Chapter 4. Reactions in Aqueous Solution. Lecture Presentation. John D. Bookstaver St. Charles Community College Cottleville, MO

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1 Lecture Presentation Chapter 4 in Solution 2012 Pearson Education, Inc. John D. Bookstaver St. Charles Community College Cottleville, MO

2 Properties of Solutions Solute: substance in lesser quantity in a solution Solvent: substance in greater quantity in a solution Solution: solute + solvent (solute is DISSOLVED in a solvent) Homogenous: type of mixture = SOLUTION Heterogenous: type of mixture but is not a solution!

3 Electrolytic Properties Ability of a solution to conduct electricity. There must be a flow of electrically charged particles (ions)! Consider the figure below. Why only 1 bulb lights? Which substance is an electrolyte? Which is a non-electrolyte?

4 Electrolytes Electrolytes: a substance that dissociates or ionize when dissolved in water. What is meant by dissociates or dissociation??? When an ionic substance dissolves in water, the solvent pulls the individual ions from the crystal and solvates them. This process is called dissociation.

5 Strong, Weak, Non-Electrolytes A strong electrolyte dissociates completely when dissolved in water into its ions. A weak electrolyte only dissociates partially when dissolved in water. soluble molecular acids and bases A non-electrolyte does not dissociate in water. Molecular = covalent compounds!!

6 Strong Electrolytes Are. Strong acids Strong bases Soluble ionic compounds Know all 7 acids and 8 bases!

7 Examples of Strong Electrolytes Strong electrolytes: ions completely dissociate Strong acid: H 2 SO 4 (aq) 2 H + (aq) + SO 4 2- (aq) Strong base: Ba(OH) 2 (aq) Ba 2+ (aq) + 2 OH - (aq) Soluble ionic salts: NaCl (aq) Na + (aq) + Cl (aq) : The single arrow indicates that the Na + and Cl ions have no tendency to recombine to form NaCl.

8 Weak Electrolytes Weak Electrolytes: partial dissociation Example: acetic acid, CH 3 COOH CH 3 COOH (aq) H + (aq) + CH 3 COO (aq) They only produce a small degree of ionization in water (e.g., weak acids & weak bases). This balance produces a state of chemical equilibrium.

9 Nonelectrolyte Nonelectrolyte: no dissociation Soluble molecular compounds (that are not acids or bases): C 6 H 12 O 6, CH 3 OH

10 Precipitation that result in the formation of an insoluble product are called precipitation reactions. A precipitate is an insoluble solid formed by a reaction in solution. An Example: Pb(NO 3 ) 2 (aq) + 2KI (aq) PbI 2 (s) + 2KNO 3 (aq) bright yellow ppt

11 Solubility The solubility of a substance at a particular temperature is the amount of that substance that can be dissolved in a given quantity of solvent at that temperature. A substance with a solubility of less than 0.01 mol/l is regarded as being insoluble. Factors that affect rate of solubility - Pulverizing, Stirring, and Heating

12 Solubility Guidelines

13 Sample Exercise (a) nickel(ii) phosphate (Ni 3 (PO 4 ) 2 ) (b) sodium carbonate (Na 2 CO 3 ) soluble (c) lead(ii) chloride (PbCl 2 ) insoluble insoluble

Exchange (Metathesis) Rxns Also called double displacement reactions. This is when 2 ionic compounds exchange (transpose) cations and anions to form new compounds.

14 Exchange (Metathesis) Rxns Also called double displacement reactions. This is when 2 ionic compounds exchange (transpose) cations and anions to form new compounds. Look out for the exceptions for these rxns: See gas forming rxns. Look at solubility rules to determine state symbols. Some examples: AgNO 3 (aq) + KCl(aq) BaCl 2 (aq) + K 2 SO 4 (aq)

15 A molecular equation: Ionic Equations AgNO 3 (aq) + KCl(aq) AgCl(s) + KNO 3 (aq) An ionic equation: more accurately reflects the species that are found in the reaction mixture Ag + (aq) + NO 3 (aq) + K + (aq) + Cl (aq) AgCl(s) + K + (aq) + NO 3 (aq) Notice: Only the aqueous salts are written in the ion form because they dissociate in water! All other states remain unchanged.

