Chaptwe 7 - Trends in the periodic table

atomic radius

Half the distance between the nuclei of two atoms of the same element that are joined together by a single covalent bond

Are groups horizontal or vertical?

vertical

Are periods vertical or horizontal?

horizontal

Why do the values of the atomic radius increase down the groups?

1. There is an extra energy level which is further away from the nucleus

2. Screening effect of inner electrons: Inner electrons in an atom block or shield the outer electrons from the pull of electrons

Why does atomic radius decrease across periods

1. Increase in effective nucleur charge: As the nucleur charge increases (number of protons) the atomic radius tends to decrease because the stronger pull from the nucleus draws the electrons closer

2. No increase in the screening effect of electrons

First ionisation energy definition

The first ionisation energy of an atom is the minimum energy required to completely remove the most loosely bound electron from a neutral gaseous atom in its ground state

Equation

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

Units for energy

kj mol -1

Ionisation energy down a group

Decreases

Bigger atoms/ atomic radius and more shielding so weaker attraction between nucleus and outer electrons

- Screening effect of inner electrons

Ionisation energy across a period

increases j

- Increasing effective nuclear charge or atomic number

- Decreasing atomic radius

exceptions to ionization energy trend

Any sublevel that is exactly have filled or fully filled has more stability therefore slightly higher ionisation

exceptions to ionization energy trend - examples

Beriliyium (BE): 1s2, 2s2 9fully filled)

Magnesium (Mg): 1s2, 2s2, 2p6, 3s2

Nitrogen: 1s2, 2s2, 2p3 (Half full)

Phosphorus: 1s2, 2s2, 2p6, 3s2, 2p3 (half filled)

Second ionisation energy

The energy required to remove an electron from an ion with one positive charge in the gaseous state

X+ (g) -> X2+ (g) + e-

Third ionisation energy equation

X2+(g) -> X3+(g) + e-

What do successive ionisation energy values provide

Successive ionisation energy values provide further evidence for the number of electrons in an atom

What happens for each successive ionisation energy level

For each successive ionisation energy level (electron removed) there is an increase in ionisation energy. i.e. the 2nd ionisation value is higher than the 1st​

reason: What happens for each successive ionisation energy level

For each electron removed then effect of the positive charge is increased meaning a greater attraction between the nucleus and remaining electrons

Small peaks

ISmall peaks in ionisation energy correspond to removing an electron from a full or half full sublevel.

Large jumps/small drops

Large ionisation jumps or drops in ionisation energy correspond to removing an electron from a change in energy level.

What does group 1 have

Group 1 has extremely low ionisation due to one electron

What does group 8 have

Group 8 has very high ionisation due to high stability

Why does the value of electronegativity decrease down the groups in the periodic table

1. Increasing atomic radius

2. Screening effect of inner electrons

Why does the value of electronegativity increase down the groups in the periodic table

1. Increasing effective nuclear charge

2. Decreasing atomic radius

Trends within alkali metals

- All very reactive as they have low first ionisation value/energy

- Reactivity increases down the group because atomic radius increases

Alkali metals: with oxygen

Potassium + Oxygen = Potassium Oxide

Reaction with water

Sodium + water = Sodium Hydroxide +Hydrogen

Trends with Hallogens

Very reactive non metals due to them needing only one more electrron to be stable

Reactivity decreases down the group due to electronegativity decreasing