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Chemistry > Module 3 - The Periodic Table And Periodicity > Flashcards

Module 3 - The Periodic Table And Periodicity Flashcards

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1
Q

What is the magnitude of an ionisation energy influenced by??

A

Nuclear charge

Atomic radius

Shielding

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2
Q

Effect of nuclear charge

A

The greater the nuclear charge the greater the attraction on the outer electrons

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3
Q

Effect of atomic radius

A

The further the electrons is from the nucleus, the weaker the attraction

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4
Q

Effect of shielding

A

Full inner shells of electrons exert a repelling effect on outer electrons. This reduces the attraction between the nucleus and outer electrons

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5
Q

Trends in first ionisation energy

Decrease down a group

A

Nuclear charge increases

Atomic radius increases so outer electron is further from the nucleus

Shielding increases as the number of full inner shells increases

Increase in distance and shielding outweighs the increased nuclear charge

Nuclear attraction on outer electron decreases

First ionisation energy decreases

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6
Q

Trends in first ionisation energy

General increase across a period

A

Nuclear charge increases

Atomic radius decreases so outer electron is closer to the nucleus

Shielding stays the same as electrons are added to the same shell

Nuclear attraction on outer shell increases

First ionisation energy increases

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7
Q

Decrease between groups 2 and 3

A

2p orbitals have slightly more than 2s

S electrons provide slightly greater shielding of the p electron

Less energy needed to remove outer electron despite increased nuclear charge

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8
Q

Decrease between groups 5 and 6

A

Spin pairing occurs in 2p subshell of oxygen

Paired electrons in a 2p orbital of oxygen repel each other

Less energy needed to remove outer electron despite increased nuclear charge

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9
Q

What is the atomic orbital

A

Region around the nucleus that can hold up to 2 electrons with opposite spin

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10
Q

General trend in the first ionisation energies of the first 11 elements

A

Nuclear charge increases
Atomic radius decreases

Outer electron is closer to the nucleus, shielding stays the same as electrons are added to the same shell

So ionisation energy increases across a period

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11
Q

Group number

A

Number of outer shell electrons

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12
Q

Period number

A

Number of electron shells

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13
Q

Periodicity

A

The repeating trend in physical and chemical properties across periods of the period table

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14
Q

Structure of a metallic lattice

A

Cations are in fixed positions giving the shape of the metal

Delocalised electrons are mobile and can move throughout the structure

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15
Q

What is a metallic bond??

A

A strong electrostatic attraction between a lattice of positive ions and a sea of delocalised electrons

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16
Q

Electrical conductivity in giant metallic lattices

A

Delocalised electrons can move and carry charge

Electrons will move towards the positive terminal, electrons will be supplied by the negative terminal

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17
Q

Melting and boiling points in giant metallic lattices

A

High due to large amount of energy needed to overcome the strong electrostatic attractions between positive ions and delocalised electrons

Metallic bond strength depends on number of outer shell electrons to be delocalised and charge density of the metal ion

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18
Q

Solubility in giant metallic lattices

A

All insoluble

Some will react with water

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19
Q

Structure of giant covalent lattices

A

Atoms held together in a network of strong covalent bonds forming a giant covalent lattice

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20
Q

Electrical conductivity of giant covalent lattices

A

Non-conductors, no delocalised electrons all outer shell electrons are used in bonding

Exception - graphite - delocalised electrons are present which can move

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21
Q

Melting points and boiling points of a giant covalent lattice

A

High due to the large amount of energy needed to break the many strong covalent bonds

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22
Q

Solubility of giant covalent lattices

A

All insoluble

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23
Q

What is a homologous series??

A

A series of organic compounds having the same functional group but with each successive member differing

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24
Q

What is the general formula??

A

The simplest algebraic formula for a member of a homologous series

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25
Q

What is the structural formula??

A

The minimal detail that shows the arrangement of atoms in a molecule

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26
Q

What is the displayed formula??

A

The relative positioning of atoms and the bonds between them

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27
Q

What is the skeletal formula??

A

The simplified organic formula showing only a carbon skeleton and associated functional groups

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28
Q

Functional group

A

A group of atoms responsible for the characteristic reactions of a compound

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29
Q

Formula for an alkyl group

A

CnH2n+1

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30
Q

What does aliphatic mean??

A

A compound containing carbon and hydrogen

Joined in straight chains, branched chains or non-aromatic rings

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31
Q

What does alicyclic mean??

A

An aliphatic compound arranged in non-aromatic rings with or without side chains

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32
Q

What does aromatic mean??

A

A compound containing a benzene ring

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33
Q

What does saturated mean??

