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OCR A Level Chemistry > Paper 3B Part1 > Flashcards

Paper 3B Part1 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 a fuel cell?

A

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

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126
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.

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127
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

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128
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

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129
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

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130
Q

Draw an alkaline fuel cell

A

check notes

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131
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

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132
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

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133
Q

Define electrochemical reaction

A

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

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134
Q

Define electrode

A

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

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135
Q

Define electrolyte

A

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

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136
Q

What is an electrochemical cell consisted of?

A

An electrochemical cell consists of 2 electrodes + 1 electrolyte

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137
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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138
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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139
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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140
Q

What is a bidentate ligand? Example

A

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

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141
Q

Draw [Ni(NH2CH2CH2NH2)3]2+

A

Check notes

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142
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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143
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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144
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
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145
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

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146
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

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147
Q

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

A

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

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148
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

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149
Q

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

A

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

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150
Q

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

A

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

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151
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

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152
Q

Colour of Sc3+

A

Colourless

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153
Q

Colour of Ti2+

A

colourless

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154
Q

Colour of Ti3+

A

lilac

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155
Q

Colour of Ti+4

A

colourless

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156
Q

Colour of Ti5+

A

colourless

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157
Q

Colour of V2+

A

lilac

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158
Q

Colour of V3+

A

green

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159
Q

Colour of V4+

A

blue

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160
Q

Colour of V5+

A

yellow

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161
Q

Colour of Cr2+

A

blue

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162
Q

Colour of Cr3+

A

green

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163
Q

Colour of Cr4+

A

colourless

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164
Q

Colour of Cr5+

A

colourless

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165
Q

Colour of Cr6+

A

orange

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166
Q

Colour of Mn2+

A

pale pink

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167
Q

Colour of Mn+3

A

colourless

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168
Q

Colour of Mn4+

A

dark purple

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169
Q

Colour of Mn5+

A

colourless

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170
Q

Colour of Mn6+

A

green

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171
Q

Colour of Mn7+

A

lilac

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172
Q

Colour of Fe2+

A

pale green

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173
Q

Colour of Fe3+

A

pale yellow

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174
Q

Colour of Fe4+

A

colourless

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175
Q

Colour of Fe5+

A

colourless

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176
Q

Colour of Fe6+

A

colourless

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177
Q

Colour of Co2+

A

Pink

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178
Q

Colour of Co3+

A

green

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179
Q

Colour of Co4+

A

colourless

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180
Q

Colour of Co5+

A

colourless

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181
Q

Colour of Ni2+

A

green

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182
Q

Colour of Ni3+

A

colorless

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183
Q

Colour of Ni4+

A

colourless

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184
Q

Colour of Cu1+

A

colourless

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185
Q

Colour of Cu2+

A

blue

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186
Q

Colour of Cu3+

A

colourless

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187
Q

Colour of Zn2+

A

colourless

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188
Q

[Cu(H2O)6]2+ colour

A

blue solution

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189
Q

Cu(OH)2 colour

A

Blue precipitate

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190
Q

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

A

Deep blue solution

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191
Q

[CuCl4]2- colour

A

Yellow solution

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192
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

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193
Q

[Fe(H2O)6]2+ colour

A

pale green solution

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194
Q

Fe(OH)2 colour

A

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

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195
Q

[Fe(H2O)6]3+ colour

A

Yellow solution

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196
Q

Fe(OH)3 colour

A

Reddy brown precipitate

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197
Q

[Mn(H2O)6]2+ colour

A

Very pale solution

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198
Q

Mn(OH)2 colour

A

Light brown precipitate which darkens in air

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199
Q

[Cr(H2O)6]3+ colour

A

Violet solution

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200
Q

Cr(OH)3 colour

A

Grey/green precipitate

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201
Q

[Cr(OH)6]3- colour

A

Green solution

202
Q

[Cr(NH3)6]3+

A

Purple solution

203
Q

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

A

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

204
Q

[Co(H¬2O)6]2+ + 6NH3

A

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

205
Q

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

A

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

206
Q

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

A

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

207
Q

Cu2+ + 2OH-  Cu(OH)2 colours

A

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

208
Q

Mn2+ + 2OH-  Mn(OH)2 colours

A

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

209
Q

Fe2+ + 2OH-  Fe(OH)2 COLOURs

A

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

210
Q

Fe3+ + 3OH-  Fe(OH)3 colours

A

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

211
Q

Cr3+ + 3OH-  Cr(OH)3 colours

A

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

212
Q

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

A

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

213
Q

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

A

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

214
Q

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

A

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

215
Q

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

A

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

216
Q

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

A

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

217
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

218
Q

The relative mass of an electron

A

1/1836

219
Q

The relative mass of a proton

A

1

220
Q

The relative mass of a neutron

A

1

221
Q

Charge of proton

A

1

222
Q

Charge of neutron

A

0

223
Q

Charge of electron

A

-1

224
Q

What is an isotope?

