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A LEVEL - Chemistry (Inorganic) > 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 are the 3 subatomic particles?

A
  • Protons
  • Neutrons
  • Electrons
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2
Q

Describe the structure of an atom.

A
  • Nucleus with protons and neutrons

* Electron shells around the nucleus

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3
Q

Compare the size of the nucleus with the size of the whole atom.

A

Its diameter is relatively tiny.

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4
Q

What subatomic particles make up the nucleus?

A

Protons and neutrons

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5
Q

Where is most of the mass of the atom found?

A

The nucleus.

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6
Q

How are the masses and charges of subatomic particles given and why?

A
  • As RELATIVE masses and RELATIVE charges

* Because the actual mass and charge values are tiny

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7
Q

What is the relative mass and charge of a proton?

A

Mass: 1
Charge: +1

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8
Q

What is the relative mass and charge of a neutron?

A

Mass: 1
Charge: 0

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9
Q

What is the relative mass and charge of an electron?

A

Mass: 0.0005
Charge: -1

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10
Q

What is a nuclear symbol?

A
  • The symbol for an element found in the periodic table that shows its mass number and atomic number.
  • e.g. 7Li3
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11
Q

What is the mass number?

A

The number of protons and neutrons in an element.

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12
Q

What is the symbol for the mass number?

A

A

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13
Q

What is the atomic number?

A

The number of protons in an element.

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14
Q

What is the symbol for the atomic number?

A

Z

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15
Q

What is another name for the atomic number?

A

Proton number

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16
Q

What identifies a particular element?

A

The atomic/proton number.

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17
Q

In a neutral atom, the number of protons equals…

A

The number of electrons.

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18
Q

How can you find the number of neutrons in an element?

A

No. of Neutrons = Mass number - Atomic number

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19
Q

What causes a negative ion to have a negative charge?

A

It has gained electrons (so it has more electrons than protons).

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20
Q

What causes a positive ion to have a positive charge?

A

It has lost electrons (so it has more protons than electrons).

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21
Q

What are isotopes?

A

Atoms with the same number of protons but different numbers of neutrons.

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22
Q

Give an example of isotopes.

A

Cl-35 and Cl-37.

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23
Q

What determines the chemical properties of an element?

A

The number and arrangement of electrons.

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24
Q

What determines the physical properties of an element?

A

Usually the mass.

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25
Q

Compare different isotopes in terms of chemical and physical properties and explain why this is.

A
  • Same chemical properties -> Same electron configuration

* Different physical properties -> Different masses

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26
Q

What is the difference between physical and chemical properties of an element?

A
  • Physical -> Can be observed without changing the element (e.g. colour)
  • Chemical -> Can only be observed by changing the chemical identity (e.g. flammability)
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27
Q

What is relative atomic mass?

A

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

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28
Q

What is the symbol for relative atomic mass?

A

Ar

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29
Q

What is relative isotopic mass?

A

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

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30
Q

What is the difference between relative isotopic mass and relative atomic mass?

A

Relative isotopic mass refers to only one isotope, while relative atomic mass takes into account all of the different isotopes of an element.

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31
Q

Are relative atomic mass and relative isotopic mass whole numbers?

A
  • Relative atomic mass - Usually not a whole number

* Relative isotopic mass - Usually a whole number

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32
Q

A natural sample of chlorine contains a mixture of Cl-35 (75%) and Cl-37 (25%). What are the relative isotopic masses and relative atomic mass?

A
  • Relative isotopic masses - 35 and 37

* Relative atomic mass - 35.5

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33
Q

What is 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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34
Q

What is the symbol for relative molecular mass?

A

Mr

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35
Q

What two things does relative molecular mass encompass and when is each used?

A
  • Relative molecular mass -> For simple molecules

* Relative formula mass -> For ionic and giant covalent compounds

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36
Q

How do you find the relative molecular mass?

A

Add up the Ar values of all of the atoms in the molecule.

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37
Q

How do you find the relative formula mass?

