Chemical Equations and Reactions Describing Chemical Reactions

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.

Chemical symbols seen in a chemical equation

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.

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.

Diatomic elements

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.

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.

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

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.

Looking at a balanced equation

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

Types of chemical reactions

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)

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

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)

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

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.

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.

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.

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)

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.

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 )

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.

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)

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

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