1. Chemical Equations and Reactions
Describing Chemical Reactions

2. Indications of a chemical reaction
A chemical reaction is the process by which one or more substances are changed into one or more different substances.
A chemical equation represents with symbols and formulas (or with words), the identities and the relative amounts of the reactants and products.
The original substances are called reactants and are shown on the left side of the equation.
The substances formed are called products and are shown on the right side of the equation.

3. Chemical symbols seen in a chemical equation

4. Indications of a chemical reaction
There are several ways to tell if a chemical reaction has occurred.
Evolution of heat and/or light
Production of a gas
Formation of a precipitate(a solid that settles to the bottom of the test tube or a cloudiness that occurs)
Color change (may or may not indicate a chemical change)
Evolution of sound (may or may not indicate a chemical change)
Formation of a new substance
Products cannot be easily changed back into reactants.

5. Characteristics of a chemical equation
The equation must represent the facts. All reactants and products must be identified by chemical analysis.
The equation must contain the correct formulas for all substances involved.
Remember to use oxidation states when writing formulas.
Remember that some elements are diatomic and must have a 2 as a subscript when written in an equation.

6. Diatomic elements

7. Characteristics of a chemical equation
The law of conservation of mass must be observed at all times. Atoms may not be created or destroyed but may be rearranged to make new substances.
To equalize the number of atoms on both sides of an equation, coefficients are used. It is placed in front of the compound.
NEVER change the subscripts in a formula.

8. Characteristics of a chemical equation
Writing a word equation is helpful to organize the facts that are known.
EX. Methane gas reacts with oxygen in the presence of a spark to make carbon dioxide and water.
The reactants are known and the products are known. The condition required for this reaction is also known.

9. Characteristics of a chemical equation
Next a formula equation can be written.
CH4 + O2 --> CO2 + H2O
Note that there are 4 H atoms on the left but only 2 H atoms on the right. A coefficient can be placed in front of the water to equal them out. That will make 2 O atoms on the left but 4 O atoms on the right. Place a 2 in front of the O2 and now check for equal numbers of atoms.
CH O2 --> CO H2O

10. Significance of a chemical equation
The coefficients indicate the relative amounts of reactants and products. The lowest whole number ratio is shown.
The relative masses can be determined from the coefficients. Once done, the law of conservation of mass can be shown to be true.

11. Looking at a balanced equation

12. Types of chemical reactions
There are five basic types of reactions. Not all reactions fall into these five categories but these are the most common kinds.
Synthesis reactions
Decomposition reactions
Single replacement reactions
Double replacement reactions
Combustion reactions
complete
incomplete

13. Types of chemical reactions

14. Synthesis reactions
Are also known as composition reactions or as direct combination reactions
Occur when 2 or more elements or small compounds combine to form 1 larger compound.
A + X --> AX (may or may not have subscripts)

15. Examples of synthesis reactions
reactions with sulfur to form sulfides
Fe + S --> Fe2S3
reactions with oxygen to form oxides
S + O2 --> SO2
reactions of metals with halogens to form salts
2Na + Cl2 --> 2NaCl
reactions of oxides of active metals with water to form hydroxides
CaO + H2O --> Ca(OH)2

16. Decomposition reactions
Occur when a single compound breaks down into two or more simpler substances.
Are the opposite of synthesis reactions.
Usually energy must be added to cause these to occur.
AX --> A + X(may or may not have subscripts)

17. Examples of decomposition reactions
decomposition of binary compounds
2H2O electricity > 2H2 + O2 (electrolysis)
decomposition of metal carbonates
CaCO3 --> CaO + CO2
decomposition of metal hydroxides
Ca(OH)2 --> CaO + H2O
decomposition of metal chlorates
2KClO3 --> 2KCl + 3O2
decomposition of acids
H2CO3 --> H2O + CO2

18. Single replacement reactions
Occur when an active element replaces a less active element.
The activity series table (p. 266) is required to predict whether a reaction will occur or not.
An element high on the table will replace an element in a compound that is lower than it is, i.e. lithium will replace lead in a compound.
Metals will replace metals; nonmetals will replace nonmetals.
Have a net ionic equation which does not include spectator ions.

19. Examples of single replacement reactions
Li + MgCl2  LiCl + Mg
Mg + ZnSO4  MgSO4 + Zn
Co + H2O  N.R.
Zn + HCl  ZnCl2 + H2
Ag + ZnBr2  N.R.

20. Double replacement reactions
Occur when the electropositive elements in two compounds switch places.
Usually occur when the reactants are in aqueous solution.
Require a driving force to determine whether or not they occur.
Formation of a precipitate (solubility table)
Formation of a gas (CO2, SO2, NO2, SO3)
Formation of water
Have a net ionic equation which does not include spectator ions.

21. Examples of double replacement reactions
MgCO3(aq) + 2HCl(aq)  MgCl2(aq) + H2O(l) + CO2(g)
2KCl(aq) + Pb(NO3)2(aq)  2KNO3(aq) + PbCl2(s)
NaNO3 + KCl  N.R. (no driving force observed)

22. Combustion reactions Occur in the presence of oxygen.
Produce carbon dioxide and water as products if one of the reactants is an organic compound and the combustion is complete.
Produce carbon monoxide and water if the combustion is incomplete. You would need to be told if this was the case.

23. Examples of combustion reactions
C7H O2  7CO2 + 8H2O
CH4 + 2O2  CO2 + 2H2O
C6H12O6(aq) + 6O2(g) → 6CO2(aq) 6H2O(l )
C2H5OH(aq) + 2O2(g) incomplete→ 2CO(g) + 3H2O(l )

24. More types of reactions
Acid-base reactions are also called neutralization reactions because they produce a salt and water.
Precipitation reactions are those which produce a precipitate as a product (double replacement reactions).
Oxidation-reduction reactions are those in which two of the reacting atoms have their charges changed. One is oxidized and one is reduced.

25. Examples of acid-base reactions
HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
H2SO4(aq) + KOH(aq)  K2SO4(aq) + H2O(l)
HC2H3O2(aq) + Ba(OH)2(aq)  Ba(C2H3O2)2(aq) + H2O(l)

26. Examples of precipitation reactions
Ba(NO3)2(aq) + Na2SO4(aq)  2NaNO3(aq) + BaSO4(s)
2KCl(aq) + Pb(NO3)2(aq)  2KNO3(aq) + PbCl2(s)

27. Examples of redox reactions
Li + MgCl2  LiCl + Mg
Li goes from 0 to +1 and is oxidized and Mg goes from +2 to 0 and is reduced
Mg + ZnSO4  MgSO4 + Zn
Mg goes from 0 to +2 and is oxidized and Zn goes from +2 to 0 and is reduced
Zn + HCl  ZnCl2 + H2
Zn goes from 0 to +2 and is oxidized and H goes from +1 to 0 and is reduced