16 Net Ionic Equation To form the net ionic equation, cross out anything that does not change from the left side of the equation to the right. The only things left in the equation are those things that change (i.e., react) during the course of the reaction. Those things that didn t change (and were deleted from the net ionic equation) are called spectator ions. Ag + (aq) + NO 3 (aq) + K + (aq) + Cl (aq) AgCl(s) + K + (aq) + NO 3 (aq)

17 Rules for Writing Net Ionic Equations 1. Write a balanced... molecular equation Refer to solubility chart for state symbols. 2. Dissociate all... strong electrolytes (the aqueous salts) 3. Cross out... spectator ions 4. Write the net ionic equation with the... species that remain

18 Write the net ionic equations for the following: K 2 CrO 4(aq) + Ba(NO 3 ) 2 (aq) BaCrO 4(s) + KNO 3(aq) AgNO 3(aq) + CaCl 2(aq) AgCl (s) + Ca(NO 3 ) 2(aq)

19 Acids Arrhenius acid: substances that increase the concentration of H + when dissolved in water Brønsted-Lowry acids: they are proton donors A proton in this case is the hydrogen ion (H + ) Molecules of different acids ionize to form different numbers of protons (H + )! Monoprotic acids: yield 1 proton per acid molecule HCl (aq) H + (aq) + Cl - (aq) HNO 3 (aq) H + (aq) NO 3 - (aq) Diprotic acids: yield 2 protons per acid molecule (a 2-step process) H 2 SO 4 (aq) H + (aq) + HSO 4 - (aq) step 1 HSO 4 - (aq) H + (aq) + SO 4 2- (aq) step 2 How many protons does acetic acid (CH 3 COOH), a weak acid donates?

Bases (Alkali) Arrhenius base: substances that increase the concentration of OH when dissolved in water. Brønsted-Lowry base: They are proton acceptors.

20 Bases (Alkali) Arrhenius base: substances that increase the concentration of OH when dissolved in water. Brønsted-Lowry base: They are proton acceptors. They want H + Bases dissociate to produce the hydroxide ion (OH - )! LiOH (aq) Li + (aq) + OH - (aq) red blue Compounds that do not contain OH - can also be bases!! They react with water to form OH - NH 3 (aq) + H 2 O (l) NH 4 + (aq) + OH - (aq)

21 Strong and Weak Acids and Bases Strong acids/ bases are strong electrolytes and are completely ionized in solution. Weak acids/bases are weak electrolytes and are partially ionized in aqueous solution.

22 Acid-Base In an acid-base reaction, the acid donates a proton (H + ) to the base.

23 Neutralization Acid + Base Salt + water Salt = an ionic compound Special type of metathesis rxn EXOTHERMIC! Some examples: HCl(aq) + NaOH(aq) NaCl(aq) + H 2 O(l) CH 3 COOH(aq) + NaOH(aq) CH 3 COONa(aq) + H 2 O(l)

24 Neutralization When a strong acid reacts with a strong base, the net ionic equation is HCl(aq) + NaOH(aq) NaCl(aq) + H 2 O(l) H + (aq) + Cl (aq) + Na + (aq) + OH (aq) Na + (aq) + Cl (aq) + H 2 O(l)

25 Neutralization When a strong acid reacts with a strong base, the net ionic equation is HCl(aq) + NaOH(aq) NaCl(aq) + H 2 O(l) H + (aq) + Cl (aq) + Na + (aq) + OH (aq) Na + (aq) + Cl (aq) + H 2 O(l) H + (aq) + OH (aq) H 2 O(l)

26 Gas Forming Exceptions to Metathesis rxns: Carbonates and bicarbonates react with ACIDS to form salt, water and CO 2 (g). HCl (aq) + NaHCO 3 (aq) NaCl (aq) + H 2 O (l) + CO 2 (g) Sulfite reacts ACIDS to form a salt, sulfur dioxide, and water: SrSO 3 (s) + 2HI(aq) SrI 2 (aq) + SO 2 (g) + H 2 O(l) Sulfides (S 2- ) reacts with ACIDS to form salt and H 2 S (g) 2 HCl (aq) + Na 2 S (aq) 2 NaCl (aq) + H 2 S (g) Rotten egg smell!

27 Knowledge about H 2 S

28 Oxidation-Reduction An oxidation occurs when an atom or ion loses electrons. A reduction occurs when an atom or ion gains electrons. One cannot occur without the other. These are called redox reactions! OIL: oxidation is loss RIG: reduction is gain

29 The Rules: Oxidation Numbers Elements in their elemental form have an oxidation number of 0. Na (s), O 2 (g), P 4 (s), S 8 (s), C (s), Ba (s) etc The oxidation number of a monatomic ion is the same as its charge. K+ = +1, S 2- = -2, Ca 2+ = +2, Fe 3+ = +3, Cl - = -1 etc Nonmetals tend to have negative oxidation numbers, although some are positive in certain compounds or ions. Oxygen has an oxidation number of 2, except in the peroxide ion, in which it has an oxidation number of 1. Hydrogen is 1 when bonded to a metal, +1 when bonded to a nonmetal. Nonmetals tend to have negative oxidation numbers, although some are positive in certain compounds or ions. Fluorine always has an oxidation number of 1. The other halogens have an oxidation number of 1 when they are negative; they can have positive oxidation numbers, however, most notably in oxyanions The sum of the oxidation numbers in a neutral compound is 0. The sum of the oxidation numbers in a polyatomic ion is the charge on the ion.