A

Single carbon carbon bonds only

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34
Q

What does unsaturated mean??

A

The presence of multiple carbon-carbon bonds

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35
Q

What is a structural isomer??

A

Compounds with the same molecular formula but different structural formula

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36
Q

Why does ionisation energy increase across a period??

A

Increasing nuclear charge

No change in shielding

Atomic radius decreases

Nuclear attraction increases

More energy needed to remove outer electron

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37
Q

Why does ionisation energy decrease down a group??

A

Atomic radius increases

More full inner shells so more shielding

These factors outway the increased nuclear charge

Nuclear attraction decreases

Less energy is needed to remove the outer electron

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38
Q

Why do successive ionisation energies increase??

A

Same nuclear charge

Fewer electrons

Electrons pulled closer to the nucleus

Nuclear attraction increases

More energy needed to remove the next electron

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39
Q

What is atomic orbital??

A

Area around the nucleus that can hold 2 electrons with opposite spin

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40
Q

What is an orbital??

A

A region around the nucleus that can hold up to 2 electrons with orbital spin

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41
Q

What is a shell??

A

Region around the nucleus which contains orbitals with the same principle quantum number

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42
Q

What is a subshell??

A

A group of orbitals with the same principle quantum number

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43
Q

How to detect carbonate ions are present

A

Add dilute nitric acid

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44
Q

Indication that carbonate ions are present when dilute nitric acid is added

A

Effervescence

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45
Q

How to prove gas produced is carbon dioxide??

A

Bubble gas through limewater which will turn cloudy

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46
Q

How to detect sulphate ions

A

Add nitric acid and barium nitrate solution

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47
Q

What will occur if sulphate ions are present when nitric acid and barium nitrate solution is added??

A

A white precipitate of barium sulphate shows the presence of sulphate ions

48
Q

Equation of the qualitative analysis of carbonate ions

A

2H+ + CO32- ~~~> CO2 + H20

49
Q

Equation of the qualitative analysis of sulphate ions

A

Ba2+ + SO42- ~~~> BaSO4

50
Q

How to detect for the presence of a halide ion

A

Add nitric acid and then silver nitrate solution to an aqueous solution of the sample

51
Q

After adding nitric acid and silver nitrate solution to a sample, what will indicate a halide ion is present

A

A white, cream or yellow precipitate will mean a halide ion is present

52
Q

Equation of the qualitative analysis of halide ions

A

Ag+ + X- ~~~> AgX

53
Q

How to determine WHICH halide ion is present

A

Add ammonia

54
Q

If a Cl- halide ion is present what will be the colour of the precipitate when ammonia is added

A

White

55
Q

If a Br- halide ion is present what will be the colour of the precipitate when ammonia is added

A

Cream

56
Q

If a I- halide ion is present what will be the colour of the precipitate when ammonia is added

A

Yellow

57
Q

What is the solubility of a Cl- ion in NH3

A

Soluble in dilute NH3

58
Q

What is the solubility of a Br- ion in NH3

A

Soluble in concentrated NH3

59
Q

What is the solubility of an I- ion in NH3

A

Insoluble in concentrated NH3

60
Q

Test for carbonates

A

Add nitric acid

61
Q

Test for sulphates

A

Add barium nitrate

62
Q

Test for halides

A

Add silver nitrate

63
Q

Which ion should be tested for first and why??

A

By testing for carbonate ions first, subsequent tests are only needed if the unknown is not a carbonate

64
Q

When does the sulphate test not work??

A

When carbonate ions are present

If carbonate ions are present, a false positive could be produced

65
Q

Qualitative analysis of ammonium ions

A

Add sodium hydroxide solution

66
Q

What happens when sodium hydroxide solution is added to a solution containing ammonium ions??

A

Ammonia gas is produced

No effervescence as the ions are very soluble

67
Q

Equation for the qualitative analysis of ammonium ions

A

NH4+ + OH- ~~~> NH3 + H2O

68
Q

How to test for ammonia gas

A

Warm the mixture to release ammonia gas

Test the gas with damp RED litmus, the alkaline gas will turn the litmus BLUE

69
Q

When is limestone used in the blast furnace??

A

In the extraction of iron

70
Q

When limestone is heated what is formed??

A

CaO (quicklime) and CO2

71
Q

What happens when limestone is firstly heated then water is added??

A

Slaked lime is formed

Ca(OH)2

72
Q

What happens when limestone is heated then excess water is added??

A

Limewater is formed

Ca(OH)2

73
Q

What happens when CO2 is added to limewater??

