CHAPTER 11

Chemical Equations

Click here to see reactions

Basic format and what happens:

Reactants → Products Bonds broken → bonds formed Atoms are not created or destroyed, but

rearranged

They can be read and written:

Word Equations – put equations in a word form

Iron + oxygen → Iron (III) Oxide Hydrogen Peroxide → Water + Oxygen  Chemical Equations – typically we use chemical symbols

and chemical formulas to write chemical equations.

Fe + O2 → Fe2O3

This would be a skeleton equation and does not show the relative amounts of reactants and products nor their states.

You might also see: Symbols (p 708) are used to indicate the Physical

State of a compound.(s) solid(l) liquid(g) gas(aq) aqueous – dissolved in water

Fe (s) + O2 (g) → Fe2O3 (s)

  Catalyst – speeds up the rate of a reaction, but is

not used up.

MnO2

H2O2 (aq) → H2O (l) + O2 (g)

Balancing a chemical reaction: To balance a reaction, first write the skeletal equation then add only coefficients so that the equation obeys the law of conservation of mass.

1) Determine the correct formulae of the reactants and products.2) Write the skeletal equation3) Determine the number of atoms of each element on both sides to determine inequalities.4) Balance one atom at a time using coefficients5) Make sure all coefficients are at their lowest ratio

Hints:*Leave polyatomic ions intact if they are on both sides*balance solo elements last*know the diatomic elements: H2, N2, O2, F2, Cl2, Br2, I2

Section 11.2

1. Combination Reactions (Synthesis) Two or more substances combine to form

a single substance Reactants are usually two elements or

two compounds Examples:

2K (s) + Cl2 (g) → 2KCl (s)

2S (s) + 3O2 (g) → SO2 (g)

2. Decomposition Reactions

A single compound is broken down into two or more products.

Rapid decomposition reactions producing gas and heat (explosions!)

  Examples:CaCO3 (s) → CaO (s) + CO2 (g)

2H2O (l) → 2H2 (g) + O2 (g)

3. Single Replacement Reactions

Atoms of one element replace the atoms of a second element in a compound

Mg (s) + 2HCl (aq) → MgCl2 (aq) + H2 (g)

Whether one metal will displace another metal is determined by the relative activities of the two metals.

In the following example, Mg must be more reactive than Cu for the reaction to take place.

Mg (s) + Cu(NO3)2 (aq) → MgNO3 (aq) + Cu (s)

4. Double Replacement Reactions

1) Involve an exchange of positive ions between 2 reacting compounds – usually between 2 ionic compounds

For a reaction to take place one of the following must occur: A precipitate must form

Na2S (aq) + Cd(NO3)2 (aq) CdS (s) + 2NaNO3 (aq) A gas must be produced

2NaCn (aq) + H2SO4 (aq) 2HCN (g) + Na2SO4 (aq) Water is produced

Ca(OH)2 (aq) + 2HCl (aq) CaCl2 (aq) + H2O (l)

5. Combustion Reactions

Oxygen is often one of the reactants CO2 and water are often products Examples:

O2 + CH4 H20 + CO2

Cellular respiration

Can you identify these reaction types?

How about these?

11.3 Describing Reactions in Aqueous Solutions

  There are three types of equations used to

describe reactions in solution.

1) Molecular equation – shows the overall reaction, but not necessarily in the actual form in solution.

Pb(NO3)2 (aq) + Na2SO4 (aq) PbSO4 (s) + 2NaNO3 (aq)

2) Complete Ionic Equation – represents all reactants or products that are soluble (aq) as ions. Pb2+ + 2NO3

- + 2Na+ + SO42- PbSO4 (s)+2Na+ +

2NO3-

   3) Net Ionic Equation – includes only those ions that react to form a solid or liquid. 

Pb2+ + SO42- PbSO4 (s)

  Spectator Ions - Na+ and NO3

- are spectator ions because they are not part of the reaction (do not form the solid or liquid) and remain dissolved.