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Chemistry - OCR A > 3.1.1 - Periodictiy > Flashcards

3.1.1 - Periodictiy Flashcards

1
Q

Define periodicity.

A

Periodicity is the repeating pattern of properties in a group

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2
Q

What are the 4 blocks that the periodic table is split into?

A
  • S block
  • P block
  • D block
  • F block
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3
Q

What is the trend in electron configuration?

A

You add 1 electron as you go across the period, when you start a new period you fill a new shell.

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4
Q

Why do elements with similar electron configurations react similarly?

A

Chemical reactions only involve outer shell electrons so they have similar chemical reactions.

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5
Q

How are elements classified?

A

Elements that have the same number of outer shell electrons are put in the same group.

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6
Q

What is ionisation energy?

A

Ionisation energy is the energy required to form a positive ion. It is a measure of how easily an atom loses an electron.

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7
Q

Define first ionisation energy?

A

The energy required to remove one electron from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous +1 ions.

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8
Q

What is the general equation for first ionisation energy?

A

X —> X+ + e-

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9
Q

What are the 3 main factors effecting ionisation energy?

A
  • atomic radius
  • nuclear charge
  • electron shielding
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10
Q

What effect does the atomic radius have on an atom?

A

The greater the distance the lead nuclear attraction, forces of attraction stop sharply with the increasing distance.

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11
Q

What effect does nuclear charge have on an atom?

A

As the nuclear charge increases there is a greater attraction between the nucleus and the outer shell electrons.

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12
Q

What effect does electron shielding have on an atom?

A

The inner electrons repel the outer shell electrons, this repulsion is called shielding. The more shells there are the more shielding there is so there is less nuclear attraction.

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13
Q

What happens to the atomic radius, nuclear charge and electron shielding as you go down the group?

A

All three increase as you go down the group.

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14
Q

What happens to the atomic radius, nuclear charge and electron shielding as you go across the period?

A
  • the nuclear charge increases
  • the atomic radius
  • the electron shielding
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15
Q

What is the trend in ionisation energy as you go down the group?

A

The nuclear attraction decreases as you go down so the ionisation energy increases as you go down.

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16
Q

What is the trend in ionisation energy going across the periodic table?

A

The nuclear attraction increases so the ionisation energy increases across each period.

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17
Q

What happens to the ionisation energy between group 2 and 3?

A

The ionisation energy decreases.

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18
Q

Why does the ionisation energy between group 2 and 3 decrease?

A

It is due to the sub shell structure:
- the outer electrons in group 3 are in the p-orbital and the p-orbitals have more energy that’s the s-orbitals
- so the electrons are further away from the nucleus
- p-orbitals also have additional shielding provided by the s-orbitals
- the increased shielding and atomic radius override the affects of the increase in nuclear charge so the ionisation energy drops

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19
Q

What is successive ionisation energy?

A

It is a measure of the energy required to remove each electron in turn.

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20
Q

What is the maximum number of ionisation energies equivalent to?

A

The number of electrons there are in an atom of an element.

21
Q

What is the general equation of second ionisation energies?

A

X+ —> X2+ + e-

22
Q

What causes the successive ionisation energy to increase?

A
  • As each electron is removed, there is less repulsion between electrons
  • So the shells draw closer causing the atomic radius to get smaller,
  • This mean that the nuclear attraction increases and the ionisation energy increases.
23
Q

What does a large increase on an ionisation energy graph mean (eg. 5-6)?

A

It correlates the the number of electrons in the outer shell (5 electrons in the outer shell).

24
Q

Define metallic bonding.

A

The electrostatic attraction between positive metal ions and delocalised electrons.

25
Q

What are the properties of a giant metallic lattice?

A
  • high melting and boiling point
  • good electrical conductivity
  • insoluble in water
  • malleability and ductility
26
Q

Why do giant metallic lattices have a high melting and boiling point?

A

The electrons are free to move but the positive ions are fixed, the attraction between them are strong. So a high temp is needed to break the metallic bonds to dislodge the ions.

27
Q

Why is a giant metallic lattice a good conductor?

A

The delocalised electrons are able to move freely through the lattice which allows them to carry a charge therefore it can conduct electricity.

28
Q

Why are giant metallic lattices malleable and ductile?

A

The delocalised electrons give the structure a degree of give which allows the atoms in the layers to slide across each other.

29
Q

Why are giant metallic lattices not soluble?

A

That atomes are not able to for hydrogen bonds so it can not dissolve/be pulled apart by water that is polar.

30
Q

What are the 2 types of lattices that a covalent compound can be formed?

A

Simple molecular or giant covalent

31
Q

What are the 3 main properties of a simple molecular lattice?

A
  • low melting and boiling point
  • non-conductors of electricity
  • solubility
32
Q

Why does a simple molecular lattice have a low melting and boiling point?

A

In a solid state the lattice is held together by weak intermolecular forces, so not much energy is required to overcome them.

33
Q

Why do simple covalent lattices not conduct electricity?

A

There are no delocalised electrons to move and carry a charge.

34
Q

Why are simple covalent molecules soluble?

A

Non-polar molecules are soluble in non-polar solvents and polar molecules are soluble in polar solvents.

35
Q

What are the 3 main properties of a giant covalent lattice?

A
  • high melting and boiling point
  • insolubility
    -electrical conductivity
36
Q

Why do giant covalent lattices tend to have high melting and boiling points?

A

There are lots of strong covalent bonds so a lot of energy is required to overcome them.

37
Q

Are giant covalent lattices soluble?

A

No they are insoluble in both polar and non-polar solvents.

38
Q

Can giant covalent lattices conduct electricity?

A

Generally no, because there are no free ions or electrons to carry a charge.

39
Q

What are 3 examples of giant covalent lattices?

A
  • diamond
  • graphite
  • silicone dioxide
40
Q

What is the structure in a diamond like?

A

Each carbon is covalently bonded to 4 other carbons in a tetrahedral arrangement.

41
Q

Can diamond conduct electricity?

A

No, it has no delocalised electrons

42
Q

How hard is diamond?

A

Very hard, there are lots of strong covalent bonds in a 3D arrangement that are not easily broken

43
Q

What is the structure of graphite like?

A

Each carbon is covalently bonded to 3 other carbons in a trigonal planar arrangement, they are bonded between layers

44
Q

Can graphite conduct electricity?

A

Yes, the 4th electron from each atom is delocalised and free to carry a charge

45
Q

Is graphite hard?

A

Very soft, there are weak London forces between the layers that are easily broken meaning that the layers can slide

46
Q

What is the structure of graphene like?

A

Each carbon is covalently bonded to 3 other carbons in a 2D trigonal planar arrangement

47
Q

Can graphene conduct electricity?

A

Yes, the 4th electron from each atom is delocalised

48
Q

Is graphene hard?

A

Very strong, as it is a single sheet and all the atoms are covalently bonded

Chemistry - OCR A (39 decks)

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