Group 2 - Alkali Earth

• Group 2 metals

  • Beryllium
  • Magnesium
  • Calcium
  • Strontium
  • Barium
  • Radium
• Group 2 metals

  • Become more reactive as you go down the group
• Atomic radius

  Increases from Mg to Ba since each atom has an additional full shell of electrons making the atom bigger
• Ionisation energy

  
  

  • Decreases from Mg to Ba as the outermost electron is further from the nucleus AND there is an additional full shell of shielding between the nucleus and the outermost electron
  • Shielding reduces the attractive force of the nucleus for the outermost electrons due to the repulsive force of the inner shell electrons.
• Metallic bonding

  
  

  • Each metal forms +2 ions with a sea of delocalised electrons. They have high melting points
  • The trend in melting points of group 2 metals decreases as you go down the group.
• Ionic radius of +2 ions increases as you go down the group


  
  

  • The ionic radius of the +2 ions increases as you go down the group because each consecutive element has an additional full shell of electrons.
  • The electrostatic attraction between the positive ions and delocalised electrons weakens as the ion size increases.
  • Metallic bonding is therefore weaker as you go down the group, and so less energy is needed to melt the metal
• Reaction with oxygen

  
  

  • Group 2 metals react at room temperature with oxygen in air to form a coating of metal oxide on the surface
  • Group 2 metals burn in oxygen to form the metal oxide
  • Coating of MO usually on surface - needs to be removed with sandpaper before reacting
• Reaction with water

  1. Mg reacts with warm water to form Mg(OH)2

  2. Other group 2 metals react with cold water to form the metal hydroxide

  3. Exothermic reaction

  4. Fizzing due to hydrogen gas

  5. Metal dissolves (faster down the group)

  6. Less soluble (Mg and Ca) hydroxides may also form a precipitate

  
  

  Alkaline solutions of hydroxides turn red litmus blue because of presence of OH- ions
• Reactions of group 2 metals with water and steam are redox reactions:

  • Metal: 0 to +2
  • Hydrogen: +1 to 0
• Reaction with steam

  1. Group 2 metals react with steam to form the metal oxide instead of the hydroxide

  2. Mg burns with a bright white flame, Mg appears as a white powder

  3. Same redox as for water
• Metal displacement reactions

  
  

  • Used to extract less reactive metals from their ores, e.g. extracting Ti from TiO2 using Mg
  • Expensive and diffuclt so need a good reason to do it
• Steps to extract Ti from TiO2

  
  

  • TiO2 is reacted with Cl2 and C:
  • TiCl4 TiO2 + 2Cl2 + 2C -> TiCl4 + 2CO
  • TiCl4 is purified by fractional distillation in argon atmosphere (it is a liquid at room temperature unlike solid ionic TiO2)
  • Ti is extracted by reaction between TiCl4 and Mg in an argon atmosphere at 500°C TiCl4 + 2Mg -> Ti + 2MgCl2
  • In this redox reaction Ti is reduced from +4 to 0 and Mg is oxidised from 0 to +2.
• Metal oxides are bases


  Can react with acids in neutralisation reactions, e.g. CaO reacting with SO2 to remove it from power station waste gases
• Solubility of group 2 hydroxides

  • Become more soluble as you go down the group
  • Mg(OH)2 is nearly insoluble
  • Ca(OH)2 is partially soluble - pH 11
  • Sr(OH)2 and Ba(OH)2 are strongly alkaline
  • Ba(OH)2 (s) -> Ba2+(aq) + 2OH-(aq)
• • Insoluble group 2 hydroxides produce a white precipitate when NaOH is added
  • Mg2+(aq) + 2OH-(aq) -> Mg(OH)2(s)
• Uses of group 2 hydroxides

  • Mg(OH)2 used in medicine to neutralise stomach acid and treat constipation
  • Ca(OH)2 used in agriculture to neutralise acidic soils and test for CO2
• Magnesium hydroxide


  • Used in medicine (in a suspension called milk of magnesia) to neutralise excess stomach acid (HCl) and to treat constipation
  • Mg(OH)2(s) + 2HCl(aq) à MgCl2(aq) + 2H2O(l)
  • Advantage over calcium carbonate is that it does not produce carbon dioxide gas that would cause bloating and wind
• Calcium hydroxide


  • Used in agriculture to neutralise acidic soils
  • Used to test for the presence of carbon dioxide
  • An aqueous solution of calcium hydroxide is called limewater. It can be used as a test for carbon dioxide because it turns cloudy due to the formation of a white precipitate of calcium carbonate
  • Ca(OH)2(aq) + CO2(g) -> CaCO3(s) + H2O(l)
• Calcium hydroxide is slightly soluble in water at room temperature. As the temperature rises, the solubility decreases.
• Determining the solubility of calcium hydroxide

