Chemical Properties of Metals

Nature of Oxides of Metals

To Understand the concept, Let us do this small experiment

  • Take an about 10 cm long magnesium ribbon and rub it with sand paper so as to expose its shining surface
  • hold the magnesium ribbon in fire tongs and introduce the other end in Bunsen flame till it catches fire
  • hold the burning magnesium over a glass plate

Observation- You will observe that the magnesium ribbon burns with a dazzling white flame and forms a white residue which falls on the glass plate

Actually the magnesium ribbon chemically reacts with oxygen present in the air to form magnesium oxide.

Step-2

Divide the residue into two parts .

  • Dissolve the first part in 5 ml of distilled water ,by shaking contents in a test tube
  • Test the solution with blue as well as red litmus Paper
  • You will observe that Red litmus turns blue indicating Magnesium oxide is basic in nature
  • Actually Magnesium oxide reacts with water to form magnesium hydroxide which is basic in nature

Take the second part of magnesium oxide in a clean test tube

  • add to it 5 ml of Hydrochloric acid
  • gently shake the contents
  • You will observe magnesium oxide disolve to form clear solution and test tube becomes warm
  • It shows magnesium oxide enters chemical reaction with hydrochloric acid

Metals reacts with oxygen to form respective oxides

Oxides dissolve with water to form respective hydro oxides which are basic in nature

  • Most metals reatcs with oxygen to form their oixdes
  • some metals react more rapidly with oxygen as compared to other metal ,i.e. the rate of reaction of different metals with oxygen is different
  • Oxides of the metals sodium,potassium,calcium and magnesium are soluble in water. they turn red litmus blue
  • Oxides of all metals dissolve in acids to form their respective metal salts and water ,hence oxides are basic in nature

Reaction of Metals with water

  • Metals form respective metal hydroxide and hydrogen gas when react with water.

Metal + Water ⇨ Metal hydroxide + Hydrogen

  • Most of the metals do not react with water. However, alkali metals react vigorously with water.

Reaction of sodium metal with water: Sodium metal forms sodium hydroxide and liberates hydrogen gas along with lot of heat when reacts with water.

Na + H2O ⇨ NaOH + H2

Reaction of potassium metal with water: Potassium metal forms potassium hydroxide and liberates hydrogen gas along with lot of heat when reacts with water.

K + H2O ⇨ KOH + H2

Reaction of calcium metal with water: Calcium forms calcium hydroxide along with hydrogen gas and heat when reacts with water.

Ca + 2H2O ⇨ Ca(OH)2 + H2

Reaction of magnesium metal with water: Magnesium metal reacts with water slowly and forms magnesium hydroxide and hydrogen gas.

Mg + 2H2O ⇨ Mg(OH)2 + H2

When steam is passed over magnesium metal, magnesium oxide and hydrogen gas are formed.

Mg + H2O ⇨ MgO + H2

Reaction of aluminium metal with water: Reaction of aluminium metal with cold water is too slow to come into notice. But when steam is passed over aluminium metal; aluminium oxide and hydrogen gas are produced.

2Al + 3H2O ⇨ Al2O3 + 2H2

Reaction of zinc metal with water: Zinc metal produces zinc oxide and hydrogen gas when steam is passed over it. Zinc does not react with cold water.

Zn + H2O ⇨ ZnO + H2

Reaction of Iron with water: Reaction of iron with cold water is very slow and come into notice after a long time. Iron forms rust (iron oxide) when reacts with moisture present in atmosphere.

Iron oxide and hydrogen gas are formed by passing of steam over iron metal.

3Fe + 4H2O ⇨ Fe3O4 + 4H2

Other metals usually do not react with water or react very slowly.

Reaction of metals with dilute Mineral Acids

  • When a metal reacts with dilute acid, salts are formed. During this reaction hydrogen gas is evolved. In other words, when a metal is added to dilute acids, salt and hydrogen gas are formed.