A

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

225
Q

In the periodic table, horizontal rows are called …

A

periods

226
Q

In the periodic table, vertical collums are called …

A

groups

227
Q

How many groups are there?

A

18

228
Q

What is group 1?

A

Alkali metals

229
Q

What is group 2?

A

Alkaline Earth Metals

230
Q

What is groups 3-12/

A

Transition metals

231
Q

What is group 15?

A

Pnictogens

232
Q

What is group 16?

A

Chalcogens

233
Q

What is group 17?

A

The halogens

234
Q

What is group 18?

A

The noble gases

235
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

236
Q

Define ionisation

A

when an atom loses an electron from its outer shell.

237
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.

238
Q

What 3 things can affect ionisation energy?

A

Shielding, nuclear charge and atomic radius

239
Q

Define shielding

A

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

240
Q

Define nuclear charge

A

The positive charge of the nucleus

241
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

242
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.

243
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.

244
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.

245
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.

246
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).

247
Q

Where does metallic bonding occur?

A

Only in metals

248
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

249
Q

Define metallic bonding

A

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

250
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

251
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

252
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)

253
Q

Do metals dissolve in solvents?

A

no

254
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

255
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

256
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

257
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

258
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

259
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

260
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

261
Q

What is graphene? properties

A

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

262
Q

What is a nanotube?

A

When graphene forms a tube like strutcure

263
Q

What is a bucky ball?

A

When carbon forms a cage like structure

264
Q

What are different physical forms of elements called?

A

allotropes

265
Q

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

A

Metallic bonding, giant covalent structures, covalent molecules

266
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

267
Q

Is silicon a metal or non-metal?

A

TRICK QUESTIONIts a metalloid

268
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)

269
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.

270
Q

How do phosphorus, sulphur and chlorine exist?

A

simple molecules, with strong covalent bonds between their atoms.

271
Q

How does argon exist?

A

Argon exists as separate atoms (it is monatomic)

272
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.
273
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

274
Q

What contains enthalpy?

A

Elements and compounds

275
Q

What is enthalphy (H)?

A

the heat that is stored in a chemical system

276
Q

What is a chemical system?

A

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

277
Q

Enthalpy cannot be measured, what can?

A

Enthalpy change

278
Q

What is enthalpy change (ΔH)?

A

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

279
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

280
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

281
Q

Burning fuel, exo or endo?

A

Exo

282
Q

Condensing a vapour, exo or endo?

A

Exo

283
Q

Evaporation, exo or endo?

A

Endo

284
Q

Neutralising an acid, exo or endo?

A

Exo

285
Q

Thermal decomposition of copper carbonate, exo or endo?

A

Endo

286
Q

Rapid oxidation of iron, exo or endo?

A

Endo

287
Q

Shop bought hand warmers, exo or endo?

A

Exo

288
Q

Sports injury cool packs, exo or endo?

A

Endo

289
Q

Enthalpy change equation

A

H(Products) – H(Reactants)

290
Q

Sign for exothermic reaction

A

negative

291
Q

Sign for endothermic reaction

A

positive

292
Q

Define activation energy

A

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

293
Q

Draw an exothermic activation energy graph

A

Check notes

294
Q

Draw an endothermic activation energy graph

A

Check notes

295
Q

Draw an exothermic enthalpy profile

A

Check notes

296
Q

Draw an endothermic enthalpy profile

A

Check notes

297
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)

298
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)

299
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)

300
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)

301
Q

What symbol is used for standard conditions?

A

ϴ

302
Q

What conditions are used for enthalpy change values?

A

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

303
Q

Two equations for calculating enthalpy changes

A

q = mcΔTΔH = q/n

304
Q

What is each part of ΔH = q/n?

