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Chemistry A Level > Periodic Table > Flashcards

Periodic Table Flashcards

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1
Q

How does the distance away from the nucleus affect the energy of an energy level?

A

An orbital has more energy when it’s further away from the nucleus

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2
Q

How many electrons can the first shell hold?

A

2

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3
Q

How many electrons can the 2nd shell hold?

A

8

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4
Q

How many electrons can the 3rd shell hold?

A

18

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5
Q

How many electrons can the 4th shell hold?

A

32

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6
Q

How many electrons can an orbital hold?

A

2

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7
Q

How many orbitals does the s sub-shell have?

A

1

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8
Q

How many orbitals does the p sub-shell have?

A

3

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9
Q

How many orbitals does the d sub-shell have?

A

5

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10
Q

How many orbitals does the f sub-shell have?

A

7

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11
Q

What shape is an s orbital?

A

A spherical shape

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12
Q

What shape is a p orbital?

A

A dumb-bell shape

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13
Q

Which sub-shell is filled after 3p?

A
  • 4s, 3d is filled after 4s as it is higher in energy than 4s
  • Electrons from 4s are removed before electrons from 3d when forming ions
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14
Q

When filling box diagrams do you fill the boxes 2 electrons at a time?

A
  • No, you fill each box with one electron each first and then backfill
  • When a sub-shell is full you can move onto the next sub-shell
  • One arrow in each box in box diagrams should be pointing up, one should be pointing down
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15
Q

What is the electronic configuration of copper?

A

[Ar] 4s^1 3d^10

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16
Q

What is the electronic configuration of chromium?

A

[Ar] 4s^1 3d^5

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17
Q

Why do shells that are further away from the nucleus have more energy?

A

Electrons need energy to climb shells meaning that electrons furthest away from the nucleus have climbed the most shells so they must have the highest energy. The shells that have the electrons with the highest energy have the highest energy

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18
Q

Why do electrons in the same orbitals have arrows pointing in opposite directions?

A

It represents the electron’s tendency to repel each other

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19
Q

Why is electrons added to and removed from the 4s orbital before the 3d orbital?

A

The 3d sub-shell has more energy than the 4s sub-shell so the 4s sub-shell is filled and has electrons removed from it before the 3d sub-shell

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20
Q

What is ionisation energy?

A

The energy required to remove an outer electron from the attractive force of the nucleus

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21
Q

What is first ionisation energy?

A

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

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22
Q

What is the trend in ionisation energy going down a group?

A

Ionisation energy decreases going down a group as the atomic radius increases, this means that the nuclear attraction experienced by the outer electrons decreases as outer electrons are further away and experience a greater shielding effect as the number of shells increases going down the group

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23
Q

Why does ionisation energy increase going across a period?

A

The amount of protons in the nucleus increases meaning that the nuclear attraction increases. The shielding effect that the outer electrons experience is similar across a period as they’re in the same shell, the atomic radius also decreases going across a period as there is a stronger nuclear attraction because the number of protons increases as you go across a group. A smaller atomic radius means a higher ionisation energy

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24
Q

Why does the ionisation energy decrease between magnesium and aluminium (and beryllium and boron) when the trend is that ionisation energy increases across a period?

A

Mg has 12 electrons, aluminium has 13. Aluminium’s valence electron is shielded more as it’s from the 3p orbital and is therefore shielded by the two electrons in the 3s orbital. Magnesium’s valence electron is in the 3s orbital so it’s shielded by less electrons. Aluminium’s valence electron also experiences less nuclear attraction as it is in the 3p orbital which has a slightly higher energy level than the 3s orbital, therefore it’s further away from the nucleus and experiences less nuclear attraction. Experiencing less nuclar attraction and more shielding means that it takes less energy to remove Aluminium’s valence electron and therefore aluminium has a lower first ionisation energy.

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25
Q

Why does the ionisation decrease between nitrogen and oxygen when the ionisation energy is meant to increase across a period?

