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OCR A Level Chemistry > deck_8929427 > Flashcards

deck_8929427 Flashcards

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1
Q

Define acid

A

dissociates in water and releases H+ ions

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2
Q

Define alkali

A

dissociates in water to release OH- ions

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3
Q

Neutralisation equation

A

H+ + OH-1  H2O

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4
Q

Bronsted Lowry Acid:

A

Proton donor

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5
Q

Bronsted Lowry Base:

A

Proton acceptor

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6
Q

pH = 1 (x 10 the concentration of H+ ions) pH =

A

2

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7
Q

What makes an acid strong?

A

It fully dissociates

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8
Q

As a Strong acid fully dissociates, pH of a strong acid can be calculated from

A

concentration of acid

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9
Q

pH =

A

-log10[H+]

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10
Q

Sulphuric acid is dibasic, what does this mean?

A

H2SO4  2H+ + SO42-

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11
Q

[H+] =

A

10-pH

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12
Q

Weak acids form an acid dissociation constant called

A

Ka

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13
Q

Ka =

A

[products] / [reactants]

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14
Q

[H+] = for weak acids

A

root([HA] Ka)

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15
Q

[HA] is what?

A

concentration of weak acid

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16
Q

HCl + H2O –>

A

H3O+ + Cl-

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17
Q

What is H3O+

A

A hydronium ion, also known as an oxonium ion

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18
Q

What is water acting as in the equation:HCl + H2O  H3O+ + Cl-

A

a base

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19
Q

What is a bronsted-lowry base?

A

A Bronsted Lowry base is a substance which accepts protons in solution

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20
Q

What is a bronsted-lowry acid?

A

A Bronsted Lowry acid is a substance which releases or donates protons in solution

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21
Q

What does a bronsted-lowry acid-base reaction involve?

A

the transfer of a proton from one base to another

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22
Q

What does ammonia do in water? include equation

A

Ammonia, in water, accepts a protonNH3 + H2O  NH4+ + OH-

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23
Q

In the equation,NH3 + H2O  NH4+ + OH-, what is water acting as?

A

An acid

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24
Q

For a bronsted-lowry base, what is the pH of an acid?

A

So for the Bronsted Lowry theory, a base does not need to have a pH>7

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25
Q

What are substances that can act as either a base or an acid called?

A

Substances that can act as either an acid or a base are amphoteric

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26
Q

What are some old defnitions of acids?

A

.The definition of acids has developed since the time of the ancient Greeks.Simpler ideas involved substances that had a sour taste, contained hydrogen, hydrogen ions or had a pH lower than 7

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27
Q

.When a chemical reacts with an acid it is the __________ from the acid which is driving the reaction forward

A

hydrogen ion, H+,

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28
Q

HCl (aq) + NaOH (aq) 

A

NaCl (aq) + H2O (l)

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29
Q

How did the Bronsted-Lowry name come about?

A

Johannes Nicolaus Bronsted and Martin Thomas Lowry did not work together but both chemists formulated the idea that acids are proton donors and bases are protons acceptors

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30
Q

The Brønsted–Lowry acid–base theory is a development of these earlier ideas and states that:

A

.An acid is a proton donor.A base is a proton acceptor

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31
Q

.How many protons, neutrons and electrons does a Hydrogen ion, H+ have?

A

.A hydrogen ion is just a proton

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32
Q

We know that HCl forms a covalent bond between the hydrogen and chlorine so what happens when it dissolves in water to become hydrochloric acid?

A

.HCl(g) + aq  H+(aq) + Cl-(aq)

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33
Q

HCl(g) + aq  H+(aq) + Cl-(aq).On closer inspection is?

A

HCl(g) + H2O(l)  H3O+(aq) + Cl-(aq)

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34
Q

.H3O+(aq) is known as the hydronium ion, also known as _______ ion

A

hydroxonium

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35
Q

.In this equation:HCl(g) + H2O(l)  H3O+(aq) + Cl-(aq)what is HCl and H2O acting as?

A

The HCl is a proton donor (a Brønsted–Lowry acid).The H2O is a proton acceptor (a B-L base)

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36
Q

In the reverse equation of HCl(g) + H2O(l)  H3O+(aq) + Cl-(aq) what happens to the roles?

A

.In the reverse equation, the roles reverse:H3O+(aq) + Cl-(aq)  HCl(g) + H2O(l)

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37
Q

once an acid has ‘donated’ a proton it would become able to ‘accept’ a proton back and hence act as a base, what is this called?

A

We call these pairs of chemicals conjugate acid-base pairs

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38
Q

Give an example of a conjugate acid-base pair?

A

HCl and Cl- = Acid and conjugate base

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39
Q

What are the acid base pairs in HCl + H2O H3O+ + Cl-

A

HCl = acid 1Cl- = base 1H2O = base 2H3O+ = acid 2

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40
Q

What are the acid base pairs in CH3COOH + H2O CH3COO- + H3O+?

A

CH3COOH = acid 1CH3COO- = base 1H2O = base 2H3O+ = acid 2

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41
Q

What are the acid base pairs in NH3 + H2O NH4+ + OH-?

A

Base 2 Acid 1 Acid2 Base 1

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42
Q

What are the acid base pairs in HCO3- + HCl H2CO3 + Cl-?

A

Base 2 Acid 1 Acid 2 Base 1

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43
Q

Complete the equation for the conjugate acid-base pair: CH3CH(OH)COOH + CH3CH2CH2COOH

A

CH3CH(OH)COO- + CH3CH2CH2COOH2+

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44
Q

List the roles of H+ in reactions

A

spectator ions:Acid + metal  salt + hydrogenSolid carbonates and soluble carbonates:Acid + carbonate  Water + carbon dioxidebase (metal oxides): Acid + base  salt + waterAlkali: H+ + OH-  H2O

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45
Q

What did Soren Sorenson do?

A

introduced simple numbers to represent the colours of indicators using an electrochemical cell to measure the hydrogen content

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46
Q

What did Soren Sorenson find?

A

.He found the ion concentration had a very large range of values of powers of 10 (10-1 to 10-14)

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47
Q

What is a strong acid? With general equation

A

.A strong acid is one which completely dissociates into ions in a solutionHA  H+ + A-

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48
Q

So for a strong acid, [H+] =

A

concentration of the acid ([HA])

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49
Q

pH =

A

-log10[H+]

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50
Q

What does the equation pH = -log10[H+] tell us?

A

.It tells us the relative hydrogen ion concentration of a given solution

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51
Q

What must we remember with the equation pH = -log10[H+]?

A

.THIS ONY WORKS FOR STRONG ACIDS

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52
Q

.An increase in H+ x10, it reduces the pH by how much?

A

1

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53
Q

What is pH of 1M HCl?

A

0

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54
Q

What is Sorenson’s pH scale?

A

.The logarithmic scale means that a shift of one pH unit means a 10x change in the acidity and alkalinity of the solution.Theoretically there is no limit to the pH scale

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55
Q

What type of acid is HCl, what does this mean?

A

monobasic = [H+]

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56
Q

What type of acid is H2SO4, what does this mean?

A

dibasic = 2[H+]

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57
Q

What type of acid is H3PO4, what does this mean?

A

tribasic = 3[H+]

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58
Q

[H+] =

A

10-pH

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59
Q

What is the [HCl] of a solution with pH 1.8?

A

0.016 moldm-3

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60
Q

What is the [H2SO4] of a solution with pH 1.8?

A

0.008 moldm-3 as dibasic acid

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61
Q

50cm^3 of 0.1M HCl is diluted to 100cm^3 with water, what is the change in pH?

A

Before dilution, pH = 10n dilution, HCl conc is halved to 0.05MAfter dilution, pH = 1.30Change = 0.3

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62
Q

What is a strong acid?

