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OCR A Level Chemistry > Paper 3B Actual > Flashcards

Paper 3B Actual Flashcards

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1
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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2
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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3
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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4
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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5
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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6
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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7
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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8
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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9
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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10
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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11
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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12
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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13
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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14
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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15
Q

Draw the born haber cycle for NaCl

A

check notes or google

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16
Q

Draw the born haber cycle for MgCl2

A

check notes or google

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17
Q

Draw the born haber cycle for CuO

A

check notes or google

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18
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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19
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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20
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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21
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

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22
Q

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

A

check notes

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23
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

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24
Q

A chemical reaction will proceed when?

A

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

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25
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

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26
Q

What happens to entropy over time?

A

Entropy must increase over time

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27
Q

When is a system in a state of high entorpy?

A

When its degree of disorder is high

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28
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

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29
Q

What does entropy always do?

A

The total entropy always increases, and the process is irreversible

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30
Q

What is the unit for entropy

A

Sin JK^-1mol^-1

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31
Q

Why is S (entropy) alwasy potisive?

A

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

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32
Q

How does the entropy of each state differ?

A

Solid has lowest entropy, and gas has the highest

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33
Q

Draw a graph of waters change in enthalpy as temperaurer increases

A

check notes

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34
Q

Systems that are more chaotic have a _____ entropy value

A

higher

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35
Q

(s)  (s) + (g) , what is delta S

A

+

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36
Q

(g)  2(g) , what is delta S

A

+

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37
Q

What is delta S?C2H5OH(l)  C2H5OH(g)

A

+

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38
Q

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

A

-

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39
Q

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

A

+

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40
Q

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

A

-

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41
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

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42
Q

Sum for delta S^theta

A

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

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43
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.

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44
Q

deltaSsurroundings is proportional to what?

A

NAME?

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45
Q

deltaSsurroundings =

A

(-deltaHsystem)/T

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46
Q

deltaStotal =

A

deltaSsystem + deltaSsurroundings

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47
Q

deltaStotal >

A

0

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48
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

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49
Q

What is gibbs free energy equation?

A

deltaG = deltaH – T deltaS < 0

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50
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

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51
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

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52
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)

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53
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
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54
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
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55
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

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56
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.

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57
Q

What elements have been used in coinage for many years?

A

copper, silver, nickel, and zinc

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58
Q

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

A

Iron

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59
Q

What element is used for electrical cables and water pipes?

A

Copper

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60
Q

Titanium has great strength, what are the elements applications?

A

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

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61
Q

Sc orbital configuration

A

[Ar] 4s2 3d1

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62
Q

Fe orbital configuration

A

Fe – [Ar] 4s2 3d6

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63
Q

Ni orbital configuration

A

Ni – [Ar] 4s2 3d8

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64
Q

Zn orbital configuration

A

Zn – [Ar] 4s2 3d10

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65
Q

Cr orbital configuration

A

Cr – [Ar] 4s1 3d5

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66
Q

Cu orbital configuration

A

Cu – [Ar] 4s1 3d10

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67
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

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68
Q

Fe2+ orbital configuration

A

Fe2+ - [Ar] 4s0 3d6

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69
Q

Fe3+ orbital configuration

A

Fe3+ - [Ar] 4s0 3d5

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70
Q

Cu2+ orbital configuration

A

Cu2+ - [Ar] 4s0 3d9

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71
Q

Cr3+ orbital configuration

A

Cr3+ - [Ar] 4s0 3d3

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72
Q

Mn2+ orbital configuration

A

Mn2+ - [Ar] 4s0 3d5

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73
Q

Mn4+ orbital configuration

A

Mn4+ - [Ar] 4s0 3d3

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74
Q

Sc3+ orbital configuration

A

Sc3+ - [Ar] 4s0 3d0

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75
Q

Zn2+ orbital configuration

A

Zn2+ - [Ar] 4s0 3d10

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76
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.

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77
Q

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

A
  • Scandium and zinc
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78
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
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79
Q

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

A

strong oxidizing agent

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80
Q

What does the observed colour of a solution depend on?

A

The observed colour of a solution depends on the wavelengths absorbed

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81
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.

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82
Q

Draw the exam colour chart

A

check notes

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83
Q

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

A

Cyan

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84
Q

What colour is cu2+

A

white

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85
Q

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

A

blue

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86
Q

Transition metals form complex ions or ?

A

coordination compounds

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87
Q

WHat do transition metals form when they make complex ions?