A

The ionisation energy decreases between oxygen and nitrogen because nitrogen has no orbitals in the 2p sub-shell that are fully filled, whereas oxygen has one orbital that is fully filled in the 2p sub-shell. The outer electron is now spin paired in the first 2p orbital. The electrons that are spin paired the ionisation energy is lower for oxygen than nitrogen

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26
Q

Why does the melting point increase across period 3 metals?

A

The atomic radius of the cations decreases while the charge increases, there are also more delocalised electrons so the strength of the metallic bonds is increased meaning that the melting point is higher

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27
Q

What are the structures of the period 3 elements?

A
  • Na, Mg, Al - giant metallic lattice
  • Si - giant covalent lattice
  • P, S, Cl - simple molecules
  • Ar - atoms
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28
Q

How does metallic bonding work?

A

There are strong electrostatic forces of attraction between cations and delocalised electrons as they’re oppositely charged

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29
Q

What is a giant covalent lattice?

A

A network of atoms joined together by covalent bonds

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30
Q

What is a simple molecule?

A

A molecule with strong covalent bonds between atoms but weak intermolecular forces between molecules meaning they have low melting and boiling points

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31
Q

What type of bonding do noble gases have?

A

Weak interatomic forces between atoms

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32
Q

Explain the trends in boiling point in period 3

A
  • Na, Mg and Al have high boiling points due to the strong electrostatic forces of attraction between cations and delocalised electrons
  • Silicon has a high boiling point due to their strong covalent bonds that have to be broken
  • P, S and Cl have low boiling points because of the weak intermolecular forces that have to be overcome (the covalent bonds don’t need to be overcome for the molecule to boil or melt
  • Ar has a low boiling point due to the weak interatomic forces that have to be overcome
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33
Q

Why are metals good conductors?

A

They have delocalised electrons that can move throughout the structure making it a good conductor

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34
Q

Is silicon a good conductor?

A

No, it has a limited number of electrons that can move through the structure meaning it’s a poor conductor

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35
Q

Are simple molecules and noble gases good conductors?

A

No, they have no free electrons meaning they’re poor conductors of electricity

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36
Q

Why is silicon insoluble in water?

A

There are no possible interactions between silicon and water molecules, the covalent bonds require a lot of energy to overcome and no compensating energy is released due to there being no interactions so they can not be broken meaning that silicon is insoluble in water

37
Q

What happens when water is added to Na?

A

A vigorous reaction occurs and NaOH and H2 are formed

38
Q

What happens when water is added to Mg?

A

Mg(OH)2 and H2 are formed (slow reaction)

39
Q

What happens to Al when water is added?

A

It is insoluble meaning there is no reaction

40
Q

What happens when water is added to Si?

A

Insoluble

41
Q

What happens when water is added to P?

A

Insoluble

42
Q

What happens when water is added to S?

A

Insoluble

43
Q

What happens when water is added to Cl?

A

Soluble (slight reaction)

44
Q

What happens to Ar when water is added?

A

It dissolves (soluble)

45
Q

Write the equation for the second ionisation of magnesium

A

Mg+(g) → Mg2+(g) + e-

46
Q

What is the equation for the reaction between metal and water?

A

Metal + water → metal hydroxide + hydrogen

47
Q

What is the difference between magnesium’s reaction with cold water and steam?

A

It reacts slowly with water and more vigorously with steam

48
Q

What happens when calcium reacts with water?

A

A white precipitate of calcium hydroxide is formed

49
Q

What happens when strontium and barium react with water?

A

A small precipitate is formed when strontium reacts with water with more dissolved product, barium reacts with water to form even less precipitate

50
Q

What happens to group 2 elements when heated in air?

A

They burn vigorously to from white metal oxides

51
Q

How do you test for carbonate ions?

A
  • Add dilute nitric acid to the sample, and then pass any gas evolved through limewater
  • A positive result is when the limewater goes milky due to a white precipitate being formed, the precipitate is a suspension of calcium carbonate
52
Q

How do you test for sulfate ions?

