Form 3 Chemistry – COMPOUNDS OF METALS

A. METAL OXIDE

            An oxide is a compound of oxygen with another element.

Preparation of oxides:

There are two ways of preparing metal oxides

 

a) DIRECT METHOD

It involves direct combination of metal with oxygen

         

1. INDIRECT METHOD
   

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It involves thermal decomposition of a salt of carbonate, hydroxide (I) nitrates

               

CLASSIFICATION OF METAL OXIDES

1. BASIC OXIDES

            Are oxides of metal when reacts with acid forms salts

            Examples are MgO, CaO, Na₂O, K₂O            

                 

2. ACIDIC OXIDES

            Are oxides of non metals and when dissolved in water form solution

            Examples are CO₂, NO₂, SO₂

           

3. AMPHOTERIC OXIDES

These are oxides which have both basic and acidic properties

(H neutralize both acidic and Alkali to form salt)

Examples are ZnO, Al₂O₃ and PbO

4. NEUTRAL OXIDES

            Are oxides which do not react with either acid or base (Have no basic/acidic properties)

            Examples are H₂O, CO, NO, N₂Cl

5. PEROXIDES

            When elements burn in excess air, peroxides are formed

            Examples are N₂O₂, H₂O₂

6. MIXED OXIDES

            These are oxides of two simple oxides

            Examples are Fe₃O₄ – Tri iron tetra oxide 

                                           (Mixture of Fe O and Fe₂O₃)

SOLUBILITY.

 

USES OF OXIDES

I. Used in preparation of salts in the laboratory

    

II. Used in formation of slag

III. Used as a drying agent

IV. Used in manufacture of motors

 V. Oxides react with acid to form salt and water

The following are the reaction equations when the following oxides react with hydroxides

B. METAL HYDROXIDE

Very active metals (K, Na and Ca) dissolves in water forming hydroxide soluble oxides dissolves in water to form hydroxide (Alkali)

Soluble oxides dissolve in water to form hydroxide (alkali)

                

Caustic soda and caustic potash are used in the manufacture of SOAP.

Lime water is used to detect CO₂ gas. The rest of the hydroxides are insoluble.

PREPARATION OF INSOLUBLE HYDROXIDES

i) Dissolve a soluble salt in water i.e. CU(N0₃)₂

ii) Dissolve a soluble hydroxide in water i.e. Ca(OH)₂

iii) Mix them together

   Cu(NO₃)₂ + Ca(OH)₂ → Cu(OH)₂ + Ca(NO₃)₂

iv) Separate the two by filtration.

NB: The filtrate is Ca(NO₃)₂ and the residue is Cu(OH)₂

PROPERTIES OF HYDROXIDE

i) Have bitter taste

ii) Neutralize acids to form salt and water

Solubility of hydroxide in water

• The nature of hydroxide of metals varies according to position of the metal in the reactivity series.
  

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The hydroxide of sodium, calcium and potassium are the strongest because they dissociate completely in water

          

Ammonium hydroxide is weak base because it does not dissociate completely in water

                               

Question:

By using diagram and details explain/ show how you would prepare

i) Zinc hydroxide

ii) Magnesium hydroxide

Solution:

i) Zinc hydroxide (Zn(OH)₂)

– Dissolve soluble salt in water (Zn(N0₃)₂).

– Dissolve soluble hydroxide in water (NaOH).

– Mix them together.

Zn(NO₃)₂ + 2NaOH  Zn(OH)_2(s)  + NaNO_3(aq)

Separate by filtration

CARBONATE AND HYDROGEN CARBONATE

• Carbonates occur in many natural forms examples are chalk, limestone and marbles. Examples of carbonates that occur naturally are the carbonates of Zinc, iron, lead and Manganese.
  