A

A strong acid is one which completely dissociates into ions in a solution

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63
Q

WHat does pH = -log10[H+] only work for?

A

ONLY WORKS FOR STRONG ACIDS

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64
Q

[H^+] equation for changing concentraation

A

[H+] = [H+]old x (old volume / new volume)

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65
Q

How are weak acids different to strong ones?

A

.Weak acids do not dissociate fully like strong acids

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66
Q

What happens when a weak acid dissociates?

A

An equilibrium is established between the ions and the acid

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67
Q

As an equibrlium is established when weak acids dissociate, what can we calculate?

A

we can calculate an equilibrium constant Ka (Acids dissociation cons tat)

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68
Q

What does Ka equal for, HA H+ + A-?

A

Ka = ( [H+] [A-] ) / [HA]

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69
Q

For the Ka equation what two assumptions need to be made?

A

.A pure acids will split evenly into equal quantities of [h+] and [A-].The concentration of the acid doesn’t change as the equilibrium lies so far to the left it can be considered equal

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70
Q

does [H+] = [H+ aq]?

A

yes

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71
Q

.A sample of ethanoic acid of concentration 0.04 moldm-3.What is the value of Ka?

A

2.5x10-9 moldm-3

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72
Q

Steps to calculating pH of weak acid

A

.Step 1: Calculate [H+] from ka and [HA].Step 2: Calculate pH

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73
Q

Are the approximations for weak acids justified?

A

.The first approximation assumes that the dissociation from water is negligible.[H+(aq)]eqm = [A-(aq)]eqm - if pH>6 then water dissociates and [H+] from water will be more significant than from dissociation of acid.This approximation breaks down for very weak acids or very dilute solutions.Second approximation assumes that the concentration of the acid [HA] is much greater than the [H+] concentration at equilibrium.[HA]start&raquo_space; [H+]eqm [HA]eqm = [HA]start – [H+]eqm so [HA]eqm = [HA]start.Not valid as acid gets stronger as [H+] becomes more significant and real difference between [HA]eqm = [HA]start – [H+]eqm.Not justified for stronger weak acids / very dilute solutions.

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74
Q

.The value of Ka for weak acids is almost always very small and difficult to compare numbers with negative indices, how do we get around this?

A

create a scale of more useable numbers we often use the value of pKa, when talking about the aciditing of weak acids

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75
Q

What is pKa often used for?

A

to compare acids in biological systems

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76
Q

pKa equation

A

pKa = -log10Ka

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77
Q

What do the values of pKa mean?

A

The higher the value of pKa, the weaker the acidThe lower the value of pKa, the stronger the acid

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78
Q

How does wine show behaviour of typical dibasic and tribasic acids?

A

Wines often contain traces of sulphurous acid, H2SO3, added as a preservative.Sulphurous acid is dibasic and its dissociation is shown as:H2SO3 H+ + HSO3 - (pKa = 1.92)HSO3 - H+ + SO3 2- (pKa = 7.18).For the first dissociation, H2SO3, acts as a weak acid.From the pKa values, HSO3, is a far weaker acid than H2SO3.The behaviour is typical of dibasic and tribasic acids

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79
Q

Where strong acids completely dissociate to release all H+ ions into solution, weak acids only partially dissociate, explain this in terms of equations

A

Strong: HA → H+ + A-Weak: HA ⇌ H+ + A-

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80
Q

As weak acids form an equilibrium their dissociation can be represented by the acid dissociation constant, Ka, whats the equation for this?

A

Ka = “[H+(aq)][A-(aq)]” /”[HA(aq)]”

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81
Q

Explain why a larger value for Ka means a lower pH

A

.The larger the value for Ka the more the equilibria lies to the right and so the lower the pH i.e. higher [H+]

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82
Q

What alters Ka?

A

Temperature

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83
Q

.As the values for Ka are very small it is easier to give their negative logarithm, pKa, instead, what is the equation for this?

A

pKa = - logKa

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84
Q

pKa = - logKa, what is the inverse equation for this?

A

The inverse for this equation is: Ka = 10-pKa

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85
Q

This means that the weaker an acid the:._____ Ka._____ pKa

A

.Smaller Ka.Larger pKa

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86
Q

.The [H+] of a weak acid depends on the value of what?

A

Ka and [HA]

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87
Q

.The equilibrium concentrations, [ ]eqm, can be used to determine a value for Ka, what is the equation for this?

A

Ka = “[H+]eqm [A-]eqm” /”[HA]eqm”

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88
Q

There are two approximations that need to be made when calculating Ka, what are they?

A

.HA dissociation forms equal [H+] and [A-].The change in [HA] is negligible so [HA]eqm = [HA]start

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89
Q

Using the approximations, Ka can be determined by what equation?

A

Ka = “[H+]eqm 2” /”[HA]start”

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90
Q

How can pH be calculated for a weak acid?

A

[H+] = √(“Ka x [HA]” ) → pH = -log[H+]

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91
Q

.A value for Ka can be determined experimentally, how?

A

by using a pH meter to get the pH of a standard solution

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92
Q

.The larger the value for Ka, the greater the _________

A

dissociation

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93
Q

When calculating values for Ka there are issues with the approximations made, what are these issues?

A

• At pH values >6 water dissociation is significantTherefore doesn’t work for very weak acids or very dilute solutions• If [H+] concentration is significant there will be a difference between [HA]eqm and [HA]startTherefore doesn’t work stronger weak acids with Ka > 10-2 mol dm-3 or very dilute solutions

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94
Q

All aqueous solutions contain which ions?

A

H+ and OH- ions

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95
Q

H2O

A

H+ + OH-

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96
Q

In ____ [H+] > [OH-]In ____ [OH-] > [H+]In ____ [H+] = [OH-]

A

In acids [H+] > [OH-]In alkalis [OH-] > [H+]Neutral [H+] = [OH-]

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97
Q

For every 500,000,000 H2O molecules, only _ dissociates

A

1

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98
Q

For every 500,000,000 H2O molecules, only 1 dissociates, what does this mean for the equilbrium?

A

the equilibrium is on the left hand side

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99
Q

Do the Kc equation for water equilibrium

A

Kc = ([H+][OH-]) / [H2O]

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100
Q

[H2O] x Kc =

A

[H+] x [OH-]

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101
Q

[H2O] is such a large excess it can be classed as a ____

A

constant

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102
Q

What is Kw?

A

ionic product of water

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103
Q

Kw = [H2O] x Kc, so replace [H2O] x Kc with Kw in the correct equation

A

Kw = [H+] x [OH-]

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104
Q

At 298K, Kw =

A

1x10-14 mol2dm-6

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105
Q

Why is the pH of pure water at 298K, 7?

A

Kw = [H+] x [OH-]1x10-14 = [H+] x [OH-]The concentrations are the same and so it can be written as:1x10-14 = [H+]21x10-7 = [H+]pH = 7

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106
Q

Effect of temperature on Kw

A

The dissociation of water is endothermicTemperature increase will move the equilibrium to the right, and so Kw will increaseTemperature decrease will move the equilibrium to the left, and so Kw will decreaseWater will always remain neutral at all temperatures, pH may change with temperature, but [H+] = [OH-] all the time and so it is neutral

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107
Q

.Water ionises very slightly, acting as both an acid and as a base – setting up an ________

A

equilibrium

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108
Q

.Water dissociates a very, very small amount according to the equation – it must do, otherwise it would not _____________

A

conduct electricity

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109
Q

.1 dm3 (1000g) of water is mainly ________ H2O

A

undissociated

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110
Q

. [H2O(l)] = 1000/18 = 55.6 mol dm-3 (a constant), why?

A

.1 dm3 (1000g) of water is mainly undissociated H2O

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111
Q

.If we know the concentration of hydroxide ions we can rearrange Kw to give us the value of what?