A

Add up the Ar values of all of the atoms in the formula.

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38
Q

Ar: C = 12, H = 1, O = 16Find the relative molecular mass of C2H6O.

A

Mr = (2 x 12) + (6 x 1) + 16 = 46

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39
Q

Ar: Ca = 40.1, F = 19Find the relative formula mass of CaF2.

A

Mr = 40.1 + (2 x 19) = 78.1

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40
Q

Compare the Ar of an atom and an ion.

A

They are the same.

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41
Q

What is isotopic abundance?

A

The amount of a particular isotope as a percentage of all of the atoms of an element.

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42
Q

How can you calculate the Ar of an element from its isotopic abundances?

A

1) Multiply each relative isotopic mass by its % abundance and add these up.2) Divide by the total % abundance (usually 100).

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43
Q

20% of all boron atoms on earth have a relative isotopic mass of 10.0, while 80% have a relative isotopic mass of 11.0. Find the Ar of boron.

A

• (20 x 10) + (80 x 11) = 1080• 1080 / 100 = 10.8

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44
Q

What is a mass spectrum?

A

A bar graph of mass to charge ratio (x) against abundance (y) of different isotopes (or other particles).

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45
Q

What device produces mass spectra?

A

Mass spectrometers.

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46
Q

What is a mass spectrometer?

A
  • A device that is used to find out what samples are made up of by measuring the masses of their components
  • Produces mass spectra
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47
Q

What can mass spectra tell us?

A
  • Relative isotopic masses and abundances

* Relative atomic masses of different elements

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48
Q

What is on the y-axis of a mass spectrum?

A

Abundance of ions (as a number or percentage).

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49
Q

What is on the x-axis of a mass spectrum?

A

Mass to charge ratio (m/z).

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50
Q

For the mass spectrum of an element, what does the height of each peak give?

A

The relative isotopic abundance of that isotope.

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51
Q

On the x-axis of a mass spectrum, what is m/z?

A

The mass to charge ratio of the ion.

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52
Q

For the mass spectrum of an element, what does the m/z value usually correspond to and why?

A

The relative isotopic mass, since the charge is usually +1.

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53
Q

How can you work out the Ar of an element from its mass spectrum?

A

1) Multiply each relative isotopic mass (x-axis) by its abundance (y-axis) and add these up.
2) Divide by the sum of the isotopic abundances.

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54
Q

Practice calculating the Ar of an element from its mass spectrum.

A

Pg 7 of revision guide.

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55
Q

Silicon can exist in three isotopes. 92.23% of silicon is Si-28 and 4.67% is Si-29. The Ar of silicon is 28.1. Calculate the abundance and isotopic mass of the third isotope.

A

1) Abundance of third isotope = 100 - 92.23 - 4.67 = 3.10%
2) Let X equal the final isotopic mass
• 28.1 = ((28 x 92.23) + (29 x 4.67) + (X x 3.10) / 100
• 28.1 = (2717.87 + (X x 3.10)) / 100
• 2810 - 2717.87 = X x 3.10• 29.719 = X
• So the final isotope is Si-30.

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56
Q

Describe how you can predict the mass spectrum for a diatomic molecule.

A

1) Express each abundance as a decimal (e.g. 75% = 0.75)
2) Make a table showing all of the possible combination of atoms in the molecule.
3) For each molecule, multiply the abundances to get the abundance for the molecule.
4) Look for any molecules that are the same and add up their abundances. (e.g. Cl-37, Cl-35 and Cl-35, Cl-37 are the same).
5) Divide all of the relative abundances by the smallest relative abundance to get the smallest whole number ratio.
6) Use these values to predict the heights of each peak on the mass spectrum.

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57
Q

Chlorine has two isotopes. Cl-35 has an abundance of 75%. Cl-37 has an abundance of 25%. Predict the mass spectrum of Cl2.

