Reactivity of Group 1 Metals increases as there are more valence electron shells

ESF: Electrostatic Forces of Attraction

Electron Trends in the Periodic Table

Number of outer shell electrons is the same
Number of complete electron shells increases by 1
Number of a group is the same as the number of electrons in the outer shell of elements in that group, except for group 0
Shielding increases
Atomic radius increases

Number of outer shell electrons increases by 1
Number of complete electron shells stays the same
The point at which a new period starts is the point at which electrons begin to fill a new shell
Nuclear Charge increases across a period
Shielding is the same across a period
Nuclear Attraction increases across a period

Valency

Valency is the combining power of an atom
Equal to number of hydrogen atoms it can combine with or displace from a compound
- Valency of hydrogen is 1
Valency is not the same as the number of valence electrons
- e.g., Nitrogen has 5 valence electrons but its valency is 3

Elements	Na	Mg	Al	Si	P	S	Cl	Ar
Protons	11	12	13	14	15	16	17	18
Nuclear Charge	+11	+12	+13	+14	+15	+16	+17	+18
Electron Configuration	2,8,1	2,8,2	2,8,3	2,8,4	2,8,5	2,8,6	2,8,7	2,8,8
Shielding (inner shell electrons)	10	10	10	10	10	10	10	10
Nuclear Attraction	+1	+2	+3	+4	+5	+6	+7	+8
Atomic Radius	0.19	0.16	0.13	0.118	0.11	0.1	0.099	0.095

Atomic Radius: Distance from the outermost valence electron to the nucleus
Atomic Radius decreases across a period
- The outer electrons are in the same shell but the increasing net nuclear attraction pulls the outer electrons closer to the nucleus
Increases down a group
- Number of protons increases, however, number of shielding electrons increases too. Effective nuclear charge therefore remains approx. constant
- Increase in radius is due to higher principle energy levels being filed, and valency energy level is located further from the nucleus
Largest atomic radius is Caesium (Cs)

Electronegativity: Ability of an atom to attract bonding pair of electrons in a covalent bond
- In a covalent bond between 2 different elements, the electron density is not shared equally
  - $C l_{2}$ has a bonding pair of electrons, shared equally
  - $H Cl$ has a bonding pair of electrons not shared equally
  - Chlorine is more electronegative
    - Has a greater ability to attract a bonding pair of electrons not shared equally than carbon
Decreases down a group
- Down a group, there is an increase in atomic radius
- Leads to weaker ESF between the nucleus and the furthermost valence electron
  - Because shielding effect increases, i.e., repulsion from core electrons
Increases across a period
- Increase in number of protons
- Increase in positive charge in the nucleus
- Increased ESF between valence electrons and the nucleus
- Decrease in atomic radius, increase in ESF
Highest Electronegativity is Fluorine (Fl)

Energy required to remove an electron from a given atom
First Ionisation Energy (of an element): Energy required to remove 1 mole of electrons from 1 mole of gaseous atoms

$M (g) ⟶ M^{+} (g) + e^{-}$

Can use ionisation energy to find reactivity of metals
- When metals react, they lose electrons (ionisation energy)
- Group 1 only has first ionisation energy, as it only has 1 valence electron
Increases across a period*
- Stronger ESF between nucleus and valence electrons due to smaller atomic radius
- Greater number of protons results in greater force of attraction between nucleus and valence electrons
  - Shielding remains the same, proton number increases, nuclear charge increases
- Energy required to overcome this ESF increases, therefore ionisation energy increases
Decreases down a period
- Increasing atomic radius
- ESF between valence electron and nucleus is weaker
- Increase in shielding effect
  - Increase in number of energy levels
- Decrease in net ESF between valence electron and nucleus
- Energy required to overcome ESF decreases, therefore ionisation energy decreases

Electrons in a covalent bond are not always shared equally
Therefore, it relies on the electronegativity of the elements involved
Covalent bond that has an unequal sharing of electrons is called a polar covalent bond
Covalent bond that has an equal sharing of electrons is called a non-polar covalent bond
Example:

$^{δ -} F -^{δ +} B$

“δ±” is the dipole, so when it says “write a dipole across the polar bonds”, do this
Bonds with the greatest permanent dipole will have the greatest difference in their electronegativity(s)