A

[H+]

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112
Q

What is the pH of 0.4 mol dm-3 NaOH?

A

Kw = [H+][OH-] = 1.00 × 10–14 mol2dm–6 (1.00 × 10_14 )/([0.4])=[H+] = 2.5 x 10-14 mol dm-3 pH = -log10[H+] = -log10[2.5 x 10-14 ] = 13.60

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113
Q

What is the pH of a solution with [OH-] = 2 x 10-2 mol dm-3 at 25oC ?

A

Step 1 : calculate [H+] from Kw and [OH-] Kw = [H+][OH-] = 1.00 x 10-14 [H+] = Kw = 1.00 x 10-14 = 5.00 x 10-13 mol dm-3 [OH-] 2.00 x 10-2Step 2 : Use calculator to find pH pH = - log [H+] = -log (5.00 x 10-13) = 12.30

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114
Q

What are the concentrations of H+ (aq) and OH- (aq) in a solution of pH 3.25 at 25oC ?

A

Step 1 : Use calculator to find [H+(aq)][H+] = 10-pH = 10-3.25 = 5.62 x 10-4 mol dm-3Step 2 : Calculate [OH-] from Kw and [H+]Kw = [H+][OH-] = 1.00 x 10-14[OH-] = Kw = 1.00 x 10-14 = 1.78 x 10-11 mol dm-3 [H+] 5.62 x 10-4

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115
Q

For pH values that are whole numbers, it is easy to work out the [H+] and [OH-] concentrations as the indices add up to what?

A

-14

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116
Q

.The pH of weak bases can be calculated via a similar method to that used for ______

A

weak bases

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117
Q

.In an aqueous solution, there will always be both H+ (aq) and OH- (aq) ions present such that ________

A

[H+ (aq)][OH- (aq)] = Kw

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118
Q

.A solution is ______ when [H+ (aq)] > [OH- (aq)] .A solution is ______ when [H+ (aq)] = [OH- (aq)].A solution is ______ when [OH- (aq)] > [H+ (aq)]

A

.A solution is acidic when [H+ (aq)] > [OH- (aq)] .A solution is neutral when [H+ (aq)] = [OH- (aq)].A solution is alkaline when [OH- (aq)] > [H+ (aq)]

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119
Q

.So a solution that is acidic will still contain __ ions, it is just that there are more __ions (and vice versa in an ______ solution)

A

OH-H+Alkaline

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120
Q

.The value of Kw controls the ______ of each ion

A

Concentrations

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121
Q

• Kw can also tell us the pH of ______

A

pure water

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122
Q

• As the water splits into equal concentrations of OH- and H+ ions we can make what assumption?

A

Kw = [H+][OH-] = 1.00 × 10–14 mol2dm–6 = [H+]2 = 1.00 × 10–14 = [H+] = 1.00 × 10–7 pH = -log10[H+] = -log10[1.00 x 10-7 ] = 7

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123
Q

When is neutral not neutral?

A

.A neutral solution is defined by a equal number of moles of H+ ions and OH- NOT by a pH of 7 as you may have been taught.Whilst this value is 7 at approximately 298K, the value of Kw increases with temperature

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124
Q

.The pOH scale measures what?

A

the concentration of hydroxide ions

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125
Q

.Therefore you can solve fpr pH, pOH and [OH-] given just ___

A

[H+]

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126
Q

pH + pOH =

A

14

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127
Q

.Kw = [H+][OH-] =

A

1x10-14

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128
Q

Define Enthalpy of reaction ΔrH

A

enthalpy change that accompanies a reaction in the molar quantities shown in a chemical equation under standard conditions, with all reactants and products in their standard states

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129
Q

Define Enthalpy of formation ΔfH

A

enthalpy change when one mole of compound is formed from its constituent elements under standard conditions

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130
Q

Define Enthalpy of combustion ΔcH

A

enthalpy change when one mole of substance is burnt completely in excess oxygen under standard conditions

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131
Q

Water slightly ionises, what equation will show this?

A

H2O(l) H+(aq) + -OH(aq)

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132
Q

Ka of :H2O(l) H+(aq) + -OH(aq)=

A

[H+][-OH] / [H2O]`

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133
Q

Ka x [H2O] =

A

Kw

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134
Q

Kw =

A

[H+][-OH]

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135
Q

What is Kw?

A

Kw is the ionic product of water, at 25’C it equals 1x10-14 mol2dm-6

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136
Q

How does ionic bonding work?

A
  1. Loss of an electron(s) by an element2. Gain electrons by a second element3. Attraction between positive and negative ions
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137
Q

Na + Cl, ionisation electron addinity and latice equations

A

Ionisation Energy – Na  e- + Na+ +496kJmol-1Electron Affinity – e- + Cl  Cl- -349kJmol-1Lattice Enthalpy – Cl- + Na+  NaCl -766kJmol-1

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138
Q

What is electron affinity?

A

.The enthalpy change when one mole of electron is added to one mole of atoms in the gaseous phase to form one mole of -1 ions

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139
Q

Is repulsion between two electrons exothermic or endothermic, why?

A

.Repulsion between two negatively charged things requires energy so exothermic

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140
Q

What is lattice enthalpy?

A

.ΔHlatt is the enthalpy change when 1 mole of ionic substance is formed from its gaseous ions under standard conditions

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141
Q

.The strength of the ionic bond is related to the lattice enthalpy, how?

A

more exothermic the greater the ionic bonding

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142
Q

Lattice enthalpy can’t be measured, why?

A

Cannot be measured directly as cannot form one mole of ionic lattice from gaseous ions

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143
Q

What are the factors affecting lattice enthalpy? How?

A

Charge – .The greater the charge on the ions, the stronger the attraction – therefore, more exothermic lattice enthalpySize – .Smaller ions can pack together more tightly, therefore there is greater attraction and more exothermic lattice enthalpy

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144
Q

Why, when moving down group two, do the temperatures required to break the lattice increase?

A

.They have decreasing charge densities.Mg2+ is a smaller ion than Ba2+ , so the +2 charge occupies a smaller volume – this means Mg2+ has a higher charge density than Ba2+.Mg2+ can distort the electron clouds within the CO32- ion (called polarisation), this weakens the covalent bonding in the ion and reduces its decomposition temperature

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145
Q

What is a buffer?

A

a solution that minimises pH change when a small amount of acid or alkali is added

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146
Q

What two types of buffers are they?

A

Weak Acid and Salt of Weak AcidExcess Weak Acid and Strong Base

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147
Q

Suggest and explain a Weak Acid and Salt of Weak Acid buffer

A

Weak Acid – Ethanoic Acid – CH3COOH(aq) CH3COO-(aq) + H+(aq)Salt of Weak Acid – Sodium Ethanoate – CH3COONa(s) + (aq) CH3COO-(aq) + Na+(aq)¬Buffer Contains – CH3COOH(aq) CH3COO-(aq) + H+(aq)When add H+, H+ reacts with conjugate base, equilibrium will move to the left to reduce the amount of H+, pH is constantAdding –OH (alkali) –-OH + H+  H2OConc of H+ decreases, equilibrium moves to the right to increase the conc of H+, pH is constant

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148
Q

Suggest and explain an excess Weak Acid and strong base buffer

A

CH3COOH(aq) + NaOH(aq) –> CH3COONa(aq) + H2O(l)CH3COONa(aq) CH3COO-(aq) + Na+(aq)CH3COOH(aq) CH3COO-(aq) + H+(aq)When add H+, H+ reacts with conjugate base, equilibrium will move to the left to reduce the amount of H+, pH is constantAdding –OH (alkali) –-OH + H+  H2OConc of H+ decreases, equilibrium moves to the right to increase the conc of H+, pH is constant

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149
Q

Excess methanoic acid is reacted with potassium hydroxide, explain how a buffer solution is produced and how pH is controlled when 5cm3 of HCl is added.