  1. Combine the two soluble salts

  2. Insoluble salt forms a precipitate

  3. Filter the solution containing the precipitate

  4. Wash the precipitate with deionised water

  5. Allow the precipitate to dry
• Solubility of G2 Sulfates

  
  

  • Group 2 sulfates become less soluble as you go down the group. MgSO4 is the most soluble, BaSO4 is the least soluble
  • When a precipitate forms these are the ionic equations:
  • Sr2+(aq) + SO42-(aq) -> SrSO4(s) (white precipitate)
  • Ba2+(aq) + SO42-(aq) -> BaSO4(s) (white precipitate)
  • An overall equation would include soluble spectator ions for example the reaction of soluble BaCl2 with soluble Na2SO4: BaCl2(aq) + Na2SO4(aq) -> BaSO4(s) + 2NaCl(aq)
• Reaction of barium with sulfuric acid

  1. Ba + H2SO4 → BaSO4 + H2

  2. The formation of the insoluble barium sulfate layer halts the reaction, which does not happen with acids such as HCl or HNO3 since the salts formed are soluble
• Barium sulfate


  • Very insoluble
  • Absorbs x-rays
  • Used in medicine as a 'barium meal' given to patients who need x-rays of their intestines
  • The barium absorbs the x-rays and so the gut (which would ordinarily be invisible to x-ray) shows up on the x-ray images allowing blockages or other problems to be identified
  • Barium compounds are toxic but the very low solubility of barium sulfate means that it can pass through the body without being absorbed into the blood
  • It would be deadly to use barium sulfate or barium chloride for the same purpose
• Testing for sulfate ions

  1. Add barium chloride acidified with HCl to a solution

  2. A white precipitate of barium sulfate forms if sulfate ions are present

  3. Acidification is necessary to remove interference from carbonate ions
• Precipitation reaction


  
  

  • A reaction where an insoluble solid is formed from soluble ions
  • For example, the formation of barium sulfate from barium ions and sulfate ions
  • Ba2+(aq) + SO42-(aq) -> BaSO4(s)
  • Whenever a pair of soluble ions which together form an insoluble solid, are present together in solution they will rapidly precipitate out from solution forming a cloud precipitate
  • This could be a fine cloudiness or a thick gel like mixture (transition metal hydroxides are very gel like).
• Preparation of an insoluble salt

  1. Combine the two soluble salts

  2. Insoluble salt forms a precipitate

  3. Filter the solution containing the precipitate

  4. Wash the precipitate with deionised water

  5. Allow the precipitate to dry
• Flame test for group 2 metal ions

  1. Dip a wire loop in nitric acid

  2. Dip in the salt you want to test

  3. Hold the loop in the hottest part of a roaring Bunsen flame

  4. Observe the colour
• Sr2+


  • Forms a white precipitate with addition of sulfate ions
  • Crimson red flame test
• Test for NH4+ the ammonium cation

  1. Add a few drops of NaOH to the test solution

  2. Warm gently

  3. Ammonia gas is produced if ammonium ions are present

  4. Test using damp red litmus paper which turns blue in the presence of ammonia gas
• Outer coating of MO (metal oxide) or M(OH)2 (metal hydroxide) inhibits reaction
• Twist the Mg strip to fit it into the test tube
• Mg will not react much visible with cold water
• Metal Hydroxides and Alkalinity

  • Mg(OH)2 is the least alkaline
  • Ba(OH)2 is the most alkaline
  • Alkalinity increases down the group
  • This is because of solubility - the more the soluble the hydroxide, the more alkaline the solution of the metal hydroxide will be due to the presence of dissosciated OH- ions in solution.
• Reaction with Steam


  • Cotton wool soaked in water at the bottom the tube
  • Chunk of metal in the middle
  • Heated
  • Extra AE provided with steam (not given with water) makes MO not M(OH)2
• Titanium


  High melting point solid metallic lattice
• TiCl4


  Molecular liquid
• Fractional Distillation


  • In an argon atmosphere
  • Not in oxygen because presence of oxidation would lead to formation of TiO2 and MgO
  • (Argon is also inert)
• TiO2


  Titanium sponge - porous with many holes
• BaSO4


  Such a low solubility that none of it will dissolve to travel into bloodstream
• Titanium 


  • Very useful metal
  • Low density
  • Corrosion resistant
  • Used to make strong and light alloys in the aerospace and medical industries
  • While its ore is abundant it is relatively difficult to extract
  • Ti cannot be extracted by carbon because it forms TiC
  • Ti cannot be extracted by electrolysis like aluminium because it has to be very pure.
  • Titanium is therefore extracted using a redox displacement reaction with the more reactive Mg.