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OCR A A-Level Inorganic Chemistry > Chapter 7 > Flashcards

Chapter 7 Flashcards

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1
Q

What do groups in the periodic table show? How does this relate to the elements?

A
  • Groups in the periodic table have the same number of electrons in their outer shell, and similar chemical properties
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2
Q

What is periodicity? Required.

A
  • Repeating trends in the physical and chemical properties of elements across a period
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3
Q

What are 4 properties that have periodicity?

A
  • Electronic configuration
  • Ionisation energy
  • Structure
  • Melting points
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4
Q

What trend do periods have in terms of electronic configuration?

A
  • Each period starts with an electron in the new highest energy level
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5
Q

What is first ionisation energy? Required.

A
  • The energy required for the removal of 1 mol of electrons from 1 mol of gaseous atoms of an element
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6
Q

Write an equation for the first ionisation energy of chlorine.

A
  • Cl (g) -> Cl+(g) +e- (it is not shown as diatomic as first ionisation energy refers to 1 mol of atoms)
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7
Q

What is second ionisation energy?

A
  • The amount of energy needed to turn 1 mol of gaseous 1+ ions into 1 mol of gaseous 2+ ions
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8
Q

What 3 factors affect ionisation energy?

A
  • Atomic radius
  • Nuclear charge
  • Electron shielding
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9
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
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10
Q

How does nuclear charge affect ionisation energy?

A
  • The more protons there are in the nucleus of an atom, the greater the attraction between the nucleus and the outer electrons
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11
Q

How does electron shielding affect ionisation energy?

A
  • Electrons are all negatively-charged and therefore repel each other, meaning that inner-shell electrons repel outer-shell electrons
  • This reduces the attraction between the nucleus and the outer electrons
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12
Q

How do successive ionisation energies compare with each other?

A
  • They increase
  • There is a big increase in the ionisation energy between different shells of an atom
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13
Q

Why do successive ionisation energies increase?

A
  • The atomic radius decreases
  • The nuclear attraction between the remaining electrons and nucleus increases (the same number of protons attract fewer electrons)
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14
Q

Describe the trends in first ionisation energy in the periodic table.

A
  • They decrease down a group (down across the different periods)
  • They generally increase across a period
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15
Q

Why does first ionisation energy decrease down a group?

A
  • The atomic radius and shielding increases
  • Nuclear attraction on outer electrons decreases
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16
Q

Why does first ionisation energy generally increase across a period?

A
  • The nuclear charge increases, which causes the atomic radius to decrease
  • There is similar shielding
  • The nuclear attraction therefore increases
17
Q

In what 2 places does the first ionisation energy not increase across a period?

A
  • Between groups 2 and 3
  • Between groups 5 and 6
18
Q

Why does the first ionisation energy decrease between groups 2 and 3?

A
  • In group 3, the p sub-shell starts to be filled
  • As it has more energy than the s sub-shell, it is easier to remove the electron in the p sub-shell
  • Group 3 therefore has a lower first ionisation energy than group 2
19
Q

Why does the first ionisation energy decrease between groups 5 and 6?

A
  • In group 6, the p sub-shell electrons start to pair up
  • The paired electrons in one orbital repeal each other, making it easier for one of them to be removed
  • Therefore, group 6 has a lower first ionisation energy than group 5
20
Q

What does the graph of first ionisation energy against atomic number look like? Give 5 details.

A
  • The highest points are the noble gases
  • The lowest points are the group 1 metals
  • First ionisation energy generally increases across each period
  • There are dips between groups 2 and 3 and 5 and 6
  • Each period has a lower first ionisation energy than the previous period (the entire line is shifted down)
21
Q

What is metallic bonding? Required.

A
  • The strong electrostatic attraction between cations and delocalised electrons
22
Q

What structure do metals have?

A
  • They all have a giant metallic lattice structure
23
Q

What are the 2 main properties of metals?

A
  • High electrical conductivity
  • High melting and boiling points
24
Q

When can metals conduct electricity, and how?

A
  • When in the solid or liquid states
  • This is because the delocalised electrons are free to move through the structure and carry charge
25
Q

Why do most metals have high melting and boiling points?

A
  • Large amounts of energy are needed to overcome the strong electrostatic attraction between the cations and electrons
26
Q

Are metals soluble? Why?

A
  • No
  • They react with solvents instead, such as group 1 metals with water
27
Q

What are giant covalent lattices? Required.

A
  • Networks of atoms bonded by strong covalent bonds
28
Q

List 2 elements that form giant covalent lattices.

A
  • Carbon
  • Silicon
29
Q

What are the melting and boiling points of giant covalent lattices like, and why?

A
  • They are high because a lot of energy is needed to break the strong covalent bonds
30
Q

Are giant covalent lattices soluble? Why?

A
  • No
  • The covalent bonds are too strong to be broken due to interactions with the solvent
31
Q

Can giant covalent lattices conduct electricity? Why?

A
  • Most aren’t as they don’t have any electrons that are free to conduct electricity
  • Graphite and graphene are exceptions to this as they only use 3 of their 4 electrons in their outer shell to form covalent bonds
32
Q

What is the difference between graphite and graphene?

A
  • Graphene is a single layer of graphite (in graphite the separate layers are bonded together by London forces)
33
Q

How do melting points vary across periods 2/3?

A
  • The melting points increase from periods 1 to 4
  • There is a sharp decrease between groups 4 and 5
  • From group 5 to 0, the melting points are low
34
Q

Why do melting points decrease between groups 4 and 5?

A
  • The elements change from giant molecular structures to simple molecular substances
35
Q

Why do melting points increase between groups 1 and 3?

A
  • The larger the charge of the cation (and therefore also the relative amount of delocalised electrons) the stronger the metallic bonds

OCR A A-Level Inorganic Chemistry (16 decks)

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