Chapter 7 Periodicity

Question 1

What is periodicity?

Answer 1

A repeating pattern in either chemical or physical properties across different periods of the periodic table

Question 2

Why does the atomic radius decrease across a period?

Answer 2

There are more protons in the nucleus increasing the nuclear charge
This means that the electrons in the shells experience a stronger electrostatic force of attraction drawing them closer to the nucleus
The number of electron shells does not increase across a period so the radius decreases

Question 3

What changes has the periodic table undergone since mendeleevs version.

Answer 3

Mendeleev ordered them by atomic mass and left out gaps due to his correct predictions of the elements properties
Since the discovery of protons the table is now arranged by atomic number

Question 4

What is ionisation energy?

Answer 4

Measures how easily an atom loses electrons to form positive ions

Question 5

What three factors affect ionisation energy?

Answer 5

Atomic Radius
Nuclear charge
Electron shielding

Question 6

How does the atomic radius affect ionisation energy?

Answer 6

The greater the distance between the nucleus and the outer electrons, the less the nuclear attraction
As the force of attraction falls it becomes easier (requiring less energy) to form positive ions

Question 7

How does nuclear charge affect ionisation energy?

Answer 7

The more protons in the nucleus of an atom, the greater the attraction between the nucleus and outer electrons

Question 8

How does electron shielding affect ionisation energy?

Answer 8

Electrons are negatively charged and so inner shells repel outer shell electrons. This repulsion, called the shielding effect, reduces the attraction between the nucleus and outer electrons

Question 9

What is the first ionisation of Na?

Answer 9

The energy required to remove 1 electron from each atom in 1 mole of gaseous Na atoms to form 1 mole of gaseous 1+ Na ions written in kJ/mol

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

Question 10

Why is each successive ionisation energy of an atom greater than the last?

Answer 10

As each electron is removed, there is less repulsion between the electrons and each shell will be drawn slightly closer to the nucleus
As the distance of each electron from the nucleus slightly, the nuclear attraction increase as more energy is needed to remove each successive electron
When the ionisation of the atom/ ion overlaps with a change in shell there will be a greater increase in ionisation energy because the distance between the electron and nucleus has decreased more significantly

Question 11

What predictions can be made from successive ionisation energies?

Answer 11

The number of electrons in the outer shell
The group of the element in the periodic table
The identity of an element

Question 12

What do trends in first ionisation energy prove?

Answer 12

The provide evidence for the existence of shells and sub shells

Question 13

What is general pattern of first ionisation energies?

Answer 13

General increase across a period as atomic radius decreases e.g. H->He, Li->Ne, Na->Ar
A sharp decrease in first ionisation energy between the Ned of one period and the start of the next period
e.g. He->Li, Ne->Na

Question 14

What is the trend in first ionisation energy down a group?

Answer 14

They decrease down a group because the nuclear charge increase is outweighed by the increase in radius (and to a lesser extent the electron shielding)

Question 15

Why does the first ionisation energy increase from Lithium to Beryllium?

Answer 15

Li and Be outer electrons are in the same sub shell so no more electron shielding takes place, however Be has more protons and therefore a greater nuclear charge to attract electrons

Question 16

Why does the first ionisation energy decrease slightly from Beryllium to Boron?

Answer 16

B outer shell electron is further from the nucleus, than Be, in a 2p subshell which is in a higher energy subshell (than the Be 2s) making it easier to remove the outermost electron from the 2p in B than the 2s in Be

Question 17

Why does the first ionisation energy increase from Boron to Carbon and Nitrogen?

Answer 17

There are more protons in C and N than B so a more attractive nuclear charge, because electrons are added to the same sub shell so no increase in electron shielding takes place

Question 18

Why does the first ionisation decrease from Nitrogen to Oxygen?

Answer 18

This marks the start of electron pairing in the p-orbitals.
In N and O the highest energy electrons are in a 2p subshell
In O the paired electrons in one of the 2p orbitals repel one another, making it easier to remove an electron from an O atom than a N atom
Therefore the first ionisation of O is lower than N

Question 19

What is metallic bonding?

Answer 19

Strong electrostatic attraction between actions and delocalised electrons

Question 20

What is the structure of a metallic compound?

Answer 20

The actions are fixed in position maintaining the structure and shape of the metal. This is in a regular arrangement
The delocalised electrons are mobile and can move throughout the structure

Question 21

What are common properties of metals?

Answer 21

Strong metallic bonds
High electrical conductivity
High mp/bp

Question 22

How do metals conduct electricity?

Answer 22

The conduct electricity in both solid and liquid states because when a voltage is applied to the metal the delocalised electrons can move through structure carrying a charge

Question 23

Are metals soluble?

Answer 23

No, not even in polar solvents because any interactions would lead to a reaction

Question 24

What effects do periodicity have on metals?

Answer 24

As you go across the period there are more delocalised electrons creating a greater force of attraction
As the number of delocalised electrons increase the electrical conductivity increases

Question 25

What are the properties of giant covalent structures?

Answer 25

High mp/bp
Cannot conduct electricity expect for graphene and graphite)
Nearly always insoluble

Question 26

Why do giant covalent lattices have a high mp/be?

Answer 26

Covalent bonds are strong

High temperatures are necessary to provide the large quantity of energy needed to break the strong covalent bonds

Question 27

Why are most giant covalent structures insoluble?

Answer 27

Covalent bonds are far too strong to be broken by interactions with solvents

Question 28

Why can’t most giant covalent lattices conduct electricity?

Answer 28

In carbon (diamond) and silicon, all four outer electrons are involved in covalent bonding, so none are available to be mobile charge carriers

Question 29

What is diamond and what is it used for?

Answer 29

Diamond is a giant covalent lattice of carbon with a tetrahedral structure because it forms four covalent bonds with neighbouring carbon atoms.
This means that it is extremely hard and is used for drill bits and jewellery
Because all outershell electrons are involved in bonding it cannot conduct electricity

Question 30

What is graphite and what is it used for?

Answer 30

Graphite is a giant covalent lattice of carbon with a hexagonal structure because it forms 3 covalent bonds with neighbouring carbon atoms
Carbons final outershell electron is not involved in bonding so can carry a charge.
Graphite is formed of 2D layers held together by weak IM forces (London forces)
Often used in pencils as lates can break and slide off of eachother

Question 31

What is graphene?

Answer 31

It’s a single layer of graphite and is a 2D covalent lattice
It’s the thinnest and strongest material ever made
It has a similar electrical conductivity to copper

Question 32

What trends in melting points are there across periods 2 and 3? (These can be applied to all periods)

Answer 32

The melting point increase from group 1 to 14 (4) because the giant metallic structures become stronger due to the increase difference in charge between the actions and high quantities of delocalised electrons. These are giant structures
There is a sharp decrease from group 14 (4) to 15 (5) because the compounds change from giant structures to simple molecules
The melting points are comparatively low from group 15 (5) to 18 (8)