the equilibrium to the left, reducing the H+ concentration to near the original value

127
Q

What happens when you add OH- to a buffer?

A

H+ ions are removed, causing HA to dissociate into more H+ ions, shifting the equilibrium to the right and returning [H+] to its original value

128
Q

How do you work out the pH of a buffer solution?

A

Use the concentrations and the Ka to work out [H+], and then use that to work out the pH

129
Q

How do you choose which indicator to use when doing a titration?

A

Pick an indicator that changes colour in the pH range where the pH curve is almost vertical (near the equivalence point)

130
Q

What is standard lattice enthalpy?

A

Enthalpy change when 1 mole of an ionic lattice is formed from gaseous ions
Na+(g) + Cl-(g) => NaCl(s)

131
Q

How does ionic radius affect lattice enthalpy?

A

Higher charge on the ions causes more negative lattice enthalpy (due to stronger electrostatic forces)

132
Q

How does charge affect lattice enthalpy?

A

Smaller ionic radii lead to more negative lattice enthalpy because they have higher charge density and sit closer together in the lattice, increasing the strength of the attraction between them

133
Q

What happens when a solid ionic lattice dissolves in water?

A

The bonds between the ions break to form gaseous ions (endothermic, the opposite of lattice enthalpy), and bonds between the ions and water are made. Soluble substances have exothermic enthalpies of solution

134
Q

What is the enthalpy change of hydration?

A

Enthalpy change when 1 mole of gaseous ions dissolve in water

135
Q

How does ionic charge affect the enthalpy of hydration?

A

Ions with a higher charge are better at attracting water molecules, so more energy is released when the bonds are made, giving them a more negative enthalpy change

136
Q

How does ionic radius affect the enthalpy of hydration?

A

Smaller ions have a higher charge density so they attract the water molecules better and have a more exothermic enthalpy of hydration

137
Q

What is entropy?

A

A measure of the number of ways that particles can be arranged and the number of ways that energy can be shared out between particles. A large, positive entropy value (JK⁻¹mol⁻¹) indicates a high level of chaos. It tells us how much energy the substance will require internally for every degree the temperature rises

138
Q

What causes entropy to increase?

A

The formation of a gas, and the formation of more particles

139
Q

Why are endothermic reactions feasible?

A

Because particles are more energetically stable when there’s more disorder

140
Q

How do you calculate the enthalpy change of a reaction?

A

Enthalpy of products minus enthalpy of reactants. When this is positive the reaction is more likely to be feasible

141
Q

What determines the tendency of a process to take place?

A

The free energy change (ΔG). This is given by the formula ΔG = ΔH - TΔS
ΔH = enthalpy change (J/mol)
ΔS = entropy (JK⁻¹mol⁻¹)
T = temperature (K)
The more negative the value of ΔG, the more feasible the reaction. If ΔG is positive, the reaction cannot occur.

142
Q

How do halogens interact with organic solvents?

A

They dissolve well in them

143
Q

How do you determine if a reaction is redox?

A

If the oxidation numbers of at least 2 of the elements change

144
Q

What is a conjugate base-acid pair?

A

A pair of molecules that only differ by one H+ ion

145
Q

[PPQ] Why does formation of a solid decrease entropy?

A

Solids have less disorder/ways of arranging molecules so have lower entropy

146
Q

[PPQ] In terms of Kp, how does increasing the pressure cause the equilibrium position to change for the reaction 2NO + O2 => 2NO2

A

The equilibrium would shift to the right because it’s the side with the fewest gaseous moles. The denominator in the kP expression would initially increase because of the pressure, but then the equilibrium shifts and the numerator increases to restore the old Kp

147
Q

What is partial pressure?

A

The pressure exerted by one gas in a mixture of gases. If two gases are produced in the same quantity, they will have the same partial pressures. It’s the mole fraction of a gas multiplied by the total pressure

148
Q

What is total pressure?

A

The sum of all the partial pressures of individual gases

149
Q

What is a mole fraction?

A

Proportion of a gas mixture that is a particular gas. It’s the number of moles of specific gas divided by the total moles of gas

150
Q

What is the formula for Kp?