A

Calcium carbonate is formed

Limewater turns milky because a precipitate is formed

74
Q

What happens when excess CO2 is added to limewater

A

Ca(HCO3) is formed

Present in hard water

75
Q

Why is limestone heated with clay in a kiln

A

To make cement

76
Q

When cement is mixed with sand, water and crushed rock what is formed??

A

Concrete

77
Q

Why is limestone used in lakes and soils

A

To neutralise acidity

78
Q

What is oxidation

A

Loss of electrons

Gain of oxygen

79
Q

What is reduction??

A

Gain of electrons

Loss of electrons

80
Q

Physical properties of group 2 elements

A
  • High melting and boiling point (higher than group 1 but lower than transition elements)
  • atomic radii increases upon descent down a group
  • increased shielding outweighs increased nuclei charge
  • ionisation energy decreases down the group
81
Q

Redox reactions of group 2 elements

A

All group 2 elements react readily with o2, h2o and dilute acids

82
Q

What is the trend in ionisation energies down group 2

A

Sum of the 1st and 2nd ionisation energies decrease down the group

83
Q

Reactions of group 2 elements with oxygen

A

All burn in oxygen

Reactivity increases down a group

84
Q

Observation of magnesium reacting with oxygen

A

Silver solid burns with a bright white light to form a white solid

85
Q

Equation of magnesium reacting with oxygen

A

Mg (s) + O2 (s) ~~~> 2MgO (s)

86
Q

Reaction of group 2 elements with water

A

Reactivity increases down a group

87
Q

Reaction of beryllium with water

A

No reaction

88
Q

Reaction of magnesium with water

A

Reacts with steam

Silver solid reacts slowly with effervescence to produce a cloudy solution

Resulting solutions are alkaline

Alkalinity increases down a group as the hydroxide becomes more soluble

89
Q

Reaction of ca/sr/Ba with water

A

Reacts with cold water

90
Q

Group 2 reactions with dilute acids

A

All react with dilute acids

Reactivity increases down the group

91
Q

Uses of group 2 compounds

A

Limestone

Quarried for building materials

Products used to neutralise acidic soil/emissions

92
Q

Uses of Mg(OH)2 and CaCO3

A

Antacids

Neutralise excess stomach acid and prevent acid indigestion

93
Q

Uses of BaSO4

A

Barium meal

Insoluble so not absorbed into the blood stream

X-rays cannot pass through so enables X-rays of the digestive tract to be taken

94
Q

Physical properties of the halogens

A

Low mp and bp (increases down the group due to increasing strength of induced dipole dipole interactions)

Atomic radio increases down the group

Solubility - non polar molecules dissolve well in non polar solvents

95
Q

Appearance of cl2 at rtp

A

Pale green gas

96
Q

Appearance of Br2 at rtp

A

Orange/brown liquid

97
Q

Appearance of I2 at rtp

A

Silver/grey solid

98
Q

Colour of cl2 in a non polar solvent

A

Pale green

99
Q

Colour of br2 in a non polar solvent

A

Orange

100
Q

Colour of i2 in a non- polar solvent

A

Purple/pink

101
Q

Colour of cl2 in water

A

Very pale green

102
Q

Colour of br2 in water

A

Orange

103
Q

Colour of i2 in water

A

Yellow/brown

104
Q

Oxidising strength in halogens

A

Decreases down a group

105
Q

Displacement reactions in halogens

A

An element higher up in the group will displace an ion lower down from its compound

106
Q

Reaction of chlorine with water

A

Chlorine disproportionates

107
Q

Equation to show the disproportionation of Cl2

A

Cl2 + H2O ~>~

108
Q

What is disproportionation

A

The same element is simultaneously oxidised and reduced

109
Q

Disproportionation of chlorine with an alkaline

A

2NaOH + Cl2 ~~~> NaCl + NaOCl + H2O

110
Q

Conditions of a reaction between chlorine and an alkali

A

Cold (below room temp)

Using dilute NaOH

111
Q

Ionic equation EXAMPLE with silver nitrate

A

X- + Ag+ ~~~> AgX

112
Q

What is used to distinguish between Cl and Br/I

A

Use Silver nitrate

113
Q

Uses of chlorine

A

Water treatment

Chlorine bleach

Bromine manufacture

Production of Organochlorine compounds such as PVC, solvents and Pesticides

114
Q

Pros of cl2 in water treatment

A

Kills bacteria

Prevents spread of disease through water supply

115
Q

Cons of cl2 in water treatment

A

Toxic so careful handling required

May react with hydrocarbons in water forming harmful substances

Chemistry (7 decks)

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