7 - Periodicity

Question 1

How did Mendeleev arrange the known 60 elements in the periodic table?

Answer 1

He arranged them in order of atomic mass.
He also lined up elements in groups with similar properties. If group properties did not fit, he swapped or left gaps assuming the atomic mass measurement was wrong.

Question 2

What common name did Mendeleev give to undiscovered elements?

Answer 2

He put “eka”infront of the name of the element above the unknown element. eg Eka-silicon.

Question 3

What does periodicity mean?

Answer 3

The repeating trend in properties of elements accross a period.

Question 4

Describe the trend in electron configuration accross period 4.

Answer 4

The 4s subshell has a lower energy than the 3d subshell so it fills up first. The 4p subshell has the highest energy so it fills up last.

Question 5

What is the common name given to each group in the periodic table?

Answer 5

1- Alkali metals 2- Alkali earth metals. 3-12 - transition elements. 15 - pnictogens 16 - Chalcogens 17 - Halogens. 18 - noble gases.

Question 6

What is the first ionisation energy?

Answer 6

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. KJ/mol

Question 7

What are the three factors effecting ionisation energy?

Answer 7

Atomic radius, atomic charge, shell sheilding.

Question 8

What is the second ionisation energy?

Answer 8

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

Question 9

What can the successive ionisation energies tell us about an atom?

Answer 9

It provides evidence for different electron energy levels in atoms. It can tell us: the number of electrons in the outer shell, the group that the element is in, and the identity of an element.

Question 10

Which ionisation energy would be a big jump above the others in a Fluorine atom.

Answer 10

The 8th ionisation energy. Fluorine has 7 electrons in its outershell so the ionisation will greatly increase when the whole outer shell has been removed and the outer electron becomes closer to the nucleous and not sheilded.

Question 11

What is the general trend in ionisation energy accross a group and why?

Answer 11

The ionisation energy increases accross a group because the amount of protons increases so the nuclear charge increases holding the electron more firmly.

Question 12

Explain in detail the trend in first ionisation energy accross period 2.

Answer 12

The first ionisation energy increases accross the period because: the nuclear charge increases, the shell sheilding is similar, so the nuclear attraction increases and the atomic radius decreases. This means it is harder to remove the outer electron.

Question 13

Explain the trend in ionisation energy down a group.

Answer 13

Although the nuclear charge increase, its effect is outweighed by increased nuclear radius and to a less extent, shell sheilding. Ionisation energy decreases.

Question 14

Although the ionisation energy increases accross period two and three, there is a drop in the ionisation energy in the same position in each group. Explain it.

Answer 14

The fall in the ionisation energy from Beryllium to Boron. The 2p subshell has a higher energy then the 2s subshell so its the 2s electron is harder to remove.

Question 15

Explain the drop in ionisation energy between Nitrogen and oxygen.

Answer 15

The drop in ionisation energy marks the start of electron pairing in the 2p orbital. In oxygen, the paired electrons in the two p orbital repel each other making it easier for one of the electrons to be removed.

Question 16

Explain why metals can conduct electricity as a solid.

Answer 16

Metallic bonding means each metal atom donates its negative outer shell electron into a shared pool of delocalised electrons. The cations left behind consist of the nucleaus and the outershell electrons and our in fixed position. Delocalised electons carrry charge.

Question 17

Give some properties of metals.

Answer 17

The giant metallic lattice means:

• Metals have a high melting point (strongly bonded)
• Electrical conductivity
• Malluable (layers slide over each other)

Question 18

Give three atoms that form giant covelant structures.

Answer 18

Non metlas: Boron, Carbon and silicon. They are held together by a network of strong covelant bonds.

Question 19

What is the bond angle in diamond.

Answer 19

Carbon forms 4 covelant bonds so it is a tetrahedral structure with a bond angle of 109.5 degrees.

Question 20

Give some properties of giant covelant structures.

Answer 20

• High melting point (strong covelant bond)
• They are insoluable in almost all solvents. (covelant bonds are too strong to be broken by interaction of solvents)
• Non conductors (except graphene and graphite).

Question 21

What is the difference between graphene and graphite?

Answer 21

Graphene is a single layer of graphite composed of hexagonally arranged carbon atoms joined by three strong covelant bond. The layers in graphite are bonded by weak london forces. Graphene is the strongest and thinnest material ever made.

Question 22

Describe the melting point accross period 2. pg103

Answer 22

The melting point increases from group 1 to group 14 (Li-C). This is because they can form giant structures and the covelant bonds gradually get stronger (more electronegative). There is then a sharp drop from group 14 to 15. 15-18 are simple molecules with much lower melting points.

Question 23

Describe the melting point accross period 3.

Answer 23

The melting point increases from group 1 to 14. These are giant structures. The elting point is then significantly lower from group 15 onwards because these are si ple molecules only held together by london forces.