A

ligand forms bonds with the central transition metal ion

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88
Q

Examples of complex ions

A

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

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89
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

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90
Q

What does monodentate mean?

A

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

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91
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)

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92
Q

Draw [Cu(H2O)6]2+

A

check notes

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93
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
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94
Q

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

A

2

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95
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
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96
Q

[Cr(H2O)4Cl2]+ name

A

¬¬Tetraaquadichlorochromium (III) ion

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97
Q

[Co(H2O)5Cl]+ name

A

Pentaaquamonochlorocobalt (II) ion

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98
Q

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

A

the suffix ‘-ate’ follows the metal

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99
Q

Name of [Fe(CN)6]4-

A

Hexacyanoferrate (II) ion

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100
Q

What does cobalt become in an anion complex ion?

A

cobaltate

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101
Q

What does aluminium become in an anion complex ion?

A

aluminate

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102
Q

What does chromium become in an anion complex ion?

A

chromate

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103
Q

What does vanadium become in an anion complex ion?

A

vanadate

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104
Q

What does copper become in an anion complex ion?

A

cuprate

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105
Q

What does iron become in an anion complex ion?

A

ferrate

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106
Q

What does nickel become in an anion complex ion?

A

nickelate

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107
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)

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108
Q

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

A

white precipitate

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109
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)

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110
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)

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111
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)

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112
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)

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113
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)

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114
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
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115
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-
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116
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)
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117
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)
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118
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
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119
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
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120
Q

WHat is oaplatin used in?

A

treatment of colorectal cancer

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121
Q

Types of cancer therapy

A
  • Surgery- Radiotherapy- Chemotherapyo Cytotoxico Targeted Anti-endocrine Novel targeted agents- Immuno-therapy- Gene therapy
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122
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
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123
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)
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124
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
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125
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
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126
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)

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127
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

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128
Q

What is a bidentate ligand? Example

A

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

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129
Q

Draw [Ni(NH2CH2CH2NH2)3]2+

A

Check notes

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130
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)

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131
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

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132
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
133
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

