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A Level Chemistry > Module 3.1 - The Periodic Table > Flashcards

Module 3.1 - The Periodic Table Flashcards

1
Q

How is the modern periodic table arranged?

A

INCREASING atomic number

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2
Q

What do periods in the modern periodic table show?

A

Gradual changes in properties across a period - periodicity.

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3
Q

What are groups in the modern periodic table?

A

Ordered elements to show similar chemical properties

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4
Q

Define periodicity

A

Trend in properties that is repeated across each period

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5
Q

Which groups are part of the s-block?

A

Groups 1, 2, H and He

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6
Q

Which groups are part of the d-block?

A

Transition metals

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7
Q

Which groups are part of p-block?

A

Groups 3 to 8 except He

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8
Q

What is the periodic trend in electron configuration?

A

The electron configuration along a period increases by one and down the group it stays the same but n increases

F —> Ne = [He]2s2 2p5 —> [He] 2s2 2p6
F —> Cl = [He] 2s2 2p5 —> [Ne] 3s2 3p5

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9
Q

What is ionisation?

A

When atom gain or lose electrons

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10
Q

What is plasma?

A

Mixture of positive and negative ions

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11
Q

What is ionisation energy?

A

The energy needed to form a positive ion (remove an electron)

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12
Q

What does first ionisation energy mean?

A

The energy required to remove one electron from each atom in one moles of the gaseous element to form one mole of gaseous 1+ ions.

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13
Q

Write an equation of Na that shows first ionisation energy.

A

Na (g) —> Na+ (g) + e-

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14
Q

Write an equation to show the second ionisation of Na

A

Na + —> Na^2+ + e^-

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16
Q

What 3 factors affect first ionisation energy?

A

> Atomic radius
Nuclear charge
Electron shielding

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17
Q

How does atomic radius affect first ionisation energy?

A

The larger the atomic radius, the smaller the nuclear attraction with outer electrons as the positive charge is further away from the outer electron. Therefore the first ionisation energy will be lower.

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18
Q

How does the nuclear charge affect ionisation energy?

A

The higher the nuclear charge, the more attracted the outer electron is to the nucleus. Therefore the higher the first ionisation energy

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19
Q

Why is each successive ionisation energy higher than the one before?

A

> As each electron is removed, there is less shielding so the shell is drawn close - atomic radius decreases
The positive nuclear charge outweigh negative charges as each electron is removed
As the distance between the nucleus and outer electrons decreases, the nuclear attraction increases. Therefore more energy is needed to remove each successive electron.

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20
Q

Why is there a massive jump between the 5th and 6th ionisation energies for nitrogen?

A

This is because the 5 electrons in the outer shell has been removed which means that the 6th ionisation energy will be a lot higher as it is an inner shell that is closer to the nucleus

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21
Q

Why does the ionisation energies increase across a period?

A

> Number of protons increases so BIGGER nuclear attraction on electrons
Same number of shells so with bigger attraction, atomic radius DECREASES
Same number of shells means SIMILAR shielding
THEREFORE more energy is needed

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22
Q

Why is there a decrease in ionisation energy from group 2 to 13?

A

> The outer electron in group 2 is in the s-orbital whereas the outer electron in group 13 is in the p-orbital.
Despite being in a higher energy level, the electron is FURTHER away so it is easier to remove the outer electron from group 13 elements.

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23
Q

Why is there a decrease in ionisation energies from group 15 to group 16?

A

> Group 15 elements have a single electron in its p-orbital, whereas group 16 elements have spin-paired electrons.
This means that the outer electron in group 16 is easier to move due to repulsion from the spin paired electrons

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24
Q

Why is there a sharp decrease in ionisation energies from the end of one period to the start of the next?

A

> As the number of shells increase, the distance from the outer electrons to the nucleus increases, therefore the nuclear attraction decrease
Also, it means an increase in electron shielding
Therefore the outer electron is easier to remove at the start of the next period when compared to the end of the last.

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25
Q

Why does ionisation energy decrease down a group?

A

> N.o of shells increases, so the distance from the nucleus to the outer electron increases, causing a weaker nuclear attraction on the outer electrons
More inner shells, so more shielding, so easier to remove outer electron
Increase in protons < n.o of shells, shielding
Therefore less energy is required

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26
Q

What is the biggest factor to affect the ionisation energies moving across a period?

A

Increased nuclear charge

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27
Q

What are the biggest factors to affect the ionisation energies moving down a group?

A

Increasing distance and shielding

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28
Q

How does electron shielding affect first ionisation energy?

A

Inner shells repel outer shell electrons as they are the same charge. The larger the number of inner shells there are, the larger the shielding effect and the smaller the nuclear attraction to the outer electrons. Therefore the lower the first ionisation energy.

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29
Q

Describe a giant metallic lattice

A

Lattice of fixed positive ions surrounded by a sea of delocalised electrons

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30
Q

What is metallic bonding?

