CHEMICAL REACTIONS: PRECIPITATION and OXIDATION - REDUCTION

Learning Objectives.

By the end of this module you should be able to:

• Classify chemical reactions as decomposition, combination, displacement or exchange

• Recognise spectator ions in a chemical reaction and write ionic equations

• Use solubility rules to identify the products of precipitation reactions

• Describe the dissolving process and account for solubility in terms of solute and solvent interactions

• Recognise oxidation-reduction reactions by identifying electron transfer between reactants

• Write balanced half-equations and use these to write balanced equations for redox reactions

• Identify oxidants and reductants in redox reactions

• Determine and use oxidation numbers to identify redox processes

Reference: ESA Chapters 5, 21, 4

CLASSIFYING CHEMICAL REACTIONS

Chemical reactions can be classified in two ways; according to the way the reactant particles behave in relation to each other or according to changes that occur to the reacting particles.

If we focus on the number and nature of the reactants and products and on the type of transformation needed to convert the reactant(s) into the product(s), reactions can be classified as:

•     Combination (Synthesis) - elements or less complex compounds come together to form a single more complex compound

•    Decomposition - a compound breaks apart into either elements or less complex compounds

•    Displacement - a single element replaces another one in a compound

•    Exchange – ions from one compound switch places with ions from another compound to form two new compounds

Examples of different types of reactions

Decomposition

•      This is the simplest type of reaction.

•      A single compound is broken down into several different compounds.

Example 1:  When green crystals of copper(II) carbonate are heated, a gas (which can be identified as carbon dioxide) is released and black copper(II) oxide is left as a residue.

i.e.             CuCO3(s)  →  CuO(s)  +  CO2(g)

Since one reactant has visibly become two products, decomposition must have occurred.

Combination

•      This is the opposite of decomposition

•      It involves the joining of two or more elements or compounds to form a single compound.

Example: Hydrogen and chlorine react to produce hydrogen chloride H2(g)  +  Cl2(g)  →  2HCl(g)

In a combination reaction, there will be several reactants but only one product.

Displacement Reactions

•      In a displacement reaction, one element is displaced from a compound by another element in its elemental form.

•      The equation describing the reaction must have an element and a compound as reactants and a different element and compound as products.

Example: If zinc metal is placed in a solution of copper(II) sulfate, the blue solution is decolourised. Colourless zinc sulfate is formed and copper metal forms a brown deposit.

Zn(s)  +  CuSO4(aq)    →    Cu(s)  +  ZnSO4  (aq)

Exchange Reactions

•         An exchange reaction is one in which atoms or groups of atoms are interchanged between two compounds.

•         For ionic compounds, an exchange reaction can be identified when two reactants give two products by a change ofpartners between the anions and cations.

Example:     When solutions of barium chloride and sodium sulfate are mixed, a precipitate of barium sulfate

is formed and sodium chloride is left in solution.

BaCl2(aq)  +  Na2SO4(aq)  →  BaSO4(s)  +  2NaCl(aq)   Written as an ionic equation:  Ba2+(aq)  +  SO42-(aq) →  BaSO4(s)

Activity 5.1 Types of reactions

Questions

1.    List the number of reactants and the number of products found in the combination reaction (as shown in the table) and link this to the definition of a combination reaction.

2.    List the number of reactants and the number of products found in the decomposition reaction (as shown in the table) and link this to the definition of a combination reaction

3.    Explain how you would compare a displacement reaction to an exchange reaction.

Exercise 5.1

Identify the type of reaction shown in each of the following chemical equations. Justify your answers.

(a)        2KClO3(s)  →     2KCl(s) + 3O2(g)

(b)        AgNO3(aq) + KCl(aq) →  AgCl(s) + KNO3(aq)

(c)        2H2(g) + O2(g)  →   2 H2O(g)

(d)        2AgNO3(aq) + Cu(s)  →  Cu(NO3)2(aq) + 2Ag(s)

(e)        CaCO3(s)   → CaO(s)   + CO2(g)

(f)         2NaI(aq) + Cl2(g)  →   2NaCl(aq)   + I2(s)

Ionic Equations and Spectator Ions

When ionic compounds are dissolved in water, the anions and cations dissociate and so essentially exist independently. For example, when solid sodium chloride, NaCl(s), is dissolved in water we can write:

This implies that the ions have separated and dispersed through the solution so they will be surrounded by water molecules. A solution of sodium chloride is usually written as NaCl(aq) which is misleading as it does not show the separated ions. However, recognising the dissociation of ions can simplify the way equations for some reactions are written.

For example: For the reaction of sodium chloride and silver nitrate solutions

NaCl(aq)  +  AgNO3(aq)  →  AgCl(s)  +  NaNO3(aq)

We can consider the dissociation of the salts in solution and write the equation as:

Na+(aq)  +   Cl-(aq)  +  Ag+(aq)  +  NO3-(aq)  →  AgCl(s)  +  Na+(aq)   +  NO3-(aq)

Since nothing has happened to the Na+ and NO3- ions during the reaction they are known as spectator ions so they can be eliminated from the equation as follows:

Ag+(aq)  +  Cl-(aq)   →  AgCl(s)

This is known as a net ionic equation, or more often, an ionic equation. (Note that the formation of solid

AgCl implies a new compound as the ions have “changed partners” from those in the original solutions and the formation of the solid removes the ions from the solution).

Activity 5.2 – ionic equations

This activity is designed to help you understand ionic equations

Questions

1.    Consider the reactions modelled in the beakers drawn in the box

(a)   (i)    Which beakers contain solid, insoluble substances? Give reasons for your answers.

(ii)   Which beakers contain solutions of ionic substances? Give reasons for your answers.

(iii) Which of the reactions produces a gas? How do you know this?

(b)   Indicate which of the equations (i) to (iii) corresponds to the reactions 1, 2 or 3 above.

(i).     Zn (s) + 2HCl (aq) → ZnCl2  (aq) + H2 (g)

(ii).    Zn (s) + Cu(NO3)2  (aq) → Zn(NO3)2  (aq) + Cu (s)

(iii)   Zn(NO3)2  (aq) + Na2CO3  (aq) → ZnCO3  (s) + 2NaNO3  (aq)

(c)   In each of the reactions 1 to 3, there are ions present in the solutions that do not participate in the chemical reaction. In other words, they exist in the same form. both before and after the reaction.  These substances are called spectator ions. Use the equations (i) to (iii) to identify the spectator  ions for each reaction.

2.    In the table below, the equation for one of the previous reactions is written in three different ways.

Molecular Equation     Zn(s) + Cu(NO3)2(aq) → Zn(NO3)2(aq) + Cu(s)

Ionic Equation            Zn(s) + Cu2+ (aq) + 2NO3 −(aq) → Zn2+ (aq) + 2NO3 −(aq) + Cu(s)

Net Ionic Equation       Zn(s) + Cu2+ (aq) → Zn2+ (aq) + Cu(s)

(a)   (i)    Which equation represents the ionic substances as bonded together in the solution?

(ii)  Which equation represents the ionic substances as separate ions in solution?

(iii) Which equations is a better representation of how the species take part in the reaction?

(iv)  Compare the net ionic equation to the other two equations.

a.  What chemical species is missing in the net ionic equation?

b.  Explain why it is valid to remove this species from the equation.

(b)        Write ionic and net ionic equations for the remaining reactions in 1: Zn(s) + 2HCl (aq) → ZnCl2  (aq) + H2  (g)

Zn(NO3)2  (aq) + Na2CO3  (aq) → ZnCO3  (s) + 2NaNO3 (aq)