Module 3: chapter 7 periodicity

How did Mendeleev create the periodic table

• Arranged 60 known elements in order of atomic mass

• Lined up elements with similar properties.

• Swapped any elements that don't the properties ( e.g. Te, I)

• Left gaps for undiscovered elements + predicted properties of them

How is the periodic table arranged now

• • 114 elements arranged in order of atomic number

• Groups are columns going down (no. Of e- in outer shell)

• Periods are rows that go across (no. Of e- shells)

What is periodicity

Across each period there is a repeating trend in properties

Examples of periodicity

Metals → non metals, election configuration, ionisation energy, structure, melting point,

What is the trend in electron configuration across a period

Each period starts with an electron in a new higher energy shell

What is the trend in electron configuration across period 2

Fill s subshell then P subshell

What is the trend in electron configuration across period 3

Fill s subshell then p subshell

What is the trend in electron configuration across period 4

Fill s subshell then 3D subshell then P subshell

How is the periodic table divided into blocks

Matching their highest energy level ( s PDF)

What is ionisation energy

Measures how easily an atom loses elections to form a positive ion

First ionisation energy

The energy required to remove one election from each atom in one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions

What are the factors affecting ionisation energy

1. Atomic radius

2. Nuclear charge (number of protons)

3. Electron shielding

Now many ionisation energies will an element have

As many as there are electrons

Do successive ionisation energies increase or decrease

Increase

What do successive ionisation energies provide evidence about

For the different election energy levels. Can make predictions about the number of electrons in the outer shell the group of the element and the identity of the element

Graph for trends in 1st ionisation energies

Structure of metals

Layers of positively charged ions in a sea of delocalised electrons

What are the properties of metals

• Conductors of electricity

• Malleable

• High melting point

• Insoluble in water → electrostatic attraction too strong for water to break

Why does magnesium have a higher melting point than sodium

Mg has a 2 + charge so attraction between ion and election is stronger and requires more energy to break

Which elements are giant covalent lattices in elemental form

Boron, Carbon, silicon

What are the properties of giant covalent structures

• High melting and boiling point

• Insoluble in water (bonds too strong for water to break)

• Generally don't conduct electricity except graphene / graphite

What are some carbon allotropes

Diamonds → each carbon is bonded to 4 others ( tetrahedral )

Graphite → each carbon only bonded to 3 others electron delocalised between layers hexagonal rings

What does the trend in melting points for period 2 look like

Increases from lithium to Carbon as giant covalent and increase in charge

Then drops as simple molecules so weaker bonds and monatomic so only breaking London forces

What does the trend in boiling points across period 3 look like

Increases from sodium until silicon as charge increases

Silicone has highest as giant covalent and sharing electrons so elections localised and held strongly

Then decreases except s is higher than P as There are more electrons in S8 so stronger London forces