  • Process
    Sulfate salts and chloride salts are formed when metals react with dilute sulfuric acid and hydrochloric acid.
  • Some metals reacts vigorously with dilute sulfuric acid or hydrochloric acid for example, potassium, sodium, lithium and calcium.
    Examples
    2Na + 2HCl —— > 2NaCl + H2
    Magnesium, zinc, iron, tin, lead and aluminum react normally with dilute acids. For example,
    Mg + H2SO4 —— > MgSO4 + H2
  • Which metals do not react with dilute acids ?
    Gold and silver
  • What is a dilute acid?
    Acids contain a large number of hydrogen atoms, but a dilute acid has less acidic characters because of the amount of hydrogen atoms. It does not cause much harm as compared to acids. It may cause a slight burn or irritation on touching the skin.
  • When alkali (base) reacts with metal, it produces salt and hydrogen gas.

Alkali + Metal ⇨ Salt + Hydrogen

Example: Sodium hydroxide gives hydrogen gas and sodium zincate when reacts with zinc metal.

2NaOH + Zn ⇨ Na2ZnO2 + H2

  • Sodium aluminate and hydrogen gas are formed when sodium hydroxide reacts with aluminium metal.

2NaOH + 2Al + 2H2O ⇨ 2NaAlO2 + 2H2

 

Displacement Reaction

Our Objective

To study a single displacement reaction with the help of iron nails and copper sulfate solution.

Theory

What is a displacement reaction?

Displacement reaction is a chemical reaction in which a more reactive element displaces a less reactive element from its compound.
Both metals and non-metals take part in displacement reactions.

Example : Reaction of iron nails with copper sulphate solution.

  • How is the chemical reactivity of metals linked with their position in the electrochemical series?
  • Chemical reactivity of metals is linked with their relative positions in the activity series.
  • Certain metals have the capacity to displace some metals from the aqueous solutions of their salts.
  • A metal placed higher in the activity series can displace the metal that occupies a lower position from the aqueous solution of its salt.
  • The displacement reaction is not limited to metals only. Even non-metals can take part in these reactions.

Examples are halogens. The activity series of halogen is F > Cl > Br > I.

Note the point:

An important thing to remember with single displacement reaction is that elements that form cations can only displace cations and elements that form anions can only displace anions.

Classification of single displacement reaction:

  • Anion replacement reaction
  • Cation replacement reaction

(A) Cation Replacement Reaction

In this reaction, one cation replaces another one from its solution. A cation is a positively charged ion or metal. All metal displacement reactions are cation replacement reactions.

Let us demonstrate some examples of cation replacement reactions.

  • If a strip of magnesium metal is placed in copper sulphate solution, the blue colour of copper sulphate disappears and the magnesium metal turns brown as the displaced copper is deposited on it. In the reactivity series, the position of magnesium is above that of copper. So it is more reactive than copper and it would displace copper from copper sulphate solution.

  • Also metals like Zn, Al, Pb, Fe, etc., displace Cu from copper salt solution. The chemical reaction can be written as:

  • Iron displaces lead from the aqueous solution of lead nitrate.

  • Similarly, metals above hydrogen in the reactivity series displace hydrogen from dilute acids.

For example :The metals such as potassium, sodium, lithium, etc. react more vigorously with dilute acids forming metal salts and hydrogen gases.

  • Metals such as aluminium, zinc, magnesium, iron, etc., react safely with dilute acids.

For example,

The reaction Zn with dil. H2SO4 is often used in the laboratory for the preparation of hydrogen gas.

If a less reactive metal is added to a salt solution of more reactive metal, nothing will happen.

(B) Anion Replacement Reactions

In this reaction, one anion replaces another one from its solution. An anion is a negatively charged ion or non-metal. For example, more reactive halogen replaces less reactive halogen from its solution.

  • Chlorine displaces bromine from an aqueous solution of sodium bromide.

  • Bromine Displaces iodine from potassium iodide solution.

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