• Sea animals have carbonates in their shells in the forms of calcium carbonates. The remains of these animals sink to the bottom of sea leading to chalk formation after thousands of years, through pressure the chalk hardens to form limestone
  

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General properties of carbonates:

       1. All carbonates are insoluble in water except those of sodium, potassium and ammonium.

       2. All carbonates give carbon dioxide on heating except those of sodium and potassium

          

       3. All carbonates give carbon dioxide with dilute acids

              

PREPARATION OF SODIUM/ POTASSIUM CARBONATES AND SODIUM/ POTASSIUM HYDROGEN CARBONATES:

a) A solution of sodium hydroxide (caustic soda) is saturated with carbon dioxide, giving a solution of sodium bicarbonates

                          

If a solution of potassium hydroxide (caustic potash) is saturated with carbon dioxide, a solution of potassium bicarbonate is formed

                               

a) This solution is divided into two equal portions, one portion is allowed to evaporate at room temperature, when crystals of sodium hydrogen carbonate (sodium bicarbonate) or potassium hydrogen carbonate (potassium bicarbonate) are obtained which are be filtered and dried.

b) An equal volume of the original sodium hydroxide or potassium hydroxide solution is added to the second portion. Sodium carbonate or potassium carbonate is formed

                                                       

The solution is evaporated down and allowed to cool crystals of sodium carbonate is formed 

 DIFFERENCES BETWEEN SODIUM CARBONATE AND SODIUM BICARBONATE

                               (Na₂CO₃)                                                                                (NaHCO₃)

 
 

All carbonates and hydrogen carbonates react with acids to form carbon dioxide, water and salt.

Dilute sulphuric acid react very little with calcium carbonate because the product calcium sulphate is insoluble, form a layer on the outside of the carbonates and stop further reaction.

Also dilute sulphuric acid and hydrochloric acid react little with lead II Carbonate because the product, the lead sulphate and lead II Chloride are insoluble, they form layer on the outside of the carbonates which cause to stop further reaction

• Nitric acid react readily with all carbonates because all nitric are soluble.
  

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TEST FOR CARBONATES AND HYDROGEN CARBONATE.

a) Add dilute nitric acid to a solution of a substance. Carbon dioxide gas is evolved, water and salts are formed.

b) i) Add magnesium sulphate to a substance, if carbonates are present white precipitates are formed.

ii) If hydrogen carbonates are present no precipitates are formed

c) Action of heat

Potassium carbonate and sodium carbonate do not decompose when heated. Other carbonate decompose to form an oxide and carbon dioxide

          

 

Preparation of sodium carbonate by SOLVARY PROCESS

1. Calcium carbonate is heated in a limekiln

                      

2. The calcium oxide is dissolved in water

           

3. The carbon dioxide is reacted with ammonia brine

           

As the solution cools, sodium hydrogen carbonates the least soluble combination of ions, precipitates out and removed

4. The sodium hydrogen carbonate is heated

           
5. The Ammonium chloride and calcium hydroxide are reacted

           

            NITRATES

Nitrates are important chemicals for industries purpose

Examples;

• potassium nitrate is used in gun powder
  
• Sodium nitrates occur in nature as salt petres
  
• Aluminium nitrate is useful in the manufacture of sulphuric acid as a fertilizer. 
  

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PREPARATION OF NITRATES

1. By neutralizing nitric acid with an alkali or base

2. by the action of nitric acid on a carbonate

3. by the action of metal on nitric acid copper

            (Only for magnesium, zinc, lead and copper)

           

PREPARATION OF NITRATES

1. All nitrates are soluble and form crystalline solids

2. Action of heat on nitrates

a) Sodium and potassium nitrates give nitrites

            

b) Many metal nitrates give the oxide

Example; i) Effect of heat on copper nitrate crystals when the blue crystals of copper nitrates are heated in a boiling tube they dissolve in their water of crystallization, prolonged heating gives brown fumes of nitrogen dioxide and the residue, copper oxide is a black.

           

        ii) Effect of heat on lead nitric crystals

• When heated they decrepitate (expand with a crackling noise) and melt. Prolonged heating gives brown fumes of nitrogen dioxide and a yellow residue of lead monoxide (litharge)
  

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c) Silver and mercury nitrates give the metal

                

d) Ammonium nitrates give nitrous oxide

           

3. The action of concentrated sulphuric acid

     – Concentrated sulphuric acid, heated with a nitrate liberates nitric acid

4. All nitrates except those of sodium and potassium are hydrated, sodium nitrate and calcium nitrates are deliquescent, and other nitrates are not.

Summary of chemical properties of Nitrates

         

TEST FOR NITRATES

1. Add concentrated sulphuric acid to the solid nitrate and warm Acid fumes (nitric acid) are involved. Add 3 – 4 copper turning and arm the mixture, the colour of the gas turns to deep brown.