134
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

135
Q

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

A

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

136
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

137
Q

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

A

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

138
Q

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

A

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

139
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

140
Q

Colour of Sc3+

A

Colourless

141
Q

Colour of Ti2+

A

colourless

142
Q

Colour of Ti3+

A

lilac

143
Q

Colour of Ti+4

A

colourless

144
Q

Colour of Ti5+

A

colourless

145
Q

Colour of V2+

A

lilac

146
Q

Colour of V3+

A

green

147
Q

Colour of V4+

A

blue

148
Q

Colour of V5+

A

yellow

149
Q

Colour of Cr2+

A

blue

150
Q

Colour of Cr3+

A

green

151
Q

Colour of Cr4+

A

colourless

152
Q

Colour of Cr5+

A

colourless

153
Q

Colour of Cr6+

A

orange

154
Q

Colour of Mn2+

A

pale pink

155
Q

Colour of Mn+3

A

colourless

156
Q

Colour of Mn4+

A

dark purple

157
Q

Colour of Mn5+

A

colourless

158
Q

Colour of Mn6+

A

green

159
Q

Colour of Mn7+

A

lilac

160
Q

Colour of Fe2+

A

pale green

161
Q

Colour of Fe3+

A

pale yellow

162
Q

Colour of Fe4+

A

colourless

163
Q

Colour of Fe5+

A

colourless

164
Q

Colour of Fe6+

A

colourless

165
Q

Colour of Co2+

A

Pink

166
Q

Colour of Co3+

A

green

167
Q

Colour of Co4+

A

colourless

168
Q

Colour of Co5+

A

colourless

169
Q

Colour of Ni2+

A

green

170
Q

Colour of Ni3+

A

colorless

171
Q

Colour of Ni4+

A

colourless

172
Q

Colour of Cu1+

A

colourless

173
Q

Colour of Cu2+

A

blue

174
Q

Colour of Cu3+

A

colourless

175
Q

Colour of Zn2+

A

colourless

176
Q

[Cu(H2O)6]2+ colour

A

blue solution

177
Q

Cu(OH)2 colour

A

Blue precipitate

178
Q

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

A

Deep blue solution

179
Q

[CuCl4]2- colour

A

Yellow solution

180
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

181
Q

[Fe(H2O)6]2+ colour

A

pale green solution

182
Q

Fe(OH)2 colour

A

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

183
Q

[Fe(H2O)6]3+ colour

A

Yellow solution

184
Q

Fe(OH)3 colour

A

Reddy brown precipitate

185
Q

[Mn(H2O)6]2+ colour

A

Very pale solution

186
Q

Mn(OH)2 colour

A

Light brown precipitate which darkens in air

187
Q

[Cr(H2O)6]3+ colour

A

Violet solution

188
Q

Cr(OH)3 colour

A

Grey/green precipitate

189
Q

[Cr(OH)6]3- colour

A

Green solution

190
Q

[Cr(NH3)6]3+

A

Purple solution

191
Q

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

A

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

192
Q

[Co(H¬2O)6]2+ + 6NH3

A

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

193
Q

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

A

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

194
Q

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

A

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

195
Q

Cu2+ + 2OH-  Cu(OH)2 colours

A

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

196
Q

Mn2+ + 2OH-  Mn(OH)2 colours

A

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

197
Q

Fe2+ + 2OH-  Fe(OH)2 COLOURs

A

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

198
Q

Fe3+ + 3OH-  Fe(OH)3 colours

A

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

199
Q

Cr3+ + 3OH-  Cr(OH)3 colours

A

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

200
Q

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

A

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

201
Q

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

A

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

202
Q

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

A

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

203
Q

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

A

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

204
Q

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

A

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

205
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

206
Q

The relative mass of an electron

A

1/1836

207
Q

The relative mass of a proton

A

1

208
Q

The relative mass of a neutron

A

1

209
Q

Charge of proton

A

1

210
Q

Charge of neutron

A

0

211
Q

Charge of electron

A

-1

212
Q

What is an isotope?

A

Isotopes are atoms of the same element, with a different number of neutrons (same number of protons/electrons)

213
Q

In the periodic table, horizontal rows are called …

A

periods

214
Q

In the periodic table, vertical collums are called …

A

groups

215
Q

How many groups are there?

A

18

216
Q

What is group 1?

A

Alkali metals

217
Q

What is group 2?

A

Alkaline Earth Metals

218
Q

What is groups 3-12/

A

Transition metals

219
Q

What is group 15?

A

Pnictogens

220
Q

What is group 16?

A

Chalcogens

221
Q

What is group 17?

A

The halogens

222
Q

What is group 18?

A

The noble gases

223
Q

What are metalloids/semi metals? Why do they have that name?

A

Elements that touch the line between metal and non-metals, they have a combination of properties of metals and non-metals

224
Q

Define ionisation

A

when an atom loses an electron from its outer shell.

225
Q

Define first ionisation

A

The energy required to remove 1 mole of electrons from one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions.

226
Q

What 3 things can affect ionisation energy?

A

Shielding, nuclear charge and atomic radius

227
Q

Define shielding

A

Shielding is when inner electrons screen the outer electrons from the pull from the nucleus

228
Q

Define nuclear charge

A

The positive charge of the nucleus

229
Q

Define atomic radius

A

Atomic radius is radius of an atom, we measure atomic radius by measuring the distance between 2 nuclei of touching atoms and halving the distance

230
Q

Describe and explain the trend of atomic radius size as you go along period 3

A

There is an increase in nuclear charge as you go along the period. The stronger the nuclear charge the more it can pull the electrons closer to the nucleus. The shielding is the same as you go along the period. The general trend for atomic radius is it decreases as you go along period 3 as the electrons are pulled closer to the nucleus.

231
Q

Describe and explain the trend of period 3 first ionisation energy

A

First ionisation energy enthalpy increases across a period, the nuclear charge increases across a period, attraction of electrons to the nucleus increases, atomic radius decreases (so electrons are closer), it takes more energy to remove that first electron.

232
Q

Why does first ionisation energy dip between magnesium and aluminium?

A

There is a decease between magnesium and aluminium as magnesium has a full sub-shell stability, as the highest occupied sub-shell is complete (3s). Aluminium has one electron in a higher sub-shell (3p), this one electron is removed more easily as it is further away from the nucleus, and therefore the first ionisation energy is lower than magnesium.

233
Q

Why does first ionisation energy dip between phosphorus and sulphur?

A

There is a decrease between phosphorus and sulphur because phosphorus has a half subshell stability, sulphur has one 3p orbital that contains a pair of electrons. These paired electrons repel each other, so one of these electrons is easier to remove therefore sulphur has a lower fist ionisation energy then phosphorus.

234
Q

Describe and explain the trend in first ionisation energy as you go down group 3

A

First ionisation energy decreases down a group.This is because more inner shell electrons as you go down the group, so there is shielding of outer electrons, attraction of electrons to the nucleus decreases, atomic radius increases (so electrons are further away).

