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A-level Chemistry Yr 12 Mine > 7 - Periodicity > Flashcards

7 - Periodicity Flashcards

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1
Q

What does Periodicity mean?

A

Across each period there is a repeating trend in properties of the element is called periodicity.

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2
Q

How is the periodic table arranged by elements?

A
  • By increasing atomic (proton) number.
  • In periods showing repeating trends in physical and chemical properties (periodicity).
  • In groups having similar chemical properties.
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3
Q

What is the trend down the group in electron configuration?

A

Elements in each group will have the same number of electrons in each sub shell.

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4
Q

What does each period start with?

A

An electron in a new highest energy shell.

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5
Q

What is the trend across period 2 in electron configuration?

A

The 2s sub-shell fills with two electrons, followed by the 2p sub-shell with the six electrons.

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6
Q

What is the trend across period 3 in electron configuration?

A

The 3s sub-shell is filled with two electrons, followed by the 3p sub-shell with six electrons.

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7
Q

What is the trend across period 4 in electron configuration?

A

The 4s sub-shell is filled first, due to the fact a new highest sub shell is needed to occupied. Then 3d sub shell is filled with ten electrons and then 4p sub shell with six electrons.

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8
Q

What group are in the s-blocks for electron configuration?

A

Group 1 and 2

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9
Q

What group are in the p-blocks for electron configuration?

A

Group 13 to group 18

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10
Q

What group are in the d-blocks for electron configuration?

A

Group 3 to 12 (transition metals)

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11
Q

The old group for halogens were 7, what group are they called now?

A

Group 17

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12
Q

How is the periodic table broken into electron configuration blocks?

A

The highest energy sub-shell.

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13
Q

What does ionisation energy measure?

A

How easily an atom loses electrons to form positive ions.

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14
Q

What is the definition for the First Ionisation Energy?

A

The First Ionisation Energy is the energy required to remove one electron from each atom in one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions.

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15
Q

Write the equation for the first ionisation energy for Sodium/Na?

A

Na (g) —> Na+ (g) + e-

Remember (g) as it has to be a gas

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16
Q

Why has the element got to be in a gaseous state for ionisation energy?

A

So intermolecular forces do not affect the energy required.

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17
Q

When presenting the energy required for ionisation energy, why must you have a + infront of the temperature (i.e +496 kgmol)?

A

It is an endothermic process as energy is put in so the bonds are broken.

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18
Q

What are the three factors affecting ionisation energy?

A
  • Nuclear Charge
  • Atomic Radius
  • Electron Shielding
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19
Q

How does Nuclear Charge affect attraction?

A

The more protons in the nucleus, the more positively charged the nucleus is the stronger the attraction for the negative (electrons).

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20
Q

How does atomic radius affect attraction?

A

Attraction falls off rapidly with distance, an electron closer to the nucleus will be much more strongly attracted.

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21
Q

How does electron shielding affect attraction?

A

Electrons are negatively charged and so inner-shell electrons repel outer shell electrons. This repulsion is called the shielding effect and reduced the attraction between the nucleus and outer electrons.

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22
Q

In first (or any) ionisation energy is an electron lost or gained?

A

An electron is lost as it creates a 1+ ion.

23
Q

What is the definitions for Second Ionisation Energy?

A

The second ionisation energy is energy required to remove one electron from each ion in one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions.

24
Q

What is the equation for first ionisation energy and second ionisation energy for Helium/He?

A

He (g) —> He+ (g) + e-

He+ (g) —> He2+ (g) + e-

25
Q

Why is the second ionisation energy greater than first ionisation energy?

A

After the first electron is lost, in first I.E, the outermost electron is pulled closer to the helium nucleus. The nuclear attraction on the renaming electron increases due to less repulsion and more I.E energy is needed to remove this second electron.

26
Q

This element is in period 3, the 1st I.N (ionisation number) has I.E (ionisation energy) is 578.
2nd I.N has I.E is 1,817
3rd I.N has I.E is 2,745
4th I.N has I.E is 11,577
5th I.N has I.E is 14,842
What is the element and why?

A

Aluminium

We know this element is in group 13, has 3 electrons in the outermost shell. As the large difference in I.E is between 3rd and 4th I.N.

27
Q

What is trend in first ionisation energy down a group and why?

A

Decreases
The atomic radius increases and also more inner shells so electron shielding increase down the group. This causes nuclear attraction on outer electrons to decrease and therefore the first ionisation energy decreases as well down the group.
Nuclear charge does increase due to more protons but it is outweighed by the points above.

28
Q

What is trend in first ionisation energy across a period and why?

A

Increases
Nuclear charge increase due to more protons and there in the same shell so similar shielding. This causes for nuclear attraction to increase and the atomic radius then decreases. Causing first ionisation energy to increase.

