Chem 110 General Principles of Chemistry

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1 Chem 110 General Principles of Chemistry Chapter 3 (Page 88) Aqueous Reactions and Solution Stoichiometry In this chapter you will study chemical reactions that take place between substances that are dissolved in solutions. Here we will focus on aqueous solutions in which the reactants are dissolved in water. These are the most common solutions you will encounter in this course. We begin by defining terms that are used to describe solutions and why solutions are so often used as a medium in which to carry out reactions. You will learn what happens to ionic compounds when they dissolve in water and you will learn more about acids and bases and the reactions between them. One of the goals here is to learn what causes chemical reactions to occur in solution, and you will discover that in many cases we can actually predict the outcome of such reactions. You will also study another important class of reactions called oxidation-reduction, which can be viewed as reactions in which electrons are transferred from one substance to another. LEARNING OBJECTIVES You should keep the following goals in mind. 1. To learn the meaning of the terms: solute, solvent dilute solution, saturated solution, unsaturated solution, supersaturated solution, solubility, electrolytes, non-electrolytes and reaction terminology. 1

2 2. To learn what happens to ionic solutes when they dissolve in water and, to learn how to write chemical equations to represent the changes that take place. 3. To learn how molecular acids and bases form ions by reacting with water and how a dynamic equilibrium is able to account for the low concentrations of ions in solutions of weak acids and bases. 4. To learn how to write chemical equations for the reactions of electrolytes in aqueous solutions. 5. To learn which factors cause a metathesis reaction and how to predict the products of a metathesis reaction. 6. To revise assigning oxidation numbers to atoms in a chemical formula. 7. To be able to identify oxidation and reduction, and to identify oxidizing and reducing agents in a reaction. 8. To learn how to work problems that deal with the stoichiometry of reactions in solution, especially those that deal with the experimental technique of titration. Terms Here are some of the terms that you should be familiar with in order to understand the contents of Chapter 3 properly. The meaning of these terms can be found in the Appendix E (Glossary), on Page 1152 at the end of the text. The words/expressions are: solvent solute concentration solubility dilute solution concentrated solution 2

3 unsaturated solution saturated solution supersaturated solution dissociation nonelectrolyte electrolyte strong electrolyte weak electrolyte metathesis reaction double displacement reaction double replacement reaction precipitation reactions acid-base neutralization salt net ionic equation molecular equation ionic equation spectator ion acids strong acid acidic anhydride strong base bases basic anhydride weak acid weak base dynamic equilibrium chemical equilibrium oxidation oxidizing agent redox reaction reduction reducing agent oxidation state oxidation-reduction reaction oxidation number disproportionation reaction Water plays an active role in dissolving ionic compounds because it consists of polar molecules that are attracted to the ions. When an ionic compound dissolves in water, the ions dissociate from each other and become solvated by water molecules. Because the ions are free to move, their solutions conduct electricity. Water also dissolves many covalent substances with polar bonds. It interacts with some H-containing molecules so strongly it breaks their bonds and dissociates them into H + (aq) ions and anions. In water, the H + ion is bonded to water forming H 3 O + (hydronium ). A molecular equation for an aqueous ionic reaction shows undissociated substances. A total ionic equation shows all soluble ionic compounds as separate, solvated ions. Spectator ions appear unchanged on both side of the equation. By eliminating the spectator ions you see the actual chemical change in a net ionic equation. 3