235
Q

Where does metallic bonding occur?

A

Only in metals

236
Q

Describe and explain the metallic structure

A

.Giant structure.The atoms are in tightly packed layers, which form a regular lattice structure.Each atom in a solid metal structure has donated its negative electrons from its outer shell forming an ion.The outer electrons become delocalised and create a ‘sea of free electrons’.The positive metal ion (cation) is fixed in a portion maintaining the structure of the metal.The delocalised electrons are mobile and can move through the structure

237
Q

Define metallic bonding

A

.Metallic bonding is the strong electrostatic attraction between the metal ions (cations) and the delocalised electrons

238
Q

Where do metals conduct? Why?

A

.Most only conduct in solid and liquid states.The delocalised electrons can move freely anywhere with the metal lattice allowing them to conduct electricity

239
Q

Why do metals have large bp/mp?

A

.Strong electrostatic attraction between the metal ion and the delocalised electrons.Needs a lot of energy to break

240
Q

What does the strength of the metal depend on? What does this also affect?

A

.The strength of metal depends on the charge of the metal ion in the structure (the greater the charge, the greater the number of delocalized electrons, the stronger the attraction, the stronger the bond).This also affects the size of the ion (the smaller the ion, the closer the electrons to the nucleus, the stronger the bond)

241
Q

Do metals dissolve in solvents?

A

no

242
Q

Why are metals good conductors of heat?

A

.When a metal is heated the delocalised electrons gain kinetic energy and move faster.This movement transfers the gained energy throughout the metal

243
Q

Why are pure metals soft?

A

When a metal is hit, the layers of metal ions are able to slide over each other, and so the layers do not shatter

244
Q

What is an alloy and why are they harder than pure metals?

A

.An alloy is a mixture of two metals, thereby distorting the layers so they can’t slide over each other and make a harder new metal

245
Q

diamond structure

A

.Giant covalent bonding.Forms a lattice.Made of carbon only.Strong covalent bond between each atom.Each carbon is bonded to 4 other carbons

246
Q

Diamond properties

A

.Hard – a giant covalent structure and has lots of strong bonds.Doesn’t conduct electricity – no delocalised electrons that can move.High melting point – many strong covalent bonds between atoms , it needs a lot of energy to break

247
Q

Graphite structure

A

.Giant covalent bonding.Made of carbon only.Hexagonal structure.Strong covalent bonds between carbon atoms.In parallel layers.Weak intermolecular forces between layers.Each carbon is bonded to 3 other carbons.Delocalised electrons in each layer

248
Q

Graphite properties

A

.Soft/slippery – the layers can slide over each other because there are weak forces between the layers.Conduct electricity – delocalised electrons that can move through the whole structure.High melting point – graphite has a giant structure with lots of strong covalent bonds between atoms, it needs lots of energy to break bonds

249
Q

What is graphene? properties

A

.One layer of graphite.Highly conductive and strong and flexible

250
Q

What is a nanotube?

A

When graphene forms a tube like strutcure

251
Q

What is a bucky ball?

A

When carbon forms a cage like structure

252
Q

What are different physical forms of elements called?

A

allotropes

253
Q

On a melting point across a period graph, how can you section the different points?

A

Metallic bonding, giant covalent structures, covalent molecules

254
Q

Why does the melting point increase across period 3?

A

.The charge in the metal ions increase.The number of delocalised electrons increases.So the strength of the metallic bonding increases.Need more energy to break stronger metallic bonds so the melting points and boiling points increase

255
Q

Is silicon a metal or non-metal?

A

TRICK QUESTIONIts a metalloid

256
Q

What is silicons structure like?

A

It has a giant covalent structure exactly the same as carbon in diamond (each silicon is bonded to four other silicon’s)

257
Q

Why does silicon have a high melting point?

A

you have to break lots of string covalent bonds in order to melt it, and this requires a lot of energy to break.

258
Q

How do phosphorus, sulphur and chlorine exist?

A

simple molecules, with strong covalent bonds between their atoms.

259
Q

How does argon exist?

A

Argon exists as separate atoms (it is monatomic)

260
Q

Why is Cl, Ar, S8 and P4’s melting points so low?

A
  • When these four substances melt or boil, it is the London forces between the molecules which are broken. These are very weak bonds. So little energy is needed to overcome them.
261
Q

In what order do the melting points of Cl2, Ar, S8 and P4 decrease? Why?