A

aA(g) + bB(g) + dD(g) + eE(g)
Kp = (p(D)^d * p(E)^e)/(p(A)^a * p(B)^b)
p(x) = partial pressure of x
Expressions for Kp only included gases

151
Q

What is the only thing that can change the equlibrium constant (Kc or Kp)?

A

Changing the temperature. Whether Kc/Kp increases or decreases depends on whether the reaction is endothermic or exothermic

152
Q

What happens to Kc and Kp when the concentration is changed?

A

If the concentration of one side increases, in order to balance this out, so does the concentration of the other side. This keeps the constant the same.

153
Q

What happens to Kc and Kp when the pressure is changed?

A

It produces more of the side with fewer gas molecules, which raises pressure back to the original level and keeps the constant the same.

154
Q

How do you balance a half equation?

A

e.g.: H2SO4 + 8H+ 8e- => H2S + 4H2O

Step 1: Look at how many H2Os are required on the right side (H2SO4 => H2S requires 4H2O)
Step 2: Add as many H+ ions as necessary to the left side
Step 3: Add an equal number of e-

155
Q

How do you set up an electrochemical cell?

A

Two different metals dipped in salt solutions of their own ions and connected by a metal wire. The two beakers are also connected by a salt bridge, a piece of filter paper soaked in KNO3 which dips into both beakers

156
Q

What happens in an electrochemical cell (zinc and copper)?

A

Zinc loses electrons more easily than copper, so in the zinc beaker, zinc is oxidised (Zn => Zn2+ + 2e-) and electrons are released into the circuit. In the less reactive beaker (copper), the same number of electrons is taken from the circuit, reducing Cu2+ into copper atoms (Cu2+ + 2e- => Cu). Electrons always flow from the most to the least reactive metal

157
Q

What is electrode potential?

A

A measure of how readily an element is oxidised. More easy to oxidise = more negative electrode potential

158
Q

How is the electrode potential measured?

A

Voltage measured when the half-cell is connected to a standard hydrogen electrode

159
Q

Describe the standard hydrogen half cell

A

Inert platinum electrode in a 1 mol acid solution, at 298K and 100 kPa. There is also a glass tube with a hole to let bubbles of H2 escape. The half reaction is 2H+ + 2e- => H2

160
Q

Halogen redox reactions

A

Cl2 + 2Br- => Br2 + 2Cl- (orange/brown) in organic
Cl2 + 2I- => I2 + 2Cl- (yellow/orange) in organic
Br2 + 2I- => I2 + 2Cl-
Lighter species gets the electrons

161
Q

[PPQ] Why is the H-O-H bond angle greater in ice than in steam?

A

Ice has hydrogen bonding, while steam does not. H2O has 2 bonding pairs and 2 lone pairs, and since lone pairs repel more than bonded pairs and the lone pairs in ice are used up for the hydrogen bonds, their effect on the bond angle is smaller, increasing the bond angle

162
Q

What happens when you add iodine and starch to water?

A

It turns blue

163
Q

How are electrochemical cells drawn?

A

Reduced form | Oxidised form || Oxidised form | Reduced form

Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)

164
Q

What are the conventions for drawing cells?

A

The cell with the more negative potential goes on the left

165
Q

What is the overall cell potential?

A

E cell = E more positive - E more negative

166
Q

How can E values be used to determine if a metal will react with the aqueous ions of another metal?

A

Look at the two half equations. The one with the more positive half reaction is written forwards in the whole equation, which represents the feasible direction of the two half-reactions. Example:

Zn2+ + 2e- <=> Zn (E = -0.76V)
Cu2+ + 2e- <=> Cu (E = 0.34V)

Zn + Cu2+ => Cu + Zn2+
So zinc reacts with copper ions

167
Q

When can predictions using E values be wrong?

A

When the conditions are not standard or the concentration is changed
When the reaction kinetics are not favourable (the reaction is so slow that it never seems to happen, or the high activation energy stops it from happening)

168
Q

How do fuel cells generate electricity?