2. The brown ring test is normally used to confirm the nitrates.

To a solution of nitrates are added a few drops of freshly prepared ferrous sulphate solution.

A few drops of concentrated sulphuric acid are added carefully and brown ring obtained at the junction between the liquid.

It is nitrate ferrous sulphate (FeSO₄NO) formed that

     

The nitrogen monoxide then reacts with more iron (II) sulphate to give the brown compound, which appears as a brown ring

USES OF NITRATES

1. Potassium nitrate is used as a food preservative; it is also used to make slow burning fuses.

2. Sodium nitrate is mainly used as a fertilizer

3. Ammonium nitrate is mixed with chalk, the mixture is known as NITROCHALK, widely used as fertilizer.

It is also used in manufacturing explosives.

4. Silver nitrate is used to make silver bromide and silver iodides chemicals used to make photographic film.

CHLORIDES

Chlorine can react with all metals to form chloride. Most common metals are attacked by dil. Hydrochloric acid to form chlorides

Example is Mg, Ca, Al, Zn, and Fe.

           

Those elements (metals) which are not attached by dilute carbonates or oxides of these metals react with dil.HCl hydrochloric acid to produce chloride          

   PREPARATION OF CHLORIDE

• Metallic chlorides can be prepared by direct or indirect or indirect methods.
  

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1. BY DIRECT METHOD

– Chloride can be prepared by direct action of chlorine on metals

                            

Iron (iii) chloride can be prepared by passing Cl₂ over a heated metal.

An iron wire is placed in hand glass tube and a steam of dry Cl₂ passes over it.

The reaction continues without further application of heat source and iron (iii) chlorides collect in a bottle as shown.

PREPARATION OF SOLUBLE CHLORIDE:

Soluble chlorides can be prepared by mixing dilute hydrochloric acid (Dil.HCL)

i) Oxides

ii) Hydroxides

iii) Carbonates

iv) Metals

Methods for preparation of soluble chloride are summarized below 

K          prepared by action of the oxide, hydroxide or carbonate

Na        dilute hydrochloric acid

Ca

Mg       Prepared by Action of the metal, oxide or carbonate on

Al         Dilute hydrochloric acid

Zn       

Fe

NB: Lead and silver chloride are only common insoluble chlorides

CHEMICAL PROPERTIES OF CHLORIDES

• To identify a chlorine in a liquid mixture add nitrates or silver nitrate (A white precipitate formed)
  

                 
  

• To identify a chloride from solid mixture, add concentrated sulphuric acid
  

  The gas that produces thick white fumes with ammonium solution is produced
  

                          
  

• Mix chloride with oxidizing agent (MnO₂)
  

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             A green wish yellow gas is involved 

                       

METAL SULPHATES

i) Preparation of soluble metal sulphates

Soluble sulphates can be prepared in the laboratory by dissolving a metal, a carbonate, a hydroxide or an oxide in dilute H₂SO₄

2. Preparation of insoluble sulphates

Insoluble sulphates can be prepared by adding dilute sulphuric acid to lead or barium ions

CHEMICAL PROPERTIES

1. All sulphates react with barium chloride to give a white precipitate which is insoluble in dilute HCl

Note. Sulphites and carbonates give precipitate which is soluble in dilute HCl.

2. The sulphates or iron (II), copper (II) and zinc (called green, blue and white vitriol respectively) decompose when heated strongly to give the oxide of the metal, sulphur dioxide, sulphur trioxide and water

    

Note. In case of copper (II) and Zinc Sulphates, no sulphur dioxide is produced

          

3. The sulphates of sodium, potassium, magnesium and calcium don’t decompose on heating

USES OF SULPHATES

1. Sodium sulphate

    Is used to treat wood pulp in the manufacture of paper

2. Anhydrous (CaSO₄) is used as a source of sulphur dioxide for the manufacture of sulphuric acid.

3. Ammonium sulphate is an important fertilizer

4. Aluminium sulphate

    Is used in the purification of sewage as it precipitates proteins

5. CuSO₄ (copper (II) sulphate) is used as fungicide i.e. destroyer of fungi on plants

Also it is used as an electrolyte for electroplating copper.