Patterns in the Periodic Table

Patterns in the Periodic Table

The periodic table is made up of columns (groups) and horizontal rows – (periods).

The elements in Groups have similar chemical properties – due to the same number of outer electrons .

Metals are on the left hand side. Non-metals are on the right hand side.

Atomic SizeAtomic Size

As we move from left to right along a period the atomic size decreases.

This is because we are adding in more electrons in the same energy level.

There will be a corresponding increase in the number of protons in the nucleus.

Therefore, more + and – charges which are attracted to each other and so the atom will be smaller.

Atomic sizeAtomic size

As we go down a group the Atomic Size increases.

This is because we are adding more energy levels.

The electrons on the outside are “shielded” from the nucleus.

Therefore, less attraction and the atom is bigger.

Ionisation EnergyIonisation Energy

The First Ionisation Energy is – the energy required to remove 1 electron from each atom in 1 mole of an element in the gaseous state!

ExampleK(g) —>K+(g) + e- ΔH = + 425 kJ mol/lAs we go down a group the First Ionisation Energy

decreases due to the “ shielding” effect of the extra energy levels.( It’s easier to remove an electron! )

Patterns in the tablePatterns in the table

The First Ionisation Energy increases as you go left to right along a period in the periodic table.

This is due to increasing number of electrons within the same energy level and so a corresponding increase in protons in nucleus, resulting in a greater electrostatic attraction. Therefore it is more difficult to remove an electron!

Descending a groupDescending a group

As you go down a group in the periodic table the First Ionisation Energy decreases due to the increasing energy levels – the outer electrons are “shielded” from the nucleus by the extra energy levels and so are easier to remove!

Removing more than one electron!ExampleMg+ (g) —>Mg 2+ (g) + e- ΔH = + 1460 kJThis is the Second Ionisation Energy

ElectronegativityElectronegativity

This is a measure of the Attraction for Bonding ElectronsAtoms of different elements will have a different

attraction for bonding electrons i.e. a different electronegativity!

Electronegavitvity increases as you go from left to right along a period.

Electronegativity deceases as you go down a group.

BondingBonding

Covalent Bonding When non metal elements share electron

clouds, to achieve the stable electron arrangement of the noble gases.

Both + nuclei are attracted to the shared cloud and so the atoms are held by an “ electrostatic” attraction. A molecule is a group of atoms held together by a covalent bond.

Ionic BondsIonic Bonds

Bond formed between + metal ions and – non metal ions. Metal atoms lose electrons to achieve a stable electron arrangement, non metal atoms gain electrons to achieve a stable electron arrangement.

A large network of + and – ions form the crystal lattice. Ionic bonds are very strong.

Metallic bondingMetallic bonding

The atoms of metal elements are held together by metallic bonding. The electrons are free moving i.e. “ delocalised”. The move from one nucleus to the next, allowing metals to conduct electricity. There are the same number of protons in the nucleus as surrounding electrons – therefore metals are neutral.

Metallic bonds are very strong, since a lot of energy is required to overcome the attraction between the delocalised electrons and the nuclei. The greater the number of delocalised electrons the greater charge on the ions and the stronger the metallic bond.

Metallic strength increases from group 1 to group 2 etc.

Intra molecular ForcesIntra molecular Forces

These are the forces of attraction within molecules or compounds.

Covalent, Ionic and metallic bonding are examples of Intra molecular forces.

Intermolecular ForcesIntermolecular Forces

These are forces of attraction between DIFFERENT molecules.

ExampleVan der Waal’s forcesHydrogen Bonding

Polar Covalent BondsPolar Covalent Bonds

When H bonds with another H they form a covalent bond.

This is because they both have the same number of protons in the nucleus and so both have an equal attraction I.e. same electronegativity, for the shared electron cloud.

When H and Cl bond it is polar.The Cl has more protons in nucleus and so will have a

greater “pull” on the shared electron cloud. Cl has a greater electronegativity than H.

DipoleDipole

When there is a difference in electronegativity between atoms in a molecule a dipole can be created.

This is when each end of the molecule becomes “ slightly charged”. One end will be “slightly negative” = δ- . The other end will be “slightly positive” = δ+..

