3.1.1 Periodicity

Definitions

Term	Definition
Periodicity	A repeating trend in physical and chemical properties of the elements across the periodic table.
Groups	A vertical column in the periodic table. Elements in a group have similar chemical properties and their atoms have the same number of outer shell electrons.
Periods	A horizontal row in the periodic table. Elements show trends in properties across a period.
Shielding effect	The repulsion between electrons in different inner shells. Shielding reduces the net attractive force between the positive nucleus and the outer shell electrons.
Metallic bonding	The strong electrostatic attraction between the regularly arranged metal cations and the delocalised valence electrons between them.
Delocalised electrons	Electrons shared between more than two atoms / ions.
Giant metallic lattice	A three dimensional structure of positive ions and delocalised electrons, bonded together by strong metallic bonds.
Giant covalent lattice	A three dimensional structure of atoms, bonded together by strong covalent bonds.

The periodic table

History

Then
- Mendeleev arranged the elements in order of atomic mass
- Swapped elements to arrange them into groups of similar properties
- Gaps left where he thought elements would be found
- Predicted properties for missing elements
- Newly discovered elements filled in the gap and matched the predicted properties
Now
- Arranged in increasing atomic number
- In vertical columns (groups) with same number of outer electrons + similar properties and horizontal rows (periods) giving number of highest energy electron shell

Arrangement

In the order of increasing atomic number
Periodicity: in periods showing repeating trends in physical and chemical properties e.g. metals \(\rightarrow\) non-metals
In groups with similar properties

Electron configuration pattern

Across period
- Each period starts with an electron in a new highest energy shell
- Period 2: \(2s\) fills \(\rightarrow\) \(2p\) fills
- Period 3: \(3s\) fills \(\rightarrow\) \(3p\) fills
- Period 4: only \(4s\) and \(4d\) occupied in shell

Blocks

s/p/d/f-block meaning: the highest energy electron is in a s/p/d/f-orbital
S, p, d and f block

Name of groups

Group number	Name
1	Alkali metal
2	Alkaline earth metals
3-12	Transition elements
15	Pnictogens
16	Chalcogens
17	Halogens
18	Noble gases

Ionisation energy

First ionisation energy

Energy required to remove one electron from each atom in one mole of gaseous atoms of an element, forming one mole of gaseous 1+ ions
Unit = \(kJ \cdot mol^{-1}\)
Equation: \(X(g) \rightarrow X^+(g) + e^-\)

Factors affecting ionisation energy

Atomic radius
- Greater distance between nucleus and outer electrons = less nuclear attraction
- Large effect on ionisation energy as force of attraction falls sharply with increasing distance
Nuclear charge (weakest effect, outweighed)
- More protons in nucleus (greater nuclear charge) = greater attraction between the nucleus and the outer electrons = increase in ionisation energy
Electron shielding
- Shielding effect: electrons are negatively charged so inner shell electrons repel outer-shell electrons
- Reduces the attraction between nucleus and outer electrons \(\rightarrow\) reduce ionisation energy

First ionisation energy trends across a period

Increases across a period
- Nuclear charge increases
- Same number of shells so similar shielding
- Atomic radius decreases
- Nuclear attraction increases \(\rightarrow\) first ionisation energy increases
Falls when the p sub-shell is starting to be filled (e.g. \(Li \rightarrow Be\))
- \(2p / 3p\) sub-shell has a higher energy than \(2s / 3s\) sub-shell so the electron is easier to remove
- Still larger than IE before the decrease
Falls when pairing of electrons in p sub-shell starts (e.g. \(N \rightarrow O\))
- Paired electrons in one of the p orbitals repel one another so it is easier to remove an electron from the atom
- Still larger than IE before the decrease

First ionisation energy trend down a group

Decrease down a group
Atomic radius increases
More inner shells so shielding increases
Increase in atomic radius and shielding outweighs the increasing nuclear charge
Nuclear attraction on outer electrons decreases \(\rightarrow\) first ionisation energy decreases