A

The solution now contains HCOOK(aq) HCOO- (aq) + K+(aq) and HCOOH(aq) HCOO-(aq) + H+ equilibriums in solution – the buffer. When 5cm3 of HCl is added, it dissociates into H+ and Cl-, the H+ ions increase the concentration of the H+ already in the solution, so the HCOOH(aq) HCOO-(aq) + H+ point of equilibrium moves to the right – decreasing the conc of H+ in solution, and so keeping the pH constant.

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150
Q

Ka x ( [HA]/[A-] ) = ? what does each part mean?

A

[H+] = Ka x ( [HA]/[A-] )[HA] = concentration of weak acid[A-] = concentration of conjugate base

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151
Q

50cm3 of 1.2M NaOH reacts with 250cm3 of ethanoic acid (1M), Ka = 1.74x10-5, what is the pH?

A

CH3COOH CH3OO- + H+CH3COOH + NaOH  CH3COO- + Na+ + H2O1. Moles weak acid = 250x1 / 1000 = 0.25mol2. Moles NaOH = 50x1.2 / 1000 = 0.06mol3. Moles A- = 0.06mol4. Moles HA = 0.19mol5. [H+] = Ka ([HA] / [A-]) = 1.74x10-5 x (0.19/0.06) = 5.51x10-56. pH = 4.26

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152
Q

Equation to work out lattice enthalpy from born haber cycle

A

L.E = FORMATION – sum(ATOM + I.E + E.A)orL.E = FORMATION – sum(REST)

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153
Q

What are Born Haber Cycles used for?

A

Born Haber cycles can be used to calculate a measure of ionic bond strength based on experimental data

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154
Q

Draw the born haber cycle for NaCl

A

check notes or google

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155
Q

Draw the born haber cycle for MgCl2

A

check notes or google

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156
Q

Draw the born haber cycle for CuO

A

check notes or google

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157
Q

Blood must contain a pH of what?

A

Blood must contain a pH of 7.40 +- 0.05

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158
Q

What is the most important buffer in blood?

A

the carbonic acid – hydrogencarbonate

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159
Q

Carbonic acid dissociation equation

A

H2CO3 H+ + HCO3-

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160
Q

If blood becomes too acidic (acidosis), what presents?

A

fatigue, shortness of breath, shock, death

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161
Q

If blood becomes too alkaline (alkalosis), what presents?

A

spasms, light-headed, nausea

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162
Q

What happens if the body produces more acidic products?

A

.The H+ concentration in the blood will increase, the equilibrium of the hydrogencarbonate dissociation will shift to the left, the H+ ions will be used up to form more hydrogencarbonate, the H+ concentration will return to normal, the pH will have had minimal change.If the acid level continued to rise, fatigue, shortness of breath, shock, or death would eventually set in.This would occur when we have ran out of HCO3- ions

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163
Q

Carbon dioxide dissolves in water to form what? What is the equation for this?

A

carbonic acid:CO2 + H2O H2CO3

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164
Q

Inhalation of high levels of CO2 means what?

A

Inhalation of high levels of CO2 mean that the equilibrium in the blood must shift to the right to form more H2CO3, this means that the level of H2CO3 in the blood increase, so in the equilibrium H2CO3 H+ + HCO3- the equilibrium shifts to the right to counter this, this forms more H+ ions in the blood, making it more acidic, and resulting in acidosis which would lead to fatigue, shortness of breath, shock, and eventual death. Increased [H+] means enzymes would also be denatured.

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165
Q

Define enthalpy of solution, with an example

A

DeltasolutionH is the enthalpy change when one mole of ionic compound is completely dissolved in water under standard conditions.e.g. NaCl(s)  Na+(aq) + Cl-(aq)

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166
Q

What is enthalpy of hydration with example

A
  • Delta Hhydration¬ is the enthalpy which takes place when one mole of gaseous ions is dissolved in water forming one mole of aqueous ions under standard conditions- Na+(g) + aq  Na+(aq)
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167
Q

Why is enthalpy of hydration always exothermic?

A

These are exothermic as bonds are formed between the ions and water molecules

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168
Q

Factors Affecting Enthalpy of Hydration

A

Charge – the higher charge on the ion, the greater the attraction for the H2O molecules, therefore a more exothermic hydrationSize – smaller ions have a greater charge density compared to the larger ions, this creates a greater attraction for H2O molecules, therefore a more exothermic hydration

169
Q

Draw the general cycle of enthalpy of solution, hydration and lattice enthalpy

A

check notes

170
Q

Equation that links lattice enthalpy, enthalpy of solution, and enthalpy of hydration

A

lattice enthalpy + enthalpy of solution = sum of the enthalpy of hydrations

171
Q

Draw the titration curve for when HCl is added to NaOH

A

check notes

172
Q

Draw the titration curve for when NaOH is added to HCl

A

check notes

173
Q

Draw strong acid strong base titration curve

A

check notes

174
Q

Draw strong acid weak base titration curve

A

check notes

175
Q

Draw weak acid strong base titration curve

A

check notes

176
Q

Draw weak acid weak base titration curve

A

check notes

177
Q

The shape of the titration curve depends on what?

A
  • Substance being titrated i.e. acid or base- Type of acid or base i.e. strong or weak
178
Q

Describe and explain parts of titration curve

A
  1. Excess of base: pH decreases as acid added2. Vertical section: acid base concentration similar, pH alters rapidly3. Excess of acid: pH decreases slightly as acid is added4. The equivalence point: halfway up the vertical section
179
Q

What is the equivalence point?

A

The point halfway between the two horizontal lines on the titration curveThe equivalence point is the volume required to have completely reacted the acid and base toetehr, stoichiometric volumesThe pH at the equivalence point depend son the type of acid and base

180
Q

Where is the equivalence point on each titration curve?

A

Equivalence point = 7 on SA-SBEquivalence point < 7 on SA-WBEquivalence point > 7 on WA-SBNo equivalence point on WA-WB

181
Q

Why is the equivalence point 0 on SASB titration curves?

A
  • H+ ions completely neutralised by OH- ions, so only water and salt present, neutral
182
Q

Why is the equivalence point < 7 on SAWB titration curves?

A
  • A weak base, like NH3, will have a strong conjugate acid, NH4+ , which will react with water to produce H3O+ so the pH is less than 7
183
Q

Why is the equivalence point > 7 on WASB titration curves?

A
  • A weak acid, HA, will have a strong conjugate base, A-, which can react with water to produce OH- ions, so the pH is more than 7
184
Q

What are acid-base indicators?

A

Acid-Base indicators are weak acids that have different coloured conjugate basesIn acidic conditions the indicator equilibrium is shifted towards the weak acid (HIn)As the system becomes more basic the equilibrium shifts towards the conjugate base (In-), altering the colour

185
Q

What is the end point?

A

The end point is when equal [HIn] and [In-] are present, the colour will therefore lie between both extremesEach indicator will have a different pH value for the end point as they each have different ka values

186
Q

HIn

A

H+ + In-

187
Q

How is an acid-base indicator chosen?

A

The indicator required for a titration must have a colour change that lies within the vertical section of a titration curve

188
Q

A chemical reaction will proceed when?

A

A chemical reaction will proceed if the products are energetically more stable than the reactants

189
Q

What is entropy?

A

Entropy is a measure of the dispersal of energy in a system, the more disordered a system the greater the dispersal of energy = higher entropywe define entropy as a measure of disorder

190
Q

What happens to entropy over time?

A

Entropy must increase over time

191
Q

When is a system in a state of high entorpy?

A

When its degree of disorder is high

192
Q

As order within a system increases, its entropy decreases, why?

A

This can be explained in terms of probability: disordered states are simply more likely to exist (or emerge) than ordered states.The spontaneous direction of change is from a less probable to a more probable state

193
Q

What does entropy always do?

