3.1.1 Periodicity

Question 1

How is the periodic table structured?

Answer 1

• Arrangement of elements
• Organised by increasing atomic number
• In periods which show repeating trends in physical and chemical properties
• In groups which have similar chemical properties

Question 2

What are the blocks in the periodic table?

Answer 2

Group 1 and 2 = s-block
Group 3 to 0 = p-block
Transition metals = d-block
Lanthanides and actinides = f-block

Question 3

How are elements in the same block similar?

Answer 3

Elements in the same block have their highest electron in the same sub-shell

Question 4

What is the first ionisation energy?

Answer 4

The energy required for the removal of 1mol of electrons from one mol of gaseous atoms (kJ mol-1)

Question 5

What is the ionic equation for first ionisation energy?

Answer 5

X (g) -> X+ (g) + e-

Question 6

How does shielding affect ionisation energy?

Answer 6

• Inner shells of electrons repel the outer shells
• More shells means more shielding
• This means weaker attraction between the nucleus and the outer electrons
• Less energy is required

Question 7

How does nuclear charge affect ionisation energy?

Answer 7

• More protons means a larger nuclear charge
• There is a greater attractive force between nucleus and outer electrons
• More energy is required

Question 8

How does atomic radius affect ionisation energy?

Answer 8

• Bigger atom means a greater atomic radius
• Smaller attractive force between nucleus and outer electrons
• Easier to remove electrons

Question 9

How does ionisation energy change as you go down a group?

Answer 9

• Decreases as you go down a group
• Atomic radius increases so the outer electrons are further from the nucleus and the attractive force is weaker
• Shielding increases as there are more shells between the nucleus and outer shell electrons so attractive force is weaker
• Energy required to remove an electron decreases

Question 10

How does ionisation energy change as you go across period?

Answer 10

• Generally increases across a period
• There are dips between the 2nd and 3rd element and the 5th and 6th element in period 2 and 3

Question 11

How does ionisation energy change across period 1?

Answer 11

• Hydrogen has a high first IE as its electron is closest to the nucleus and there is no shielding
• Helium has a higher value as it has an extra proton in the nucleus so there is a stronger attraction

Question 12

How does ionisation energy change across period 2?

Answer 12

• Lithium has shielding from 1s electrons and an increase in atomic radius which cancels out the increase in nuclear charge so IE is lower than helium
• Beryllium has a higher IE due to the increased nuclear charge and no extra shielding (still in 2s)
• Boron has a drop in value as the electron has moved into a 2p orbital so the increased shielding makes the electron easier to remove
• Carbon and nitrogen have a higher IE due to increased nuclear charge
• Oxygen has a drop in IE due to the extra electron being paired in a 2p orbital so the repulsive force means less energy is required to remove the electron
• Fluorine and neon have a higher IE due to increased nuclear charge and no extra shielding

Question 13

How does ionisation energy change across period 3?

Answer 13

• Substantial drop to sodium due to extra shielding
• Similar pattern to period 2

Question 14

What are successive ionisation energies?

Answer 14

The removal of more than one electron from the same atom

Question 15

How do successive ionisation energies change?

Answer 15

• Increase as atomic radius decreases
• A large jump in energies indicates which group an element is in as a certain number of electrons can be lost easily, then it becomes difficult
• Atomic radius decreases as the electrostatic attraction between the nucleus and outer electrons increases when electrons are lost

Question 16

What is metallic bonding?

Answer 16

The strong electrostatic attraction between cations and delocalised electrons

Question 17

What structure to metals have?

Answer 17

Giant metallic lattice

Question 18

What are the properties of metals?

Answer 18

• Good electrical and thermal conductors as delocalised electrons can move freely
• Malleable as the layers can easily slide over each due to metal ions all being the same size
• High melting and boiling points as lots of energy is required to break the strong electrostatic forces
• Will have a higher melting point if it can donate more electrons to the delocalised system
• Insoluble as the metallic bond is too strong to be broken

Question 19

What is the structure of a giant covalent substance?

Answer 19

Giant covalent lattices which are a network of atoms bonded by strong covalent bonds

E.g. carbon, silicon

Question 20

What are the properties of diamond?

Answer 20

• Each carbon is bonded 4 times in a tetrahedral shape
• Very high melting point due to the strong covalent bonds which require lots of energy to break
• Very hard and insoluble as the covalent bonds are too strong to break
• Doesn’t conduct electricity as there are no delocalised electrons
• Can be cut or make gemstone or used in drills and glass cutters

Question 21

What are the properties of graphite?

Answer 21

• Each carbon is bonded 3 times with the 4th electron being delocalised
• Very high melting point due to the strong covalent bonds
• Insoluble as the covalent bonds are too strong to break
• Soft as the layers can slide easily due to the weak forces between them
• Conducts electricity as the delocalised electrons between the layers are free to move and can carry a charge

Question 22

What are the properties of graphene?

Answer 22

• One layer of graphite
  -1 atom thick and made of hexagonal carbon rings
• Good electrical conductor as it has delocalised electrons
• Lightweight, transparent and flexible
• Used in aircraft shells, super computers, high speed computing and smart phone screens

Question 23

What are the properties of silicon (IV) oxide?

Answer 23

• Adopts same structure as diamond
• Giant covalent lattice made of tetrahedral units bonded by covalent bonds
• Empirical formula is Si2O

Question 24

How does melting point change across a period?

Answer 24

• Melting point increases from group 1 to group 4 as they are all giant structures
• Groups 1 to 3 have giant metallic structures
• The metallic bonds get stronger as the metal ions have an increasing positive charge, more delocalised electrons and a smaller atomic radius so more energy is required
• Silicon and carbon have the highest melting point as they both have giant covalent structures
• There is a sharp decrease between group 4 and 5 which marks the change from giant to simple molecular structures
• Less energy is required to overcome weak intermolecular forces
• The larger the simple molecular structure, the larger the intermolecular forces so more energy is required (e.g. sulfur)
• Group 0 exist as individual atoms so have the lowest melting point due to the weakest intermolecular forces