Example – Hydrogen chloride – H δ+ Cl δ-

Van der Waal’s ForcesVan der Waal’s Forces

This is a very weak, temporary force of attraction between different atoms or different molecules.

Electrons are constantly moving round the nucleus. Sometimes there may be an unequal distribution of the electrons and this may end up in a temporary dipole being created.

The atoms with this temporary dipole attract other atoms with an oppositely charged temporary dipole and so Van der Waal’s force of attraction is created.

Permanent dipole – permanent dipole

attractions!

Permanent dipole – permanent dipole

attractions!These are additional intermolecular forces

between polar covalent molecules.E.g. H2O , H ClThese are stronger than the temporary Van der

Waal’s forces.They result in an increase in the MP and BP of

polar covalent compounds.

Hydrogen BondingHydrogen Bonding

Hydrogen has a very low electronegativity.If it bonds to an element with a higher electronegativity

e.g. 0, N or F - the molecule will be highly polar.This strong intermolecular force is called a Hydrogen

bond.A hydrogen bond is the strongest intermolecular force

but it is not as strong as a covalent intra molecular bond.Hydrogen bonding increases the BP of the compounds.

WaterWater

The density of water increases as the T falls to 4oC.But it starts to decrease between 4oC and 0oC.The water molecules start to move further apart and form

an open structure held together by Hydrogen bonds. As a result ice is less dense than water and so floats on

water.In pools the water freezes from the surface downwards,

insulating the water beneath it.

Polar MoleculesPolar Molecules

A molecule may have an unequal electronegativity within atoms – creating a polar bond.

However the polarity of the whole molecule is related to symmetry.

In polar molecules the bonds are not arranged in symmetry, asymmetrical– e.g. HCl , NH3.,H2O

Non polar molecules have bonds arranged in symmetry – the permanent dipole – permanent dipole attraction cancel each other out – e.g. CCl4, CO2

Structure and Properties of Elements!

Structure and Properties of Elements!

Element Structure M.P. Conductor?

Metals Solids High Yes

B, C, Si Covalent network solids

High No

P,S Discrete covalent molecules (solids)

Low No

H,N,O,F,Cl Diatomic gases Low No

Ne, Ar, He Monatomic gases Low No

Covalent Network SolidsCovalent Network Solids

Boron, Carbon and Silicon form Covalent Network Solids!

This is a large lattice of covalently bonded atoms.They have High MP and BP.ExamplesDiamond – a large network of tetrahedrally

arranged carbon atoms. It has a rigid 3D structure making it very hard!

GraphiteGraphite

Carbon atoms held in planar hexagonal rings ( 6 atoms joined together)

The rings slide over one another – powder appearance.

Each carbon atom has 3 bonds – 1 electron is delocalised allowing Graphite to conduct electricity.

Other Covalent NetworksOther Covalent Networks

SiliconSilicon has a similar crystal structure as diamond.

Example - Silicon carbide( carborundum ), Silicon dioxide – has a SiO4 tetrahedral structure

Silicon is a semi conductor!BoronForms interlocking B12 atoms.It can be as hard as diamond.

Covalent Molecular Solids - Discrete

Covalent Molecular Solids - Discrete

FullerinesBuckminster Fullerine – C5 and C6 rings form a dome

shaped solid ( football!) with a total of C60 atoms. Other fullerines – C70 dome structures.S

Fullerines can form addition reactions with halogens.Some are used as catalysts – (palladium containing

compounds.)SPhosphorous – P4 tetrahedral structure.Sulphur – S8 puckered rings.

SolubilitySolubility

Ionic compounds dissolve in water because water is polar!

The H δ+ and O δ- pull apart the ionic lattice.The H δ+ part of the water molecule surrounds the

negative ion and the O δ- , of the water molecule, surrounds the positive ion.

Electrostatic attractions from between the ions and the polar ends of the water molecule – this overcomes the electrostatic attraction between the ions in the lattice.

Solubility patternsSolubility patterns

Ionic or polar substances dissolve in polar solvents e.g. water.

Covalent, non polar substances dissolve in non polar solvents e.g. ethanol, tetrachloromethane.

Ionisation in waterPolar compounds can ionise when dissolved in

water I.e. separate into separate ions – this allows them to conduct electricity.