A

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

305
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)

306
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

307
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

308
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)

309
Q

Define average bond enthalpy

A

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

310
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.

311
Q

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

A

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

312
Q

What can average bond enthalpy also be called?

A

Mean bond enthalpy

313
Q

Where is average bond enthalpy taken from?

A

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

314
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)

315
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

316
Q

Define standard enthalpy change of combustion

A

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

317
Q

Define standard enthalpy change of formation

A

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

318
Q

Enthalpy change of formation of elements is always …

A

zero

319
Q

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

A

Using delta H f it’s p-r!

320
Q

Formula for hess’ law cycles involving enthalpies of formation

A

sum of ΔHf products - sum of ΔHf reactants

321
Q

Why does oxygen have no enthalpy of combustion value?

A

Oxygen can’t be combusted

322
Q

What does Hess’ law state?

A

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

323
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

324
Q

Does breaking bonds require or release energy?

A

Require

325
Q

Does forming bonds require or release energy?

A

Release

326
Q

Define exothermic

A

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

327
Q

Define endothermic

A

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

328
Q

Define average bond enthalpy

A

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

329
Q

Equation for enthalpy change

A

enthalpy of bonds broken – enthalpy of bonds formed

330
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
331
Q

Draw a general diagram for Hess’ law

A

CHECK NOTES

332
Q

Enthalpy change of reaction –

A

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

333
Q

Enthalpy change of formation –

A

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

334
Q

Enthalpy change of combustion –

A

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

335
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.

336
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

337
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

338
Q

What is Hess’ law used for?

A

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

339
Q

Where does Hess’ law come from?

A

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

340
Q

For elements, enthalpy change of formation is ALWAYS ____

A

zero

341
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

342
Q

Definition of Relative Atomic Mass (Ar)

A

Weighted mean mass of an atom relative to 1/12th the mass of an atom of carbon 12

343
Q

Definition of Relative Formula Mass (Mr)

A

The term used when working out the calculation for compounds with giant structures

344
Q

Name 3 types of giant structures

A

Giant IonicGiant CovalentGiant Metallic

345
Q

Definition of Relative Molecular Mass

A

The term used when working out the calculation for compounds that are simple molecules

346
Q

Name a type of simple molecular structure

A

Covalent Compounds

347
Q

What is the formula for finding out how much of a compound is made up of a particular element?

A

% of element = ((number of atoms of element x relative atomic mass of the element) / relative formula mass of compound) x 100

348
Q

What is the molar mass?

A

.The mass per mol of a substance in g mol^-1.Same number as Mr

349
Q

What is the amount of substance + unit?

A

.A means of counting the number of particles in a substance.Unit is the mol

350
Q

What is Avogadro’s Constant?

A

(NA) 6.02x10^23 mol-1

351
Q

What is everything measured relative to?

A

Carbon-12

352
Q

Define the mole

A

The amount of any substance containing as many elementary particles as there are carbon atoms in 12 grams of carbon-12 (6.02x10^23) particles

353
Q

number of mols (mol) =

A

mass (g) / molar mass (g mol-1)

354
Q

Smallest mass of an atom

A

1.67x10^-27

355
Q

The largest mass of an atom

A

4.52x10^-25

356
Q

What do we use the smallest mass of an atom to do?

A

Produce a relative scale called the unified atomic mass unit ‘u’ = 1.67x10^-27

357
Q

If hydrogen-1 = 1u, what does carbon-12 =?

A

12u

358
Q

Define relative isotopic mass

A

The mass of an atom of an isotope relative to 1/12th the mass of a carbon 12 atom

359
Q

What is 1/12th the mass of a carbon-12 atom?

A

1u

360
Q

Define relative atomic mass

A

The relative atomic mass is the ‘weighted mean’ mass of an atom relative to 1/12th the mass of a carbon-12 atom

361
Q

formula to find out relative atomic mass/weighted mean?

A

((mass x abundance) + (mass x abundance)) / 100

362
Q

Define molecular formula

A

The number and type of atoms of each element in a molecule

363
Q

Define empirical formula

A

Shows the simplest whole number ration of atoms of each element in a compound

364
Q

State the steps for finding the empirical formula

A
  1. find the mass2. calculate the moles3. divide all the results by the smallest value to get the ration4. adjust the ratio to get whole numbers5. workout the empirical formula
365
Q

What is water of crystallisation?