4 PRECIPITATION REACTIONS (page 92) Precipitation reactions involve the formation of an insoluble ionic compound from two soluble ones. They occur because electrostatic attraction among certain pairs of solvated ions are strong enough to cause their removal from solution. Such reactions can be predicted by noting whether any possible ion combinations are soluble. Using Solubility Guidelines (Table 3.1- Page 94) to Predict Precipitation Reactions. Three steps can be used to predict whether a precipitate will form: 1. Note the ions present in the reactants 2. Consider the possible cation anion combinations 3. Decide whether any of these combinations is insoluble Refer to page 94 and attempt Sample Exercise 3.1 before you look at the solution and follow it up with the Practice Exercises. ACID BASE REACTIONS (Page 97) Acid-base (neutralization) reactions occur when an acid (an H + yielding substance) and a base (an OH - yielding substance) reacts and the H + and OH - ions form a water molecule. Strong acids and bases (Table 3.2) dissociate completely in water whereas weak acids and bases dissociate only slightly (equilibrium reactions). In a titration, a known concentration of one reactant is used to determine the concentration of the other. An acid-base reaction can also be viewed as the transfer of a proton from an acid to a base. 4

5 Acid-Base Reactions with Gas Formation (Page 101) An ionic gas forming reaction is an acid-base reaction in which an acid transfers a proton to a polyatomic ion (carbonate, bicarbonate, sulfite, bisulfite), forming a gas that leave the reaction mixture. OXIDATION REDUCTION REACTIONS (Page 103) When one reactant has a greater attraction for electrons than another, there is a net movement of electron charge, and a redox reaction takes place. Electron gain (reduction) and electron loss (oxidation) occur simultaneously. The redox process is tracked by assigning oxidation numbers (Section 3.4 on page 103) to each atom in a reaction. The species that is oxidized (contains an atom that increases in oxidation number) is the reducing agent; the species that is reduced (containing an atom that decreases in oxidation number) is the oxidizing agent. Note: You should already be familiar with the assignment of oxidation numbers. Some Examples of Common Types of Redox Reactions 1. Combination Reaction Combination model equation A + B AB Two reactants form one product 5

6 Examples of combination reaction S(s) + O 2 (g) SO 2 (g) The overall net charge on both sides of the equation is zero. Note: Not all combination reactions are redox reaction. When SO 2 dissolves in water to produce H 2 SO 3, one of the acids responsible for acid rain, there is no change in oxidation state of any of the elements. SO 2 (g) + H 2 O(l) H 2 SO 3 (aq) 2. Decomposition Reactions Decomposition model equation XY X + Y One reactant goes to two or more products 2KClO 3 (s) 2KCl(s) + 3O 2 (g) Many explosive reactions are decomposition reactions. 2NH 4 ClO 4 (s) N 2 (g) + Cl 2 (g) + 2O 2 (g) + 4H 2 O(g) Ammonium perchlorate is a rocket fuel. Read about the air bag, a safety feature in motorcars. Note: Not all decomposition reactions are redox reactions for example the decomposition of calcium carbonate does not result in the change of oxidation numbers. CaCO 3 (s) CaO(s) + CO 2 (g) Read about the uses of CaCO 3, CaO and CO 2 in the Chemical industry. 6

7 3. Metathesis reactions (Exchange reactions) page 95 These reactions generally involve swapping ions two compounds in solution react to form two new compounds. Here are some examples: AgNO 3 (aq) + NaCl(aq) NaNO 3 (aq) + AgCl(s) BaCl 2 (aq) + K 2 SO 4 (aq) 2KCl(aq) + BaSO 4 (s) 3Ca(NO 3 ) 2 (aq) + 2H 3 PO 4 (aq) 6HNO 3 (aq) + Ca 3 (PO 4 ) 2 (s) 4. Single Displacement Reactions Single displacement model equation E + AB EB + A Two reactants two products One element + one compound goes to one compound + one element See the activity series of the metals on page 8 (Table 1) in this worksheet. 5. Disproportionation Reactions In disproportionation reactions the same substance is oxidized and reduced. In the reaction (a) 3NO(g) + heat N 2 O(g) + NO 2 (g) The nitrogen in NO is oxidized to and is also reduced to (b) 3Cl 2 (g) + 6OH - (aq)+ heat ClO 3 - (aq) + 5Cl - (aq) + 3H 2 O(l) The Cl in chlorine is oxidized to and is also reduced to 7

