Chemical Equations and Reactions

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. • CH4 + 2 O2 --> CO2 + 2 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 • 2H2Oelectricity > 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 • C7H16 + 11O2  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