A

Water molecules that are bonded into a crystalline structure of a compound

366
Q

What doe solid compounds formed from aqueous solutions have trapped in their crystal structures?

A

The crystals that are formed have water molecules trapped in the crystal structure

367
Q

What does copper sulphate exist as?

A

Blue crystals

368
Q

How does copper sulphate lose its water?

A

Heat

369
Q

What colour is anhydrous copper sulphate?

A

White

370
Q

What is the formula of copper sulphate?

A

CuSO4.xH2Ox = number of water molecules

371
Q

Anhydrous definition

A

Contains no waters of crystillisation

372
Q

Hydrated definition

A

A crystallised compound containing water molecules

373
Q

Where is the water in copper sulphate?

A

Water forms part of the crystal sturcture

374
Q

Four steps of the hydrated and anhydrous practical

A
  1. Weigh an empty crucible2. Add the hydrated salt into the weighed crucible, weigh the crucible and the hydrated salt3. Using a pipe clay triangle, support the crucible containing the hydrated salt on a tripod. Heat the crucible and contents gently for about one minute. Then heat it strongly for a further three minutes.4. Leave the crucible to cool. Then weigh the crucible and anhydrous salt
375
Q

What is a binary compound?

A

Binary compounds contain 2 elements only

376
Q

How to name a binary compound with one example

A

To name it, take the first element then change second elements name to –ide (for ionic compounds the metal always comes first) for example magnesium oxide (MgO), sodium chloride (NaCl), Calcium sulphide (CaS)

377
Q

What is a polyatomic ion?

A

Ions comprised of more than one atom

378
Q

What are brackets in formulas used for?

A

Polyatomic ions need to go into brackets

379
Q

What is a redox reaction?

A

When reduction and oxidisation happen at the same time

380
Q

What is OILRIG

A

OxidisationIsLossReductionIsGain

381
Q

What assumptions were made during the copper sulphate anhydrous and hydrated salt experiment? Any extra infomation (how to prevent/any problems that could arise)

A

.That all the water has been lost – could heat the mass then weigh and repeat till it stays the same to counteract this.No further decompositions – copper oxide might have been made if we heated it too much

382
Q

When are you allowed to use half values for equation balancing?

A

In combustion equations, specifically on the oxygen

383
Q

In what order do you balance combustion equations (hydrocarbons)?

A

Carbon first then the hydrogen before finally the oxygen

384
Q

What does aqueous mean?

A

Dissolved in water

385
Q

Equipment needed for the determination of the formula for magnesium oxide practical

A

.access to balance accurate to two decimal places.Crucible and lid.bunsen burner.Tripod stand.heat-proof mat.clay pipe triangle.Tongs

386
Q

Health and Safety for the determination of the formula for magnesium oxide practical

A

.Wear eye protection at all time (safety goggles).Take care not to touch any apparatus that is hot.Take particular care at steps 3 and 4, do not look at any bright light given off by the reacting magnesium while it is being heated.Do not place the magnesium ribbon directly in the Bunsen flame

387
Q

Method for the determination of the formula for magnesium oxide practical

A
  1. Measure the mass of crucible and lid2. Put the magnesium ribbon to the crucible. You will need to coil the magnesium so that it fits. Reweigh the crucible and lid.3. Arrange the equipment with a tripod on a heatproof mat, a clay pipe triangle on the tripod, a crucible in the triangle and a bunsen burner under it. Raise the crucible lid slightly using tongs to control the reaction4. When the reaction is nearly complete, place the crucible lid on the heatproof mat and heat the crucible strongly for 5 minutes. During this time, tap the magnesium oxide gently with tongs to break up the residue5. Allow the crucible to cool and reweigh the crucible, its contents, and lid
388
Q

What is molar gas volume?

A

The molar gas volume is the volume per mole of gas molecules at the stated temperature and pressure

389
Q

What are the conditions at RTP?

A

Conditions: RTP, Room temperature and pressure20’C101KPa (1atm) pressure

390
Q

At RTP what does 1 mole of gas have a volume of?

A

24dm^3 (24000cm^3)

391
Q

What is the volume and molar gas volume equation?

A

amount n (mol) = Volume V / molar gas volume V

392
Q

How is a gas a ‘perfect’ or ‘ideal’ gas?

A

If it obeys the ideal gas equation

393
Q

When is a gas most close to obeying the ideal gas equation and why?