A

Strong ELECTROSTATIC attraction between cations and negative delocalised electrons

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31
Q

Explain the high melting/boiling points of giant metallic lattices

A

> Attraction between positive ions and electrons are strong

> A lot of energy is required

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32
Q

Describe the trend in melting points in period 2 and 3

A

> Increase with giant metallic structure;
- Period 2: Li and Be
- Period 3: Na, Mg and Al
Continued increase with giant covalent structures;
- Period 2: B, C
- Period 3: Si
Decrease to simple covalent structures then flat;
- Period 2: N2, O2, F2, Ne
- Period 3: P4, S8, Cl2, Ar

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33
Q

Why is there an increase in melting points between Li to Be in period 2 and Na to Al in period 3?

A
  1. Li to Be in period 2 and Na to Al in period 3 are giant metallic structures
  2. Ionic charges increase
  3. Ionic radius decreases
  4. N.o of outer shell electrons increase
  5. Stronger attraction
  6. M.P. increases across period 2 group 1 to 2 and 3 from group 1 to 2 to 13
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34
Q

Why is there an increase in melting points between Be to C in period 2 and Al to Si in period 3?

A
  1. After Be in period 2 and Al in period 2, they become giant covalent structures
  2. Successive giant covalent structures have all covalent bonds
  3. All these strong covalent bonds must be broken
  4. A lot of energy required
  5. Higher M.P
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35
Q

Why is there a sudden decrease from C to N2 in period 2 and Si to P4 in period 3?

A
  1. After N2 in period 1 and P4 in period 2, the elements are simple covalent structures
  2. Only need to break the weak IMF
  3. Lower m.p than the giant covalent structures
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36
Q

What happens to the melting points between N2 to Ne in period 2 and P4 to Ar in period 3?

A

Stays low

37
Q

Describe the structure of diamond

A

Forms a lattice where each carbon atom is bonded to four other carbon atoms around it. This makes it extremely hard with a high melting point.

38
Q

Describe the structure of graphene

A

A two-dimensional giant lattice, one carbon atom thick, of interlocking hexagonal carbon rings. It is extremely strong, light and can conduct electricity.

39
Q

Describe the structure of graphite

A

A layered structures with delocalised electrons between layers, so it can conduct electricity. The layers slide over each other easily as it bonded by weak IMFs.

40
Q

Describe the structure of silicon

A

Network of atoms all bonded by strong covalent bonds. Similar structure to diamond, it has a giant covalent lattice.

41
Q

Describe the high melting/boiling point of giant metallic lattices

A

The attraction between the fixed positive ions and the delocalised electrons are strong. High temperatures are needed to overcome the metallic bonds

42
Q

Describe the good electrical conductivity of giant metallic lattices

A

Delocalised electrons can carry a current, allowing metals to conduct electricity even whilst solid.

43
Q

Why are metals malleable and ductile?

A

The delocalised electrons allows atoms or layers to slide past each other.

44
Q

What electron configuration do all group 2 elements have? What does this mean?

A

S2 - means that they LOSE 2 electrons to become 2+ ions

45
Q

Fill in the blanks:

As Group 2 elements lose 2 electrons to become 2+ ions, they are ______ agents and electron ________ .

A

Reducing agents / electron donors

46
Q

Describe the reaction of group 2 elements with water. Include an equation

A

All group 2 elements except beryllium react with water. Mg reacts slowly but each metal reacts more vigorously down the group. Fizzing and dissolving of the solid is to be observed.

M (s) + 2H2O (l) —> M(OH)2 (aq) + H2 (g)

47
Q

Describe the reaction of group 2 metals with oxygen. Include an equation

A

Group 2 metals react VIGOROUSLY with oxygen. The product is an ionic oxide.

2M (s) + O2 (g) —> 2MO (s)

48
Q

Why are group 2 elements more reactive further down the group? Use Mg and Ca as examples

A

Ca has a bigger ionic radius, more shells, more shielding which means that the outer electrons are easier to lose. Therefore group 2 elements are more reactive down the group.

49
Q

Describe the reaction of group 2 elements with dilute acids

A

All group 2 elements except beryllium react with dilute acids to form a salt and hydrogen gas. The reaction is more vigorous down the group.

M (s) + 2HCl (aq) —> MCl2 + H2

50
Q

Describe the trend in reactivity of group 2 metals in terms of first and second ionisation energies.

A

Each successive element down the group has its outer electron in a higher energy level, meaning they have a larger atomic radius and therefore more shielding. This means that the two outer electrons are lost more easily going down the group to barium as first and second ionisation energies decrease down the group. Therefore the reactivity increases as you go down the group.

51
Q

Describe the trend in alkalinity of group 2 hydroxides

A

The solubility of the hydroxides increases as you go down the group. This means that each successive element down the group with release more OH- ions. Therefore alkalinity INCREASE down the group and pH INCREASES down the group

52
Q

What is a use for Ca(OH)2?

A

Neutralising acidic soils

53
Q

What is a use for Mg(OH)2?

A

Indigestion remedies - neutralises acidic stomachs

54
Q

What is a use for CaCO3?

A

Building and construction uses - present in limestone and marble

55
Q

What type of molecules are halogens?

A

Diatomic

56
Q

Explain the trend in boiling points for Halogens

A

Each successive element has an extra shell of electrons which leads to a higher level of London forces (induced-induced dipole forces). Therefore boiling point increases down the group

57
Q

What is the general electron configuration of halogens?

