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chemistry a - level > Topic 1 - Atomic Structure and the Periodic Table > Flashcards

Topic 1 - Atomic Structure and the Periodic Table Flashcards

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1
Q

What is the definition of relative isotopic mass?

A

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

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2
Q

What is the definition of relative atomic mass?

A

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

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3
Q

What is the equation to calculate the Ar value of an element?

A

Ar = average mass per atom of an element x 12/ mass of one atom of carbon-12

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4
Q

What is used as the standard measure of relative isotopic mass?

A

Carbon-12

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5
Q

How do you calculate Relative Atomic Mass from a mass spectrum?

A

Σ (m/z value x relative or % abundance) / Σ % abundances

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6
Q

What will the mass spectrum of a diatomic molecule e.g Cl2, look like?

A

Chlorine has 2 isotopes 35Cl and 37Cl with a ratio of 3:1
Lines from the two isotopes of Cl+ will form at 35 and 37 on the x -axis
Only some molecules of Cl2 will fragment upon ionisation so therefore you will also get Cl2+ peaks at 70,72 and 74 on the x - axis and these will be in the ratio 9:6:1 which you can figure out from the relative abundance and ratio of the fragments

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7
Q

Why do molecular ions fragment?

A

As they tend to be unstable

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8
Q

What is a free radical?

A

An atoms or group of atoms which contain a single unpaired electron e.g X*

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9
Q

What is produced when a molecule fragments?

A

A positive ion and an uncharged free radical

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10
Q

What will an uncharged free radical not produce?

A

A line on the mass spectrum

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11
Q

What is good about mass spectrometry?

A

It is very precise and produces almost a fingerprint of the molecule

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12
Q

What is the definition of ionisation energy?

A

The amount of energy required to remove one mole of electrons from one mole of atoms of elements in the gaseous state to form one mole of gaseous ions

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13
Q

What are ionisation energies measured under?

A

Standard conditions

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14
Q

Write a general equation for the 1st, 2nd and 3rd ionisation energies for a metal

A

M(g) → M+ (g) + e-
M+ (g) → M²+ (g) + e-
M²+ (g) → M³+ (g) + e-

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15
Q

Write a general equation for the first and second ionisation energies together for a metal

A

M (g) → M²+ (g) + 2e-

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16
Q

Which sign is given for endothermic?

A

+

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17
Q

Why are ionisation energies always endothermic?

A

As energy always needs to be taken in so the electrostatic forces of attraction can be overcome between the nucleus and the electrons

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18
Q

What is the unit for ionisation energy?

A

kjmol^-1

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19
Q

Does the successive ionisation energy of an element increase or decrease?

A

Increase

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20
Q

Why do the successive ionisation energies of elements increase?

A

As each electron is removed, the attractive forces between the outermost electron and the nucleus increase due to decreased shielding and an increasingly positive ion, therefore each successive electron is harder to remove.

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21
Q

What are successive ionisation energies evidence for?

A

The existence of electron shells

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22
Q

Why are there big jumps (increases) in successive ionisation energies?

A

As this marks the removal of an electron from the next quantum shell at a lower energy level that is closer to the nucleus

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23
Q

What are you able to deduce from the successive ionisation energies of an element? How?

A

The group of an unknown element, as where the large jump occurs marks the jump to a lower energy level and the outer shell be empty.

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24
Q

How do you sketch a graph for the successive ionisation energies of an element?

A

x-axis - Order of electron removed
y- axis - ionisation energy

  • Write the electron arrangement for the element
  • the ionisation energy will start low and increase slowly for the outer shell electrons
  • there will be a large jump when moving to the next shell
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25
Q

What is periodicity?

A

A trend in a characteristic across a period of the periodic table

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26
Q

What are the 3 factors that affect ionisation energy?

A
  • Nuclear charge
  • Distance form the nucleus/ Atomic radius
  • Shielding from filled inner electron shells
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27
Q

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

A

It decreases

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28
Q

Why does the first ionisation energy decrease as you go down a group?

A

Whilst the nuclear charge increases, the distance of the outermost electron from the nucleus, due to more shells, and the shielding from the filled inner shells increases, therefore less energy is required to remove an electron

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29
Q

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

A

It generally increases (however, there are exceptions)

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30
Q

Why does ionisation energy generally increase across a period?

A

The shielding from filled inner remains (pretty much) constant across a period. The nuclear charge increases causing the atomic radius to decrease, therefore the electrostatic forces of attraction between the nucleus and outermost electrons increases.