A

The total entropy always increases, and the process is irreversible

194
Q

What is the unit for entropy

A

Sin JK^-1mol^-1

195
Q

Why is S (entropy) alwasy potisive?

A

All substances process some degree of disorder because particles are always in constant motion

196
Q

How does the entropy of each state differ?

A

Solid has lowest entropy, and gas has the highest

197
Q

Draw a graph of waters change in enthalpy as temperaurer increases

A

check notes

198
Q

Systems that are more chaotic have a _____ entropy value

A

higher

199
Q

(s)  (s) + (g) , what is delta S

A

+

200
Q

(g)  2(g) , what is delta S

A

+

201
Q

What is delta S?C2H5OH(l)  C2H5OH(g)

A

+

202
Q

What is delta S?C2H2(g) + 2H¬2(g)  C2H6(g)

A

-

203
Q

What is delta S?NH4Cl(s) + aq  NH4Cl(aq)

A

+

204
Q

What is delta S?4Na(s) + O2(g)  2Na2O(s) deltaS

A

-

205
Q

Define the standard entropy change

A

The standard entropy change is the entropy change that accompanies a reaction in the molar quantities expressed in the equation, under standard conditions

206
Q

Sum for delta S^theta

A

delta S^theta = sum(S^theta products) – sum(S^theta reactants)

207
Q

Why are exothermic reactions more preferable in nature?

A

as the products are more stable than reactantsthe key is not the decrease in energy but the associated increase in entropy of the surroundings.

208
Q

deltaSsurroundings is proportional to what?

A

NAME?

209
Q

deltaSsurroundings =

A

(-deltaHsystem)/T

210
Q

deltaStotal =

A

deltaSsystem + deltaSsurroundings

211
Q

deltaStotal >

A

0

212
Q

What can we learn about the entropy of the sun by the photosynthesis equation?

A

6CO2(g) + 6H2O(l)  UV light  C6H12O6(s) + 6O2(g)Negative entropy changeEntropy change in sun must be so positive it outweighs every plant on earth

213
Q

What is gibbs free energy equation?

A

deltaG = deltaH – T deltaS < 0

214
Q

How do you get to gibbs free energy equation?

A

deltaStotal¬ = deltaSsystem + deltaSsurroundings > 0deltaStotal¬ = deltaSsystem – (deltaHsystem)/T > 0T deltaStotal = T deltaSsystem – deltaHsystem > 0-T deltaStotal = -T deltaSsytem + deltaHsystem < 0deltaG = deltaH – T deltaS < 0

215
Q

Using deltaG explain when a reaction is and is not feasible

A

deltaG must be negative (<0) for a reaction to be feasible i.e. proceedIf deltaG is positive (>0) then a reaction is not feasible

216
Q

What is deltaG at the points of feasibility?

A

At the point of feasibility, we can say deltaG = 0 (assume that deltaH and deltaS don’t vary with temperature)

217
Q

At low temperatures what does deltaGsystem equal? why?

A
  • At low temperatures, deltaGsystem = deltaH (-T deltaS becomes negligible) so for a reaction to occur it needs to be exothermic
218
Q

At high tempertaures deltaGsystem ewuals what? why?

A
  • At high temperatures, deltaGsystem = -T deltaS (deltaH becomes negligible) so for a reaction to occur it needs to have a positive deltaS as – T deltaS needs to be less than 0
219
Q

Limitations to deltaG equation

A

just because the value is negative and so feasible it doesn’t mean it occurs, the reaction rate might be incredibly slow or the activation energy too high

220
Q

What is an oxidation number?

A

An oxidation number shows the charge of an atom if all of its bonds were considered totally ionic

221
Q

What are the oxidation number rules?

A
  1. Elements in their natural state = 02. The total oxidation states in a molecule = 03. The total oxidation states in an ion = the chargeGroups 1, 2, 3 - State +1, +2, +3Fluorine - State -1Hydrogen -State +1-1 in metal hydridesOxygen - State -2-1 in peroxidesChlorine -State -1+ (varies) in chlorates
222
Q

What is an oxidising agent?

A

takes electrons from what is oxidised

223
Q

What is a reducing agent?

A

gives electrons to what is reduced

224
Q

Write the half equations and full equation for 4Na + O2  2Na2O

A

4Na  4Na+ + 4e-O2 + 4e-  2O2-4Na + O2  4Na+ + 2O2-

225
Q

Write the half equations for Cu2O + H2SO4  Cu + CuSO4 + H2O

A

Cu+  Cu2+ + e-2Cu+  Cu + Cu2+

226
Q

Steps of a redox titration using iron sulphate and sulphuric acid and potassium manganate (VII)

A
  1. Make up a 150cm3 standard solution using 7g of iron sulphate by weighing by difference in a volumetric flask, record your mass measurements in a table2. a. Pipette 25cm3 of the standard solution into a conical flaskb. Add 10cm3 of 1M H2SO4 to the conical flaskc. Fill the burette with a standard solution of 0.02M potassium manganate (VII)d. Carry out the titrations until the ed point is reached, the first permanent pink colour, record your results in a tablee. Repeat until you have two concordant resultsf. Calculate the mean titre
227
Q

Reduction: MnO4- + 8H+ + 5e-  Mn2+ + 4H2OOxidation: Fe2+  Fe3+ + e- or 5Fe2+  5Fe3+ + 5e-MNO4- + 8H+ + 5Fe2+  Mn2+ + 4H2O + 5Fe3+What is the oxidising and reducing agent?

A

Fe2+ is the reducing agentMnO4- is the oxidizing agent

228
Q

Mass of impure hydrated iron II sulphate = 6.97g25cm3 of hydrated iron (II) sulphate used taken from a 250cm3 volumetric flask full of the dissolved iron tabletsMean titre of 0.02M potassium manganate (VII) = 23.4cm3Calculate the percentage purity of the impure sample of FeSO4.7H2O

A

93.30%

229
Q

What are the redox equations for the iodine/thiosulphate titrations?

A

I2 + 2e-  2I-2S2O32-  S4O62- + 2e-I2 + 2S2O32-  2I- + S4O62-

230
Q

What colour is iodine?

A

Iodine is a yellow brown solution

231
Q

What colour is iodide?

A

Iodide is a straw colour

232
Q

How do you know when an iodine/thiosulphate titration is complete?

A

Iodine is a yellow brown solutionIodide is a straw colourThe above colour change is the indication for the redox reaction to be complete

233
Q

Is an anode positive or negative?

A

Positive

234
Q

Is a cathode positive or negative?

A

Negative

235
Q

Draw a general cell

A

check notes

236
Q

What is a cell? What is it made of? What does it do?

A
  • A cell has two half cells- The two half cells must be connected with a salt bridge- Simple half cells will consist of a metal (acts as an electrode) and a solution of a compound containing that metalo For example, Cu and CuSO4- These two half cells will produce a small voltage if connected into a circuit (become a battery or cell)
237
Q

Draw the standard hydrogen half-cell

A

Check notes

238
Q

What is the standard hydrogen half-cell?

A
  • H+/H2 half-cell is chosen to produce standard potentials- H+ + e-  ½H2 - 1/2H2  H+ + e– Conditionso 298Ko 1atmo 1M H+- Both equations have a voltage of 0.0 volts
239
Q

Why is platinum used in the standard hydrogen half-cell?

A

Platinum is inert and so does not take part in the reaction

240
Q

How can you find the value of any electron potential?

A

If the standard hydrogen half-cell is connected to another cell, you can find the value of any electron potential (as the standard is0.0 volts).

