Topic 3

What do rows represent?

What do rows represent?

The number of valence electrons

What do periods represent?

The outer energy level occupied by electrons

What is atomic radius?

The distance between the nucleus and the outermost electron shell

Explain the atomic radius trends

• The larger the nuclear charge, the greater the attraction, the smaller the atom

• increases down a group

  • increase in shells = more shielding = weaker nuclear attraction = larger atom

• Decreases as atomic number increases across period

  • more positive nuclear charge, and electrons are added to the same level

Explain the ionic radius trend

• increases with negative charge across period

• Increases down a group

  • more shells

Define ionisation energy

The energy required to remove 1 mol of electrons from 1 mol of an atom to make 1 mol of a gaseous ions

Define first ionization energy

the energy required to remove 1 mol of electrons from 1 mol of gaseous atoms

Explain the ionization energy trends

• Decreases with paired electrons

  • pairs repel = easier to remove

• decreases as shielding increases

  • more shielding energy = less attraction = less energy

• Decreases as distance increases

  • further from nucleus = less attraction = lower energy

• increases as nuclear charge increases with atomic number

  • greater force of attraction = more energy required to overcome

Why does ionization energy decrease between periods?

The addition of a new energy level results in more shielding and increased radius

Why is there a decrease in ionization energy between Be and B?

B looses from P subshell, which is further away from the nucleus and easier to remove than S.

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

O has paired electrons in P subshell, which are easier to remove.

Why does successive ionization energy increase?

Less shielding = harder to remove

Define electron affinity

energy released when 1 mol of electrons is gained by 1 mol of atoms to form 1 mol of gaseous ion

Why is second electron affinity endothermic?

Due to having to overcome repulsive forces between negative ion and electron

How can you describe the first electron affinity?

exothermic

Why does electron affinity decrease down a group?

the larger the atom, the lesser the attraction

Why does Fl have a very small electron affinity?

it is very small, and adding electrons experiences much repulsion

Define electronegativity

the ability to attract electrons to itself in a covalent bond

Explain electronegativity trends

• increase in nuclear charge = increase in electronegativity

• increase in atomic radius = less attraction = decrease in electronegativity

• increases across a period

  • nuclear charge increases

• decreases down a group

  • increased shielding

Define amphoteric

A substance that can act as both an acid and a base

What oxides are amphoteric?

Aluminum Al2O3

What oxides are basic?

• Na2O

• MgO

What oxides are acidic?

• SO2

• SO3

• P4O10

• SiO2

Explain group 1 trends

• reactivity increase down the group

• MP and BP decrease down the group

  • atomic radius increase = weaker metallic bond