A

The strength of the intermolecular forces increases with the number of electrons, so the melting point decreases in this order:S8 > P4 > Cl2 > Ar

262
Q

What contains enthalpy?

A

Elements and compounds

263
Q

What is enthalphy (H)?

A

the heat that is stored in a chemical system

264
Q

What is a chemical system?

A

The system refers to the atoms, molecules and ions making up the chemicals

265
Q

Enthalpy cannot be measured, what can?

A

Enthalpy change

266
Q

What is enthalpy change (ΔH)?

A

the difference in enthalpy between the products and reactants in a reaction

267
Q

What happens in an exothermic reaction?

A

.In an exothermic chemical reaction heat energy is transferred from the system to the surroundings.Any energy loss by the system is balanced by the energy gain by the surroundings.The temperature of the surroundings increases, so we see a temperature increase

268
Q

What happens in an endothermic reaction?

A

.In an endothermic chemical reaction heat energy is absorbed from the surroundings into the system.Any energy gain by the system is balanced by the energy loss of the surroundings.The temperature of the surroundings decreases, so we see a temperature decrease

269
Q

Burning fuel, exo or endo?

A

Exo

270
Q

Condensing a vapour, exo or endo?

A

Exo

271
Q

Evaporation, exo or endo?

A

Endo

272
Q

Neutralising an acid, exo or endo?

A

Exo

273
Q

Thermal decomposition of copper carbonate, exo or endo?

A

Endo

274
Q

Rapid oxidation of iron, exo or endo?

A

Endo

275
Q

Shop bought hand warmers, exo or endo?

A

Exo

276
Q

Sports injury cool packs, exo or endo?

A

Endo

277
Q

Enthalpy change equation

A

H(Products) – H(Reactants)

278
Q

Sign for exothermic reaction

A

negative

279
Q

Sign for endothermic reaction

A

positive

280
Q

Define activation energy

A

The minimum amount of energy required to start a reaction by breaking the bonds

281
Q

Draw an exothermic activation energy graph

A

Check notes

282
Q

Draw an endothermic activation energy graph

A

Check notes

283
Q

Draw an exothermic enthalpy profile

A

Check notes

284
Q

Draw an endothermic enthalpy profile

A

Check notes

285
Q

Define Enthalpy change of reaction (ΔHr) with example

A

the enthalpy change when the reaction occurs in the molar quantities shown in the chemical reactione.g. Zn(s) + Cu2+(aq)  Cu(s) + Zn2+(aq)

286
Q

Define Enthalpy change of formation (ΔHf) with example

A

the enthalpy change when 1 mole of compound is formed from its elementse.g. C(s) + O2(g)  CO2(g)

287
Q

Define combustion change of reaction (ΔHc) with example

A

the enthalpy change when 1 mole of substance is burnte.g. CH4(g) + O2(g)  CO2(g) + H2O(g)

288
Q

Define Enthalpy change of neutrilisation (ΔHneut) with example

A

the enthalpy change when solutions of an acid and an alkali react together under standard conditions to produce 1 mole of water. It is always measured per mole of water formed.e.g. HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)

289
Q

What symbol is used for standard conditions?

A

ϴ

290
Q

What conditions are used for enthalpy change values?

A

Standard conditions:Standard pressure = 100KPaStandard temperature = 298K (25’C)Standard concentration = 1 moldm-3

291
Q

Two equations for calculating enthalpy changes

A

q = mcΔTΔH = q/n

292
Q

What is each part of ΔH = q/n?

A

q = heat energy (J)n = number of moles (mol)

293
Q

What is each part of q = mcΔT?

A

q = heat energy (J)m = mass of substance heated or cooled (g)c = specific heat capacity of water (4.18 Jg-1K-1)ΔT = change in temperature (‘C or K)

294
Q

What is the final step of calculating enthalpy change equations?

A

Multiply the answer by the co-efficient of the reagent not in excess

295
Q

When doing enthalpy change calculations, and you dont have a weight for the solution, what do you use?

A

The volume of it in cm^3, Cm^3 = g

296
Q

When doing enthalpy change calculations with an equation and bond energy values, how do you get the final answer after finding the values of either side of the equation?

A

bonds broken - bonds made(left - right)

297
Q

Define average bond enthalpy

A

the average enthalpy change when one mole of gaseous covalent bonds is broken

298
Q

Why might the calculated average bond enthalpy different from the actual?

