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A Level Chemistry > Topic 1 - Atomic structure > Flashcards

Topic 1 - Atomic structure Flashcards

1
Q

Flame spectrometry

A

Flame spectrometry is a method that allows the determination of the elemental composition of a sample by analysis of the emission spectra that are produced after heating a sample.

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2
Q

Emission spectra

A
  • Emission spectra are produced when electrons are excited
  • They absorb energy to move from a lower energy level, such as the ground state, to a higher energy level.
  • When the electrons move back to a lower energy level they release a ‘quantum’ of energy (fixed amount) in the form of light of a specific frequency. This light is one line of an emission spectrum.
  • This provides evidence that electrons are in shells, rather than moving randomly moving anywhere around an atom.
  • The distance moved is proportional to the FREQUENCY of light.
  • Short distance moved, low frequency, longer wavelength
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3
Q

Ionisation energy

A
  • Ionisation energy is a measure of the amount of energy needed to remove electrons from atoms to form positively charged ions
  • They represent an oxidation reaction as the species on the left of the arrow loses an electron
  • Positively charged ions are always formed
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4
Q

Explain why the ionisation energy is always endothermic

A
  • As electrons are negatively charged and protons in the nucleus are positively charged, there will be an attraction between them.
  • The closer the electron is to the nucleus, the greater the electrostatic attraction and the more energy is required to overcome this attraction.
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5
Q

1st ionisation energy of sodium definition

A

The energy required to remove ONE MOLE of electrons (to infinity) from ONE MOLE of gaseous atoms to form ONE MOLE of gaseous sodium 1+ ions.

Na (g) → Na+ (g) + e-

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6
Q

2nd ionisation energy of sodium definition

A

The energy required to remove ONE MOLE of electrons (to infinity) from ONE MOLE of gaseous sodium +1 ions to form ONE MOLE of gaseous sodium +2 ions.

Na+ (g) → Na2+ (g) + e-

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7
Q

3rd ionisation energy of sodium

A

The energy required to remove ONE MOLE of electrons (to infinity) from ONE MOLE of gaseous sodium +2 ions to form ONE MOLE of gaseous sodium +3 ions.

Na2+ (g) → Na3+ (g) + e-

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8
Q

What energy change is represented by this equation?

Ca+ (g) → Ca2+ (g) + e-

A

The second ionisation energy of calcium

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9
Q

What energy change is represented by this equation?

Ca (g) → Ca2+ (g) + e-

A

The first AND second ionisation energy of calcium

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10
Q

Factors that affect the ionisation energy

A
  • Number of electron shells (electron shielding)
  • Atomic or ionic radius
  • The number of electrons within a shell
  • Number of protons in the nucleus
  • More shells = smaller I.E. More electron repulsion
  • Larger radius = smaller I.E. Greater distance between the nucleus and outer electron, less electrostatic attraction
  • More electrons within a shell (e.g. a lone pair). Smaller I.E. as there is repulsion
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11
Q

Effective nuclear charge definition

A

The effective nuclear charge is the net positive charge experienced by an electron in a polyelectronic atom (atom with more than one electron).

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12
Q

Trend in effective nuclear charge across a period

A

Effective nuclear charge increases left to right across a period for electrons in the same shell as there are more protons in the nucleus and the number of shells remains constant.

Effective nuclear charge decreases as electrons are removed from shells further away from the nucleus as there is more electron shielding.

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13
Q

Trend in 1st ionisation energy across a period

A

First ionization energy increases left to right across a period:

  • Number of protons increases
  • Electrons are added to the same shell
  • The amount of electron shielding stays the same
  • Thus the effective nuclear charge increases
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14
Q

Trend in 1st ionisation energy down a group

A

First ionisation energy decreases going down a group:

  • The atomic radius increases
  • So the outer shell is further from the nucleus
  • The number of shells increases
  • So the electron shielding increases
  • Thus the effective nuclear charge decreases
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15
Q

Explanation for why group 3 elements have a lower first ionisation than group 2

A
  • The 1st ionization energy of boron is less than beryllium for period 2 and the 1st ionization energy of aluminium is less than calcium for period 3.
  • Boron has the electronic configuration [He]2s22p1 and aluminium has the electronic configuration [Ne]3s23p1.
  • The p subshell is a little further from the nucleus than the s subshell, so the effective nuclear charge is less than for s subshell electrons.
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16
Q

Explanation for why group 6 elements have a lower first ionisation than group 5

A
  • The 1st ionization energy of oxygen is less than nitrogen for period 2 and the 1st ionization energy of sulfur is less than phosphorus for period 3.
  • Oxygen has the electronic configuration [He]2s22p4 and nitrogen has the electronic configuration [He]2s22p3.
  • Nitrogen has its 3 p subshell electrons in three separate orbitals. Oxygen has one of its three orbitals with a pair of electrons. They repel, lowering the 1st ionization energy.
  • Phosphorus has the electronic configuration [Ne]3s23p3 and nitrogen has the electronic configuration [Ne]3s23p4
  • Phosphorus has its 3 p subshell electrons in three separate orbitals. Sulfur has one of its three orbitals with a pair of electrons. They repel, lowering the 1st ionization energy.
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17
Q

