Periodicity

Question 1

How did Mendeleev arrange the periodic table?

Answer 1

By atomic mass

Question 2

How is the periodic table arranged now?

Answer 2

By atomic/proton number

Question 3

What are the horizontal rows of the periodic table called?

Answer 3

Periods

Question 4

What does the number of the period show?

Answer 4

The number of the highest energy electron shell in an elements atoms

Question 5

What are the 4 blocks on the periodic table?

Answer 5

s, p, d and f

Question 6

Where is the s-block?

Answer 6

Left hand side of the table - groups 1 and 2

Question 7

Where is the p-block?

Answer 7

Right hand side of the table - groups 3, 4, 5, 6, 7 and 0

Question 8

Where is the d-block?

Answer 8

In the middle - the transition elements

Question 9

Where is the f-block?

Answer 9

Right at the bottom - the high mass new elements

Question 10

What is first ionisation energy?

Answer 10

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

Question 11

How is sodiums first ionisation energy written?

Answer 11

Na(g) -> Na +(g) + e-

Question 12

What factors affect ionisation energy?

Answer 12

• Atomic radius
• Nuclear charge
• Electron shielding

Question 13

How does atomic radius affect ionisation energy?

Answer 13

The greater the distance between the nucleus and the outer electrons, the less the nuclear attraction. Therefore, the lower the ionisation energy the greater the atomic radius

Question 14

How does nuclear charge affect ionisation energy?

Answer 14

The more protons there are in the nucleus of an atom, the greater the attraction between the nucleus and the outer electrons. Therefore, the greater the ionisation energy the greater the nuclear charge

Question 15

How does electron shielding affect ionisation energy?

Answer 15

Electrons are negatively charged and so inner-shell electrons repel outer-shell electrons. This repulsion, called the shielding effect, reduces the attraction between the nucleus and the outer electrons. So the greater the electron shielding, the lower the ionisation energy

Question 16

Why may a successive ionisation energy be greater than the previous?

Answer 16

Because there are less electrons to the same charge of the nucleus - so there is a greater nuclear attraction on the remaining electrons so the ionisation energy increases

Question 17

What is second ionisation energy?

Answer 17

The energy required to remove one electrons from one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions

Question 18

What would cause a large jump in successive ionisation energies?

Answer 18

The successive ionisation energy being removing an electron from a different/new shell. The atomic radius is smaller and the electron shielding is less so the forces of attraction will be greater and the ionisation energy will be greater

Question 19

What happens to first ionisation energy down a group?

Answer 19

• Atomic radius increases
• More inner shells so shielding increases
• Nuclear attraction on outer electrons decreases
• First ionisation energy decreases

Question 20

What happens to first ionisation energy across a period?

Answer 20

• Nuclear charge increases
• Same shell so similar shielding
• Nuclear attraction increases
• Atomic radius decreases
• First ionisation energy increases

Question 21

Why is there a fall in ionisation energy between beryllium and boron?

Answer 21

Marks the start of the filling of the 2p sub-shell

In boron, the 2p electron is easier to remove than one of the 2s electrons in beryllium. Therefore, the first ionisation energy of boron is less than the first ionisation energy of beryllium

Question 22

Why is there a fall in ionisation energy between nitrogen and oxygen?

Answer 22

In nitrogen and oxygen the highest energy electrons are in a 2p-subshell

In oxygen, the paired electrons in one of the 2p sub-shells repel each one another, making it easier to remove an electron from an oxygen atom than a nitrogen atom. Nitrogen has no paired electrons

Question 23

What are all metals at room temperature?

Answer 23

Solids - except mercury which is a liquid

Question 24

What is metallic bonding?

Answer 24

The strong electrostatic attraction between cations and delocalised electrons

Question 25

How does metallic bonding work?

Answer 25

In a solid metal structure, each atom has donated its negative outer-shell electrons to a shared pool of electrons, which are delocalised throughout the whole structure. The positive ions (cations) left behind consist of the nucleus and the inner electron shells of the metal ions. The cations are fixed in position, maintaining the structure and shape of the metal. The delocalised electrons are mobile and able to move throughout the structure.

Question 26

How are atoms held together in a metal structure?

Answer 26

In a giant metallic lattice

Question 27

What are the properties of metals?

Answer 27

• Strong metallic bonds
• High electrical conductivity
• High melting and boiling points

Question 28

Why can metals conduct electricity?

Answer 28

Because they have delocalised electrons than can move through the structure as mobile charge carriers

Question 29

Why do metals have high melting and boiling points?

Answer 29

High temperatures are necessary to overcome the strong electrostatic attraction between the cations and electrons to get the metal melting/boiling

Question 30

Are metals soluble?

Answer 30

No

Question 31

What are the structures of non-metals boron, carbon and silicon?

Answer 31

Giant covalent lattice - many billions of atoms are held together by a network of strong covalent bonds

Question 32

What is a giant covalent lattice?

Answer 32

A structure where each carbon/silicon/boron atom is bonded to 4 others with bond angles of 109.5 degrees

Question 33

What are the melting and boiling points like for giant covalent lattices?

Answer 33

High. The high temperatures are necessary to provide the large quantity of energy needed to break the covalent bonds

Question 34

What is the solubility of giant covalent lattices like?

Answer 34

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

Question 35

What is the electrical conductivity like for giant covalent lattices?

Answer 35

Non-conductors of electricity (expect graphite and graphene). Typically, all 4 outer shell electrons are involved in covalent bonding so there are no free charge carriers

Question 36

What is graphite?

Answer 36

Composed of parallel layers of hexagonally arranged carbon atoms. The layers are bonded by weak London forces. The bonding in the hexagonal layers only uses 3 of the carbons 4 outer-shell electrons. The spare electron is delocalised between the layers, so electricity can be conducted as in metals.

Question 37

What is graphene?

Answer 37

A single layer of graphite, composed of hexagonally arranged carbon atoms, linked by strong covalent bonds. It has the same electrical conductivity as copper. It is the thinnest and strongest material ever made

Question 38

What is the trend in melting points across Periods 2 and 3?

Answer 38

• The melting point increases from Group 1 to 4
• There is a sharp decreases in melting point between Group 4 and 5
• The melting points are comparatively low from group 5 to group 0

Question 39

Why is there a sharp decrease in melting points between Groups 4 and 5 for Periods 2 and 3?

Answer 39

Because it shows how the structures go from giant to simple molecules

Question 40

What are the structure of the Period 2 elements?

Answer 40

Ni and Be = Giant metallic structures
B and C = Giant covalent structures
N2, O2, F2 and Ne = simple molecular structures

Question 41

What are the structures of the Period 3 elements?

Answer 41

Na, Mg and Al = Giant metallic structures
Si = giant covalent structure
P4, S8, Cl2 and Ar = simple molecular structure