A

1) Table of abundances:
• Cl-35, Cl-35 = 0.75 x 0.75 = 0.5625
• Cl-35, Cl-37 = 0.75 x 0.25 = 0.1875
• Cl-37, Cl-35 = 0.25 x 0.75 = 0.1875
• Cl-37, Cl-37 = 0.25 x 0.25 = 0.0625
2) Cl-35, Cl-37 is the same as Cl-37, Cl-35. So Cl-35, Cl-37 = 0.1875 + 0.1875 = 0.375
3) Smallest abundance = 0.0625
• Cl-35, Cl-35 -> 0.5625 / 0.0625 = 9 (Mr = 70)
• Cl-35, Cl-37 -> 0.375 / 0.0625 = 6 (Mr = 72)
• Cl-37, Cl-37 -> 0.0625 / 0.0625 = 1 (Mr = 74)
4) Draw the graph of relative abundance against m/z.
(See diagram pg 8 of revision guide)

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58
Q

How can the mass spectrum of a molecule be used to identify its Mr?

A

By looking at the m/z value of the peak with the highest m/z.
(NOTE: Ignore any very small peaks above this - due to carbon-13)

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59
Q

During mass spectrometry, when the molecules in a sample are bombarded with electrons, what is formed?

A

A molecular ion (M+).

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60
Q

When identifying the Mr of a molecule by its mass spectrum, why are there very small peaks after the M+ peak and why can they be ignored?

A

These are caused by the presence of carbon-13 atoms and are uncommon.

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61
Q

When looking at a mass spectrum for an unknown alcohol, how would you identify the alcohol?

A
  • Look at the peak with the highest m/z value (ignoring any tiny peaks after it)
  • The m/z value is the Mr of the alcohol
  • Calculate the Mr of the first few alcohols until you find the corresponding alcohol
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62
Q

Explain how mass spectrometry works.

A

1) Ionisation - The substance is bombarded with electrons, which knock out electrons to form positive ions
2) Acceleration - The ions are accelerated through an electric field.
3) Deflection - A magnetic field deflects the ions by different amounts according to their mass to charge ratio.
4) Detection - The number of ions arriving at each m/z value is detected by a computer.

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63
Q

What are some uses of mass spectrometry?

A 
Used to analyse samples in:
• Radioactive dating
• Space research
• Detection of illegal drugs in sport
• Pharmaceutical industry
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64
Q

Remember to revise mass spectra.

A

Pg 7 + 9 of revision guide

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65
Q

Describe the idea of electronic structure.

A
  • Electrons move around in quantum shells -> e.g. 2nd shell
  • The shells contain different types of subshell -> e.g. 2nd shell contains 2s and 2p subshell
  • Each subshell has a different number of orbitals, which can each hold 2 electrons -> e.g. 2p subshell has 3 orbitals, so 6 electrons
  • Shells/Energy levels -> Subshells -> Orbitals
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66
Q

What is the numbering of quantum shells called?

A

They are given numbers called “principal quantum numbers”.

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67
Q

What are quantum shells also known as?

A

Energy levels

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68
Q

How are the shells numbered?

A

The further from the nucleus, the greater the energy level, starting at 1.

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69
Q

How many electrons can an orbital hold?

A

2

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70
Q

What are the four subshells?

A

s, p, d, f

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71
Q

How many orbitals and electrons in an ‘s’ subshell?

A

Orbitals: 1

Max. electrons: 2

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72
Q

How many orbitals and electrons in a ‘p’ subshell?

A

Orbitals: 3

Max. electrons: 6

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73
Q

How many orbitals and electrons in a ‘d’ subshell?

A

Orbitals: 5

Max. electrons: 10

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74
Q

How many orbitals and electrons in an ‘f’ subshell?

A

Orbitals: 7

Max. electrons: 14

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75
Q

Name the four subshells and the number of electrons in each.

A
  • s - 2
  • p - 6
  • d - 10
  • f - 14
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76
Q

Name the subshells in the 1st energy level.

A

1s

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77
Q

Name the subshells in the 2nd energy level.

A

2s 2p

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78
Q

Name the subshells in the 3rd energy level.