Trend in melting points across period 3

A
  • Melting point increases from group 1 to group 3
  • Strength of metallic bonding increases from groups 1 to 3
  • The ionic radius decreases
  • The charge increases
  • Thus, charge density increases
  • More electrons delocalize per atom
  • Group 4 elements form macromolecules
  • With lots of strong covalent bonds
  • That require a lot of thermal energy to break
  • Group 5 to group 7 elements form simple molecules
  • With weak intermolecular forces between the molecules
  • That require little thermal energy to break
  • Group 8 elements exist as single atoms, not molecules
  • They do not have intermolecular forces
  • They have weak interatomic forces
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18
Q

Metallic bonding definition

A

The electrostatic attraction between cations and the sea of delocalised electrons

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19
Q

Explanation for why magnesium has a higher melting point than sodium

A
  • Magnesium has stronger metallic bonds than sodium
  • Magnesium has a smaller atomic radius than sodium
  • Magnesium has a greater charge than sodium, +2 vs +1
  • Thus, charge density increases
  • One more electron delocalizes per magnesium atom than sodium
  • So the magnesium cations have a greater electrostatic attraction to the sea of delocalised electrons than sodium
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20
Q

Explanation for why aluminium has a higher melting point than magnesium

A
  • Aluminium has stronger metallic bonds than magnesium
  • Aluminium has a smaller atomic radius than magnesium
  • Aluminium has a greater charge than magnesium, +3 vs +2
  • Thus, charge density increases
  • One more electron delocalizes per aluminium atom than magnesium
  • So the aluminium cations have a greater electrostatic attraction to the sea of delocalised electrons than magnesium
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21
Q

Periodicity definition

A

Repeating trends in physical properties across a period

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22
Q

Group of elements with the highest 1st ionisation energy

A

Noble gases

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23
Q

Group of elements with the lowest 1st ionisation energy

A

Group 1 metals

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24
Q

Group of elements with the highest 2nd ionisation energy

A

Group 1 metals

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25
Q

Group of elements with the highest 3rd ionisation energy

A

Group 2 elements

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26
Q

Z

A

The symbol for the atomic number of an element

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27
Q

Atomic number

A

The number of protons in the nucleus of an atom

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28
Q

Mass number

A

The number of protons and neutrons in the nucleus of an atom

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29
Q

Charge on a ion in mass spectrometry

A

1+

30
Q

Relative atomic mass

A

The weighted mean mass of atoms of an element compared to 1/12th the mass of an atom of carbon-12

31
Q

Relative isotopic mass

A

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

32
Q

Isotopes

A

Atoms with the same number of protons and different number of neutrons. They have identical chemical properties because they have the same number of electrons in the outer shell.

33
Q

A sample of bromine Br2 is analysed by mass spectrometry. It has the isotopes 79Br and 81Br.

Predict the number of peaks on a mass spectrum and identify the species responsible for the peaks.

A
  • 5 Peaks
  • Peak at m/z = 79 indicates 79Br+
  • Peak at m/z = 81 indicates <span>81</span>Br+
  • Peak at m/z = 158 indicates 79Br-79Br+
  • Peak at m/z = 160 indicates 79Br-<span>80</span>Br+
  • Peak at m/z = 162 indicates <span>81</span>Br-<span>81</span>Br+
34
Q

A sample of oxygen O2 is analysed by mass spectrometry. It has the isotopes 15O, 16O and 17O.

Predict the number of peaks on a mass spectrum and identify the species responsible for the peaks.

A
  • 8 Peaks
  • Peak at m/z = 15 indicates 15O+
  • Peak at m/z = 16 indicates 16O+
  • Peak at m/z = 17 indicates 17O+
  • Peak at m/z = 30 indicates 15O-15O+
  • Peak at m/z = 31 indicates 15O-16O+
  • Peak at m/z = 32 indicates 16O-16O+
  • Peak at m/z = 33 indicates 16O-17O+
  • Peak at m/z = 34 indicates 17O-17O+
35
Q

A sample of nitrogen N2 is analysed by mass spectrometry. It has the isotopes <span>13</span>N, 14N and 15N.

Predict the number of peaks on a mass spectrum and identify the species responsible for the peaks.

A
  • 8 Peaks
  • Peak at m/z = 13 indicates 13N+
  • Peak at m/z = 14 indicates 14N+
  • Peak at m/z = 15 indicates 15N+
  • Peak at m/z = 26 indicates 13N-13N+
  • Peak at m/z = 27 indicates 13N-14N+
  • Peak at m/z = 28 indicates 14N-14N+
  • Peak at m/z = 29 indicates 14N-15N+
  • Peak at m/z = 30 indicates <span>15</span>N-<span>15</span>N+
36
Q

A sample of bromine Br2 is analysed by mass spectrometry. It has the isotopes 79Br and 81Br.

