Chapter 7 - The periodic table Flashcards Preview

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Flashcards in Chapter 7 - The periodic table Deck (23):
1

Explain the meaning of the term periodicity and give an example using the elements in period three

Periodicity is the gradual change in a property across a period that is repeated across each period. For example, the atomic radius of atoms across period three gradually falls.

2

Briefly explain what is meant by ionisation energy

The ionisation energy is the energy required to remove an outer electron from the attractive force of the nucleus.

3

What is the first ionisation energy?

The first ionisation energy is the energy required to remove an electron from each of one mole of gaseous atoms.

4

Why does the first ionisation energy decrease down the group?

Further down the group the outer electron is further from the nucleus and shielded by electrons in more inner shells. The electrostatic attraction to the nucleus is reduced.

5

Write an equation for the ionisation energy of chlorine

Cl(g) --> Cl+(g) + e-

6

Explain whether chlorine or bromine has the lower first ionisation energy. (4 marks)

Bromine atoms are larger, so the outer electrons are further from the nucleus and shielded by more inner electrons. The electrostatic attraction between bromine's outer electron and the nucleus is less, so the first ionisation energy is lower.

7

Why does the first ionisation energy increase across period 2 and 3?

The number of protons in the nucleus increases, the outer electron is in the same shell with similar shielding. The electrostatic attraction of the outer electron to the nucleus increases.

8

Why do Boron and Aluminium have lower first ionisation energies than Beryllium and magnesium? Which is an exception to the general trend.

The outer electrons of boron and aluminium are in a p orbital, higher in energy.

9

Why does Oxygen and Sulphur have lower first ionisation energies than nitrogen and phosphorus - it's an exception to the general trend.

The outer electrons of oxygen and sulphur have paired electrons in a p-orbital. These repel each other so one electron is more easily removed.

10

What is the structure of the metallic elements?

The metallic elements have a giant metallic lattice structure.

11

What is the structure of the non-metallic elements?

They have either a giant covalent lattice or are composed of simple molecules or atoms.

12

Why is there strong bonding in metals such as lithium beryllium, sodium and magnesium.

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

13

What bonding is there in carbon and silicon?

Giant covalent lattice's - The covalent bonds between the atoms are strong, these bonds are throughout the structure.

14

What are the three structures of pure carbon?

Diamond graphite and graphene.

15

What is the bonding like in simple molecules like oxygen, fluorine and nitrogen?

They are molecules with strong covalent bonds between atoms, but weak intermolecular forces between the molecules.

16

What is bonding like in noble gases?

There are weak interatomic forces between the atoms

17

Explain the boiling points in metals, carbon and silicon and other non-metals.

Metals have high boiling points due to the strong electrostatic attraction between the cations and delocalised electrons.
Carbon and silicon have high boiling points due to the strong covalent bonds that have to be broken.
The other non-metals have low boiling points due to weak intermolecular forces.

18

Explain conductivity in in giant covalent structures

Giant covalent lattices are non-conductors of electricity, the only exceptions are graphene and graphite.

19

Explain the solubility of metals

Metals do not dissolve.

20

Explain the electrical conductivity of metals

Metals have delocalised electrons that can move so they are good conductors. Graphite and graphene are good conductors because they have delocalised electrons.

21

Explain the solubility of giant covalent structures

Giant covalent lattices are insoluble in almost all solvents. The covalent bonds holding together the atoms in the lattice are far too strong to be broken by interaction with solvents.

22

Explain why the general trend is that the first ionisation energy increases across period 2

The first ionisation energy increases across period 2 because each subsequent atom has one extra proton. The electron being removed is in the same shell so that shielding from inner electrons is similar. Overall, the electrostatic attraction of the outer electron to the nucleus is increased meaning that more energy is required.

23

Explain why the first ionisation energy of oxygen is lower than that of nitrogen?

In nitrogen there are no paired electrons in the 2p orbitals. In oxygen, electrons are paired in one of the 2p orbitals. These electrons repel each other.