A
  • Add drops of barium nitrate to a solution of the sample
  • A positive result is a white precipitate of barium sulfate
  • Ba2+ (aq) + SO42-(aq) → BaSO4 (s)
53
Q

How do you test for halide ions?

A

Add silver nitrate to a solution of the sample, white (milk) precipitate if chloride ions are present, cream precipitate if bromide ions are present, yellow (butter) precipitate if iodide ions are present
- Confirm the result with ammonia (Cl- dissolve in dilute or concentrated ammonia, Br- dissolve in concentrated ammonia, I- insoluble in dilute and concentrated ammonia)

54
Q

What are the ionic equations for the reactions between silver nitrate and the different halogens?

A
  • Ag+(aq) + Cl-(aq) → AgCl(s)
  • Ag+(aq) + Br-(aq) → AgBr(s)
  • Ag+(aq) + I-(aq) → AgI(s)
55
Q

How do you test for ammonium ions?

A
  • Add sodium hydroxide to sample, heat and test any gas evolved with damp red litmus paper (turns blue if ammonia is present), there will also be a strong smell
  • NH4+(s) + OH-(aq) → NH3(g) + H2O(l)
56
Q

What order should the tests for carbonate, sulfate, ammonium and halide ions be done in and why?

A

1) Carbonate ions
2) Sulfate ions
3) Halide ions
4) Ammonium ions
- This is because carbonate ions would give a precipitate in tests 2 and 3 as well if they hadn’t already been detected in test 1 (alternatively you could add dilute nitric acid in tests 2 and 3 as dilute nitric acid removes carbonate ions)
- Sulfate ions would produce a white precipitate in test 3 if not detected already in test 2

57
Q

How many ionisation energies does a group 3 element have?

A

3

58
Q

How many ionisation energies does a group 2 element have?

A

2

59
Q

Why is each successive ionisation energy greater than the previous ionisation energy?

A
  • As each electron is removed there is less repulsion between the remaining electrons and each shell is drawn in slightly closer to the nucleus
  • As the distance of each electron to the nucleus decreases slightly as the shell is drawn in, the nuclear attraction increases, this means that more energy is required to remove each electron
60
Q

What happens to the boiling point of a metal as its atomic number increases?

A
  • As the atomic number of a metal atom increases, the nuclear charge increases and the sea of electrons becomes more negatively charged
  • This means there is a stronger attraction between the negatively charged sea of electrons and the positively charged nucleus, this means the ionic size decreases
  • This means that more energy is required to break the strong electrostatic forces of attraction between metal ions and the sea of electrons and therefore, boiling point increases as atomic number increases
61
Q

Why does the ionisation energy decrease between groups 5 and 6 (between N and O, and also between P and S)

A

In group 5, elements have one electron in each of their p-orbitals. In group 6 the outermost electron is spin paired in the p-orbital. Electrons that are spin paired experience some repulsion, this makes the valence electron slightly easier to remove. Therefore the ionisation energy is lower for group 6 elements than for group 5 elements.

62
Q

What is the trend in ionisation energy from one period to the start of another period?

A
  • Ionisation energy sharply decreases as you move from the end of one period to the start of another one, this is because anew shell is added as you move from one period to the next
  • As a new shell is added, the valence electron is further away from the nucleus meaning it experiences less nuclear attraction and because the valence electron is in a new shell, it experiences a greater shielding effect, this means that the ionisation energy decreases dramatically
63
Q

Can silicon conduction electricity?

A

Yes but very poorly

64
Q

How is diamond’s lattice formed?

A

Each carbon atom is surrounded by 4 carbon atoms

65
Q

How is graphene’s lattice formed?

A

Graphene forms a 2D lattice, one carbon atom thick of interlocking hexagonal rings

66
Q

Describe the structure of a giant metallic lattice

A
  • Delocalised electrons are spread throughout the structure
  • These electrons can move throughout the structure
  • It is impossible to tell which electron originated from which particular ion
  • Over the whole structure the charges of metal ions and electrons balance out giving the molecule an overall negative charge
67
Q

Why do metals have high melting and boiling points?

