Chemical reactions Chapter 11

Section Overview 11.1: Describing Chemical Reactions 11.2: Types of Chemical Reactions 11.3: Reactions in Aqueous Solution

Describing Chemical reactions Section 11.1

Describing Chemical Reactions Chemical reactions are necessary to sustain life. In a chemical reaction, one or many substances change (reactants) into one or more new substances (products). Chemists use a chemical equation to convey as much as possible about what happens in a chemical reaction. Word Equations: Reactants are on the left and products on the right with an arrow between them that means yields, or gives, or reacts to produce. Reactants  Products Iron + oxygen  iron (III) oxide Hydrogen peroxide  water + oxygen

Describing Chemical Reactions Chemical Equations: A representation of a chemical reaction using the formulas for the reactants and products instead of words. Fe + O2  Fe2O2 Equations that just show the formulas and not the amount of each of the reactants and products are called skeleton equations. To add more information, you can also add the state of the of substances in parentheses. Fe(s) + O2(g)  Fe2O2(s) Most chemical reactions require a catalyst, which speeds up the reaction and is written above the arrow.

Writing a Skeleton Formula Example Problem: Hydrochloric acid and solid sodium hydrogen carbonate are shown before being placed into a beaker to react. The products formed are aqueous sodium chloride, water, and carbon dioxide gas. What is the skeleton formula for this reaction?

Writing a Skeleton Equation Example Problem: Hydrochloric acid and solid sodium hydrogen carbonate are shown before being placed into a beaker to react. The products formed are aqueous sodium chloride, water, and carbon dioxide gas. What is the skeleton formula for this reaction? Solution: Write the formulas for everything given and put into a skeleton formula. Hydrochloric acid: HCl(aq) Water: H2O(l) Carbon dioxide gas: CO2(g) Sodium chloride: NaCl(aq) Sodium hydrogen carbonate: NaHCO3(g) NaHCO3(s) + HCl(aq)  NaCl(aq) + H2O(l) + CO2(g)

Balancing Chemical Equations A balanced equation is a chemical equation in which both sides have equal numbers of atoms of each element. To balance an equation, you place coefficients, small whole numbers that are placed in front of the formulas in an equation in order to balance it. To write a balanced chemical equation, first write the skeleton equation. Then use coefficients to balance the equation. Occasionally, a skeleton equation may already be balanced.

Balancing Chemical Equations Determine the correct formulas for all the reactants and products. Write a skeleton equation. Determine the number of atoms each side of the equation. Count a polyatomic ion as a single unit if it appears unchanged on both sides. Balance the elements one at a time using coefficients. Begin with elements that only appear once on each side. Never balance by changing subscripts! Check each atom or polyatomic ion to be sure they are balanced. Make sure the coefficients are in the lowest possible ratio.

Balancing Chemical Equations Example Problem 1: Hydrogen and oxygen react to form water. Write a balanced equation for the reaction.

Balancing Chemical Equations Example Problem 1: Hydrogen and oxygen react to form water. Write a balanced equation for the reaction. Solution: Skeleton Formula First H2 + O2  H2O H = 2 H=2 O = 2 O=1 Skeleton isn’t balanced, so start adding coefficients. 2H2 + O2  H2O H=4 H=2 O=2 O=1 Still un balanced 2H2 + O2  2H2O H=4 H=4 O=2 O=2

Balancing Chemical Equations Example Problem 2: Balance the following equation Al + O2  Al2O3

Balancing Chemical Equations Example Problem 2: Balance the following equation Al + O2  Al2O3 Solution: Al + O2  Al2O3 4Al + O2  2Al2O3 Al =1 Al=2 Al=4 Al=4 O=2 O=3 O=2 O=6 2Al + O2  Al2O3 4Al + 3O2  2Al2O3 Al=2 Al=2 Al=4 Al=4 O=2 O=3 O=6 O=6 2Al + O2  2Al2O3 Al=2 Al=4 O=2 O=6

Types of Chemical reactions Section 11.2

Classifying Chemical Reactions The five types of chemical reactions are combination, decomposition, single-replacement, double-replacement, and combustion. Combination Reactions: A chemical change in which one or more substances react to form a new single substance. R + S  RS 2K + Cl2  2KCl Decomposition Reaction: A chemical change in which a single compound breaks down into two or more simpler compounds. RS  R + S 2HgO  2Hg + O2

Classifying Chemical Reactions Single-Replacement Reaction: A chemical change in which one element replaces a second element in a compound. T + RS  TS + R 2K + 2H2O  2KOH + H2 Whether one metal will replace another depends upon the reactivities of the two metals. A reactive metal will replace any metal listed below it in an activity series.