A

real gases obey the equation very closely at low pressure (no more than atmospheric pressure) and high temperature (room temperature). Under these conditions a gas is most like a gas and least like a liquid.

394
Q

What 5 assumptions does the kinetic theory make about gas molecules?

A

• The particles are moving in straight lines at random.• We can neglect the volume of the particles themselves in comparison with the total volume of the gas (occupy negligible volume).• The particles do not attract one another (exert no force on one another).• The kinetic energy of the particles is proportional to the temperature of the gas.• No energy is lost in collisions between particles.

395
Q

What is the ideal gas equation?

A

pV = nRT

396
Q

In the ideal gas equation what is ‘p’ and what is it measured in?

A

Pressure, measured in Pa

397
Q

In the ideal gas equation what is ‘n’ and what is it measured in?

A

Amount of gas, measured in mol

398
Q

In the ideal gas equation what is ‘T’ and what is it measured in?

A

Temperature, measured in K

399
Q

In the ideal gas equation what is ‘V’ and what is it measured in?

A

Volume, measured in m^3

400
Q

In the ideal gas equation what is ‘R’ and what is it measured in?

A

Ideal gas constant, measured in J/mol/K

401
Q

What is the ideal gas constant?

A

8.31 J/mol/K

402
Q

How to convert from cm^3 to m^3

A

x10^-6

403
Q

How to convert from dm^3 to m^3

A

x10^-3

404
Q

How to convert from ‘C to K?

A

273

405
Q

How to convert from KPa to Pa?

A

x10^3

406
Q

Max number of electrons in the first shell?

A

2

407
Q

Max number of electrons in the second shell?

A

8

408
Q

Max number of electrons in the third shell?

A

18

409
Q

Max number of electrons in the fourth shell?

A

32

410
Q

What is an electron shell?

A

A group of atomic orbitals with the same principal quantum number, n

411
Q

What’s the principal quantum number?

A

Represented as n, a number representing the overall energy level of the orbital. The bigger the number, the further the distance between the energy level and the atomic nucleus

412
Q

Formula to work out how many electrons in the shell?

A

2(n^2)

413
Q

What is an orbital?

A

A region of high probability within an atom that can hold 2 electrons with opposite spin

414
Q

What are shells made up of?

A

Orbitals

415
Q

What did Pauli discover in 1924?

A

.Orbitals only hold 2 electrons.Electrons carry a negative charge.Spin on-axis - generate a magnetic field.Spin clockwise or anti-clockwise, represented by arrows.Electrons in the same orbital must spin in different directions

416
Q

What does an s-orbital look like?

A

A sphere

417
Q

Which shells have s-orbitals?

A

From n=1 onwards, each shell contains one s-orbital (max two electrons)

418
Q

What does a p-orbital look like?

A

Dumbbell shaped, like a balloon squashed in the middle

419
Q

Which shells have p-orbitals?

A

From n=2 onwards, each shell contains three p-orbitals (max 6 electrons)

420
Q

Which shells have d-orbitals?

A

From n=3 upwards, each shell has 5 d-orbitals (max 10 electrons)

421
Q

Which shells have f-orbitals?

A

From n=4 onwards, each shell has 7 f-orbitals (max 14 electrons)

422
Q

What does each part of the notation 1s^2 represent?

A

1 = energy levels = type of orbital2 = number of electrons in orbital

423
Q

What is Aufbau’s principle?

A

Electrons fill the lowest energy orbitals in sequence

424
Q

What is the order that the orbitals fill up?

A

1s 2s 2p 3s 3p 4s 3d 4p 5s

425
Q

What is an electron energy level made up of?

A

An electron energy level is made up of atomic orbitals with the same principal quantum number

426
Q

What is a sub-shell?

A

Within each shell, orbitals of the same type are grouped together as a sub-shellsEach sub-shell is made up of only one type of orbital only, so there are s, p and d sub-shells

427
Q

What is Hunds Rule

A

Electrons singly occupy orbitals before pairing up

428
Q

When ionising, using the orbital model, which electrons are lost?

A

The ones in the outer shell, so even if 3d has electrons in it they will be lost from 4s when ionising

429
Q

Which two elements don’t follow Aufbau’s principle?

A

Copper and Chromium

430
Q

How does Copper not follow Aufbau’s principle?