A

3s 3p 3d

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79
Q

Name the subshells in the 4th energy level.

A

4s 4p 4d 4f

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80
Q

How many electrons fit into each of the first four energy levels?

A
  • 1st - 2
  • 2nd - 8
  • 3rd - 18
  • 4th - 32
81
Q

What is an orbital?

A

A region of space in which an electron spends most of its time.

82
Q

Compare the energy of different orbitals.

A

Those within the same subshell have the same energy.

83
Q

How is the movement of electrons in an orbital related and what is this called?

A
  • They spin in opposite directions to each other.

* Spin-pairing

84
Q

Describe the shape of s-orbitals.

A

Spherical.

85
Q

Describe the shape of p-orbitals.

A
  • Dumbbell shaped.
  • Px orbital goes along the x-axis.
  • Py orbital goes along the y-axis.
  • Pz orbital goes along the z-axis.
86
Q

Remember to revise the shape of orbitals.

A

Pg 10 of revision guide.

87
Q

How can boxes be used to represent subshells?

A
  • Each subshell is made up of the right number of boxes, which each represent an orbital.
  • Two arrows can be draw in each box to represent electrons.
88
Q

When drawing orbitals as boxes, what is important about the direction of the arrows in each box?

A

There is an up and a down arrow in each box, representing the electrons spinning in opposite directions.

89
Q

How can you work out the electronic configuration of an element/ion?

A

1) Work out the number of electrons
2) Fill up the lowest energy subshells first
3) Electrons fill orbitals singly before they start pairing up

90
Q

What us the order in which the subshells fill up?

A

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

91
Q

How does subshell notation work?

A
  • The energy level is given first (e.g. “1”s2)
  • The subshell is given next (e.g. 1”s”2)
  • The number of electrons is last (e.g. 1s”2”)
92
Q

What must you remember about how the electron subshells fill up?

A
  • The 4s subshell fills up before the 3d subshell

* Electrons always fill each orbital in a subshell singly before pairing up

93
Q

Practice drawing out the electron box diagram for nitrogen and oxygen.

A

See diagram pg 10 of revision guide.

94
Q

Describe the electron box diagram for nitrogen (7).

A
  • Up and down arrow in 1s box
  • Up and down arrow in 2s box
  • Up arrow in each of the 2p boxes
95
Q

Explain the shorthand method sometimes used to write electronic configurations.

A
  • Noble gas symbols are sometimes used to replace the start of an electronic configuration
  • E.g. Calcium (1s2 2s2 2p6 3s2 3p6 4s2) can be written as [Ar]4s2
96
Q

What are the different blocks in the periodic table?

A
  • s-block
  • d-block
  • p-block
97
Q

What is the block on the left of the periodic table called?

A

s-block

98
Q

What is the block in the middle of the periodic table called?

A

d-block

99
Q

What is the block on the right of the periodic table called?

A

p-block

100
Q

What defines s-block elements?

A

The highest energy/outermost electron is in the “s” subshell.

101
Q

What defines p-block elements?

A

The highest energy/outermost electron is in the “p” subshell.

102
Q

What is the electronic configuration of Ca? (Atomic number: 20)

A

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

103
Q

What is the electronic configuration of Ca2+? (Atomic number: 20)

A

1s2 2s2 2p6 3s2 3p6

104
Q

What is the electronic configuration of Cl? (Atomic number: 17)

A

1s2 2s2 2p6 3s2 3p5

105
Q

What is the electronic configuration of Cl-? (Atomic number: 17)

A

1s2 2s2 2p6 3s2 3p6

106
Q

How do you work out the electronic configuration of an ion?

A
  • Work out the number of electrons in the ion
  • Write out the configuration for this number of electrons
  • Electrons are taken from/added to the highest energy subshell.
107
Q

What defines d-block elements?

A

The highest energy electron is in the “d” subshell.
NOTE: This is not the outermost electron, since there are electrons in the 4s subshell fill up first, but have lower energy.

