Practice:1. The electronic structure of an atom of an element in Group 6 of the Periodic Table could be:A 1s2 2s2 2p2

B 1s2 2s2 2p4

C 1s2 2s2 2p6 3s2 3p6 3d6 4s2

D 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6

2.

Electron configuration & chemical properties

Key Words:• Periodic law• Groups• Periods• Transition

metals• Metalloids• Lanthanides• Actinides

Objectives:

- electronic structure determines the chemical properties of an element

- periodic table is divided into blocks

Outcomes:D: recall that chemical properties are related to electronic structure- Know the blocks of the periodic tableA-C:- Know the chemical properties of:

- s-block elements- d-block elements- p-block elements

Periods & groups:• The reactivity of an element, and how it

combines with other elements, is determined by its arrangement of electrons in its outer shell

• The periodic table arranges elements in order of their atomic number

• Groups: the vertical columns in the periodic table

• Periods: the horizontal rows in a periodic table

• All the elements in a period have the same number of electron shells.

• So, the elements in each group and period show particular characteristics and trends in their chemical and physical properties

Periodic Law: the properties of the elements are a function of their atomic numbers

Blocks

Block Groups Subshell:s 1 + 2 Outer electrons in s subshells

p 3+4+5+6+7+0 Outer electrons in p subshells

d Transitional metals

Outer electrons in d subshells

f Lanthanides + actinides

Outer electrons in f subshells

s-block elements

• Reactive metals• Lower melting temperature• Lower boiling temperature• Lower density• Conduct electricity• Include hydrogen and helium – but usually

treated as a separate group.

Than other metals

d-block elements• Called Transitional metals• Less reactive that Group 1+ 2 metals – this is

because the inner d orbital is being filled while the outer s orbital is full

• All conduct electricity and heat• Are shiny, and hard• Ductile – pulled into shape• Malleable – hammered into shape• Mercury is the only exception – low melting

temperature liquid at room temperature

f-block elements• Lanthanides – are all similar• Actinides – all radioactive– Only the actinides up to uranium are naturally

occurring– The others have all been synthesises by scientists

and have extremely short half-lives

p-block elements• All the non-metals and metalloids• Include Tin and Lead– Form positive ions– Form ionic bonds with non-metals

• Many metals in p block do not have strong metallic characteristics– All conduct heat and electricity– Called post transitional metals generally

unreactive

• Metalloids occur in a diagonal block• Mostly like non-metals• Conduct electricity – but poorly• Silicon and germanium are responsible for

microchips

• Non-metals all form covalent bonds with other non-metals & ionic bonds with metals

• Majority do not conduct electricity• Some elements form giant covalent structures

Practice:

Trends in the Periodic Table

Key Words:• Atomic radius• Ionisation

energy• Melting

temperature

Objectives:

- Understand trends in the periodic table

Outcomes:

D: understand and describe the trends in the periodic table

A-C: Explain the trends in the periodic table - - -- ionization energy based on given data or recall of the shape of the plots of ionization energy versus atomic number using ideas of electronic structure and the way that electron energy levels vary across the period.

- melting temperature of the elements based on given data

Key Words• Ionization energy: the amount of energy it takes to

strip away the first electron• Electronegativity: a measure of how tightly an atom

holds onto its outer shell electrons• Nuclear charge: the attractive force between the

positive protons in the nucleus and the negative electrons in the energy levels. The more protons, the greater the nuclear charge.

• Shielding: inner electrons tend to shield the outer electrons from the attractive force of the nucleus. The more energy levels between the outer electrons and the nucleus, the more shielding.

• Atomic radius: measure of the size of atoms, usually measured from the nucleus to the outer shell

• Ionic radius: the size of ions

Key points:• Across periodic table: elements gain electrons

• Down a group : elements gain electron shells.

This changes the diameter of atoms which affects their physical and chemical properties

• The atomic radius generally decreases across a period:

The nuclear charge becomes increasingly positive as the number of protons in the nucleus increases.

The number of electrons also increases BUT they are all in the same shellThis means that they are attracted more strongly to

the nucleus – so reducing the atomic radius across a period

• The atomic radius generally increases down a group:

The outer electrons enter new energy levels down a group

So, even though the nucleus has more protons, the electrons are further away and they are screened by more electron shells.So, they are not held so tightly and the atomic radius

increases

Atoms to ion:

• The atomic radius changes when atoms form ions

• Positive ions always have a smaller ionic radius that the original atom.– Because: the loss of electron(s) means that the

remaining electrons each have a greater share of the positive charge of the nucleus so are more tightly bound

– And when an ion in formed, a whole ion shell is usually lost

• Negative ion has a larger atomic radius than that of the original atom

even though the extra electrons are in the same electron shell, the addition of the negative charge means that the electrons are less tightly bound to the nucleusSo the atomic radius is larger

Periodic Trends in Ionisation Energy:• Then more tightly held the outer electrons, the higher the ionisation

energy

3 main factors affecting ionisation energy of an atom:

The attraction between the nucleus & the outermost electron – decreases as the distance between them increases reducing the ionisation energy

The size of the positive nuclear charge - a more positive nucleus has a greater attraction for the outer electron so higher ionisation energy

Inner shells of electrons repel the outer electron, screening or shielding it from the nucleus - the more electron shells there are between the outer electrons and the nucleus, the less firmly held the outer electron is lower ionisation energy

Ionisation energy & periods:• Ionisation energy increases across a period• It becomes harder to remove an electron

• This is because:Increasing positive nuclear charge across the period

o Without the addition of extra electron shells to screen the outer electrons

The atomic radius gets smaller & electrons are held more firmly – so it requires more energy to make ionisation happen

The end of each period is marked by the high ionisation energy of a noble gas – this is a result of a stable electronic structure & indicates their unreactive natures

• (b) shows that First ionisation energies do not increase smoothly across a period

• This is because of subshells within each shell• E.g: the first ionisation energy of Be is larger

than B, Mg has a larger first ionisation energy than Al – why?– For Be or Mg, an electron must be removed from a

full s-shell. – Full subshells are particularly stable – so it requires

more energy than removing a single p electron from B or Al

• Nitrogen & phosphorous have unexpectedly high first ionisation energies:– They both have a half-full outer p subshell.– Half full subshells seem to have greater stability – So requires more energy

Ionisation energy decreases down a group – it becomes easier to lose an electron

Patterns in physical properties• The physical properties are closely linked to the

structure and bonding of atoms

• Melting temperature: the temperature at which the pure solid is in equilibrium with the pure liquid, at atmospheric pressure.– this is affected by the packing & binding of atoms

within a substance– It changes as you go across a period

• The relatively high melting temperatures of the metals (e.g. Li, Mg, Al) are due to their metallic structure.– The atoms are held tightly together is a ‘sea of

electrons’– It takes a lot of energy to separate them

• Giant molecular structures (metalloids-silicon, carbon-in form of diamond):– Strong covalent bonds between atoms which hold

them tightly in a crystal structure– Very difficult to remove individual atoms– So very high melting temperature

• Simple molecular structures:– Most non-metals found on right of periodic table– Small, individual molecules– Strong covalent bonds within molecules– But, molecules are held together by weak

intermolecular forces– Can be separated easily– Low melting temperature

Practice