Periodicity (copy)

First ionisation energy definition

energy required to remove one mole of electrons from one mole of atoms of an element in the gaseous state to form one mole of gaseous ions

First ionisation energy equation

E (g) —> E+ (g) + e-

First ionisation energy _____ across a period

increases

First ionisation energy ______ down a group

deacreases

What 3 factors influence ionisation energy?

1- Atomic radii

2- Nuclear charge

3- Electron shielding

How does atomic radii affect first ionisation energy?

Electrons in shells further from nucleus are less attracted to nucleus.

∴ the further the outer electron shell is from nucleus, the lower the ionisation energy

How does nuclear charge affect first ionisation energy?

Nuclear charge increases w. increasing atomic n. means there are greater attractive forces between the nucleus and outer electrons

∴ more energy is required to overcome the attractive forces when removing an electron

How does electron shielding effect first ionisation energy?

shielding effect- when electrons in full inner shells repel electrons in outer shells preventing them to feel the full nuclear charge

∴ greater shielding of outer electrons by inner electron shells = lower ionisation E

Why does ionisation energy decrease down a group despite increase nuclear charge?

• Atomic radius increases

• Shielding (by inner shell electrons) increases

• ∴ attraction between nucleus + outer electrons decreases

Why does ionisation energy across a period increase?

• Across nuclear charge increases

• Distance between nucleus and outer electrons remains reasonably constant no significant change in atomic radii

• Shielding by inner shell electrons remains the same

Why is there a rapid decrease in ionisation E between last element in one period and 1st element in next period

• Increased distance between nucleus and outer electrons

• Increased shielding by inner electrons

• Which outweigh increased nuclear charge

Why is there a slight decrease in first ionisation energy between Beryllium and boron?

In boron- 5th electron in 2p subshell (higher energy) - further from nucleus than 2s subshell of beryllium therefore has less attraction to nucleus and requires less energy to remove

Why is there a slight decrease in first ionisation energy between N and O?

Oxygen- paired electrons in 2p subshell - repel each other - makes it easier to remove an electron in O than N

Across a period and down a group first ionisation factors- table

Successive ionisation energies of an element ______

Increase

Why do successive ionisation energies increase?

Removing an electron from a pos ion is more difficult than from a neutral atom

cuz as electrons are removed- attractive force increase- cuz of decreasing shielding + increase in proton to electron ratio

The big jumps on the graph show the change of ____ and the small jumps are the change of _____

shell

sub-shell

Diamond

• Giant cov lattice of carbon atoms

• Each carbon is cov bonded to 4 others

• Tetrahedral arrangement- bond angle 109.5

• Giant lattice structure- strong bonds in all directions

• Hardest known structure

Diamond uses

drills and glass-cutting tools

Graphite

• Each carbon bonded to three others in layered structure

• Layers made of hexagons- bond angle 120

• Spare electrons are delocalised ∴ occupy the space between layers

• Same layer atoms held by strong cov bonds

• Layers held by weak IMF allow layers to slide over each other

Graphene

• Infinite lattice of cov bonded atoms in 2D only to form layers

• Made of single layer of carbon atoms that are bonded in repeating pattern of hexagons

Silicon(IV) oxide

• Also known as silicon dioxide

• Same structure as diamond- giant cov lattice/ macromolecular structure made of tetrahedral units all bonded by strong cov bonds

Empirical formula of Silicon(IV) oxide

SiO2

Solubility of metallic substances

• Metals do not dissolve

• Interaction between polar solvents and charges in the metallic lattice but these lead to reactions, rather than dissolving

M.p. across period 2 and period 3

M.p increase group 1-4:

• Group 1-3 have metallic bonding which increases in strength due to increased forces of attraction between more electrons in outer shell that are released to the sea of electrons and smaller pos ion

• Group 4 has giant cov structure w. many strong cov bonds requiring a lot of E to overcome

  Sharp increase in m.p from group 4 to 5:

• Group 5 to 0 have simple molecular structure w. weak London forces between molecules requiring little energy to overcome

periodicity

• Pattern in the change in the properties of a row of elements
  OR
  Trend in the properties of elements across a period

• Repeated in the next row

Why does He have largest first ionisation energy?

cuz of small distance ionisation energy increases, higher nuclear charge but no shielding ∴ takes more E to remove outer electron

Electronegativities

Ability of atom to attract the pair of electrons in a cov bond

what is m.p a measure of?

Change of state that occurs between solid and liquid. Idea of strength of bond holding it together

If charge in ionic bond increases what happens?

More attraction therefore stronger bonding

What has the highest m.p of period 3?

• Silicon is a metalloid with a giant covalent structure held together by strong covalent bonds which require a lot of energy to break.

• Silicon is the only element in period 3 that is a giant covalent structure and so it has the highest melting point at 1.414^oC

Why is the melting point of sulfur (S₈) is greater than that of phosphorus (P₄)

• S₈ molecules are bigger than P₄ molecules

• Sulfur and phosphorus both have weak London forces between molecules

• The strength of London forces depends on the number of electrons the molecule has

• Hence larger molecules will experience stronger London forces 

Similarities between graphite and copper

They both have layers / planes / sheets of atoms or ions that can slide over one another

What group has highest m.p

group 4

The difference in melting points between Na and Mg is due to:

• Magnesium ions have a greater charge

• Magnesium has more delocalised/outer electrons

• Magnesium has greater attraction between ions and electrons
  OR
  Magnesium has stronger metallic bonds