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31
Q

Describe how nuclear charge affects ionisation energy?

A

The nucleus contain 2 particles, protons and neutrons. Protons have a positive charge and neutrons have no charge, therefore the overall charge on the nucleus is positive. Thus, the more protons within the nucleus the stronger the attractive forces due to a greater nuclear charge.

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32
Q

Describe how distance from the nucleus affects ionisation energy?

A

The negatively charged electrons are attracted to the positively charged nucleus. As the distance increases the attraction falls off rapidly.

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33
Q

Describe how shielding affects ionisation energy?

A

Inner shells that are filled with electrons will shield outer shells from the nucleus. Therefore, shielding reduces the forces of attraction between the nucleus and outer electrons.

34
Q

What are the 4 exceptions in period 2 and 3 that don’t follow the increase in ionisation energy as you for form left to right?

A

Be → B
N → O
Mg → Al
P → S

35
Q

Why is there a dip in first ionisation energy as you move from Be → B and Mg → Al?

A

Be has an electronic configuration of 1s2 2s2 and B has an electronic configuration of 1s2 2s2 2p1. Although there is an increase in nuclear charge, the outer electron of B is in a 2p orbital, which has a slightly higher energy, whereas in Be this is a 2s orbital. Therefore, less energy is needed to remove the outer electron. Additionally, the outermost electron in B experiences slightly more shielding from filled inner shells and is further away from the nucleus so the forces of attraction are slightly weaker. This is the same reasoning for Mg → Al.

NOTE: Although there is a dip in first ionisation energy, it is not lower than Li in period 2 and Na in period 3.

36
Q

Why is there a dip in first ionisation energy as you move from N → O and P → S?

A

In Oxygen, the extra electron (4th electron) goes into a p orbital that already contains one electron. Whilst, there is an increase in nuclear charge, this is outweighed by the spin pair repulsion between the two electrons in the p orbital. Therefore, less energy is needed to remove the outer electron from Oxygen. This is the same reasoning for P → S .

NOTE: Although there is a dip in first ionisation energy, it is not lower than C in period 2 and Si in period 3.

37
Q

What is the definition of nuclear charge?

A

The charge due to the protons in the nucleus.

38
Q

What is the definition of effective nuclear charge?

A

The effectiveness of the nuclear charge after passing through any filled inner shells.

39
Q

Why does atomic radius decrease across a period?

A

As shielding is roughly constant and the nuclear charge of the atom increases, there are stronger electrostatic forces of attraction between the outer electron and the nucleus. This results in the electrons being pulled together so a smaller atomic radius.

40
Q

How do you figure out the radius of and atom?

A
  • Atoms don’t have a fixed radius
  • In metallic and covalent structures the atoms are bonded strongly and therefore overlap giving a smaller atomic radius
  • In structures where only weak London Forces are present the atoms are just touching therefore giving a more accurate representation of their atomic radius
  • use the distance between the atoms/2 to figure out the atomic radius
41
Q

Does electronegativity increase or decrease across a period? Why?

A

The electronegativity increases as the increasing nuclear charge causes more attraction of the electron pair in the bond to the nucleus

42
Q

Does electronegativity increase or decrease down a group? Why?

A

The electronegativity decreases due to the bonded electrons being further way from the nucleus ,due to more filled shells, and increased shielding. Therefore, the electron pair in the covalent bond is less strongly attracted.

43
Q

Why do the melting and boiling points increase across Na, Mg and Al in period 3?

A

They all have metallic structures. As you move from Na to Al there is an increasing number of electrons contributing to a larger and more negative sea of delocalised electrons. Additionally, the proton number increases giving the metals a higher nuclear charge so the nuclei of the atoms are getting increasingly positive. Additionally, the sea of delocalised electrons is getting progressively nearer to the nuclei and so are more strongly attracted. Due to these three points the melting and boiling points increase.

44
Q

Why does silicon have a high melting and boiling point?

A

It is a non-metal and has a giant covalent structure. Therefore you have to break strong covalent bonds that require very large amounts of energy to overcome in order to melt/boil it.

45
Q

Does melting point increase or decrease down a group? Why?

A

The melting point decreases as metallic bonding strength decreases as the size of the atoms increases. Therefore, the electron cloud isn’t as effective at holding the ions together so they are overcome by lower energies.

46
Q

What do melting a boiling points largely depend on?

A

The structure and bonding existing in elements

47
Q

Why do phosphorus, sulphur, chlorine and argon, in period 3, have such low melting and boiling points?

