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Chemistry A-Level > Atomic and Electron Structure, Periodic Table and Ionisation Energy (2.1, 2.2, 5.1, 7.1, 7.2) > Flashcards

Atomic and Electron Structure, Periodic Table and Ionisation Energy (2.1, 2.2, 5.1, 7.1, 7.2) Flashcards

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1
Q

What are isotopes?

A

Atoms of the same element with different numbers of neutrons and different masses but the same number of protons

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2
Q

Describe the atomic structure

A

Consists of a nucleus made up of two subatomic particles: protons and neutrons
Third subatomic particle: electron arranged around the nucleus in shells

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3
Q

What is the relative isotopic mass?

A

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

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4
Q

What is the relative atomic mass?

A

The weighted mean mass of an atom relative to 1/12th of the mass of an atom of carbon-12

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5
Q

What is the relative molecular mass?

A

The average mass of a molecule or formula unit, compared to 1/12th of the mass of an atom of carbon-12

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6
Q

What are the numbers of shells (energy levels) called?

A

Principal quantum numbers

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7
Q

How does energy change in the shells?

A

Shells further from the nucleus have a higher energy than shells closer to the nucleus

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8
Q

What are the names of the different sub-shells?

A

s, p, d, f

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9
Q

What does each sub-shell have?

A

A different number of orbitals

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10
Q

How many electrons can one orbital hold?

A

2

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11
Q

How many orbitals does an s sub-shell have?

A

1

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12
Q

How many electrons can an s sub-shell hold?

A

2

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13
Q

How many orbitals does a p sub-shell have?

A

3

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14
Q

How many electrons can a p sub-shell hold?

A

6

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15
Q

How many orbitals does a d sub-shell have?

A

5

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16
Q

How many electrons can a d sub-shell hold?

A

10

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17
Q

What is an atomic orbital?

A

A region around the nucleus that can hold up to two electrons, with opposite spins

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18
Q

What are the sub-shells in the 1st shell?

A

1s

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19
Q

How many electrons can the 1st shell hold?

A

2

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20
Q

What are the sub-shells in the 2nd shell?

A

2s, 2p

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21
Q

How many electrons can the 2nd shell hold?

A

8

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22
Q

What are the sub-shells in the 3rd shell?

A

3s, 3p, 3d

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23
Q

How many electrons can the 3rd shell hold?

A

18

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24
Q

What are the sub-shells in the 4th shell?

A

4s, 4p, 4d, 4f

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25
Q

How many electrons can the 4th shell hold?

A

32

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26
Q

How many orbitals does a f sub-shell have?

A

7

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27
Q

How many electrons can a f sub-shell hold?

A

14

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28
Q

How much do you love lebae?

A

100%

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29
Q

What is the periodic table?

A

The arrangement of elements

  • by increasing atomic number (left to right)
  • in periods (horizontal rows) showing repeating trends in physical and chemical properties of elements
  • in groups (vertical columns) having similar chemical properties and the same number of electrons in outer shell
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30
Q

What is spin-pairing?

A

If there are two electrons in an orbital, they have to ‘spin’ in opposite directions

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31
Q

What is the shape of s-orbitals?

A

Spherical

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32
Q

What is the shape of p-orbitals?

A

Dumbbell - there are 3 p-orbitals and they are at right angles to one another

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33
Q

What is the order of sub-shells in increasing energy?

A

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 4d, 4f

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34
Q

How do electrons fill orbitals with the same energy?

A

They occupy each orbital singly before sharing (i.e. 1 in each orbital first)

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35
Q

What do all elements within a period have in common?

A

They have the same number of electron shells

36
Q

What do all elements within a group have in common?

A

They have the same number of electrons in their outer shell and therefore similar chemical properties

37
Q

What are the 4 blocks in the periodic table?

A

s-block, d-block and p-block and f-block

38
Q

What elements are in the s-block?

A

The elements in the first two columns (i.e. the left side)

39
Q

What elements are in the d-block?