The relative height of the peaks is in a ratio of 3 : 1 , 79Br : 81Br. Determine the relative atomic mass of the sample.

A
  • There is 75% of 79Br and 25% of 81Br.
  • Relative atomic mass = (75/100 x 79) + (25/100 x 81)
  • Relative atomic mass = 59.25 + 20.25
  • Relative atomic mass = 79.5
37
Q

A sample of bromine N2 is analysed by mass spectrometry. It has the isotopes 13N and 14N.

The relative height of the peaks is in a ratio of 3 : 2 , 13N : 14N. Determine the relative atomic mass of the sample.

A
  • There is 60% of 13N and 40% of 14N.
  • Relative atomic mass = (60/100 x 13) + (40/100 x 14)
  • Relative atomic mass = 7.8 + 5.6
  • Relative atomic mass = 13.4
38
Q

Emission spectra provide evidence for…

A

Quantum shells

39
Q

Successive ionisation energies provide evidence for…

A

Subshells and the number of shells

40
Q

An element has the following successive ionisation energies. Identify which group the element is in.

736 , 1450 , 7740 , 10500 , 13600

A

Group 2

There is a large jump from the 2nd to 3rd ionisation energies and so the 3rd electron is a removed from a full shell.

41
Q

An element has the following successive ionisation energies. Identify which group the element is in.

1060 , 1900 , 2920 , 4960 , 6280 , 21200

A

Group 5

There is a large jump from the 5th to 6th ionisation energies and so the 6th electron is a removed from a full shell.

42
Q

An element has the following successive ionisation energies. Identify which group the element is in.

786 , 1580 , 3230 , 4360 , 16000

A

Group 4

There is a large jump from the 4th to 5th ionisation energies and so the 5th electron is a removed from a full shell.

43
Q

s-block elements

A

Group 1 and 2 elements, AND helium

44
Q

p-block elements

A

Groups 3 to 8, EXCLUDING helium

45
Q

d-block elements

A

Transition metals

46
Q

f-block elements

A

Lanthanides and actinides

47
Q

Shape of a s orbital

A

Spherical

48
Q

Shape of a p orbital

A

Bi-lobed

49
Q

Number of orbitals in a p subshell

A

3

50
Q

Order of filling subshells

A

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

51
Q

Number of orbitals in a d subshell

A

5

52
Q

Maximum number of electrons in an s orbital

A

2

53
Q

Maximum number of electrons in an d subshell

A

10

54
Q

Maximum number of electrons in an s subshell

A

2

55
Q

Maximum number of electrons in an p orbital

A

2

56
Q

Maximum number of electrons in an p subshell

A

6

57
Q

Electronic configuration of a sodium atom

A

1s2 2s2 2p6 3s1

or [Ne] 3s1

58
Q

Electronic configuration of a sodium ion

A

1s2 2s2 2p6

or [Ne]

59
Q

Electronic configuration of a fluorine atom

A

1s2 2s2 2p5

or [He] 2s2 2p5

60
Q

Electronic configuration of a fluorine ion

A

1s2 2s2 2p6

or [Ne]

61
Q

Electronic configuration of a nickel atom

A

Nickel has 28 electrons

1s2 2s2 2p6 3s2 3p6 4s2 3d8

or [Ar] 4s2 3d8

62
Q

Electronic configuration of a copper atom

A

Copper has 29 electrons

1s2 2s2 2p6 3s2 3p6 4s1 3d10

or [Ar] 4s1 3d10

63
Q

Electronic configuration of a copper Cu+ ion

A

Copper Cu+ has 28 electrons

1s2 2s2 2p6 3s2 3p6 3d10

or [Ar] 3d10

The single 4s1 electron is removed

64
Q

Electronic configuration of a copper Cu2+ ion

A

Copper Cu2+ has 27 electrons

1s2 2s2 2p6 3s2 3p6 3d9

or [Ar] 3d9

The single 4s1 electron and a 3d electron are removed

65
Q

Electronic configuration of an iron atom

A

Iron has 26 electrons

1s2 2s2 2p6 3s2 3p6 4s2 3d<span>6</span>

or [Ar] 4s2 3d<span>6</span>

66
Q

Electronic configuration of an iron Fe2+ ion

A

Iron Fe2+ has 24 electrons

1s2 2s2 2p6 3s2 3p6 3d<span>6</span>

or [Ar] 3d<span>6</span>

The two 4s2 electrons are removed

67
Q

Electronic configuration of an iron Fe3+ ion

A

Iron Fe3+ has 23 electrons

1s2 2s2 2p6 3s2 3p6 3d<span>5</span>

or [Ar] 3d<span>5</span>

The two 4s2 electrons and a 3d electron are removed

68
Q

Hund’s rule

A
  • Used to fill up electron orbitals
  • Orbitals are filled singly until all of the orbitals in a subshell are full
  • Then electrons are paired up with opposite spin
69
Q

Box notation

A

Each orbital in a subshell is represented by a box

70
Q
A
71
Q
A

A Level Chemistry (2 decks)

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