A

By reacting a fuel (usually hydrogen) with an oxidant, usually oxygen. Water and electricity is produced (basically combustion but more efficient because no energy is wasted as light or sound)

169
Q

Describe a hydrogen/oxygen fuel cell

A

At the anode the platinum catalyst splits the H2 fuel into 2H+ and 2e-. The electrons flow through a wire to the cathode. The electrolyte membrane allows the H+ to flow to the cathode, where it combines with the O2 and the electrons to form water

170
Q

Advantages and disadvantages of fuel cells

A

More efficient than combustion engines, no pollution
Toxic chemicals required to make the cells
Flammable

171
Q

Combining the half equations of a storage cell

A

Fe + 2OH- => Fe(OH)2 + 2e-
NiO(OH) + H2O + e- => Ni(OH)2 + OH-

Fe + 2H2O + 2NiO(OH) => 2Ni(OH)2 + Fe(OH)2
Ionic parts cancel out (OH- and e-)

172
Q

Why do chemicals in a mixture separate in chromatography?

A

They differ in the extent to which they are soluble in the mobile phase (and the stationary phase if it’s a liquid)
They differ in the extent to which they adsorb/stick to the stationary phase

173
Q

How does thin layer chromatography work?

A

Thin layer of Al2O3 or SiO2 (stationary phase) supported on a glass plate is inserted into a container of solvent with a lid to keep the solvent evaporating. The height that each substance reached is recorded, and used to calculate the Rf (distance moved by spot divided by distance moved by solvent). UV light can be used to view invisible compounds

174
Q

What is gas chromatography?

A

A sensitive technique used for analysing complex mixtures which vaporise when heating. This method also tells you how much of each compound is present

175
Q

How does gas chromatography work?

A

A sample is vapourised and injected into the head of the chromatographic column. The column is full of an unreactive gas (the carrier gas), and a solid or liquid adsorbed into an inert solid. If the stationary phase is a liquid, separation depends on the relative solubility of the tested compound in the stationary and mobile phases. If the stationary phase is a solid, it depends on the adsorption of the components into the stationary phase. The components reach the end of the column one by one and pass into a detector, which plots a graph of time and recorded chemical response. The area under each peak gives an indication of the relative amount of that component

176
Q

What is retention time?

A

Time taken between the injection of a sample and emergence from the column

177
Q

Disadvantages of gas chromatography

A

Similar compounds such as methanal and ethanal often have overlapping peaks. Identification of unknown compounds is hard because reference times vary based on rate of flow of carrier gas and temperature. These limitations are overcome by pairing it with mass spectrometry

178
Q

What is the definition of a transition element?

A

A d-block element that can form at least one stable ion with an incomplete d sub-shell

179
Q

Describe the electron configuration of a transition element

A

1s2 2s2 ………. 3p6 3d_ (4s2 ……..

When forming ions, the 4s electrons are lost first, before the 3d ones

180
Q

What are the properties of transition elements?

A

Exist in variable oxidation states, make good catalysts, form coloured ions. They form different oxidation states because the energy levels of the 4s and 3d orbitals are close, meaning multiple electrons can be lost with not much energy. They can all form 2+ ions, and their maximum oxidation state can not exceed the number of 3d and 4s electrons

181
Q

Why do transition metals make good catalysts?

A

They can change oxidation states by gaining or losing electrons in their d shells, so they can transfer electrons to speed up reactions.

182
Q

What form do transition metals take in water?

A

[M(H2O)6]x+ (also written as Mx+ (aq))

183
Q

Reactions of Cu²⁺ with water and ammonia

A

[Cu(H2O)₆]²⁺(aq) + 2OH⁻(aq) => Cu(OH)₂(H2O)₄ + 2H₂O(l)
[Cu(H2O)₆]²⁺(aq) + 2NH₃(aq) => Cu(OH)₂(H2O)₄ + 2NH₄⁺(aq)

Pale blue solution to blue precipitate

184
Q

Reactions of Fe³⁺ with water and ammonia

A

[Fe(H2O)₆]³⁺(aq) + 3OH⁻(aq) => Fe(OH)₃(H2O)₃ + 3H₂O(l)
[Fe(H2O)₆]³⁺(aq) + 3NH₃(aq) => Fe(OH)₃(H2O)₃ + 3NH₄⁺(aq)

Yellow solution to orange precipitate

185
Q

Other precipitate colours

A

Manganese: Pale pink solution to pink precipitate
Iron II: Pale green solution to green precipitate
Green solution to grey-green precipitate

186
Q

What is “enthalpy of formation” in relation to the Born-Haber cycle for NaCl to Na+(g) + e- + Cl(g)?