A

The elements in the next ten columns (i.e the middle part)

40
Q

What elements are in the p-block?

A

The elements in the last 6 columns (i.e. the right side)

41
Q

What outer shell electronic configuration do s-block elements have?

A

s1 or s2

42
Q

What electronic configuration do d-block elements have?

A

They’re d sub-shells are being filled

43
Q

What outer shell electronic configuration do p-block elements have?

A

s2p1 - s2p6

44
Q

What is the first ionisation energy?

A

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

45
Q

What kind of process is ionisation?

A

Endothermic

46
Q

What is the equation for the first ionisation energy?

A

Na(g) —> Na+(g) + e-

47
Q

What must you have in an ionisation energy equation?

A

State symbols (g)

48
Q

What is the 2nd ionisation energy?

A

Na+(g) —> Na2+(g) + e-

49
Q

What factors affect ionisation energy?

A

Atomic radius
Nuclear charge
Electron shielding

50
Q

How does atomic radius affect ionisation energy?

A
  • The greater the distance between the nucleus and the outer electrons, the less the nuclear attraction (attraction between nucleus and outer electron)
  • Therefore, ionisation energy decreases as atomic radius increases
51
Q

How does nuclear charge affect ionisation energy?

A
  • The more positively charged (the more protons there are in) the nucleus is, the greater the attraction between the nucleus and the outer electron
  • Therefore, ionisation energy increases as nuclear charge increases
52
Q

How does electron shielding affect ionisation energy?

A
  • Electrons are negatively charged and therefore inner-shell electrons repel outer shell electrons (shielding effect)
  • As the number of electrons between the outer electrons and the nucleus increases, the outer electrons have a weaker nuclear attraction
  • Therefore, ionisation energy increases as electron shielding decreases
53
Q

What is electron shielding?

A

The decrease of nuclear attraction between the outer electrons and the nucleus as a result of inner electrons producing a shielding effect

54
Q

What is the shielding effect?

A

Electrons are negatively charged and so inner-shell electrons repel outer-shell electrons
This reduces the attraction between the nucleus and the outer electrons

55
Q

How many ionisation energies does an element have?

A

As many as there are electrons

56
Q

Why do successive ionisation energies become larger?

A
  • After the first electron is lost, the second electron is pulled closer to the nucleus
  • The nuclear attraction on the remaining electron increases and more ionisation energy will be needed to overcome the attraction and remove the electron
57
Q

What to remember when looking at successive ionisation energy graphs?

A

1) look for any large jumps in ionisation energy - this marks the change from one shell to another
2) the electrons with the largest ionisation energy are from the shell closest to the nucleus
3) ignore sub shells
4) when drawing a graph, a new shell = jump upwards in graph

58
Q

What do successive ionisation energies allow predictions about the element to be made about?

A

1) The number of electrons in the outer shell (from where the increase in ionisation energy is)
2) The group of the element in the periodic table
3) The identity of an element (if period is known)

59
Q

What are the two key patterns in the trend of first ionisation energies of the first 20 elements?

A

1) A general increase in the first ionisation energy across each period
2) A sharp decrease in the first ionisation energy between the end of one period and the start of the next period

60
Q

What happens to the first ionisation energies as you go down a group?

A

They decrease

61
Q

Why do first ionisation energies decrease down a group?

A
  • Atomic radius increases
  • Number of (inner) shells increases so shielding increases
  • This increased shielding effect outweighs the increase in nuclear charge
  • Therefore, nuclear attraction on the outer electrons decreases
  • So first ionisation energy decreases
62
Q

What happens to the first ionisation energies as you go across a period?

A

They increase

63
Q

Why do first ionisation energies increase across a period?

A
  • Nuclear charge increases
  • Same shell so similar shielding
  • Nuclear attraction increases
  • Atomic radius decreases
  • So first ionisation energy increases
64
Q

In the trend of first ionisation energies in a period where are there falls in two places despite the overall increase?

