Chapter 7: Periodicity (7.2,7.3) Flashcards

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1
Q

What is ionisation energy ?

A

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

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2
Q

What is the trend in first ionisation energies?

A

1) There is a general increase in first ionisation energies across each period
2) There is a sharp decrease in first ionisation energies between the end of a period and the start of a new one

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3
Q

Explain why first ionisation energy increases across a period?

A

1) The atomic radius decreases meaning that there is a larger nuclear attraction to the outermost electron.
2) The nuclear charge also increases due to an increase in the number of protons in the nucleus
3) They all have the same number of shells which mean that they all have similar nuclear shielding

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4
Q

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

A

1) The atomic radius increases due to there being more shells which means that there is a smaller nuclear attraction the outermost electron
2) There are more inner shells, which means that the shielding would increase
3) Although nuclear charge increases this outweighed by the other factors

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5
Q

What are the 3 factors effecting ionisation energy?

A

1) Atomic radius
2) Nuclear charge
3) Electron shielding

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6
Q

Why does ionisation energy increase as you remove more electrons?

A

After the 1st electrons are lost the 2nd electrons are pulled closer to the nucleus due to there being more protons remaining than electrons, so it requires more ionisation energy.

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7
Q

How do you make predictions using ionisation energies?

A

1) Count the number of ionisation energies before a big jump tells you the number of electrons in the outer shell
2) The large increase in ionisation energies tells you the electron has been removed from a shell closer to the nucleus

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8
Q

Melting points in giant covalent structures

A

They have high melting and boiling points because high temperatures are required to supply the large amount of energy needed to break the strong covalent bonds.

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9
Q

Solubility of giant covalent structures

A

They are insoluble in almost all solvents - the covalent bonds are too strong to be broken down by interactions with solvents

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10
Q

Electrical conductivity of giant covalent structures

A

Giant covalent structures cannot conduct electricity apart from graphene and graphite. For example in carbon (diamond) and silicon all 4 outer shell electrons are involved in covalent bonding so are not available to conduct electricity, however in graphite and graphene bonds are formed where 1 electron is available for conducting.

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11
Q

Key features of graphene:

A

1) Planar hexagonal layers with bond angles of 120 degrees
2) Strong covalent bonds
3) Only 3 bonds are formed so one electron remains - this goes into the shared pool of delocalised electrons

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12
Q

Key features of graphite

A

1) Parallel layers of hexagonally arranged carbon atoms
2) Weak forced between layers
3) Bonding to hexagonal layers, again, only uses 3 electrons

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13
Q

Describe the bonding and structure in metals

A

The strong electrostatic forces of attraction between delocalised electrons and the positive metal ions.

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14
Q

Electrical conductivity of metals

A

Metals conduct when solid and liquid. When a voltage is applied across a metal, the delocalised electrons can move, carrying charge.

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15
Q

Melting and boiling points of metals

A

For most metals high temperatures are required to provide the large amounts of energy needed to overcome the strong electrostatic forces of attraction between the cations and electrons.

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16
Q

Solubility of metals

A

Metals do not dissolve in any solvents

17
Q

Trend in boiling point across period 2

A

Boiling point of metals increases as you go across the period this is due to them having metallic bonding and the intermolecular forces being stronger as you go across, until carbon where it peaks with the highest melting point, due to it having a giant covalent structure with requires high temperatures to supply energy to break covalent bonds and then boiling point decreases again because low termperatures are required to overcome the weak London forces between covalent molecules.

18
Q

Why does Boron have a significantly higher melting point than aluminium?

A

Because it has a giant covalent structure