periodicity

• Period: A row across the periodic table.
• Group: A column down the periodic table.
• Periodicity:Trends down groups and across periods.

Blocks in periodic table

1) Period 2 (Li to Ne) and Period 3 (Na to Ar) show similar patterns.

Electron configuration: 

Na:

Mg:

Al:

Si:

P:

S:

Cl:

Ar:

2) First ionisation energy INCREASES across the period with a dip between group 2 and 3 and between group 5 and 6.

Definition: The energy requried to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions. 

Mg -> Mg+ + e-

Explanation (across a period):

• Number of protons increase
• Shielding is the same (electrons are added to the same energy shell)
• Attraction between nucleus and outer electrons increases.
• Therefore, outer electron is harder to remove.

• Shielding: same
• Distance from nucleus: decreases
• Charge of nucleus: increases

In general there's an increase in 1st IE across the period because the nuclear charge increases and the distance to outer electrons decreases.

3) Atomic radius DECREASES across the period.

Explanation:

• Number of protons increases
• Shielding is the same
• Attraction between nucleus and outer electrons increases so electrons are pulled in more tightly.

Nuclear charge increases, outer electrons are in the same shell so are held in tighter due to it being more positive. 

4) Ionic radius

Positive ions (cations) are smaller than their atoms because:

• Outer shell is lost so outer electrons have been removed. E.g. Na (11e-) Na+ (10e-)

Negative ions (anions) are larger than their atoms because:

• Electron(s) added to the same shell.
• Same number of protons attracting more electrons. More electrons = more repulsion
• Same nuclear charge.

E.g. Oxygens ion (O2-) 8protons, 10electrons

5) Melting and boiling point

Trend: Increases from group 1 to group 4 and then decreases sharply to group 5 and then group 0.

Explanation: 

This trend is due to the change in structure and bonding across the period.

• Na to Al (Structure = Giant metallic lattice. Bonding = Metallic)

As we go across the period the positive metal ions become more positive, the charge on the ion increases, there are more delocalised electrons so more energy is needed to break the bonds. M.P increases.

5) Si

Structure = Giant covalent lattice

Bonding = Covalent bonds

• High MP as strong covalent bonds require lots of energy to break. 

5) P , S , Cl ,

Structure = Simple covalent molecules

Explanation:

• Molecules held together by weak VDW's forces. 
• VDW's created on the surface of the molecule.
• More surface area, more points of contact, more VDW's therefore more energy to break. 

5) Ar

Structure = Monoatomic

• Very small surface area.
• Very few points of contact.
• Very few VDW's forces.
• Little energy required to overcome them so low MP & BP.

Metallic

• The electrostatic attraction between positive metal ions and a sea of delocalised electrons. 
• Strength - Strong.

Covalent

• A shared pair of electrons.
• Strength - Strong.

VDW's

• Temporary charge in electron distribution causing oscillating dipoles on the surface transferred onto the neighbouring molecule.
• Strength - Very weak.

Electron configuration

Be

Mg

Ca

Sr

Ba

1) Melting points: HIGH because of the strong electrostatic forces in metallic bonding.

The melting points decrease down the group because: the ion gets bigger and the no. of delocalised electrons is the same so metallic bonds get weaker. There's a lower charge:SA ratio. Less attraction to electrons so less energy required to break the bonds. 

2) Density: Group 2 ions are smaller than group 1 ions so can pack together more tightly. Group 2 has stronger metallic bonds than group 1 so more electrons in the sea resulting in stronger metallic bonds.

3) Reactivity: They are all reactive metals with similar chemical reactions(all want to lose 2 electrons). Increases down the group. All have 2 outer electrons which are lost during reactions. The most reactive can lose electrons easiest (Lowest 1st IE). E.g. Mg lowest 1st IE, more shielding, outer further away from nucleus which outweighs increase in charge. 

4) Appearance of group 2 compounds: Silver metals, shiny. Tarnish when left in air due to becoming oxidised.

5) First IE decreases down the group. (E.g. at no. 3)

1) Atomic radius INCREASES

• Extra shells of electrons - Outer electrons get further away from the nucleus.
• Extra shielding outweighs the increase in nuclear charge.
• So atoms get bigger as you go down the group.

2) Trends in reactivity down group 2:

Reactivity INCREASES down the group

• 1st IE decreases so outer electrons lost more readily.
• This is due to: 
• outer electron further away from the nucleus.
• more shielding
• extra shielding outweighs increase in nuclear charge

The reaction of water on group 2 oxides: Group 2 (II) oxides dissolve to from alkali solutions. 

Equation: MO (s) + H2O (l) --> M(OH)2 (aq)

pH of resulting solution: Above 7 for alkali. OH - strong alkali so 12-13 pH. 

Trends in alkalinity of group 2 hydroxides: Become more alkali as you go down the group as they become more soluble in water. 

Things only alkali if they're dissolved in water

Thermal decomposition of group 2 carbonates

Carbonates decompose to give a metal oxide (MO) and carbon dioxide (CO2)

Equation:

Trend in ease of decomposition: Harder to decompose as you go down the group. So MgCO3 decomposes the most readily.