A

copper steals a 4s electron to gain a full 3d orbitalIt’s actually 1s2, 2s2, 2p6, 3s2, 3p6, 4s1, 3d10

431
Q

How does chromium not follow Aufbau’s principle?

A

chromium steals a 4s electron to be able to put an electron in every 3d orbitalIt’s actually 1s2, 2s2, 2p6, 3s2, 3p6, 4s1, 3d5

432
Q

How do copper and chromium gain stable structures?

A

Cr and Cu get stable structures from full and half full 3d sub shells

433
Q

What do the big numbers in equations show?

A

The molar ratio

434
Q

What are atoms trying to achieve when they bond?

A

A full outer shell

435
Q

4 key things about ionic bonding

A

Between metals and non-metalsIts to do with the loss and gain of electronsMetals form positive chargesNon-metals form negative charges

436
Q

How does an atom gain a positive charge?

A

It loses electrons, therefore there are more protons than electorns

437
Q

How does an atom gain a nagative charge?

A

It gains electrons, therefore there are less protons than electrons

438
Q

What is an ion?

A

A charged particle

439
Q

How do ionic bonds stay together?

A

The electrostatic attraction between the positive and negative ions

440
Q

What do you use the crosses and dots for on ionic bonding diagrams?

A

use crosses for the metals electrons and dots for the non-metals electrons

441
Q

What is a lattice?

A

A lattice is a regular repeated three-dimensional arrangement of atoms, ions, or molecules in a metal or other crystalline solution.

442
Q

What is the structure of ionic compounds like?

A

A giant ionic lattice, the attraction between the oppositely charged ions acts equally in all directions, which leads to the formation of a giant ionic lattice in three dimensions

443
Q

Moles equation with concentration and volume

A

Mol = (cm^3 x moldm^-3) / 1000

444
Q

Rearrange the moles equation with concentration and volume for concentration =

A

moldm^-3 = (mol x 1000) / cm^3

445
Q

What are almost all ionic compounds at room temperature?

A

solids

446
Q

Why are almost all ionic compounds solid at room temperature?

A

At room temperature there is insufficient energy to overcome the strong electrostatic forces of attraction between the oppositely charged ions in the giant ionic lattice, high temperatures are needed to provide this energy

447
Q

Why is the melting points higher for ionic lattices containing ions with greater ionic charges? What else does ionic attraction depend on?

A

The melting points are higher for lattices containing ions with greater ionic charges, as there is a stronger attraction between ions.The ionic attraction also depends on the size of the atom

448
Q

Melting point of NaF

A

993’C

449
Q

Melting point of CaF2

A

1423’C

450
Q

Melting point of Na2O

A

1275’C

451
Q

Melting point of CaO

A

2614

452
Q

What do many ionic compounds dissolve in?

A

Polar solvents like water

453
Q

Why might an ionic compound not be soluble?

A

.Polar water molecules break down the lattice and surround each ion in the solution.But in a compound made of ions with large charges, the ionic attraction may be too strong for water to be able to break down the lattice structure, and the compound will then not be soluble

454
Q

Solubility of NaCL at 20’C

A

6.1 mol dm^-3

455
Q

Solubility of CaCl2 at 20’C

A

0.67 mol dm^-3

456
Q

Solubility of Na2CO3 at 20’C

A

2.0 mol dm^-3

457
Q

Solubility of CaCO3 at 20’C

A

1.3x10^-4 mol dm^-3

458
Q

What 2 processes does solubility require?

A

.The ionic lattice must be broken down.The water molecules must attract and surround the ions

459
Q

What does the solubility of an ionic compound in water depend on?

A

on the relative strengths of the attractions within the giant ionic lattice and the attractions between ions and water molecules

460
Q

As the solubility decreases in ionic compounds…

A

ionic charge increases

461
Q

When does and when doesn’t an ionic compound conduct electricity?