108
Q

Why is the electronic configuration of d-block elements more difficult to work out?

A

The 4s subshell fills before the 3d subshell.

109
Q

What is the electronic configuration of vanadium? (Atomic number: 23)

A

1s2 2s2 2p6 3s2 3p6 4s2 3d3

110
Q

What two elements have an unusual electronic configuration?

A
  • Copper (Cu)

* Chromium (Cr)

111
Q

What makes copper and chromium unusual in terms of electronic configuration? Why do they do this?

A
  • They donate one of their 4s electrons to the 3d subshell.

* This is because they’re more stable with a full or half-full d-subshell.

112
Q

What is the electronic configuration of chromium? (Atomic number: 24)

A

1s2 2s2 2p6 3s2 3p6 4s1 3d5

113
Q

What is the electronic configuration of copper? (Atomic number: 29)

A

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

114
Q

For EM radiation, as frequency increases…

A

Wavelength and energy decrease

115
Q

Give the order of the EM spectrum with increasing frequency.

A
  • Radio waves
  • Microwaves
  • Infrared
  • Visible light
  • UV
  • X-rays
  • Gamma rays
116
Q

What is the lowest energy of an electron called?

A

The ground state.

117
Q

Describe what happens when electrons in an atom are heated.

A
  • Electron takes in energy -> Moves to a higher energy level
  • Electron then releases energy by dropping to a lower energy level
  • This emits a frequency of light which corresponds to the energy difference between the two energy levels
118
Q

What is the ground state?

A

The lowest possible energy of an electron.

119
Q

What is excitation?

A

When electrons gain energy from their surroundings and move to a higher energy level.

120
Q

What is the notation for the ground state?

A

n = 1

121
Q

How are the different energy levels symbolised on a diagram?

A

n = 1, n = 2, n = 3, etc.

122
Q

How can the energy of each energy level be described?

A
  • Discrete

* This means that they have fixed values

123
Q

What is produced when electrons in a specific element de-excite?

A

An emission spectrum.

124
Q

Describe an emission spectrum.

A

Coloured lines of particular colours on a dark background.

125
Q

Why is the emission spectrum for each element different?

A
  • Each element has a different electron arrangement
  • So the electrons which de-excite experience different energy differences between energy levels
  • So the frequencies of radiation emitted are different
126
Q

Is it the excitation of electrons that causes light to be emitted from an element?

A

No, it is the de-excitation.

127
Q

What causes the different lines on an emission spectrum?

A
  • Electrons falling to different energy levels and from different energy levels.
  • This results in various energy differences and thus different photon frequencies
128
Q

Explain why on an emission spectrum there may be groups of lines.

A

Each group is due to electrons dropping to a different energy level.

129
Q

As the frequency increases, what happens to the lines in each set on an emission spectrum?

A

They get closer together.(See diagram pg 12 of revision guide)

130
Q

Describe and explain the emission spectrum of hydrogen.

A
  • 3 sets of lines
  • First set in the UV range -> Caused by electrons dropping to the n = 1 state.
  • Second set in the visible range -> Caused by electrons dropping to the n = 2 state.
  • Third set in the infrared range -> Caused by electrons dropping to the n = 3 state.
  • In each set, the lines are closest together on the left (where the frequency is higher).
131
Q

On an energy level diagram for the emission spectrum of an element, what happens to the energy level lines?

A
  • They get closer together.

* i.e. n = 1 and n =2 are further apart than n = 2 and n = 3.

132
Q

Remember to revise emission spectra.

A

Pg 12 of revision guide.

133
Q

Why does the de-excitation of electrons release EM radiation of a fixed frequency?

A
  • Each shell has a fixed energy
  • Moving down energy levels causes electrons to lose a fixed amount of energy
  • This energy corresponds to a particular frequency of photon
134
Q

What is some evidence for electrons existing in quantum shells?

A
  • Emission spectra - > Show that electrons always release fixed amounts of energy when moving between shells.
  • Ionisation energy decreasing down a group
135
Q

What are the 4 principles of electron shells?

