chemistry- periodicty draft

Periodicity definition

Repeating trends of physical or chemical properties

What is a period

Horizontal rows of elements. Period one starts at H, period starts at Li…

All elements in the same period have the same number of principle energy levels and there are trends in the properties as w move along a period

What is a group

A vertical column

All elements in a group have the same number of outer electrons and similar chemical properties.

How does atomic radius change down a group and why

Increases down a group

As more shells

So more shielding

So outer electrons are further from the nucleus and have a weaker attraction to the nucleus

How does atomic radius change down a period and why

Atomic radius decreases down a period

As number of protons increases and shielding stays the same

So stronger nuclear charge and stronger attraction between outer electrons and nucleus

What effect is more significant shielding/ distance or number of protons

Shielding/ distance

What is the first ionisation energy

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

Write an equation showing the process of removing the first ionisation energy

Describe and explain the trend in the first ionisation energy in group 1 elements

Ionisation energy decreases.

As the electron is removed from a higher principle energy leaves.

So distance between the outer electrons and the nucleus increases.

Shielding increases. Weaker attraction between nepucleus and outer electron so less energy required to remove electron.

Describe and explain the general trend of ionisation enerypgy across a period.

Ionisation energy increases as number of protons increase and atomic radius decreases. Shielding remains the same so there is a stronger attraction between the nucleus and the outer electrons, so more energy is required to remover electron.

Draw a graph to show ionisation energy across a period

Increases. Except for between group 2 and group 3 elements where it decreases slightly, and between group 5 and 6 where it also decreases but by less.

Why is AL first IE lower than Mg

Mg: 1s2, 2s2, 2p6, 3s2

Al: 1s2, 2s2, 2p6, 3s2, 3p1

Electron is removed from a higher energy p sub shell so there is a weaker attraction between nucleus and outer electrons. Less energy required to remove Al.

Why is S first IE lower than P

P: 1s2, 2s2, 2p6, 3s2, 3p3

S: 1s2, 2s2, 2p6, 3s2, 3p4

There is a pair of electrons in P orbital which repel each other so less energy required to remove electron.

What are successive ionisation energy’s

Measured when additional electrons are removed, always removing outermost electrons.

You can as many successive ionisation energies as there are electrons in an atom

Explain why successive ionisation energies increase

Successive IEs will always increase

As positive charge in ion increases

So ionic radius decreases .

So nuclear attraction increases.

So more energy is required to remove outer electron

Write the first four SIE for Al

How can ionisation energies be used as proof for principle energy shells.

IE will always increase by a larger margin when an electron is being removed from a lower shell that is closer to the nucleus..

What is metallic bonding

String electrostatic forces of attraction between positive ions and delocalised electrons

How many delocalised electrons are in a metallic lattice

The number of delocalised electrons is proportional to those lost by mega, atoms in their attempt to gain a noble gas configuration. E.g. Na will donate 1 per atom, Mg will donate 2

What factors influence the strength of a metallic bond and why

1. Ionic charge on the metal (higher charge stronger bond)

2. Number of delocalised electrons (more e- stronger bond)

3. Ionic radius (smaller radius stronger bond)

This is because the positive ions will be more strongly attracted to the delocalised sea of electrons

Explain the conductivity of metals

Good conductors as the delocalised electrons are free to move and carry a charge throughout the structure.

Explain the melting point of metals

High melting a boiling points due to strong electrostatic forces of attraction between positive ions and delocalised electrons

Explain the malleability and ductility of metals

They can be hammered into shape as the layers on ions can slide over one and other

What are the giant covalent structures in Period 2 and 3

Diamond, graphite, graphene, silicon, silicon dioxide

What is the melting point of giant covalent structures

High as string covalent bonds between atoms require a lot of energy to break

Describe the bonding and conductivity of diamond

Each carbon atom covalently bonded to 4 other carbon atoms.

Not conductive as no delocalised electrons

Describe the bonding and conductivity of graphite

Each carbon atom covalently bonded to three other C atoms forming hexagonal layers. Weak London forces between layers.

Yes as one delocalised electron per atom free to carry charge by moving and flowing

Describe the bonding and conductivity of graphene

1 layer of graphite, each carbon atom covalently bonded to 3 other forming a single hexagonal layer.

Conductive as delocalised electrons are free to move and flow.

Describe the bonding and conductivity of silicon

Each silicon atom covalently bonded to four other silicon atoms.

No conductivity as no delocalised electrons.

Describe the bonding and conductivity of silicon dioxide

Each Si atom bonded to four O atoms.

No as no delocalised electrons

Period 2 what is bonding and crystal structure of Li and Be

Bonding - metallic

Crystal structure - giant metallic

Period 2 what is bonding and crystal structure of C

Bonding - covalent

Crystal structure- giant covalent (diamond, graphite, graphene)

Period 2 what is bonding and crystal structure of N, O and F

Bonding, covalent

Crystal structure, simple molecular. All form diatomic and are non polar so only London forces.

Period 2 what is bonding and crystal structure of Ne

Bonding, monoatomic (single atom)

Crystal structure, covalent bonds, nonpolar atoms

Period 3 what is bonding and crystal structure of Na and Mg and Al

Bonding, metallic

Crystal structure, giant metallic lattice

Period 2 what is bonding and crystal structure of Si

Bonding, covalent

Crystal structure, giant covalent

Period 3 what is bonding and crystal structure of P, S Cl

Covalent

Simple molecular

Period 3 what is bonding and crystal structure of Ar

Monoatomic

Simple molecular, non polar atoms so only London forces between molecules.