Successive ionisation energy pattern

Equation: e.g. \(Mg^+(g) \rightarrow Mg^{2+}(g) + e^-\)
Larger than the previous one
- After the first electron is lost the remaining electrons are pulled closer to the nucleus
- Nuclear attraction on the remaining electrons increases so more energy needed
Large increase when shell change
- Shell closer to the nucleus as atomic radius drops
- Less shielding present as there are now less inner shells between electron & nucleus
- Stronger nuclear attraction so more energy needed
- Going to a more inner shell = extremely large increase
Smaller but still large increase when going to a new sub-shell / sub-shell become half filled
Can be used to work out the number of electrons in each shell + group number of the element

Metallic bond

Metallic bonding structure

Regularly arranged metal cations sitting in a sea of delocalised electrons
Each atom donate its negative outer shell electrons to a shared pool of electrons which are delocalised throughout the whole structure
Cations left behind = nucleus + inner shell electrons
Cations are fixed in position
Delocalised electrons are mobile and free to move throughout the structure
Forms a giant metallic lattice

Properties of metals

All conduct electricity
- Delocalised electrons can move through the structure and carry charge through the structure when a voltage is applied across a metal
- More delocalised electrons \(\rightarrow\) more electrons can move \(\rightarrow\) better conductivity
- Conducts electricity both in solid state and when molten
Most have high melting and boiling points
- Depends on the strength of metallic bonds
  - Greater cation charge = stronger attractive forces as more electrons are delocalised and forces between electrons + cations are stronger
  - Larger ions = weaker attractive forces due to larger atomic radius decreasing the charge density
- High temperature needed to provide the large amount of energy to overcome strong electrostatic attraction between the cations and the electrons
- Melting and boiling points decrease down the group
Dissolve in liquid metals only
- Similar force between particles
- Any interaction between polar / non-polar solvent + solute lead to a reaction rather than dissolving
  - Forces between particles are too large so it is not energetically favourable for them to mix

Giant covalent structures

Boron, carbon allotropes, and silicon (\(Si\), \(SiO_2\), \(SiC\))
A network of atoms bonded by strong covalent bonds to form a giant covalent lattice

Diamond / silicon

4 outer shell electrons of each atom form 4 covalent bonds with other carbon / silicon atoms
Tetrahedral structure
109.5° bond angle due to electron-pair repulsion
High melting and boiling points
- Atoms held together by strong covalent bonds
- High temperatures are needed to provide the large quantity of energy needed to break the strong covalent bonds
Non-conductors of electricity
- All 4 outer-shell electrons involved in covalent bond so no charged particles or mobile ions are available for conducting electricity

Graphite

Flat 2D sheets of hexagonally arranged carbon atoms (trigonal planar 120°)
Layers bonded by weak London forces \(\rightarrow\) they can slide over each other easily so graphite is soft
High melting and boiling points
- Atoms held together by strong covalent bonds
- High temperatures are needed to provide the large quantity of energy needed to break the strong covalent bonds
Can conduct electricity
- One electron from each carbon atom is delocalised and is available for conductivity

Graphene

Single layer of graphite
Hexagonally arranged (trigonal planar 120°) carbons
Very hard as there are no points of weakness in the structure
One of the thinnest + strongest material in existence (atoms held together by strong covalent bond)
High melting and boiling points
- Atoms held together by strong covalent bonds
- High temperatures are needed to provide the large quantity of energy needed to break the strong covalent bonds
Can conduct electricity
- One electron from each carbon atom is delocalised
- They can move and conduct electricity

Applications of graphene

Electronics
- Flexible displays
- Wearables
- Other next-generation electronic devices

Periodic trends in properties

Atomic radii trend across a period

Atomic radii decreases across the period
Positive charge in nucleus and negative charge in the outer shell both increases
Shielding remains similar as the number of shells doesn't change
The attraction between the nucleus and the outer electrons increases

Melting / boiling point trend across a period

Increases from Group 1 to 14
Sharp decrease between Group 14 to 15 - change from giant to simple molecular structures
Comparatively low from Group 15 to 18
- The exact boiling points depend on the type of covalent bonding
- Giant covalent bonding = very high melting and boiling points
- Simple covalent bonding = depends on strength of intermolecular forces (London forces) which depends on the mass of the nucleus
- Smaller molecular radius = lower boiling points (hence Group 18 has the lowest boiling points)