8 Table 1: Activity series of the Metals Li Most reactive K Ba More active metals are more easily oxidised Ca than less reactive metals Na Mg Reactivity decreases as you go down the series Al Zn Fe Cd Co Ni Sn Pb H 2 All metals above hydrogen displace hydrogen from acids (eg Zn(s) + 2HCl(aq) ZnCl 2 (aq) + H 2 (g)) Cu Hg Ag Au Least reactive More reactive metals are more easily oxidized than less reactive metals. More reactive metals displace less reactive metals from the salts of less reactive metals. For example Mg is more reactive than Cu therefore Mg will displace Cu from soluble copper salts eg Mg (more reactive metal) + CuSO 4 MgSO 4 + Cu (less reactive metal displaced from CuSO 4 ) 8

9 SOLUTION STOICHIOMETRY AND CHEMICAL ANALYSIS (Page 114) A common laboratory technique applicable to precipitation, acid-base and redox reactions is titration. The key point in a titration is the equivalence point, which is assessed with the aid of an indicator. Titration data can be used to establish the molarity of a solution, called standardization of a solution or to provide other information about samples analysed. Preparing Solutions by Dilution (page 112) Whenever water is added to a solution, it becomes diluted. i.e. the molarity (or mass %) decreases. The amount of solute remains constant as the solution is diluted. Thus we have: Moles of solute before dilution = Moles of solute after dilution n before dilution = n after dilution But M = n/v n = M x V MiVi = MfVf i= initial solution f= final solution Acid Base Titration Using titration data to establish concentration of acids and bases. Study Sample Exercise 3.15 on page 117. Attempt the Practice Exercises. Redox Titration: Attempt the Practice Exercise at the top of page

10 Additional Problems You are expected to go through the following problems from your textbook: Chapter 3. Nos. 3.9, 3.10, 3.11, 3.12, 3.18, 3.20, 3.32, 3.37, 3.39, 3.45, 3.66, Past Exam Questions Multiple Choice Questions 1. Identify a neutralization reaction: A) AgNO 3 (aq) + HCl(aq) AgCl(s) + HNO 3 (aq) B) 2HI(aq) + H 2 O 2 (aq) I 2 (s) +2H 2 O(l) C) NaOH(aq) + HCl(aq) NaCl(aq) + H 2 O(l) D) MgO(s) + H 2 O (l) Mg(OH) 2 (s) (1) 2. The oxidation state for the element in the reactant that is oxidised in the equation below is: 4Al(s) + 3O 2 (g) 2Al 2 O 3 (s) A) 0 B) -2 C) +2 D) +3 (1) 10

11 3. Identify the reducing agent in the following reaction: NO 3 - (aq) + S 8 (s) NO 2 (g) + SO 4 2- (aq) A) NO 3 - B) S 8 C) NO 2 D) SO 4 2- (1) 4.1 Study the reaction between chromium chloride and sodium hydroxide and answer the questions that follow. CrCl 3 (aq) + 3NaOH(aq) Cr(OH) 3 + 3NaCl i) Using the rules for solubility, predict the solubility of the products. (1) ii) Write the net ionic reaction for this equation. (1) 4.2 The following reaction occurs in a basic medium: Cl 2 (g) ClO - (aq) + Cl - (aq) Balance the redox equation using the half reaction method. Show your working. (4) 11

Chap. 4 AQUEOUS RXNS. O H δ+ 4.1 WATER AS A SOLVENT 4.2 AQUEOUS IONIC REACTIONS. Page 4-1. NaOH(aq) + HCl(g) NaCl(aq) +H 2 O

Chap. 4 AQUEOUS RXNS. O H δ+ 4.1 WATER AS A SOLVENT 4.2 AQUEOUS IONIC REACTIONS. Page 4-1. NaOH(aq) + HCl(g) NaCl(aq) +H 2 O Chap. AQUEOUS RXNS.1 WATER AS A SOLVENT Describe solution composition in terms of molarity Describe strong and weak electrolyte solutions, including acids and bases Use ionic equations to describe neutralization