241
Q

Why a salt bridge? Using potassium chloride and copper as an example

A
  • The salt bridge is used to connect up the circuit- The free moving ions conduct the charge- A salt bridge is usually made from a piece of filter paper (or material) soaked in a salt solution, usually potassium nitrate- The salt should be unreactive with the electrodes and electrode solutionso For example, potassium chloride would not be suitable for copper systems as chloride ions can form complexes with copper ions, a wire is not used because the metal wire would set up its own electrode system with the solutions
242
Q

What is standard electrode potential?

A

The emf of a half cell compared with a standard hydrogen electrode at standard conditions of 298K, concentration 1.0M and 1 atm of pressure.

243
Q

What is emf?

A

The emf is the difference in the positive and negative sides of the cell.

244
Q

Describe a two ion cell system?

A

Both of the ions in the equation will be at a concentration of 1M, platinum electrode is used as it is INERT.2 aqueous ions in a half equation, means they both have to be 1.0M solutions in the same half cell

245
Q

Draw a manganese/lead system

A

check notes

246
Q

General properties of d-block elements

A

The d-block elements have high melting and boiling points.The d-block elements are good conductors of both electricity and heat.

247
Q

What elements have been used in coinage for many years?

A

copper, silver, nickel, and zinc

248
Q

What element is used extensively in construction and production of tools?

A

Iron

249
Q

What element is used for electrical cables and water pipes?

A

Copper

250
Q

Titanium has great strength, what are the elements applications?

A

It has many aerospace and medical applications (for example joint replacement).

251
Q

Sc orbital configuration

A

[Ar] 4s2 3d1

252
Q

Fe orbital configuration

A

Fe – [Ar] 4s2 3d6

253
Q

Ni orbital configuration

A

Ni – [Ar] 4s2 3d8

254
Q

Zn orbital configuration

A

Zn – [Ar] 4s2 3d10

255
Q

Cr orbital configuration

A

Cr – [Ar] 4s1 3d5

256
Q

Cu orbital configuration

A

Cu – [Ar] 4s1 3d10

257
Q

How do copper and chromium minimize repulsions?

A

Copper and chromium minimize repulsions by being half full or full, chromium has 4s and 3d orbitals half full, copper has 3d full

258
Q

Fe2+ orbital configuration

A

Fe2+ - [Ar] 4s0 3d6

259
Q

Fe3+ orbital configuration

A

Fe3+ - [Ar] 4s0 3d5

260
Q

Cu2+ orbital configuration

A

Cu2+ - [Ar] 4s0 3d9

261
Q

Cr3+ orbital configuration

A

Cr3+ - [Ar] 4s0 3d3

262
Q

Mn2+ orbital configuration

A

Mn2+ - [Ar] 4s0 3d5

263
Q

Mn4+ orbital configuration

A

Mn4+ - [Ar] 4s0 3d3

264
Q

Sc3+ orbital configuration

A

Sc3+ - [Ar] 4s0 3d0

265
Q

Zn2+ orbital configuration

A

Zn2+ - [Ar] 4s0 3d10

266
Q

Define a transition element

A

A transition element is a d-block element that forms at least one ion with an incomplete d sub-shell.

267
Q

Which ‘d’ block elements do not fit the transition element definition?

A
  • Scandium and zinc
268
Q

What are some characteristic properties of transition metals and there compounds, with examples. (not conduction, melting, boiling etc.)

A
  • They form compounds in which the transition element has different oxidative stateso Fe2+ = +2o Fe3+ = +3- They form colored compounds- The elements and their compounds can act as catalystso Fe in Haber processo Ni in hydrogenation of alkenes
269
Q

A species containing a transition element in its highest oxidation state is often a what?

A

strong oxidizing agent

270
Q

What does the observed colour of a solution depend on?

A

The observed colour of a solution depends on the wavelengths absorbed

271
Q

Why does copper sulphate solution appear blue?

A

Copper sulphate solution appears blue because the energy absorbed corresponds to red and yellow wavelengths, wavelengths corresponding to blue light aren’t absorbed.

272
Q

Draw the exam colour chart

A

check notes

273
Q

If the colour we observe is red, what colour has been absorbed?

A

Cyan

274
Q

What colour is cu2+

A

white

275
Q

What colour is [Cu(H2O)6]2+

A

blue

276
Q

Transition metals form complex ions or ?

A

coordination compounds

277
Q

WHat do transition metals form when they make complex ions?

A

ligand forms bonds with the central transition metal ion

278
Q

Examples of complex ions

A

[Cr(H2O)6]3+ [CuCl4]2-

279
Q

What is a ligand?

A

a molecule or ion that can donate a pair of electrons with the transition metal ion to form a coordinate bond

280
Q

What does monodentate mean?

A

Monodentate ‘one tooth’ means each ligand donates just one pair of electrons.

281
Q

Examples of ligands and how many loan pairs they have

A

H2O (oxygen has two loan pairs) NH3 (nitrogen has a loan pair) Cl- (chloride has a loan pair)CN- (carbon has a loan pair) OH- (oxygen ha a loan pair)

282
Q

Draw [Cu(H2O)6]2+

A

check notes

283
Q

Describe the structure and bonding in [Cu(H2O)6]2+

A
  • The central ion is Cu2+- The ligands are water moleculeso Each molecule donates a pair of electrons from the O atom to the Cu2+ to form a co-ordinate bond- The co-ordination number is 6o This indicates the number of coordinate bonds to the central metal ion
284
Q

What is the oxidation of Co in [Co(H2O)5Cl]+

A

2

285
Q

How to name complex ions?

A
  • The name gives the metal ions and its oxidation state last, and the name/number of ligands before- Pre-fixes di, tri, tetra, penta, hexa used- Ligands are listed alphabetically, with prefixes not allowed to alter this order
286
Q

[Cr(H2O)4Cl2]+ name

A

¬¬Tetraaquadichlorochromium (III) ion

287
Q

[Co(H2O)5Cl]+ name

A

Pentaaquamonochlorocobalt (II) ion

288
Q

If the complex ion is an anion, what do you do?

A

the suffix ‘-ate’ follows the metal

289
Q

Name of [Fe(CN)6]4-

A

Hexacyanoferrate (II) ion

290
Q

What does cobalt become in an anion complex ion?

A

cobaltate

291
Q

What does aluminium become in an anion complex ion?

A

aluminate

292
Q

What does chromium become in an anion complex ion?

A

chromate

293
Q

What does vanadium become in an anion complex ion?

A

vanadate

294
Q

What does copper become in an anion complex ion?

A

cuprate

295
Q

What does iron become in an anion complex ion?

A

ferrate

296
Q

What does nickel become in an anion complex ion?

A

nickelate

297
Q

What does the value of pKa show?

A

The higher the value of pKa, the weaker the acidThe lower the value of pKa, the stronger the acid

298
Q

Equation for finding hydrogen ion concentration in buffer

A

[H+] = ka x ([HA]/[A-])

299
Q

In an electrochemical cell, where does oxidation happen?

A

the half-cell with the most negative standard electrode potential

300
Q

What is the relationship between electrode potential and metal reactivity?

A

More negative electrode potential = more reactive the metal

301
Q

Which equation links E (cell), E(positive electrode), and E(negative electrode)

A

E (cell) = Eo (positive electrode) – Eo (negative electrode)

302
Q

The more reactive a metal is, the more likely it is to what?

A

to lose electrons and form a positive ion

303
Q

More reactive metals have more negative electrode potentials and vice versa for what?

A

Non-metals

304
Q

in general if the electrode potential for the reaction is positive it is regarded as being what?

A

feasible

305
Q

What is the problem with predicting the feasibility of electrode reactions?

A
  • The value of E says something about the feasibility of the reaction under standard conditions only- The value of E says something about the feasibility of the reaction, but does not say anything about the rate of the reaction
306
Q

What is the anticlockwise rule?