A

Natural states to ionic lattice (e.g. Na + 0.5Cl2 => NaCl)

187
Q

What is “atomisation energy” in relation to the Born-Haber cycle for NaCl to Na+(g) + e- + Cl(g)?

A

Splitting an element into its atoms (e.g. Cl2(g) => 2Cl(g))

188
Q

What is “electron affinity” in relation to the Born-Haber cycle for NaCl to Na+(g) + e- + Cl(g)?

A

An atom taking a lone electron (e.g. Na+(g) + e- + Cl(g) => Na+(g) + Cl-(g))

189
Q

What is “lattice enthalpy” in relation to the Born-Haber cycle for NaCl to Na+(g) + e- + Cl(g)?

A

Gaseous ions to ionic lattice (e.g. Na+(g) + Cl-(g) => NaCl(s))

190
Q

What is “hydration enthalpy” in relation to the Born-Haber cycle for NaCl to Na+(g) + e- + Cl(g)?

A

Dissolving gaseous ions in water

191
Q

What is “solution” in relation to the Born-Haber cycle for NaCl to Na+(g) + e- + Cl(g)?

A

Energy released when dissolving an ionic lattice in water

192
Q

What enthalpies add up to make the energy change for NaCl => Na+(g) + Cl(g) + e-?

A
  • enthalpy of formation of NaCl + atomisation Cl + atomisation Na + ionisation Na
  • lattice enthalpy - Cl e- affinity
193
Q

Writing half equations for things such as Cr₂O₇²⁻ => Cr³⁺

A

Step 1: Look at how many H2Os are required on the right side (7 oxygens lost, so 7H2O)
Step 2: Add as many H+ ions as necessary to the left side (7 * 2 = 14)
Step 3: Add an equal number of e⁻ on the left side

194
Q

Combining two half equations

A

Multiply the equations until the e⁻s are equal, and then simply merge the equations together

195
Q

How does haemoglobin work?

A

The structure is a Fe²⁺ ion, surrounded by one globin molecule, 4 nitrogen ligands, and either a water or an oxygen molecule. This complex carries oxygen around the body. If carbon monoxide is inhaled, it replaces the oxygen and isn’t easily replaced, which is a problem as it prevents oxygen from reaching cells

196
Q

[Important] Describe the left branch of the Born-Haber cycle for MgCl2 (bottom to top)

A 
Minus enthalpy of formation
2x atomisation energy of chlorine
Atomisation energy of magnesium
First ionisation of magnesium
Second ionisation of magnesium

Formation, atomisation, ionisation

197
Q

[Important] Describe the right branch of the Born-Haber cycle for MgCl2 (top to bottom)

A

2x First electron affinity of chlorine
Hydration enthalpy of magnesium
2x Hydration enthalpy of chlorine
Minus solution enthalpy

The last 3 steps add to make the lattice enthalpy

198
Q

If ΔH = -92 kJ/mol for 3H2(g) + N2(g) => 2NH3, how do you calculate the energy released per mole of hydrogen gas?

A

Divide -92 by 3 (it’s the molar ratio of hydrogen), so 1 mol of hydrogen reacting = 30.7kJ released

199
Q

What causes entropy to increase?

A

Solid becoming a liquid, liquid becoming a gas, temperature rising, or a solid dissolving. The entropy also rises when solids reactants form gaseous products, or the products contain more gaseous molecules than the reactants

200
Q

How do you calculate ΔS (change in entropy) for a solution?

A

The sum of the entropy of the products, minus the sum of the entropy of the reactants

201
Q

Formula for lattice enthalpy and enthalpy of formation

A

ΔLE H = ΔFH - everything else

202
Q

How do you represent the entire Born-Haber cycle mathematically?