A

They are simple molecular substances that are only held together by weak London Forces that require little energy to overcome. The sizes of their melting and boiling points are governed completely by the sizes of the molecule.

48
Q

What is an emission spectrum?

A

An emission spectrum is a display of the many possibilities of electron jumps between energy levels. Each line correspond to a specific energy value which infers the electrons only have a limited choice of allowed energies.

49
Q

Why do we observe an emission spectrum when we heat an atom of an element?

A

When an atom is heated the electrons within the atom are excited from their ground state into higher energy levels. When the electron returns to its ground state it emits energy as an electromagnetic wave of a specific frequency. If this frequency falls within the visible spectrum we observe a colour.

50
Q

What are atoms made up of?

A

Quantum energy levels

51
Q

What are Quantum energy levels made up of?

A

Quantum Sub-shells

52
Q

What do the electrons in a given quantum cells all share?

A

Similar energies

53
Q

What are the sub-shells in Quantum shell 1?

A

1s

54
Q

What are the sub-shells in Quantum shell 2?

A

2s, 2p

55
Q

What are the sub-shells in Quantum shell 3?

A

3s, 3p, 3d

56
Q

What are the sub-shells in Quantum shell 4?

A

4s, 4p, 4d, 4f

57
Q

How many orbitals are there in an s sub-shell?

A

1

58
Q

How many orbitals are there in a p sub-shell?

A

3

59
Q

How many orbitals are there in a d sub-shell?

A

5

60
Q

How many orbitals are there in an f sub-shell?

A

7

61
Q

How many electrons can an s sub-shell hold?

A

2

62
Q

How many electrons can a p sub-shell hold?

A

6

63
Q

How many electrons can a d sub-shell hold?

A

10

64
Q

How many electrons can an f sub-shell hold?

A

14

65
Q

What are the three important rules for filling sub-shells?

A

1.Of the orbitals available, the added electrons occupies the orbital of the lowest energy first
2. Each orbital can hold a maximum of two electrons, but they must have opposite spin
3. If a number of orbitals of equal energies is available , the added electron will go into a vacant orbital, keeping spins parallel, before 2 electrons occupy the same orbital.

66
Q

What is the electron number equal to?

A

The atomic number (bottom one)

67
Q

What is the arrangement of electron sub shells in increasing energy up to 4d?

A

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d

68
Q

Why is the 4s sub-shell filled before the 3d sub-shell?

A

As it is of a lower energy level

69
Q

What are the 2 exceptions for the way that you fill sub-shells?

A

Chromium - 24e-
Copper - 29e-

70
Q

Why are Chromium and Copper filled in a different way?

A

As electrons always want to be organised in the most stable arrangement possible. For Chromium this means for the 3d and 4s sub shells there is one electron in each orbital. For Copper this means the 3d sub-shell is filled entirely before the 4s orbital gains one.

Normally, the 4s sub shell is fully filled prior to the 3d one as it is of a lower energy, however, in these two cases it is energetically favourable to fill the 3d one as well.

71
Q

What is the electronic configuration of Cr?

A

24e-
1s2,2s2,2p6,3s2,3p6,3d5,4s1

72
Q

What is the electronic configuration of Cu?

A

29e-
1s2,2s2,2p6,3s2,3p6,3d10,4s1

73
Q

What block are group 1 and 2 in?

A

S-block, this means their outermost electron will be in an s sub-shell.

74
Q

What bock are the transition metals in?

A

D-block, this means their outermost electron will be in a d sub-shell.

75
Q

What block are groups 5→8 in?

A

P-block, this means their outermost electron will be in a p sub-shell.

76
Q

What is the shape of an s orbital?

A

A circular shape, 2 on either side of the nucleus at (0,0,0), along the x,y or z axis

77
Q

What is the shape of a p orbital?

A

An hourglass shape along the x, y or z axis

78
Q

When d block metals form ions where is the at electron lost from? Why?

A

The 4s sub-shell as it behaves as the outermost, highest energy orbital

79
Q

Write the electronic configuration of Fe(3+)

A

1s2,2s2,2p6,3s2,3p6,3d5

80
Q

Write the electronic configuration for Cr(3+)

A

1s2,2s2,2p6,3s2,3p6,3d3

81
Q

Write the electronic configuration for O(2-)

A

1s2,2s2,2p6

82
Q

Are the trends in melting and boiling points the same across periods 2 and 3?

A

Yes, both follow the same rules and are dependant on structure and bonding in the element

chemistry a - level (26 decks)

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