A

Arrange the redox half equations so that:- The electrons are on the left- The largest negative electrode potentials are at the top- The reaction will then take place anticlockwise around the half equations- E (cell) = E (bottom) – E(top)

307
Q

When does predicting feasibility of electrode reactions go wrong?

A

When the conditions are not standard, this includes:- A change in concentration- If the half equations are in equilibriumIf the kinetics are not favorable:- The rate of reaction may be slow so reaction does not appear to happen- If the reaction has a high activation energy

308
Q

Define catalyst

A

a substance that increases the rate of a chemical reaction by providing an alternative reaction pathway of lower activation energy (Ea)

309
Q

What happens when Al3+, Ca2+, Mg2+ Add NaOH or KOH or NH4OH

A

white precipitate

310
Q

Description of solution, observation with NaOH, and equation for Copper Cu2+

A

Transparent blue Pale blue precipitateInsoluble in excess Cu2+(aq) + 2OH-(aq)  Cu(OH)2(s)

311
Q

Description of solution, observation with NaOH, and equation for Iron(II) Fe2+

A

Pale green Dark green precipitateTurns brown on contact with airInsoluble in excess Fe2+¬(aq) + 2OH-(aq)  Fe(OH)2(s)

312
Q

Description of solution, observation with NaOH, and equation for Iron(III) Fe3+

A

Orange/brown Orange/brown precipitateInsoluble in excess Fe3+(aq) + 3OH-(aq)  Fe(OH)3(s)

313
Q

Description of solution, observation with NaOH, and equation for Chromium(III) Cr3+

A

Violet Grey-green precipitateSoluble in excess giving dark green solution Cr3+ + 3OH-(aq)  Cr(OH)3(s)Cr(OH)3(s) + 3OH-(aq)  [Cr(OH6]3-(aq)

314
Q

Description of solution, observation with NaOH, and equation for Manganese(II) Mn2+

A

Pale pink Off white precipitateRapidly turning brown on contact with airInsoluble in excess Mn2+(aq) + 2OH-(aq)  Mn(OH)2(s)

315
Q
  • Cu2+(aq) + 2e- Cu(s)If we increase the concentration of Cu2+ ions then:
A
  • Equilibrium moves to oppose the charge- Electrons are removed from the system- The electrode potential becomes more positive
316
Q

What shape are four coordinate complex ions usualy? with examples

A
  • Tetrahedral is the most common shape- E.g. [CuCl4]2- and [CoCl4]2-
317
Q

When not tertrahedral, what shape are 4 coordinate complexes, with examples

A
  • Some 4 co-ordinate complex ions are square planar in shape, with the ligands arranged at the corners of a square- E.g. [Ni(NH3)2Cl2] (cis and trans)
318
Q

When does a square planar shape take place in complex ions? example

A
  • These occur in complexes with 8-d electrons in the d subshell.- E.g. Pt(II), Pd(II), Au(III)
319
Q

What is cis platin used in?

A
  • Cancer treatment in testicular cancer and useful for ovarian, head and neck, and lung cancer- Extremely toxic
320
Q

Why did they change from cisplatin to carboplatin?

A
  • Improved chemical stability relative to cisplatin due to chelation by cyclobutane dicarboxylic acid- Essentially equivalent antitumour activity to cisplatin
321
Q

WHat is oaplatin used in?

A

treatment of colorectal cancer

322
Q

Types of cancer therapy

A
  • Surgery- Radiotherapy- Chemotherapyo Cytotoxico Targeted Anti-endocrine Novel targeted agents- Immuno-therapy- Gene therapy
323
Q

What is mustard Gas? What came about because of mustard gas?

A
  • Potent vesicant agent that burns eyes, skin and respiratory tractMustard Gas = war gasNitrogen Mustard = anticancer drug
324
Q

How does cisplatin treat cancer?

A
  • The cisplatin binds to DNA and causes a critical structural change n the DNA – a bend of 45 degrees- This stops cell replication and leads to apoptosis (cell death)
325
Q

Key points of lactic acid

A
  • 2-hydroxy propanoic acid- Also known as lactic acid- The second carbon is a chiral center- The mirror image is non-superimposableo One is found in sour milko The other is found in anaerobic respiration
326
Q

What is a fuel cell?

A

a fuel cell is a device that converts chemical energy into electrical energy, water, and heat through electrochemical reactions.

327
Q

How ddoes a fuel cell work? Are there many or one usually and why?

A

• Fuel and air react when they come into contact through a porous membrane (electrolyte) which separates them• This reaction results in a transfer of electrons and ions across the electrolyte from the anode to the cathode• If an external load is attached to this arrangement, a complete circuit is formed and a voltage is generated from the flow of electrical currentThe voltage generated by a single cell is typically rather small (< 1 volt), so many cells are connected in series to create a useful voltage.

328
Q

Differences between a fuel cell and a battery

A

Hydrogen Fuel Cell – • Open system• Anode and cathode are gases in contact with a platinum catalyst• Reactants are externally supplied, no recharging requiredGalvanic Cell (Battery) – • Closed system• Anode and cathode are metals• Reactants are internally consumed, need periodic recharging

329
Q

Differences between a fuel cell and an internal combustion engine

A

Fuel Cell – • Output is electrical work• Fuel and oxidant react electrochemically• Little to no pollution producedInternal Combustion Engine – • Output is mechanical work• Fuel and oxidant react combustively• Use of fossil fuels can produce significant pollution

330
Q

Similarities between a fuel cell and in internal combustion engine

A

• Both use hydrogen-rich fuel• Both use compressed air as the oxidant• Both require cooling

331
Q

Draw an alkaline fuel cell

A

check notes

332
Q

What are the half equations, volatages, cell potential, and overall equation for an alkaline fuel cell?

A

Half Equations – 2H20 (l) + 2e-  H2 (g) + 2 OH- (aq) E = -0.83V½ O2 (g) + 2e-  2 OH- (aq) E = +0.40VCell Potential – = 0.4- -0.83 = 1.23VOverall Equation – H2 + ½ O2  H2O

333
Q

Why Methanol not Hydrogen?

A

Some new fuel cells use methanol rather than hydrogen as the fuel because- Liquid methanol is easier to store then hydrogen gas- Methanol can be generated from biomass

334
Q

Define electrochemical reaction

A

A reaction involving the transfer of electrons from one chemical substance to another

335
Q

Define electrode

A

An electrical terminal that conducts an electric current into or out of a fuel cell (where the electrochemical reaction occurs).

336
Q

Define electrolyte

A

A chemical compound that conducts ions from one electrode to the other

337
Q

What is an electrochemical cell consisted of?

A

An electrochemical cell consists of 2 electrodes + 1 electrolyte

338
Q

What is rhe equation for the number of optical isomers

A
  • The number of isomers follows the equation: 2n, where n is the number of chiral centres
339
Q

Why are optical isomers called such?

A

The isomers are called optical isomers as they can rotate plane-polarized light (light which only travels in one plane)

340
Q

WHat makes a 6 co-ordinate compound cis or trans?

A

The bond angle between the two ligands which are different, e.g. 90’ cis, 180’ trans

341
Q

What is a bidentate ligand? Example

A

Bidentate – ‘two tooth’ ligandsMost common is ethane-1, 2-diamine

342
Q

Draw [Ni(NH2CH2CH2NH2)3]2+

A

Check notes

343
Q

Describe and explain EDTA4-

A

A Hexadentate Ligand- EDTA4- has 6 lone pairs, each of which can form a co-ordinate bond- 1 EDTA ion reacts with 1 metal ion- Ethylenediaminetetraacetic acid (EDTA)

344
Q

What is ligand substitution?

A

The addition of another ligand to a solution containing the aqua transition metal ion results in a substitution reaction

345
Q

What occurs for ligand substitution?