A

Fall from beryllium (Be) to boron (B) (or Mg to Al)

Fall from N to O (or P to S)

65
Q

Why is there a fall in first ionisation energy from beryllium to boron?

A
  • Because of the filling of the 2p sub-shell
  • The 2p sub-shell in B has a higher energy that the 2s sub-shell in Be
  • Therefore, in B the 2p electron is easier to remove than one of the 2s electrons in Be
  • So the first ionisation energy of B is less than the first ionisation energy of Be
66
Q

Why is there a fall in first ionisation energy from nitrogen to oxygen?

A
  • Because of the start of electron pairing in the p-orbitals of the 2p sub-shell
  • In N and O the highest energy electrons are in a 2p sub-shell
  • In O, the paired electrons in one of the 2p orbitals repel one another, making it easier to remove an electron from an O atom than an N atom
  • Therefore, the first ionisation energy of O is less than the first ionisation energy of N
67
Q

Describe the structure of the outer electrons in nitrogen

A
  • Three 2p electrons: 2px1, 2py1, 2pz1
  • One electron in each 2p orbital
  • Spins are at right angles - equal repulsion as far apart as possible
68
Q

Describe the structure of the outer electrons in oxygen

A
  • Four 2p electrons; 2px2, 2py1, 2pz1
  • Two electrons in one 2p orbital
  • 2p electrons start to pair
  • The paired electrons repel
69
Q

Why does aluminium have the strongest metallic bonding in period 3?

A

Furthest along metal in period

More electrons for each nucleus so a stronger nuclear attraction

70
Q

Why does the second ionisation energy of oxygen have a greater value than the first?

A

Higher nuclear attraction
Less electrons the positive charge has to be distributed over
Electron repulsion is less in O+ than O
Ionic radius of O+ is smaller than the atomic radius of O

71
Q

What is the relative mass of an electron?

A

1/1836

72
Q

What is the mass (nucleon) number? (A)

A

The number of protons plus the number of neutrons

73
Q

What is the atomic number? (Z)

A

The number of protons

74
Q

What is an ion?

A

A charged atom

75
Q

What is a cation?

A

Atoms with fewer electrons than protons - they have an overall positive charge

76
Q

What is an anion?

A

Atoms with more electron than protons - they have an overall negative charge

77
Q

What are shells?

A

Energy levels

78
Q

Describe the property of spin in electrons

A
  • Electrons are negatively charged and repel one another
  • Electrons have a property called spin - either up or down
  • The two electrons in an orbital must have opposite spins - the opposite spins help to counteract the repulsion between the negative charges of the two electrons
79
Q

What is periodicity?

A

Repeating trends in physical and chemical properties of elements

80
Q

What does each period start with?

A

An electron in a new highest energy shell

81
Q

What is the trend in electron configuration across period 2?

A

Across period 2, the 2s sub-shell fills with 2 electrons, followed by the 2p sub-shell with 6 electrons

82
Q

What is the trend in electron configuration across period 3?

A

Across period 3, the 3s sub-shell fills with 2 electrons, followed by the 3p sub-shell with 6 electrons

83
Q

What elements are in the f-block?

A

The ones in the bottom two rows

84
Q

What predictions could you make about fluorine from its successive ionisation energies?

A
  • There is a large increase between the 7th and 8th ionisation energy
  • This suggests that the 8th electron must be removed from another shell, closer to the nucleus with less shielding and that there are 7 electrons in the outer shell
85
Q

What is the second ionisation energy?

A

The energy required to remove one electron from each atom in one mole of gaseous 1+ ions to form one mole of gaseous 2+ ions

86
Q

Explain the rises in ionisation energy across period 2

A
  • Rise from Li to Be - due to the filling of the 2s sub-shell
  • Rise from B to N - due to the adding of one electron to each 2p orbital
  • Rise from O to Ne - due to the pairing of the 2p electrons

Chemistry A-Level (27 decks)

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