A

.The bond enthalpies used in the calculations are averages from different compounds, the exact bond enthalpy depends on the particular compound in which it is either formed or broken.The bond enthalpies are not determined under standard conditions. Bond enthalpies are determined by molecules in a gaseous state. But at 298K not all compounds are a gas, such as water.

299
Q

In terms of enthalpy, what is breaking of bonds always? Why?

A

Always endothermic, as you need to add energy to break bonds

300
Q

What can average bond enthalpy also be called?

A

Mean bond enthalpy

301
Q

Where is average bond enthalpy taken from?

A

the average is taken over a wide range of compounds containing that type of bond

302
Q

State the standard conditions

A

Standard pressure = 100kPa (1 atmosphere)Standard temperature = 298K (25’C)Standard concentration = 1moldm-3Standard states of all substances (carbons is a solid, hydrogen is a gas and water is a liquid)

303
Q

Define standard enthalpy change of reaction

A

The enthalpy change for a reaction as shown by the molar quantities in the chemical equationsMeasured under standard conditions

304
Q

Define standard enthalpy change of combustion

A

Enthalpy change when 1 mole of a substance reacts completely with oxygen under standard conditions`

305
Q

Define standard enthalpy change of formation

A

The enthalpy change when 1 moles of substance is formed from its elements under standard conditions

306
Q

Enthalpy change of formation of elements is always …

A

zero

307
Q

What to remember for Hess’ Law Cycles Involving Enthalpies of Formation

A

Using delta H f it’s p-r!

308
Q

Formula for hess’ law cycles involving enthalpies of formation

A

sum of ΔHf products - sum of ΔHf reactants

309
Q

Why does oxygen have no enthalpy of combustion value?

A

Oxygen can’t be combusted

310
Q

What does Hess’ law state?

A

Hess’ law states that the overall enthalpy change of the two routes is the same

311
Q

Equation for Hess’ Law Cycles involving enthalpies of combustion

A

sum of enthalpy of combustion of reactants - sum of enthalpy of combustion of products

312
Q

Does breaking bonds require or release energy?

A

Require

313
Q

Does forming bonds require or release energy?

A

Release

314
Q

Define exothermic

A

the energy required to break bonds is less than the energy given out when new bonds form

315
Q

Define endothermic

A

the energy required to break bonds is more than the energy given out when new bonds form

316
Q

Define average bond enthalpy

A

the average enthalpy change when one mole of gaseous covalent bonds is broken

317
Q

Equation for enthalpy change

A

enthalpy of bonds broken – enthalpy of bonds formed

318
Q

Why might calculated bond enthalpies not be right?

A
  1. The bond enthalpies used in the calculations are averages from different compounds, the exact bond enthalpy depends on the particular compound in which it is either formed or broken2. The bond enthalpies are not determined under standard conditions. Bond enthalpies are determined by molecules in a gaseous state. But at 298K not all compounds are a gas, such as water
319
Q

Draw a general diagram for Hess’ law

A

CHECK NOTES

320
Q

Enthalpy change of reaction –

A

This is the enthalpy change when the reaction occurs in the molar quantities shown in the chemical equation

321
Q

Enthalpy change of formation –

A

This is the enthalpy change when 1 mole of compound is formed from its elements

322
Q

Enthalpy change of combustion –

A

This is the enthalpy change when 1 mole of substance is burned

323
Q

Enthalpy change of neutralisation –

A

The enthalpy change when solutions of an acid and an alkali react together under standard conditions to produce 1 mole of water. It is always measured per mole of water formed.

324
Q

Hess’ Law

A

If a reaction can take place by more than one route, and the initial and final concentrations are the same, the total enthalpy change is the same regardless of the route taken

325
Q

Things to remember when using Hess’ law

A

If you have more than one compound, then you have to add the enthalpy of formation for each compound togetherEnthalpy of formation is given for 1 mole of compound formed. If there is more than one mole, you multiply it

326
Q

What is Hess’ law used for?

A

Hess’ Law determines enthalpy changes indirectly, for when determining them directly isn’t possible

327
Q

Where does Hess’ law come from?

A

Hess’ Law comes from the idea of the conservation of energy

328
Q

For elements, enthalpy change of formation is ALWAYS ____

A

zero

329
Q

Enthalpy change of reaction equation =

A

sum of the enthalpy change of formation of products - sum of the enthalpy change of formation of reactants OR sum of the enthalpy change of combustion of reactants - sum of the enthalpy change of combustion of products

OCR A Level Chemistry (17 decks)

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