A
  • Electrons can only exist in fixed orbits, or shells, and not anywhere in between.
  • Each shell has a fixed energy.
  • When an electron moves between shells, EM radiation is emitted or absorbed.
  • The radiation emitted or absorbed will always have a fixed frequency.
136
Q

What do the clear lines on an emission spectrum indicate?

A
  • Energy levels are discrete (not continuous)

* Electrons jump between energy levels, not move

137
Q

What is ionisation?

A

The removal of 1 or more electrons.

138
Q

Define the first ionisation energy.

A

The energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms to form 1 mole of gaseous 1+ ions.

139
Q

When writing ionisation energy equations, what is it important to remember?

A

Use the (g) state symbol for everything except the electrons.

140
Q

Is ionisation endothermic or exothermic?

A

Endothermic

141
Q

Write the equation for the first ionisation energy of oxygen.

A

O(g) -> O+(g) + e-

142
Q

Is the enthalpy change for ionisation positive or negative?

A

Positive

143
Q

What causes a high ionisation energy?

A
  • A strong attraction between the electron and the nucleus

* So more energy is needed to overcome the attraction and remove the electron.

144
Q

What 3 factors affect ionisation energy?

A

1) Nuclear charge
2) Electron shell (+ subshell!)
3) Shielding

145
Q

Explain how nuclear charge affects ionisation energy.

A
  • The more protons there are in the nucleus, the more positively charged the nucleus is
  • So there is a stronger attraction for the electrons.
146
Q

Explain how the electron shell affects ionisation energy.

A
  • Attraction decreases with distance
  • An electron in an electron shell close to the nucleus experiences a stronger attraction than one in a shell further away.
147
Q

What is shielding?

A

The decreased attraction of an outer electron by the nucleus due to more inner electrons.

148
Q

Explain how shielding affects ionisation energy.

A

• More shielding electrons -> Weaker attraction of outer electron

149
Q

Explain how first ionisation energy changes down a group.

A

• Decreases going down a group
Because:
• Each consecutive element has an extra electron shell
• Outer electrons are further from the nucleus -> Weaker attraction
• More inner shells for shielding -> Weaker attraction

150
Q

Define the second ionisation energy.

A

The energy needed to remove 1 mole of electrons from 1 mole of gaseous 1+ ions to form 1 mole of gaseous 2+ ions.

151
Q

Write the equation for the second ionisation energy of oxygen.

A

O+(g) -> O2+(g) + e-

152
Q

How many ionisation energies are there for an element?

A

As many as there are electrons.

153
Q

Explain the graph for the successive ionisation energies for an element.

A

• Within each shell -> Slight increase in ionisation energies -> Since the attraction from the nucleus is shared between fewer electrons -> Shallow gradient
• Going into a more inner shell -> Large increase in ionisation energy -> Since the electron is closer to the nucleus -> Steep gradient
(See diagram pg 15 of revision guide)

154
Q

Describe the graph for the successive ionisation energies of sodium.

A
  • Large jump between 1st and 2nd electron
  • Small jumps between 2nd and 9th electron ( 8 electrons)
  • Large jump between 9th and 10th electron
  • Small jump between 10th and 11th electron
155
Q

Remember to revise ionisation energy graphs.

A

Pg 15 of revision guide.

156
Q

How can you tell the group of an element by its successive ionisation energy graph?

A

Count how many electrons are removed before the first big jump.

157
Q

How can you tell the electronic configuration of an element by looking at its successive ionisation energy graph?

A
  • Work from right to left, counting how many electrons are removed before each large jump.
  • This is the number of electrons in each successive shell.
158
Q

What is periodicity?

A

The trends of elements going across the periodic table.

159
Q

What are the vertical columns in the periodic table called?

A

Groups

160
Q

What are the horizontal rows in the periodic table called?

A

Periods

161
Q

What are the elements in the periodic table arranged by?

A

Proton number

162
Q

What links all of the elements in a period?

