Chapter 4: Aqueous Reactions and Solution Stoichiometry

1. Chapter 4: Aqueous Reactions and Solution Stoichiometry Prof. Dr. Nizam M. El-Ashgar Chemistry Department, IUG

2. General Properties of Aqueous Solutions • A solution is a homogeneous mixture of two or more substances. • A solution in which water is the dissolving medium is called an aqueous solution.

3. Solute vs. Solvent • The substance present in the greatest quantity is called the solvent. • Water is considered the universal solvent because of its ability to dissolve many substances. • The other dissolved substances are called the solutes. • A solvent dissolves a solute. • When water is the solvent, the solution is called an aqueous solution.

4. Electrolytes • A substance whose aqueous solutions contain ions is called an electrolyte because it will allow electric current to flow through it. Example:NaCl • A substance that does not form ions in solution is called a nonelectrolyte. Example: C12H22O11

5. Ionic Compounds in Water • When ionic compounds dissolve in water the ions separate and are surrounded by water molecules. • Once the ions have separated, they can move and conduct electricity. In the crystalline form they cannot move or conduct electricity. • The solvation process helps stabilize the ions in solution and prevents the cations and anions from recombining.

6. Ionic Compound in Water • Water molecules are polar and have a d+ and d- side. • The more electronegative oxygen atoms are attracted to the positive ions. • The less electronegative hydrogen atoms are attracted to the negative ions.

7. Ionic Compounds: undergo dissociation: processby which many ionic substances dissolve in water, the solvent pulls the individual ions from the crystal and solvates them. - {NaCl+ H2O} Polar water molecule + + H2O NaCl(s) H2O Na+(aq) +Cl-(aq)

8. Hydration of sodium chloride

9. Molecular Compounds in Water • When a molecular compound dissolves in water, the solution usually consists of intact molecules (neutral) dispersed throughout the system. • Consequently, most molecular compounds are nonelectrolytes. A few exceptions include strong acids like HCl. Ionic (electrolyte) Molecular (non electrolyte)

10. Molecular (covalent) Compounds: mostlyinsolublegases, exceptpolar organicliquids containing O & N (polar: acids, bases, alcohols, etc.) Insoluble gases: NO2, CH4, CO2, O2, P2O5, N2, CO, etc. Polar Covalent {carbon (C) chains containing H,O or N}: CH3OH, C6H12O6, C6H5OH, etc. H2O C6H12O6(s) C6H12O6(aq) H2O CH3OH(l)CH3OH (aq) Dissolve without dissociating into ions {Methanol+Water}

11. Electrolytes and Nonelectrolytes • An electrolyteis a substance that dissociates into ions when dissolved in water. (eg. NaCl in water) • A nonelectrolytemay dissolve in water, but it does not dissociate into ions when it does so. (eg. sucrose in water)

12. Strong and Weak Electrolytes • Strong electrolytes are those solutes that exist in solution completely or nearly completely as ions. • Weak electrolytes are those solutes that exist in solution mostly in the form of molecules with only a small fraction in the form of ions. • A nonelectrolytedoes NOT dissociate in water. (tend to be molecular compounds)

13. Electrolytes: Strong and Weak • A strong electrolytedissociates completely when dissolved in water. HCl(g)H2O H+(aq) + Cl-(aq) • A weak electrolyteonly dissociates partially when dissolved in water. NH4OH(aq)H2O NH4+(aq) + OH-(aq) Acetic Acid, HC2H3O2 (aq) H2O H+(aq)+ C2H3O2-(aq) ions molecules

14. Electrolytes: a visual

15. Strong Electrolytes Are • Strong acids • Strong bases • Soluble Ionic Salts

16. Example Potassium Iodide, glucose, and hydrofluoric acid are examples of a _________, ________, and _________ electrolytes, respectively. • Weak, non-, strong • Strong, non-, weak • Strong, weak, strong • Strong, weak, weak • Weak, weak, and strong

17. Chemical Equilibrium • Chemical equilibrium is the state of balance in which the relative numbers of each type of ion or molecule in the reaction are constant over time. • Equilibrium is represented by a double arrow for the yield sign. This shows that the reaction is occurring in both directions. • Dynamic Equilibrium

18. Strong Electrolytes • Remember that soluble ionic compounds are strong electrolytes. • We consider ionic compounds as those composed of metals and nonmetals – such as NaCl, FeSO4 and Al(NO3)3 – or compounds containing the ammonium ions, NH4+ - such as NH4Br and (NH4)2CO3.