Uses of group 2 elements and their compounds

Magnesium hydroxide - indigestion tablets - neutralise stomach acid (HCl)

Calcium hydroxide - agriculture to treat soil - reduce acidity of soil

• Halogens are in group 7 (VII): they have 7 outer electrons
• They are in the p block because their outer electrons are in the p sub shell
• They are diatomic elements (X2) F2, Cl2, Br2, I2

1) Explanation of trends

Atomic radius INCREASES down the group

This is because as you go down the group the atoms have extra shells of electrons. The outer electrons are further away from the nucleus. The increased shielding outweighs the increase in nuclear charge. 

The ionic radius (x-) is bigger than the atomic radius (x) because:

The protons in the ions are the same as in the atoms, there's more electron (due to negative charge on ion) so more repulsion so the ion is slightly bigger. More electrons than protons so nucleus exerts smaller force of attraction. 

Boiling point INCREASES down the group. The halogens become LESS volatile. This is because:

• Halogen molecules are held together by VDW's. The VDW's between the molecules get stronger as the no. of electrons in the molecules increases.
• As you go down the group, the molecule gets bigger, more surface area, more points of contact so more VDW's therefore more energy required to break them.

Volatility: The ease with which a liquid turns into a gas. (Volatility increases as boiling point decreases) 

Appearance of halogens

• Cl2 - green, gas at room temp, colourless in aq solution, colourless in cyclohexane
• Br2 - orange/brown, liquid at room temp, yellow/orange in aq solution, yellow in cyclohexane
• I2 - purple/silver, solid at room temp, orange/brown in aq solution, pink in cyclohexane

The halogens are more soluble in non-polar solvents than in water because they're non polar.

The solubility of the halogens in water decreases down the group. 

Oxidising powers of the halogens

An oxidising agent is a reagent that gains electrons in order to oxidise another species. 

Halogens are oxidising agents because they gain electrons to form negative ions.

The ability of the halogens to act as oxidising agents DECREASES down the group. 

Explanation as you go down the group:

• More shielding
• Outer electrons are further away from the nucleus so less attraction
• These outweigh the nuclear charge

When they react they gain an electron, the electron enters a shell that has greater shielding so is further away from the nucleus. It becomes harder to attract and gain an electron. 

Displacement Reactions 

• This series of experiments provides evidence for the trend in oxidising powers of the halogens. 
• Add a solution of each halogen to a solution of each potassium halide in turn.
• Note any colour changes and decide whether a reaction has occurred. 
• Add cyclohexane (a non-polar solvent) and note the colour in the organic (top) layer. This will tell you which halogen is now present. 

Relative reactivity of the halogens: 

Cl2 more reactive than Br2 

Br2 more reactive than I2

If X is higher up the group than Y:

X2 (aq) + 2Y- (aq) --> 2x- + Y2

Explanation:

The more reactive halogen X2 is better at accepting electrons, therefore takes them from the Y ion. 

E.g. Outer electron is fluorine is closer to the nucleus. More shielding in Cl. So higher up group there's more attraction to outer electrons. 

 Disproportionation is a redox reaction in which the same element is both oxidised and reduced.

The reaction of chlorine with water

An equilibrium is established between chlorine, water, hydrochloric acid and chloric (I) acid:

Chlorine is both oxidised and reduced

Observations

Chlorine water is green due to the presence of chlorine. The solution turns blue litmus paper red and then white. This is because HCl is a strong acid: HClO is a bleach. 

 Cl2 (Green gas) HCl (Turns blue litmus red) HClO (Bleaches litmus-turns white)

Reaction of chlorine with sodium hydroxide

• Sodium chloride, sodium chlorate (I) and water are produced. (Makes bleach)

The green colour of chlorine fades and the smell is less pungent.

Again chlorine is both oxidised and reduced.

Uses of chlorine in water treatment

Benefits - kills bacteria. 

 Risks - toxic, can form chloro organic compounds that are harmful

Examples of halides

• Sodium Chloride (NaCl) Ions present Na+ & Cl-
• Calcium Fluoride (CaF2) Ions present Ca2+ & 2F-
• Potassium Iodide (KI) Ions present K+ I-

 Test for halide ions:

• Make a solution of the substance to be tested 
• Add dilute nitric acid to remove carbonates (which interfere with the test) 
• Add silver nitrate solution (AgNO3) -> forms the precipitate
• Observe the colour of the precipitate 

Halide ion Colour of precipitate Name & formulae of precipitate Chloride White precipitate Silver Chloride (AgCl) solid
Bromide Cream precipitate Silver Bromide (AgBr) solid
Iodide Yellow precipitate Silver Ioded (AgI) solid

Silver Nitrate + Sodium Chloride

Silver Nitrate + Sodium Bromide

Silver Nitrate + Sodium Iodide

General ionic equation: Ag+ (aq) + X- (aq) --> AgX(s)

Use of silver halides 

Silver halides darken on exposure to light, producing silver metal. They are used in photography paper. 

E.g. AgBr   Ag+ + e- --> Ag

Trends in solubility of the silver halides in ammonia

• To each precipitate add dilute ammonia solution and observe if it dissolves 
• If not, add concentrated ammonia solution and see if it dissolves. 

Silver halide Solubility in ammonia Trend in solubility: Less soluble as go down

AgCl Dissolve in dilute NH3+conc NH3 

AgBr  Dissolve in conc NH3 but not in dilute NH3

AgI Doesn't disolve in NH3

Moles = Mass/Mr 

  Mass = Moles x Mr Mr = Mass/Moles

Moles = Conc x Vol

  Conc = Mol/Vol Vol = Mol/Conc

Mol=Vol of gas/24

Vol=Mol x 24