A

They have the same number of electron shells.

163
Q

How many shells do the elements in each period have?

A
  • Period 1 -> 1 shell
  • Period 2 -> 2 shells
  • Period 3 -> 3 shells
  • etc.
164
Q

What links all of the elements in a group?

A

They have the same number of electrons in their outer shell.

165
Q

What properties repeat going across each period?

A

Chemical and physical

166
Q

What properties do the elements in a group share?

A

Chemical

167
Q

Which block in the periodic table do hydrogen and helium belong to?

A

s-block

168
Q

Describe the electronic configuration of group 1-3 ions.

A

They have the configuration of the previous inert gas.

169
Q

Describe the electronic configuration of group 4-7 ions.

A

They have the configuration of the next inert gas.

170
Q

How do d-block elements form ions?

A
  • Usually lose s and d electrons

* Form positive ions.

171
Q

Explain how electronic configuration determines chemical properties.

A

The number of outer electrons determines the way in which the element bonds and how it forms ions.

172
Q

What happens to atomic radius as you go along a period and why?

A
  • It decreases
  • Because the nucleus’ positive charge increases, pulling the electrons closer to the nucleus
  • The outer electron doe not contribute to shielding, since it is added to the outer shell
173
Q

Why does shielding not increase across a period?

A

The extra electrons are added to the outer shell, so they do not impact shielding.

174
Q

Describe the trend of the first ionisation energy going across a period.

A
  • General increase

* Small drops between groups 2-3 and 5-6

175
Q

Why does the first ionisation energy generally increase going across a period?

A
  • Number of protons is increasing -> Stronger nuclear attraction -> Smaller atomic radius
  • All outer electrons at roughly the same energy level, even if they’re in slightly different orbital types -> Little extra shielding or extra distance
176
Q

Going across periods 2 and 3, between which groups are there small drops in the first ionisation energy?

A

Groups 2-3 and 5-6

177
Q

Going across periods 2 and 3, why is there a drop in first ionisation energy between groups 2 and 3?

A

• The outer electron in the group 2 element is in the “s” subshell, while in the group 3 element it is in the “p” subshell
So the group 3 element has a higher 1st ionisation energy because:
• The electron is further from the nucleus
• There is extra shielding from the “s” subshell
(NOTE: These two effects override the effect of the extra nuclear charge)

178
Q

Going across periods 2 and 3, why is there a drop in first ionisation energy between groups 5 and 6?

A

• The outer electron is in the same subshell, so shielding is identical
BUT:
• Group 5 element has singly-filled subshells, while the group 6 element has its outer electron in an orbital with 2 electrons
• The repulsion between the two electrons in an orbital means that it is easier to remove the electron from the full orbital in group 6
(NOTE: This effect overrides the effect of the extra nuclear charge)

179
Q

Describe and practice drawing out the bar chart for first ionisation energy across period 2 or 3.

A 
• Increase between groups 1 and 2
• Fall to group 3
• Increase to group 5
• Fall to group 6
• Increase to group 8
(See diagram pg 17 of revision guide)
180
Q

Which has a higher first ionisation energy, Mg-12 or Al-13, and why?

A

Mg = 1s2 2s2 2p6 3s2
Al = 1s2 2s2 2p6 3s2 3p1
• Aluminium’s outer electron is in the next subshell, so it receives more shielding and is further from the nucleus
• This overrides the effect of the extra nuclear charge and makes Al’s outer electron easier to remove
• So Mg-12 has a higher first ionisation energy

181
Q

Which has a higher first ionisation energy, P-15 or S-16, and why?

A

P = 1s2 2s2 2p6 3s2 3p3
S = 1s2 2s2 2p6 3s2 3p4
• Phosphorus’ outer subshell is singly filled, while sulfur’s outer electron is in a full orbit
• So sulfur’s outer electron experiences a greater repulsion
• This overrides the effect of the extra nuclear charge and makes S’s outer electron easier to remove
• So P-15 has a higher first ionisation energy
(See diagram pg 17 of revision guide)

182
Q

Remember to revise periodicity trends for first ionisation energy.