A
  • One or more ligands is exchanged for another- A change in colour of the solution is observed- Sometimes the complex ion changes shape/coordinate number
346
Q

Colour of [Cu(H2O)6]2+, and its colour with dropwise and excess, ammonia and hydrochloric acid

A

Complex Ion [Cu(H2O)6]2+Pale blue solution Addition of Ammonia Dropwise: Pale blue precipitation of copper(II) hydroxideExcess: Blue precipitate redissolves, forming a deep blue solutionAddition of Concentrated HClDropwise: Begins to turn greenExcess: Begins to turn yellow

347
Q

Colour of [Cr(H2O)6]3+, and its colour with dropwise and excess, ammonia

A

Complex ion [Cr(H2O)6]3+Violet solutionAmmonia Dropwise: Grey/green precipitateExcess: Precipitate redissolves to produce a purple solution

348
Q

Equation for addition of NaOH to [Cu(H2O)6]2+

A

[Cu(H2O)6]2+ + 2OH-  [Cu(H2O)4(OH)2]2+ + 2H2O

349
Q

Equation for addition of NH3 to [Cu(H2O)6]2+ in dropwise and excess

A

[Cu(H2O)6]2+ + 2NH3  [Cu(H2O)4(OH)2]2+ + 2NH4+[Cu(H2O)6]2+ + 4NH3  [Cu(NH3)4(H2O)2]2+ + 4H2O

350
Q

Equation for addition of HCl to [Cu(H2O)6]2+

A

[Cu(H2O)6]2+ + 4Cl- [CuCl4]2- + 6H2O

351
Q

Equation for addition of ammonia to [Cr(H2O)6]3+

A

[Cr(H2O)6]3+ + 6NH3  [Cr(NH3)6]3+ + 6H2O

352
Q

What is the stability constant?

A

Kstab – the equilibrium constant existing between a transition metal ion surrounded by water ligands and the complex formed when the same ion has undergone a ligand substitution reactionLike KC but for equilibrias including complex ions

353
Q

Colour of Sc3+

A

Colourless

354
Q

Colour of Ti2+

A

colourless

355
Q

Colour of Ti3+

A

lilac

356
Q

Colour of Ti+4

A

colourless

357
Q

Colour of Ti5+

A

colourless

358
Q

Colour of V2+

A

lilac

359
Q

Colour of V3+

A

green

360
Q

Colour of V4+

A

blue

361
Q

Colour of V5+

A

yellow

362
Q

Colour of Cr2+

A

blue

363
Q

Colour of Cr3+

A

green

364
Q

Colour of Cr4+

A

colourless

365
Q

Colour of Cr5+

A

colourless

366
Q

Colour of Cr6+

A

orange

367
Q

Colour of Mn2+

A

pale pink

368
Q

Colour of Mn+3

A

colourless

369
Q

Colour of Mn4+

A

dark purple

370
Q

Colour of Mn5+

A

colourless

371
Q

Colour of Mn6+

A

green

372
Q

Colour of Mn7+

A

lilac

373
Q

Colour of Fe2+

A

pale green

374
Q

Colour of Fe3+

A

pale yellow

375
Q

Colour of Fe4+

A

colourless

376
Q

Colour of Fe5+

A

colourless

377
Q

Colour of Fe6+

A

colourless

378
Q

Colour of Co2+

A

Pink

379
Q

Colour of Co3+

A

green

380
Q

Colour of Co4+

A

colourless

381
Q

Colour of Co5+

A

colourless

382
Q

Colour of Ni2+

A

green

383
Q

Colour of Ni3+

A

colorless

384
Q

Colour of Ni4+

A

colourless

385
Q

Colour of Cu1+

A

colourless

386
Q

Colour of Cu2+

A

blue

387
Q

Colour of Cu3+

A

colourless

388
Q

Colour of Zn2+

A

colourless

389
Q

[Cu(H2O)6]2+ colour

A

blue solution

390
Q

Cu(OH)2 colour

A

Blue precipitate

391
Q

[Cu(NH3)4(H2O)2]2+ colour

A

Deep blue solution

392
Q

[CuCl4]2- colour

A

Yellow solution

393
Q

HCl + Cu2+(aq) reaction colour and why

A

the complex can look green as the reaction is reversible so both blue and yellow species present

394
Q

[Fe(H2O)6]2+ colour

A

pale green solution

395
Q

Fe(OH)2 colour

A

Green precipitate, if left in air a reddy brown colour appears (Fe2+ oxidizes to Fe3+)

396
Q

[Fe(H2O)6]3+ colour

A

Yellow solution

397
Q

Fe(OH)3 colour

A

Reddy brown precipitate

398
Q

[Mn(H2O)6]2+ colour

A

Very pale solution

399
Q

Mn(OH)2 colour

A

Light brown precipitate which darkens in air

400
Q

[Cr(H2O)6]3+ colour

A

Violet solution

401
Q

Cr(OH)3 colour

A

Grey/green precipitate

402
Q

[Cr(OH)6]3- colour

A

Green solution

403
Q

[Cr(NH3)6]3+

A

Purple solution

404
Q

[Cu(H2O)6]2+ + 4NH3

A

[Cu(H2O)6]2+ + 4NH3  [Cu(NH3)4(H2O)2]2+ + 4H2OBlue solution Deep blue solution

405
Q

[Co(H¬2O)6]2+ + 6NH3

A

[Co(H¬2O)6]2+ + 6NH3  [Co(NH3)6]2+ + 6H2O

406
Q

[Cu(H2O)6]2+ + 4Cl-

A

[Cu(H2O)6]2+ + 4Cl-  [CuCl4]2- + 6H2OBlue solution Yellow/green solution

407
Q

[Co(H2O)6]2+ + 4Cl-

A

[Co(H2O)6]2+ + 4Cl-  [CoCl4]2- + 6H2OPink solution Blue solution

408
Q

Cu2+ + 2OH-  Cu(OH)2 colours

A

Cu2+ + 2OH-  Cu(OH)2Blue solution  blue precipitate

409
Q

Mn2+ + 2OH-  Mn(OH)2 colours

A

Mn2+ + 2OH-  Mn(OH)2Very pale pink solution  pale brown precipitate

410
Q

Fe2+ + 2OH-  Fe(OH)2 COLOURs

A

Fe2+ + 2OH-  Fe(OH)2Green solution  green precipitate

411
Q

Fe3+ + 3OH-  Fe(OH)3 colours

A

Fe3+ + 3OH-  Fe(OH)3Yellow/brown solution  brown precipitate

412
Q

Cr3+ + 3OH-  Cr(OH)3 colours

A

Cr3+ + 3OH-  Cr(OH)3Green solution  green precipitate

413
Q

[Mn(H2O)6]2+ + 2NH3

A

[Mn(H2O)6]2+ + 2NH3  Mn(H2O)4(OH)2 + 2NH4+

414
Q

[Fe(H2O)6]3+ + 3NH3

A

[Fe(H2O)6]3+ + 3NH3  Mn(H2O)3(OH)3 + 3NH4+

415
Q

Cr(H2O)3(OH)3 + 3OH-

A

Cr(H2O)3(OH)3 + 3OH-  [Cr(OH)6]3- + 3H2OGreen precipitate  green solution

416
Q

Cr(H2O)3(OH)3 + 3H+

A

Cr(H2O)3(OH)3 + 3H+  [Cr(H2O)6]3+Green precipitate  green solution

417
Q

Cr(OH)3(H2O)3 + 6NH3

A

Cr(OH)3(H2O)3 + 6NH3  [Cr(NH3)6]3+ + 3H2O + 3OH-Green precipitate  purple solution

418
Q

Cu(OH)2(H2O)4 + 4NH3

A

Cu(OH)2(H2O)4 + 4NH3  [Cu(NH3)4(H2O)2]2+ + 2H2O + 2OH-Blue precipitate  deep blue solution

OCR A Level Chemistry (17 decks)

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