30 Some Examples What is the oxidation state (or number) of the elements in bold? a) P 2 O 5 b) H 2 O 2 c) NaH d) Cr 2 O 7 2- e) S 8 f) PO 4 3- g) LiCoO 2

31 Type of Single-displacement reaction! Oxidation of Metals One substance is being oxidized, the other is being reduced! Acids oxidize metals to form salt + H 2 (g) Zn(s) + 2 HBr (aq) ZnBr 2 (aq) + H 2 (g) Metal salts reacts with metals to form a salt and a displaced metal. Mn (s) + Pb(NO 3 ) 2 (aq) Mn(NO 3 ) 2 (aq) + Pb (s)

32 Redox Rxns Which substance is being oxidized and which is being reduced? Zn(s) + 2 HBr (aq) ZnBr 2 (aq) + H 2 (g) Mn (s) + Pb(NO 3 ) 2 (aq) Mn(NO 3 ) 2 (aq) + Pb (s)

33 Helps predict whether one species can oxidize another. The Activity Series The Rule: Any metal on the list can be oxidized by the ions of the elements below it!!! An Example: Cu (s) + 2 Ag + (aq) Cu 2+ (aq) + 2 Ag (s) The thought process: The metal is Cu. The ion is Ag +. Silver (ion) is below Cu in the list, thus oxidation of the Cu will occur!

34 Can the following reaction occur? The Activity Series Cu 2+ (aq) + 2Ag(s) X Cu(s) + 2Ag + (aq) The Rule: Any metal on the list can be oxidized by the ions of the elements below it!!!

35 Molarity Two solutions can contain the same compounds but be quite different because the proportions of those compounds are different. Molarity is one way to measure the concentration of a solution: Molarity (M) = moles of solute volume of solution in liters 2012 Pearson Education, Inc.

36 Mixing a Solution Goal: To create a solution of a known molarity 1. Weigh out a known mass of solute 2. Add solute to a volumetric flask 3. Add solvent to the line on the neck of the flask

37 Dilution Procedure for preparing a less concentrated solution from a more concentrated solution. Moles of solute before dilution = Moles of solute after dilution

38 Dilution The molarity of the new solution can be determined from the equation M c V c = M d V d, where M c and M d are the molarity of the concentrated and dilute solutions, respectively, and V c and V d are the volumes of the two solutions.

39 Dilution Problems 1.How much 2.0 M NaCl solution would you need to make 250 ml of 0.15 M NaCl solution? 2.What would be the concentration of a solution made by diluting 45.0 ml of 4.2 M KOH to 250 ml?

40 Using Molarities in Stoichiometric Calculations 1. How many grams of Ca(OH) 2 are needed to neutralize 25.0 ml of 0.1 M HNO 3? 2. What volume of M HCl is needed to react completely with mols of Pb(NO 3 ) 2?

41 Titration A common way to determine the concentration of a solution is by titration. The process involves combining a solution of known concentration (standard solution) with a solution of unknown concentration.

Indicator: used to determine the endpoint or equivalence point of the reaction represented by a

42 Titration Titration is an analytical technique in which one can calculate the concentration of a solute in a solution. Standard solution: a reagent solution of known concentration. Indicator: used to determine the endpoint or equivalence point of the reaction represented by a color change. Equivalence point (endpoint): the point at which both solutions are stoichiometrically equivalent.

43 Titrations Problem 1. One commercial method used to peel potatoes is to soak them in NaOH solution for a short time, the remove and spray the peel. The [NaOH] is typicaly 3 to 6 M, and the solution must be analyzed periodlically. In one such analysis, 45.7 ml of M H 2 SO 4 is required to neutralize 20.0 ml of NaOH. What is the [base]? 2. What is the [NaOH] if 48.0 ml neutralizes 35.0 ml of M H 2 SO 4?

Chapter 4 Aqueous Reactions and Solution Stoichiometry

Chapter 4 Aqueous Reactions and Solution Stoichiometry Chemistry, The Central Science, 11th edition Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten Chapter 4 and Solution Stoichiometry John D. Bookstaver St. Charles Community College Cottleville,