A

.In the solid state, an ionic compound does not conduct electricity.But once melted and dissolved in water the ionic compound does conduct electricity

462
Q

Describe an ionic compounds electric properties in solid state

A

.The ions are in a fixed position in the giant ionic lattice.There are no mobile charge carriers, as the ions cannot move.An ionic compound is a non-conductor of electricity in the solid state

463
Q

Describe an ionic compounds electric properties in a liquid or dissolved state

A

.The solid ionic lattice breaks down.The ions are now free to move as mobile charge carriers.An ionic compound is a conductor of electricity in liquid and aqueous states

464
Q

Summarise the properties of ionic compounds

A

Most ionic compounds – .Have high melting and boiling points.Tend to dissolve in polar solvents such as water.Conduct electricity only in the liquid state or aqueous solutions

465
Q

What is the main component in teeth, bones and tooth enamel

A

hydroxyapatite, Ca5(PO4)3OH

466
Q

What allows tooth decay?

A

Acid conditions, from food, break down enamel and allow tooth decay

467
Q

What does saliva do?

A

Salvia helps to neutralise acidic food and also to replace ions

468
Q

How do fluoride ions help teeth?

A

.Fluoride ions help to replace lost ions by forming fluoropatite, Ca5(PO4)3F, which is stronger than hydroxyapatite and more resistant to acid conditions

469
Q

What do most toothpastes and some water sources contain?

A

Most toothpastes contain fluoride as sodium fluoride, your water may also contain fluoride depending on where you live

470
Q

What is the theoretical yield?

A

The maximum possible amount that can be made

471
Q

Why is the theoretical yield never achievable?

A

.The reactions may have not gone to completion.Other reactions (side reactions) may have occurred.Purification of the product may have resulted in loss of product

472
Q

Percentage yield % =

A

(actual yield mol / theoretical yield mol) x 100

473
Q

Describe the use of aqueous barium chloride in qualitative analysis

A

Test for sulphate ions

474
Q

What is atom economy?

A

A measure of the proportion of reactants included in the final useful product

475
Q

What happens in an ideal reaction, in terms of atom economy?

A

All reactant atoms end up within the useful product molecule, no waste produced!

476
Q

What do inefficient reactions have in terms of atom economy?

A

They are wasteful and have a low atom economy

477
Q

What do efficient reactions have in terms of atom economy?

A

.High atom economy.Important for sustainable development.Conserve natural resources and create less waste

478
Q

What is a bonding pair?

A

A pair of electrons in a covalent dot n cross diagram that are being used in the reaction

479
Q

What is a lone pair?

A

A pair of electrons in a covalent dot n cross diagram that are not being used in the reaction

480
Q

How is a sigma bond formed?

A

The head-on overlap of orbitals

481
Q

In hydrogen what happens in covalent bonding, in terms of orbitals

A

The two 1s orbitals overlap and they become a molecular orbital

482
Q

How is the octet rule broken in Boron trifluoride?

A

It becomes an electron deficient molecule

483
Q

How can the octet rule be broken in sulphur hexafluoride?

A

3rd shell of S can hold 18 electrons, octet rule can be broken

484
Q

What is NH4 ^+

A

Ammonium ion, a type of molecular ion

485
Q

What is a dative covalent bond?

A

A bond formed when both electrons in the share are donated by one atom

486
Q

For an exam question what should you do?

A

READ IT ALL AND DO EVERYTHING IT TELLS YOU

487
Q

Properties of simple covalent molecules

A

.Low melting and boiling points.Weak intermolecular forces.Not soluble in polar solvents – only in other non-polar liquids.Molecules are not charged so they don’t conduct electricity (no mobile charges).Weak and soft when solid

488
Q

What does VSEPR theory stand for?

A

.Valence.Shell.Electron.Pair.Repulsion.Theory

489
Q

What is the shape of a molecule or ion determined by?

A

The shape of a molecule or ion is determined by the number of electron pairs in the outer shell of the central atom, more specifically the number of electron pairs repelling as far away from each other as possible (maximum repulsion).

490
Q

What is the valence shell?

A

The outer shell

491
Q

the bond angle of a linear molecule

A

180’

492
Q

the bond angle of a triangular planar shape

A

120’

493
Q

the bond angle of a tetrahedron

A

109.5’

494
Q

the bond angle of an octahedron

A

90’

495
Q

Molecular formula of hydrochloric acid

A

HCl

496
Q

Molecular formula of sulphuric acid

A

H2SO4

497
Q

Molecular formula of nitric acid

A

HNO3

498
Q

Molecular formula of ethanoic acid

A

CH3COOH

499
Q

Molecular formula of sodium carbonate

A

Na2CO3

500
Q

Molecular formula of sodium hydroxide

A

NaOH

OCR A Level Chemistry (17 decks)

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