A

Equate the two sides, quantities going down get negative signs

Example:

2atomisation of Cl + atomisation of Mg + 1st ionisation Mg + 2nd ionisation Mg - formation = -2affinity of chlorine - lattice enthalpy

203
Q

Describe what happens to MgCl2(s) as you move up the left branch of the Born-Haber cycle

A 
MgCl2(s)
Mg(s) + Cl2(g)
Mg(s) + 2Cl(g)
Mg(g) + 2Cl(g)
Mg+(g) + 2Cl(g) + e-
Mg2+(g) + 2Cl(g) + 2e-
204
Q

Describe what happens to MgCl2(s) as you move down the right branch of the Born-Haber cycle

A

Mg2+(g) + 2Cl(g) + 2e-
Mg2+(aq) + 2Cl-(g)
Mg2+(aq) + 2Cl-(aq)
MgCl2(s)

205
Q

Creating a full redox equation from two half equations that occur in an acid solution

A

Write down the two half equations
MnO4-(aq) + H+(aq) => Mn2+(aq) + H2O(l)
I-(aq) => I2(s)

Balance them by changing the amounts of H+ and oxygen (in the H2O)
MnO4-(aq) + 8H+(aq) => Mn2+(aq) + 4H2O(l)
2I-(aq) => I2(s)

Add electrons to make sure the charges on both sides are the same
5e- + MnO4-(aq) + 8H+(aq) => Mn2+(aq) + 4H2O(l)
2I-(aq) => I2(s) + 2e-

Multiply both equations by appropriate amounts to ensure that they both have the same number of e-
10e- + 2MnO4-(aq) + 16H+(aq) => 2Mn2+(aq) + 8H2O(l)
10I-(aq) => 5I2(s) + 10e-

Merge the equations together

206
Q

What is the purpose of the salt bridge in an electrochemical cell?

A

Zinc ions are formed in the zinc half-cell, and copper ions are removed from the copper half-cell, so cations and ions move across the salt bridge to keep both half-cells electrically neutral

207
Q

What is the standard electrode potential (Eo) of a half cell?

A

The voltage measured when that half cell is connected to the standard hydrogen half cell

208
Q

How do you calculate the overall potential of a cell?

A

More positive potential minus the less positive potential. The half-cell with the more negative potential provides the electrons

209
Q

How do you tell if a redox reaction is feasible?

A

If the cell potential is positive, the reacton is feasible. Rremember to flip the sign if reversing the direction of the reaction, and then add the two values (don’t subtract)

210
Q

How does concentration affect the feasibility of reactions?

A

The equilibrium shifts a tiny bit away from the side with the higher concentration

211
Q

How are scandium and zinc different from other transition elements?

A

They only form one type of ion, they are white and non-catalytic

212
Q

Why do transition elements act as catalysts?

A

They adsorb the target species and weaken its bonds

213
Q

What does a complex with 2 ligands look like?

A

Linear

214
Q

What are bidentate ligands?

A

Ligands which form 2 coordinate bonds with the central ion

215
Q

How do you determine whether a complex ion is E or Z?

A

If the 2 ligands with the highest priorities are at 90 degrees to each other, the complex is Z, and if they are at 180, the complex is E

216
Q

What happens to the metal with the higher Eo?

A

The electrode whose metal has a higher Eo gains mass, as its ions gain electrons

217
Q

How can you tell if species A can oxidise species B?

A

If the Eo value associated with A is higher than the one associated with B

218
Q

What is used to oxidise Cr3+ into Cr2O7-?

A

Hydrogen peroxide

2Cr3+ + 3H2O2 + 10OH- => 2CrO4(2-) + 8H2O

219
Q

What is used to reduce Cu2+ to Cu+?

A

Iodide ions

220
Q

How can MnO4- be used to carry out redox titrations?

A

Add dilute sulfuric acid, followed by deep purple MnO4- solution until all the MnO4- has been reduced to pale pink Mn2+ ions

221
Q

From a list of redox equations and Eo values, how can you tell what the strongest reducing and oxidising agents are?

A

The reagent on the non-“e-“ side of the equation with the lowest Eo is the strongest reducing agent

The reagent on the e- side of the equation with the highest Eo is the strongest oxidising agent

222
Q

When should a platinum electrode be used?

A

When there are ions on both sides of the equation (e.g. Fe3+ + e- => Fe2+), or the compound is not a solid metal (e.g. iodine)

223
Q

What happens when transition metals form an ion?

A

They lose the electrons in their 4s shell first

224
Q

Determining the anode/cathode in a cell?

A

Oxidation happens at the anode. This means the cell where an electron is released is the anode

A-level Chemistry (8 decks)

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