29
Q

Explain why B (group 13) have a slightly lower first ionisation energy than Be (group 2)?

A

It’s due to Sub-shell structure.
Outer electron in group 13 is in the p orbital which has a slightly higher energy so the electron is further from the nucleus and has additional electron shielding from the s-orbitals. These factors override the increased nuclear charge.

30
Q

Explain why O (group 16) have a lower lower first ionisation energy than N (group 15)?

A

Due to p orbital repulsion.
Group 16 have an orbital containing two electrons whereas group 15 have orbitals only containing 1 electron. And the electron is removed from the double electron orbital first in group 16.
The repulsion between two electrons in an orbital mean an electron is easier to remove.

31
Q

What are metalloids?

A

Elements near to the metal and non-metal divide which can show metal and non-metal properties.

32
Q

What is metallic bonding?

A

Metallic bonding is the strong electrostatic attraction between cations (positive ions) and delocalised electrons.

33
Q

How are delocalised electrons formed?

A

In a solid metal structure, each atom has donated it’s negative outer shell electrons to a shared pool of electrons, which are delocalised (spread out) throughout the whole structure.

34
Q

What is a giant metallic lattice structure?

A
  • The cations/positive ions (formed from the lost delocalised electrons) are in fixed position maintaining the structure and shape of the metal.
  • The delocalised electrons are mobile and are able to move through out the structure.
  • It contain strong metallic bonds between the cations and delocalised electrons.
35
Q

What are the properties of metals?

A
  • Strong metallic bonds (attraction between positive ions and delocalised electrons)
  • High electrical conductivity
  • High melting and boiling points
36
Q

Are metals soluble?

A

Metals are insoluble, except in liquid metals, because of the strength of the metallic bond.

37
Q

Why have metals got a high melting/boiling points?

A

Due the strong metallic bonds between the delocalised electrons and positive ions.

38
Q

What affects the MP/BP of metals?

A
  • The number of delocalised electrons, more delocalised electrons the stronger the bonds.
  • The size of the metal ion and lactic structure.
39
Q

What is a Giant Covalent lattices?

A

Are huge networks of covalently bonded atoms.

40
Q

Why can carbon and silicon form a giant Covalent Lattice?

A

Carbon and silicon are in group 14 so have four electrons in their outermost shell. Carbon (in diamond form) and silicon use these four electrons to form four covalent bonds to four other carbon/silicon atoms.

41
Q

What are the properties of giant covalent lattices?

A
  • High melting and boiling points.
  • Insoulble in almost all solvents.
  • Generally non-conductors of electricity (except graphene and graphite).
42
Q

What are some examples of solid giant covalent lattices?

A

carbon, diamond, graphite, graphene and silicon.

43
Q

What shape and bond angle does diamond/silicon make?

A

Tetrahedral structure, bond angle of 109.5°

44
Q

Why do giant covalent lattices insoluble?

A

Due to the strong covalent bonds that are hard to be broken.

45
Q

Why do giant covalent lattices have high melting/boiling points?

A

Strong covalent bonds are hard to overcome.

46
Q

Why do carbon/silicon not conduct electricity?

A

All four electrons are involved in covalent bonding (no delocalised electrons).

47
Q

What is graphite?

A

Graphite is composed of parallel layers of hexagonally arranged carbon atoms, bonded by weak London forces.

48
Q

What is graphene?

A

A single layer of graphite.

49
Q

What is the difference between graphite/graphene and giant covalent bonds?

A
  • Can conduct electricity, has delocalised electrons.
  • Less dense, as it has layers.
  • Slippery as the layers can slide over each other.
50
Q

Explain the trend in melting point across Periods 2 and 3?

A
  • The melting point increases from group 1 to group 14.
  • There is a sharp decrease in melting point between group 14 and group 15.
  • The melting points are comparatively from group 15 to group 18.
51
Q

Why is there a sharp decrease between group 14 and group 15 in melting points across period 2 and 3?

A

Group 14 elements are giant covalent lattices which have strong covalent bonds which need to be overcome. Whereas from group 15 to group 18 have a simple molecule structure which have weak London forces between molecules which are easy to overcome.

52
Q

Why is there a melting point increase from group 1 to group 14 across group 2 and 3?

A

As group 1, 2 and sometimes 13, on period 3 (aluminium), are giant metallic structures which have strong metallic bonds which get stronger as the ionic radius decreases and then number of delocalised electrons increase.

In group 14 and sometime 13, across period 2 (Boron), are giant covalent latices with strong covalent bonds were a lot of energy is needed to overcome them.

53
Q

Why is melting point of the noble gases the lowest across period 2 and 3?

A

As held together by the weakest forces.

A-level Chemistry Yr 12 Mine (15 decks)

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