Periodicity  Elements in the PT are arranged in order of increasing atomic number.  Elements in the same group - same chemical and physical properties.  Across the period - repeating pattern of physical and chemical properties known as periodicity.

Periodic Trends Properties such as  Melting points  Electronegativity  First ionisation energy show periodicity

Ionisation Energy The first ionisation energy is the energy required to remove the valence electron from 1 mole of gaseous atom to form a positive ion. The equation for 1 st IE of element M M(g)  M + (g) + e - Example H(g)  H + (g) + e - ΔH = +1310kJmol -1 The amout of energy required to carry out this process for 1 mole of hydrogen atoms is 1310kilojoules.

Electronegativity  The electronegativity is the ability of an atom in a covalent bond to attract shared pairs electrons to itself.  The greater the electronegativity of an atom, the greater its ability to attract shared pairs of electrons to itself.  Electronegativity value is based on the Pauling scale. A value of 4.0 is give to F (most electronegative atom). The least electronegative elements, Ce and Fr both have a value of 0.7

Atomic Radii  Metallic radius is half the distance between nuclei in a metallic crystal.  Covalent radius is half the distance between like atoms that are bonded together in a molecule.  van der Waals’ radius is the effective radius of adjacent atoms which are not chemically bonded in a solid, but are presumably in "contact". An example would be the distance between the iodine atoms of adjacent I 2 molecules in crystalline iodine. (Source:

Atomic & ionic radii down a group  Atomic radii is determined by 2 opposing factors - Shielding effect by the electrons of the inner shell(s) - Nuclear charge (due to protons)  Moving down the group, both the nuclear charge and shielding effect increase. However, the outer electrons enter new shells. So, although the nucleus gains protons, the electrons are not further away, but also more effectively screened by an addtional shell of electrons.

Ionic radii for ions of the same charge also increases down a group for the same reason. Atomic radius increases down the group

Trends in first ionisation energy  The atomic radius increases down the group as additional electrons are added, causing the shielding effect to increase.  The further the outer shell is from the nucleus, the smaller the attractive force exerted by the protons in the nucleus.  More easily an outer electron can be removed, the lower the ionisation energy.

The first IE decreases down the group

Trends in electronegativity  There is an increasing distance between the nucleus and electrons down the group. Hence, the attractive force is decreased.  Although the nuclear charge increases down the group, this is counteracted by the increased shielding effect due to additional electron shells.

The electronegativity decreases down the group

Trends in melting point Group I  Metals are held together by metallic bonding.  The strength of metallic bonding decreases because the attractive forces between the delocalised electrons and the nuclues decreases owing to the increase in the distance. The increase in the nuclear charge is counteracted by the increase in shielding. Group 7  As the molecules become large, the attractive forces between them increases with the number of electrons in atoms or molecules.

The melting point decreases down group I The melting point increases down group 7

Atomic radii across a period  Across the period from left to right, the no. of protons and electrons increase by one.  The electrons are added to the same shell, hence, there is only a slight increase in the shielding effect across the period.  At the same time, additional protons are added to the nucleus, increasing the nuclear charge.  The effect of increasing nuclear charge more than outweighs the small increase and consequently all the electrons are pulled closer to the nucleus.

The atomic radius decreases across the period

Ionic radii across a period  The radii of positive ions decrease from Na + to Al 3+  The radii of positive ions decrease from P 3- (phosphide ion) to Cl -  The ionic radii increase from the Al 3+ to P 3-.

Isoelectronic Species  Isoelectronic species are atoms and ions that have the same number of electrons  The large increase in size from Al 3+ to P 3- is due to the presence of additional electron shell which causes a large increase in the shielding effect resullting in an increase in the ionic radius. SpeciesNa + Mg 2+ Al 3+ Nuclear charge Number of electrons10 Ionic configuration SpeciesP 3- S 2- Cl - Nuclear charge Number of electrons18 Ionic configuration

First ionisation across a period  When moving across the period from left to right, the nuclear charge increases but the shielding effect only increases slightly (since electrons enter the same shell).  Hence, the electron shells are pulled progressively closer to the nucleus.

The electronegativity increases across the period

Comparing the eletronegativity values

Summary of Periodic Trends

Chemical Properties Group I alkali metals  Li, Na and K contain 1 valence electron.  Reactive metals, stored under liquid paraffin to prevent them from reacting with air.  Readily lose their valence electron -good reducing agent  Reactivity increases down the group

 React with water to form an alkali solution of the metal hydroxide and hydrogen gas. (i) 2Li(s) + 2H 2 O(l)  LiOH(aq) + H 2 (g) Lithium floats and reacts quietly (ii) 2Na(s) + 2H 2 O(l)  NaOH(aq) + H 2 (g) Sodium melts into a ball which darts around on the surface (iii) 2Ks) + 2H 2 O(l)  KOH(aq) + H 2 (g) Heat generated from the reaction with potassium ignites the hydrogen.

 React readily with chlorine, bromine and iodine to form ionic salts, e.g. (i) 2Na(s) + Cl 2 (g)  2NaCl(s) (ii) 2K(s) + Br 2 (l)  2KBr(s) (iii) 2Ks) + I 2 (g)  2LiI(s)

Chemical Properties  Chlorine is a stronger oxidizing agent than bromine, so can remove the electron from bromide ions in solution to form chloride ions and bromine.  Similarly, both chlorine and bromine can oxidize iodide ions to form iodine. (i) Cl 2 (aq) + 2Br - (aq)  2Cl - (aq) + Br 2 (aq) (ii) Cl 2 (aq) + 2I - (aq)  2Cl - (aq) + I 2 (aq) (iii) Br 2 (aq) + 2I - (aq)  2Br - (aq) + I 2 (aq)

Test for halide ions  The presence of halide ions in solution can be detected by adding silver nitrate solution. Ag + (aq ) + X - (aq)  AgX(s) X = Cl,Br or I light AgCl white Ag(s) + ½ X 2 AgBr cream AgI yellow