Group 1 & 2 metals & the halogens - AS Edexcel Chemistry

• Add a few drops of concentrated hydrochloric acid to a sample on a watch glass.
• Dip a clean platinum or nichrome wire into the mixture and then hold it in a hot Bunsen burner flame. The colour allows you to identify the cation in the mixture.
• Coloured flames are produced due to the presence of metal ions.

In a flame test, the sample is converted to a chloride because chlorides are more volatile than other salts.

The outer electrons become excited and are promoted to a higher energy level before dropping back to its ground state (original energy level). In doing so, a photon of light is emitted in the visible range.

Flame colours for the Group  1 and 2 cations

*The magnesium flame can appear white because of the production of magnesium oxide, however magnesium does not have an excited electronic state corresponding to these wavelengths, so it emits no visible radiation.

With Oxygen

• form solid metal oxides
• reactivity increases down the group

2Mg(s) + O₂(g) --> 2MgO(s)

With Chlorine gas

• form solid metal chlorides
• reactivity increases down the group (not as clearly seen as with oxygen)

Mg(s) + Cl₂(g) --> MgCl₂(s)

• Mg reacts with steam to form MgO(s) and H₂(g)
• Ca, Sr and Ba react to form the hydroxide and H₂,

e.g. Ca(s) + 2H₂O(g) --> Ca(OH)₂(aq) + H₂(g)

• forms an alkaline solution
• reactivity increases down the group

With Water

• Beryllium doesn't react, Mg reacts only slightly
• Slaking lime - when cold water is added to calcium oxide, a fizzing sound is heard. Water vapour is released and calcium hydroxide produced. (HIGH pH) CaO(s) + H₂O(g) --> Ca(OH)₂ (aq) Strontium oxide and barium oxide react in a similar way to calcium oxide.

With dilute hydrochloric acid (HCl) and nitric acids  (HNO₃)

• forms a chloride salt & water:MgO(s) + 2HCl(aq) --> MgCl₂(aq) + H₂O(l)

• forms a nitrate salt & water:CaO(s) + 2HNO₃(aq) --> Ca(NO₃)₂(aq) + H₂O(l)  

Hydroxides with dilute hydrochloric acid (HCl) and nitric acids (HNO₃)

• Similar to oxides - salt and 2 moles of water produced

• Halogens are non-polar and strong oxidising agents
• Reactivity decreases down the group
• Low melting and boiling pts, but they increase down the group
• Solubility in water decreases down the group (iodine is almost insoluble), more soluble in hydrocarbon solvents
  

• All group 2 metal nitrates and metal chlorides are soluble
• Group 2 salts in which the anion has a charge of -2 (e.g. sulfates) are largely insoluble (except Mg and Ca salts)
• Solubility of the salts decreases down the group as atomic number & ionic size increases
• Solubility of sulfates decreases down the group, solubility of hydroxides increases.

Test for sulfate ions

• Add dilute HCl to sample to destroy any carbonate
• Then add barium chloride or barium nitrate solution
• A white precipitate of barium sulfate forms

Ba²⁺(aq) + SO₄^2-(aq) --> BaSO₄(s)

The stability of ionic compounds increases as:

• Cationic radius decreases
• The ionic charge increases

Carbonates

• GROUP 1 are stable (except for lithium carbonate, Li₂CO₃ which decomposes to give lithium oxide, Li₂O and  CO₂
  
• GROUP 2 decompose to form stable oxides and CO₂. (Beryllium carbonate is unstable, doesn't exist at room temp.)
  

Nitrates

• GROUP 1 form their nitrites (2NaNO₂) which are stable to heat, and oxygen
• GROUP 2 decompose to form their oxide, nitrogen dioxide and oxygen 

With metals

• become reduced to negative halide ions

GROUP 1: 2Na(s) + Cl₂(g) --> 2NaCl(s) GROUP 2: Mg(s) + Cl₂(g) --> MgCl₂(s)
Iron reacts with halogens to form iron(III) halides, e.g 2Fe(s) + 3Cl₂(g) --> 2FeCl₃(s)

With non-metals

• achieve noble gas configuration through covalent bonding.
• Chlorine reacts with phosphorus to form phosphorus(III) chloride (2PCl₃)

With iron(II) chloride solution - 2FeCl₂ (pale green)

• Halogen oxidises green(II) ions to brown(III) ions: 2FeCl₃(aq)

• react with ammonia to form ammonium halides: NH₃(g) + HCl(g) --> NH₄Cl(s)

Halogens with potassium halides

• Displacement (redox) reaction: 2KI(aq) + Cl₂(g) --> 2KCl(aq) + I₂(aq)

Sodium halides with concentrated sulphuric acid (H₂SO₄)

• Forms a hydrogen halide: NaCl(s) + H₂SO₄(l) --> NaHSO₄(s) + HCl(g)
• With NaBr and NaI, the acid acts as an oxidising agent, so the respective halogen is produced
• Hydrogen iodide is a stronger reducing agent than hydrogen bromide

Potassium dichromate: test for SO₂(g)
Lead ethanoate paper: test for hydrogen sulphide H₂S(g)

• Add dilute nitric acid to unknown halide solution (to prevent the precipitation of any other silver salts)
• Add silver nitrate solution
• If left in sunlight, partial decomposition of silver chloride into silver and chlorine turns the precipitate grey

Ag⁺(aq) + Cl^-(aq) --> AgCl(s)
AgCl(s) + 2NH₃(aq) --> [Ag(NH₃)₂]⁺(aq) + Cl^-(aq)