19. Sample Exercise 4.1 • The diagram below represents an aqueous solution of one of the following compounds: MgCl2, KCl, or K2SO4. Which solution does the drawing best represent? The diagram shows twice as many cations as anions, consistent with the formulation K2SO4.

20. Practice Exercise • If you were to draw diagrams like the previous one representing aqueous solutions of each of the following ionic compounds, how many anions would you show if the diagram contained six cations? a. NiSO46 b. Ca(NO3)2 12 c. Na3PO4 2 d. Al2(SO4)3 9

21. Types of Aqueous Reactions Precipitation A precipitate formed Acid-Base Reactions Acid and base are involved Oxidation Reduction (REDOX) Reactions These are reactions in which electrons are lost and gained during the reaction itself.

22. Precipitation Reactions • Reactions that result in the formation of an insoluble product are called precipitate reactions. • A precipitate is an insoluble solid formed by a reaction in solution.

23. Precipitation Reactions Aqueous solutions, reacting to produce a precipitate (an insoluble compound). Example: KI(aq) + Pb(NO3)2 (aq) Predict the solubility of compounds in reaction: Pb(NO3)2(aq)+ 2KI(aq) 2 KNO3(aq) PbI2(s) Precipitate (ppt) KI PbI2

24. Solubility • The solubility of a substance at a given temperature is the amount of substance that can be dissolved in a given quantity of solvent at a given temperature. • Most important solubility rule: Alkali metal ions and ammonium ions are soluble in water.

25. Solubility – ability to be dissolved • Soluble (aq) • a substance which easily dissolves in water • Slightly Soluble (s) • a substance that only dissolves a tiny bit in water • Insoluble (s) • substance that does NOT dissolve in water • substance remains separate from the H2O molecules (no interaction) • substances are insoluble when the ions attract so stronglythat they CANNOT be pulled apart by H2O molecules

26. Solubility: a visual

27. SOLUBILITY RULES: for Ionic Compounds (Salts) • All salts of alkali metals (IA) are soluble. • All NH4+ salts are soluble. • All salts containing the anions: NO3-, ClO3-, ClO4-, (C2H3O2-) are soluble. • All Cl-, Br-, and I- are solubleexcept forHg22+, Ag+, and Pb2+ salts. • All SO42- are solubleexcept forAg+, Hg22+, Pb2+, Ba2+, and Sr2+. • 6. All Sulfides, S2-, Sulfites SO32- and hydroxides OH-are insolubleexcept forGroup 1A,NH4+, Ca2+, Sr2+, Ba2+ • 7.Carbonates, CO32-and Phosphates, PO43- are insoluble except for alkali metal cations and NH4+ • 8. Most hydroxides are insoluble. Exceptions: Group IA hydroxides, Ca(OH)2, Sr(OH)2, Ba(OH)2 • ______________________________________________________________ • Note: All acids are aqueous solutions(thus soluble) (exception HF).

29. Sample Exercise 4.2 • Classify the following ionic compounds as soluble or insoluble in water: a. Sodium carbonate: Na2CO3, Soluble b. Lead (II) sulfate : PbSO4 Insoluble

30. Practice Exercise • Classify the following compounds as soluble or insoluble in water: • Cobalt (II) hydroxide: Insoluble • Barium nitrate: Soluble • Ammonium phosphate: Soluble

31. Solubility Rules - Practice • Are the following soluble or insoluble? • Be(C2H3O2)2 • MgS • BaSO4 • K3PO4 soluble insoluble insoluble soluble

32. Metathesis (Exchange) Reactions • Reactions in which positive ions and negative ions appear to exchange partners conform to the following generic equation: AX + BY  AY + BX Such reactions are called exchange reactions, double displacement reactions, or metathesis reactions.

33. Completing and Balancing Metathesis Equations Steps to follow • Use the chemical formulas of the reactants to determine which ions are present. • Write formulas for the products: cation from one reactant, anion from the other. Use charges to write proper subscripts. • Check your solubility rules. If either product is insoluble, a precipitate forms. • Balance the equation. Ways to Write Metathesis Reactions: • Molecular equation • Complete ionic equation • Net ionic equation

34. Double Displacement Reactions

35. Sample exercise 4.3 a) Predict the identity of the precipitate that forms when solutions of BaCl2 and K2SO4 are mixed. BaSO4 is insoluble and will precipitate from solution. KCl is soluble. b) Write the balanced chemical equation for the reaction.