Classifying Chemical Reactions

Classifying Chemical Reactions Double Replacement Reaction: A chemical change involving an exchange of positive ions between two compounds. R+S- + T+U-  R + U- + T+S- Na2S + Cd(NO3)2  CdS + 2NaNO3 Combustion Reaction: A chemical change in which an element or a compound reacts with oxygen, often producing energy in the form of heat and light. CXHY + (x+y/4)O2  xCO2 + (y/2)H2O CH4 + 2O2  CO2  2H2O

Writing Equations for Combination Reactions Example Problem: Complete the equation for the following reaction Cu + S  (two reactions possible)

Writing Equations for Combination Reactions Example Problem: Complete the equation for the following reaction Cu + S  (two reactions possible) Solution: Two reactions are possible because copper is a transition metal and can be Cu+ or Cu2+. Cu + S  CuS (balanced) OR Cu + S  Cu2S 2Cu + S  Cu2S (balanced)

Writing Equations for Decomposition Reactions Example Problem: Write a balanced chemical equation for the following decomposition reaction H2O 

Writing Equations for Decomposition Reactions Example Problem: Write a balanced chemical equation for the following decomposition reaction H2O  Solution: H2O  H2 + O2 H=2 H=2 O=1 O=2 2H2O  H2 + O2 H=4 H=2 O=2 O=2 2H2O  2H2 + O2 (balanced)

Writing Equations for Single Replacement Reactions Example Problem: Write a balanced equation for each reaction. Zn + H2SO4  Cl2 + NaBr 

Writing Equations for Single Replacement Reactions Example Problem: Write a balanced equation for each reaction. Zn + H2SO4  Cl2 + NaBr  Solution: Zn + H2SO4  ZnSO4 + H2 (balanced) Zinc is more reactive than hydrogen, so it takes its place. b. Cl2 + NaBr  NaCl + Br2 Chlorine is more reactive than bromine (reactivity decreases down a group). Cl2 + 2NaBr  2NaCl + Br2 (balanced)

Writing Equations for Double-Replacement Reactions Example Problem: Write a balanced chemical equation for each double replacement reaction CaBr2 + AgNO3  (a precipitate of silver bromide is formed) FeS + HCl  (hydrogen sulfide gas (H2S) is formed)

Writing Equations for Double-Replacement Reactions Example Problem: Write a balanced chemical equation for each double replacement reaction CaBr2 + AgNO3  (a precipitate of silver bromide is formed) FeS + HCl  (hydrogen sulfide gas (H2S) is formed) Solution: CaBr2 + AgNO3  AgBr + Ca(NO3)2 CaBr2 + 2AgNO3  2AgBr + Ca(NO3)2 (balanced) FeS + HCl  H2S + FeCl2 FeS + 2HCl  H2S + FeCl2 (balanced)

Writing Equations for Combustion Reactions Example Problem: Write balanced equations for the complete combustion of these compounds Benzene (C6H6) Ethanol (CH3CH2OH)

Writing Equations for Combustion Reactions Example Problem: Write balanced equations for the complete combustion of these compounds Benzene (C6H6) Ethanol (CH3CH2OH) Solution: Remember, oxygen is the other reactant. Products are CO2 and H2O. C6H6 + O2  CO2 + H2O 2C6H6 + 15O2  12CO2 + 6H2O (balanced) CH3CH2OH + O2  CO2 + H2O CH3CH2OH + 3O2  2CO2 + 3H2O (balanced)

Reactions in aqueous soultion Section 11.3

Net Ionic Equations Most chemical reactions take place in water, in which some of the reactants dissociate, or separate, into cations and anions when they dissolve in water. These ions can be used to write a complete ionic equation, an equation that shows dissolved ionic compounds as dissociated free ions. An ion that appears on both sides of the equation and is not directly involved in the reaction is called a spectator ion. When you rewrite the equation without the spectator ions, the net ionic equation remains.

Net Ionic Equations AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq) Complete Ionic Equation Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq)  AgSl(s) + Na+(aq) + NO3-(aq) The nitrate and sodium ions remain unchanged on both sides and are spectator ions therefore they can be removed. Net Ionic Equation Ag+(aq) + Cl-(aq)  AgSl(s) Be sure in the balanced equation that not only the atoms are equal, but also the charges.

Net Ionic Equations Example Problem: Aqueous solutions of iron (III) chloride and potassium hydroxide are mixed. A precipice of iron (III) hydroxide forms. Write the balanced net ionic equation for the reaction.

Net Ionic Equations Example Problem: Aqueous solutions of iron (III) chloride and potassium hydroxide are mixed. A precipice of iron (III) hydroxide forms. Write the balanced net ionic equation for the reaction. Solution: Fe3+(aq) + Cl- (aq) + K+(aq) + OH-(aq)  Fe(OH)3(s) + K+(aq) + Cl-(aq) Spectator ions = Cl and K Fe3+(aq) + OH-(aq)  Fe(OH)3(s) Fe3+(aq) + 3OH-(aq)  Fe(OH)3(s) (balanced)

Predicting the Formation of a Precipitate Sometimes mixing solutions of two ionic compounds can result in the formation of an insoluble salt called a precipitate. Whether or not a precipitate forms, depends on the solubility of the new compounds that form. You can predict the formation of a precipitate by sing the general rules for solubility of ionic compounds. If a compound is soluble, it will not form a precipitate. If a compound is insoluble, it will form a precipitate.