A

There are very strong electrostatic forces of attraction between the positively charged metal ions and the negatively charged sea of delocalised electrons, these forces of attraction require a lot of energy to overcome to dislodge the metal ions from their rigid positions in the lattice and therefore metals have a high melting and boiling point

68
Q

Why are metals good conductors of electricity?

A

The delocalised electrons are free to move anywhere within the lattice and therefore can conduct electricity even in the solid state

69
Q

Why are metals described as ductile?

A

They can be drawn out or stretched, this allows metals to be drawn into wires, they’re ductile because the layers of the metal can slide over each other due to the delocalised electrons

70
Q

Why are metals described as malleable?

A

They can be pressed into different shapes or hammered into thin sheets, they’re malleable because the layers of the metal can slide over each other due to the delocalised electrons

71
Q

What are the physical properties of the group 2 metals

A
  • Have reasonably high melting and boiling points
  • Have low densities
  • Form colourless (white) compounds
72
Q

Why does the reactivity increase going down group 2?

A

Going down the group, each successive element has a larger atomic radius meaning that its valence electron is further away from the nucleus meaning it experiences a weaker nuclear attraction. It is also in a new shell meaning it experiences a greater shielding effect, this means that the reactivity increases going down the group as outer electrons are easier to remove going down the group

73
Q

What do group 2 oxides react with water to form?

A

Metal hydroxides

74
Q

What si the trend in solubility of group 2 hydroxides?

A

Solubility increases going down the group as does alkalinity

75
Q

What are the main uses of the group two compounds?

A
  • Ca(OH)is used by farmers and gardeners to reduce the acidity levels of soil
  • Mg(OH)2 is used as an indigestion remedy as it can neutralise excess stomach acid, Mg(OH)2 + HCl → MgCl2 +2H2O
  • Building and construction uses, CaCO3 is an extremely useful building material, it is present in both limestone and marble
76
Q

What is a disadvantage of using group 2 carbonates as building materials?

A

They react with acids to form a salt, water and CO2. This is a problem as most rainwater has an acidic pH, this leads to gradual erosion of objects made of limestone or marble, such as buildings and statues

77
Q

Do the halogens have high or low melting and boiling points

A

Low melting and boiling points

78
Q

What is the trend in boiling point going down group 7?

A

Boiling point increases as the number of electrons increases meaning the strength of London forces between molecules increases and therefore more energy is required to overcome these intermolecular forces and boiling point therefore increases going down the group as the number of electrons increases going down the group

79
Q

What is the trend in reactivity going down group 7?

A

Reactivity decreases going down the group as atomic radius increases. As the atomic radius increases the nuclear attraction experienced by the incoming electrons decreases. Also going down the group, the number of shells increases meaning that incoming electrons experience a greater shielding effect and their ability to gain an electron in the p subshell to form a 1- ion decreases

80
Q

What colour is Cl2 in water and in cyclohexane?

A

Pale green

81
Q

What colour is Br2 in water and cyclohexane?

A

Orange

82
Q

What colour is I2 in water and cyclohexane?

A

Brown in water and violet in cyclohexane

83
Q

Can chlorine oxidise bromide ions into Br2?

A

Yes, Cl2(aq) + 2Br-(aq) → 2Cl-(aq) +Br2(aq)

84
Q

Can chlorine oxidise iodide ions into I2?

A

Yes, Cl2(aq) + 2I-(aq) → 2Cl-(aq) +I2(aq)

85
Q

Can bromine oxidise chloride ions into Cl2?

A

No

86
Q

Can bromine ions oxidise iodide ions into I2?

A

Yes, Br2(aq) + 2I-(aq) → 2Br-(aq) +I2(aq)

87
Q

Can iodine oxidise chloride or bromide ions?

A

No

88
Q

What is a disproportionation reaction?

A

A reaction in which the same element is both reduced and oxidised

Chemistry A Level (15 decks)

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