A

Pg 17 of revision guide.

183
Q

How does the way a subshell is filled affect its stability and what is the result of this?

A

Singly filled or full subshells -> More stable -> Higher ionisation energies

184
Q

What is some evidence for the theory of electron subshells?

A

• The drop in ionisation energy between groups 2-3 and 5-6

185
Q

Going across periods 2 and 3, how does melting/boiling point change?

A 
• Increases from group 1 to 4
• Drops to group 5
• Fluctuating low between groups 5 to 7
• Lowest at group 8
(See diagram pg 18 of revision guide)
186
Q

Going across periods 2 and 3, describe how the bonding changes?

A
  • Groups 1 to 3 -> Metallic (except boron)
  • Group 4 -> Giant covalent
  • Groups 5 to 6 -> Simple molecular
  • Group 8 -> Monatomic
187
Q

Give the formula for the following molecules:• Nitrogen• Oxygen• Fluorine• Phosphorus• Sulfur• Chlorine

A
  • Nitrogen - N2
  • Oxygen - O2
  • Fluorine - F2
  • Phosphorus - P4
  • Sulfur - S8
  • Chlorine - Cl2
188
Q

Explain how and why the boiling point of elements changes as you go across period 2.Elements: Li, Be, B, C, N, O, F, Ne

A

• Increase from Li to Be -> Metallic bonding gets stronger -> Due to more delocalised electrons + smaller atomic radius
• Increase to B and C -> Giant covalent structures
• Fall to N, O and F -> Simple molecular -> Weak intermolecular forces -> Strength of London forces depends on no. of electrons
• Fall to Ne -> Monatomic -> Very weak London forces
(See diagram pg 18 of revision guide)

189
Q

Explain how and why the boiling point of elements changes as you go across period 3.Elements: Na, Mg, Al, Si, P, S, Cl, Ar

A

• Increase from Na to Al -> Metallic bonding gets stronger -> Due to more delocalised electrons + smaller atomic radius
• Increase to Si -> Giant covalent structure• Fall to P, S and Cl -> Simple molecular -> Weak intermolecular forces -> Strength of London forces depends on no. of electrons -> NOTE: S higher than P, which is higher than F, due to electron number
• Fall to Ar -> Monatomic -> Very weak London forces
(See diagram pg 18 of revision guide)

190
Q

How does the boiling point of metals change going across a period and why?

A

• Boiling point increases
BECAUSE:
• Increasing number of delocalised electrons
• Decreasing ionic radius
So there is a stronger attraction between the metal ions and the delocalised electrons.

191
Q 
Compare the boiling point of elements which are:
• Metals
• Giant covalent structures
• Simple covalent molecules
• Monatomic
A

Highest: Giant covalent
Metals
Simple molecules
Lowest: Monatomic

192
Q

Why do giant covalent structures have such boiling points?

A
  • Have several strong covalent bonds

* A lot of energy is needed to break all of these

193
Q

Name two elements which form giant covalent structures.

A

Carbon and silicon

194
Q

What determines the boiling point of simple molecular structures?

A
  • The number of electrons in the molecule
  • The more electrons, the stronger the London forces
  • These are generally quite weak, so the boiling point is low
195
Q

Why do noble gases have the lowest boiling point?

A

They are monatomic, so they have very weak London forces.

196
Q

What type of structure does silicon have?

A

Giant covalent

197
Q

Order these by boiling point and explain why this is the case: phosphorus, sulfur, chlorine.

A

Highest: Sulfur (S8)
Phosphorus (P4)
Lowest: Chlorine (Cl2)
This is because of the number of electrons in each molecule determines the strength of the London forces. So S8 has the strongest intermolecular forces.

198
Q

What is the molecular formula of sulfur?

A

S8

199
Q

What is the molecular formula of phosphorus?

A

P4

A LEVEL - Chemistry (Inorganic) (21 decks)

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