A

nS^2 nP^5

58
Q

What type of ions do halogens formed ?

A

Halogens gain one electron to form 1- ions

59
Q

What is the trend in boiling points for halogens?

A

Increase down the group

60
Q

What is the trend in reactivity for halogens?

A

Decrease down the group

61
Q

Explain the trend in reactivity of halogens

A

> Atomic radius increases - nuclear attraction is smaller
Electron shielding increases
Therefore the ability to gain an electron in the p-orbital and from 1- ions decreases

62
Q

How can the reactivity of halogens be shown by redox reaction?

A

Reacting aqueous solution of halide ions (Br-, Cl-, I-) and aqueous solutions of halogens (Br2, Cl2, I2). A more reactive halogen will displace a less reactive halogen. For example:
> Cl2 + 2Br- —> 2Cl- + Br2
> Cl2 + 2I- —> 2Cl- + I2

Chlorine oxidises Br - and I- which means that it is the most reactive out the three

63
Q

Define disproportionation

A

The oxidation and reduction of the same element in a redox reaction

64
Q

Describe the disproportionation reactions of chlorine with cold dilute aqueous sodium hydroxide - form bleach

A

Forms household bleach at room temperature. Chlorine is reduced and oxidised.

Cl2(aq) + 2NaOH(aq) —> NaClO(aq) + NaCl(aq) + H2O (l)
0 -1 Reduced
0 +1 Oxidised

65
Q

Describe the disproportionation reaction of chlorine with water

A

Chlorine kills bacteria, making water safe to drink. Chlorine is reduced and oxidised in this reaction.

Cl2(aq) + H2O(l) —> HClO(aq) + HCl(aq)
0 -1 Reduced
0 +1 Oxidised

66
Q

What are the benefits of adding chlorine to water?

A

Kills bacteria and water-carried diseases such as cholera.

67
Q

What are the risks of adding chlorine to water?

A

> Some believe that it react with organic matter to form traces of chlorinated hydrocarbons which cause cancer.
Chlorine is a toxic gas

68
Q

How do you test for carbonate ions?

A

> Add a strong dilute acid

> Collect any gas and pass through limewater

69
Q

What are the positive test observations for the carbonate ion test?

A

> Fizzing/bubbling

> Limewater turns cloudy

70
Q

Do group 2 elements oxidise or reduce? Include equation

A

They get oxidised as they lose electrons

M —> M^2+ + 2e^-

71
Q

What is the ionic equation for carbonate ion test?

A

CO3^2-(aq) + 2H^+ —> H2O(l) + CO2 (g)

72
Q

What are the positive test observations for the sulfate ion test?

A

White precipitate

73
Q

How do you test for sulfate ions?

A

> Add dilute hydrochloric acid

> Add barium chloride

75
Q

What is the ionic equation for the sulfate ion test?

A

Ba^2+(aq) + SO4^2-(aq) —> BaSO4(aq)

76
Q

How do you test for halide ions?

A

> Dissolve in water
Add aqueous silver nitrate
If colour is not distinguishable, then add aqueous ammonia.

77
Q

How do you test for ammonium ions?

A

> Add sodium hydroxide solution
Warm gently
Test with RED litmus paper

78
Q

What are the positive test observations of the ammonium ion test?

A

Litmus turns blue

79
Q

What is the ionic equation for the ammonium ion test?

A

NH4^+(aq) + OH^-(aq) —> NH3(aq) + H2O(aq)

80
Q

What are the ionic equations for the halide ion test?

A

> Cl^- : Ag^+(aq) + Cl^-(aq) —> AgCl(s)

> Brl^- : Ag^+(aq) + Br^-(aq) —> AgBr(s)

> l^- : Ag^+(aq) + I^-(aq) —> AgI(s)

81
Q

What is the order of anion tests?

A
  1. Carbonate
  2. Sulfate
  3. Halide
82
Q

What are the positive test observations for the halide test?

A

> Cl^- : White precipitate
Br^- : Cream precipitate
I^- : Yellow precipitate

83
Q

What model does the trend in ionisation energy support? How?

A

Bohr’s model where describes a positive nucleus, with a negatively charged electrons orbiting around it in DEFINED ENERGY LEVELS.

84
Q

Describe the solubility of giant metallic structure

A

Do not dissolve in solvents

85
Q

What type of electrons do covalent bonds involve?

A

Localised electron - cannot move freely - fixed between bonded atoms.

86
Q

What type of electrons do metallic bonds involve?

A

Delocalised electrons - can move freely throughout the lattice and are effectively shared by all the positive ions in the lattice.

87
Q

Describe the reactions between group 2 oxides and water.

A

Group 2 oxides react with water to form metal hydroxides.

M (s) + 2H2O (l) —> M(OH)2 (aq) + H2 (g)

88
Q

Describe the solubility of metal hydroxides from the reaction between group 2 oxides and water. pH?

A

> The metal hydroxides are soluble in water and form alkaline solutions with water as they release OH- ions.
The typical pH of these solutions is between 10 and 12.

A Level Chemistry (7 decks)

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