36. Practice Exercise a) What compound precipitates when solutions of Fe2(SO4)3 and LiOH are mixed? Fe(OH)3 b) Write a balanced equation for the reaction. Fe2(SO4)3(aq) + 6 LiOH(aq) 2Fe(OH)3(s) +3Li2SO4(aq) c) Will a precipitate form when solutions of Ba(NO3)2and KOH are mixed? no (both possible products, Ba(OH)2 and KNO3, are water soluble)

37. Complete Ionic Equations • A molecular equation: is an equation showing the complete chemical formulas of the reactants and products without showing dissociated ions. Example: Pb(NO3)2(aq) + 2KI(aq)  PbI2(s)+ 2KNO3(aq) • A complete ionic equations: is an equation that shows all soluble strong electrolytes as ions. Example: Pb2+(aq) + 2NO3-(aq) + 2K+(aq) + 2I-(aq)  PbI2(s) + 2K+(aq) + 2NO3-(aq) • A net ionic equation: is an equation in which the spectator ions have been removed. • Example: Pb2+(aq) + 2I-(aq)  PbI2(s)

38. Ways of Expressing Precipitation Reactions There are three different: (1) Molecular Equations (2) Ionic Equations (3) Net Ionic Equations AgNO3 (aq) + KCl(aq) AgCl(s) + KNO3 (aq) Ag+ (aq) + NO3- (aq) + K+ (aq) + Cl- (aq)  AgCl(s) + K+ (aq) + NO3- (aq) Ag+(aq) + Cl-(aq) AgCl(s)

39. Sample Exercise 4.4 Write the net ionic equation for the precipitation reaction that occurs when solutions of calcium chloride and sodium carbonate are mixed.

40. Practice Exercise Write the net ionic equation for the precipitation reaction that occurs when aqueous solutions of silver (I) nitrate and potassium phosphate are mixed.

41. Acid-Base Reactions • Bases • Substances that produce OH− ions when dissolved in water (Arrhenius). • Ex NaOH Na+ + OH– • Proton Acceptor (L&B) • NH3 (g) • Acids • Substances that produce H+ ionswhen dissolved in water (Arrhenius). • Ex HNO3 H+ + NO3– • Proton donor (L&B) • HCl (g)

42. Acids Bases

43. Strong vs. Weak • Both strong & weak acids and bases exist • Strong acids & bases dissociate completely in solution • Exist as 100% ions • Strong electrolytes & conductors • Weak acids & bases only partially dissociate in solution • Few ions exits • Weak electrolytes & conductors

44. Strong vs. Weak Electrolytes Strong Acid: Weak Acid: HClHNO2

45. Bases The strong bases are : • Alkali metals (IA) Hydroxides • Heavy group 2A metal hydroxides: Barium Hydroxide Strontium Hydroxide Calcium Hydroxide • (Weaker: Ammonium, CalciumHydroxides)

46. Acids • Monoprotic acids contain one ionizable hydrogen. Example: HCl • Diprotic acids contain two ionizablehydrogens. Example: H2SO4 • Triprotic acids contain three ionizablehydrogens. Example: H3PO4 • Most multiprotic acids only lose one hydrogen to a reasonable extent. Citric acid

47. Properties of Acids • Acids taste sour. • Some acids also contain hydrogens that will not ionize. • Normally the hydrogen(s) that will ionize are listed at the front of the chemical formula. • Example: HC2H3O2 • Sometimes the hydrogen that will ionize is listed at the end of a COOH group. • Example: CH3COOH

48. Properties of Bases • Bases taste bitter. • Bases are substances that accept (react with ) hydrogen ions. • Bases produce hydroxide ions (OH-) when they dissolve in water. • Bases do not have to contain the hydroxide ion. Examples: NaOH vs. NH3 NaOH(aq) Na+ (aq) + OH- (aq)

49. Examoles of Strong Acids • You need to memorize the strong acids. HCl(aq) hydrochloric acid HBr (aq) hydrobromic acid HI (aq) hydroiodic acid HClO3 (aq) chloric acid HClO4 (aq) perchloric acid HNO3 (aq) nitric acid H2SO4 (aq) sulfuric acid

50. Strong Bases • You need to memorize the strong bases. Group 1 metal hydroxides: LiOH, NaOH, KOH, RbOH, and CsOH Group 2 metal hydroxides: Ca(OH)2, Sr(OH)2, and Ba(OH)2