1. 1 Chapter 4 Types of Chemical Reactions and Solution Stiochiometry

2. 2 Preview the contents of this chapter will introduce you to the following topics:  Water, Nature of aqueous solutions, types of electrolytes, and dilution  Types of chemical reactions: precipitation, acid base reactions and oxidation-reduction reaction  Stoichiometry of reactions and balancing the chemical equations

3. 3 4.1 Water, the Common Solvent Water is one of the most important substances on earth: –Cooling – engine, nuclear power plants, and many. –Transportation …etc. Water dissolve many different substances, e.g. salts, sugar, and many other To understand this process, we need to consider the nature of water: –As molecule, is H 2 O –Shape is V-shape with angle 105 o –Band type of each O – H is covalent band and polar [electrons are not equirdlntty shared] –Polarity is polar with +ve charges of hydrogen and -ve on oxygen. This Polarity of water gives it the greatest ability to dissolve compounds.

4. 4 4.1 Water, the Common Solvent This Polarity of water gives it the greatest ability to dissolve compounds. Figure 4.2 shows schematic ionic solid dissolving in water. This process is called "Hydration". Note: The stronger the ion-water attraction, the higher the solubility. Therefore, not all the solid have the same solubility [chapter 11].

5. 5 4.1 Water, the Common Solvent Water also dissolve non-ionic substances e.g. alcohols – "polar" and compatible structure to water: "like – dissolve – like" Figure 4.3 many substances don't dissolve in water e.g. animal fats "non-polar"

6. 6  Solution – composite of solute (substance to be dissolved) + Solvent (Major substance e.g. water).  Solute:  Solvent: 4.2 The Nature of Aqueous Solutions: Strong and Weak Electrolytes & non–Electrolyte 4 dissolves in water (or other “solvent”) 4 changes phase (if different from the solvent) 4 is present in lesser amount (if the same phase as the solvent) 4 retains its phase (if different from the solute) 4 is present in greater amount (if the same phase as the solute) SolutionSolventSolute Soft drink (l) Air (g) Soft Solder (s) H2OH2O N2N2 Pb Sugar, CO 2 O 2, Ar, CH 4 Sn

7. 7  One major property for characterizing aqueous solutions is its "Electrical conductivity" or its ability to conduct an electric current. Three solutions are observed: non electrolytes, weak, electrolytes and strong electrolyte. figure 4.4  The basis for conductivity properties of solutions was first correctly identified by Svante Arrhenius (1859 – 1927) 4.2 The Nature of Aqueous Solutions: Strong and Weak Electrolytes & non–Electrolyte nonelectrolyte weak electrolyte strong electrolyte An electrolyte is a substance that, when dissolved in water, results in a solution that can conduct electricity.

8. 8  Arrhenius postulated that the electric current depend directly on the number of ions present. 4.2 The Nature of Aqueous Solutions: Strong and Weak Electrolytes & non–Electrolyte Strong Electrolyte – 100% dissociation NaCl (s) Na + (aq) + Cl - (aq) H2OH2O Weak Electrolyte – not completely dissociated CH 3 COOH CH 3 COO - (aq) + H + (aq) Nonelectrolyte does not conduct electricity? No cations (+) and anions (-) in solution C 6 H 12 O 6 (s) C 6 H 12 O 6 (aq) H2OH2O

9. 9 4.3 The Composition of Solutions: To perform stoichiomctric calculations at any chemical reactions you must know two things: –The nature of the reaction – exact forms of chemical in solutions. –The amounts of chemicals – "concentration" concentration of a solution can be described in many different ways, %, molar, molal, mole. faction.. etc. We will consider here the unit "molar"(M) or molarity: "the method used to prepare molar solutions". Molarity (M) = moles of solute per volume of solution in liters:

10. 10 4.3 The Composition of Solutions: Notes:  For ionic systems the solution prepared will contain the number of moles prepared but for ionic species its different story e.g. 1.0M NaCl contains 1.0 mole NaCl or more accurate 1.0 mole of Na + and 1.0 mole of Cl -.  Molarity can be used to determine number of moles per certain volumes where: Moles = Liters of solution x Molarity. Example 4.3 Gives the concentration of each type of ion in the solutions of 0.50M Co(NO 3 ) 2 Example 4.4 Calculate the number of moles of Cl - ions in 1.75 L of 1.0x10 -3 M ZnCl 2

11. 11 4.3 The Composition of Solutions: Dilution: it is the procedure to get low concentrated solution (diluted) from a high concentrated one. Moles of solute after dilution = moles of solute before dilution. (M x V) after = (M x V) before. Example 4.7 What volume of 16 M sulfuric Acid must be used to prepare 1.5 L of a 0.10M H 2 sO 4 solution? Standard Solution: Solution used in chemical analysis. it has accurately known concentration

12. 12 4.4 Types of Solution Reactions Millions of possible chemical reaction needs system for grouping them into classes. The commonly used by chemists: Precipitation reactions AgNO 3 (aq) + NaCl(aq)  AgCl(s) + NaNO 3 (aq) Acid-base reactions NaOH(aq) + HCl(aq)  NaCl(aq) + H 2 O(l) Oxidation-reduction reactions Fe 2 O 3 (s) + Al(s)  Fe(s) + Al 2 O 3 (s)

13. 13 4.5 Precipitation Reactions Precipitate – insoluble solid that separates from solution molecular equation ionic equation net ionic equation Na + and NO 3 - are spectator ions PbI 2 Pb(NO 3 ) 2 (aq) + 2NaI (aq) PbI 2 (s) + 2NaNO 3 (aq) precipitate Pb 2+ + 2I - PbI 2 (s) Pb 2+ + 2NO 3 - + 2Na + + 2I - PbI 2 (s) + 2Na + + 2NO 3 -

14. 14 Simple Rules for Solubility of Aq. Solutions 1.Most nitrate (NO 3  ) salts are soluble. 2.Most alkali (group 1A) salts and NH 4 + are soluble. 3.Most Cl , Br , and I  salts are soluble (NOT Ag +, Pb 2+, Hg 2 2+ ) 4.Most sulfate salts are soluble (Except BaSO 4, PbSO 4, HgSO 4, CaSO 4 ) 5.Most OH  salts are only slightly soluble (Except NaOH, KOH are soluble) 6.Most S 2 , CO 3 2 , CrO 4 2 , PO 4 3  salts are only slightly soluble, i.e., Not soluble. 4.5 Precipitation Reactions Exercise 4.8: Using the solubility rules in table 4.1, predict what will happen when the following pairs of solutions are mixed. a.KNO 3 (aq) & BaCl 2 (aq) b.Na 2 SO 4 (aq) & Pb(NO 3 ) 2 (aq) c.KOH(aq) & Fe(NO 3 )(aq)

15. 15 Writing Net Ionic Equations AgNO 3 (aq) + NaCl (aq) AgCl (s) + NaNO 3 (aq) Ag + + NO 3 - + Na + + Cl - AgCl (s) + Na + + NO 3 - Ag + + Cl - AgCl (s) Write the net ionic equation for the reaction of silver nitrate with sodium chloride. 1.Write the balanced molecular equation. 2.Write the ionic equation showing the strong electrolytes 3.Determine precipitate from solubility rules 4.Cancel the spectator ions on both sides of the ionic equation 4.6 Describing Reactions in Solution

16. 16 4.7 Stoichiometry of Precipitation Reactions The procedures for calculating quantities of reactants and products involved in chemical reaction. The following steps summarized the procedure: Step 1: Identify the present in the combined solution, and determine what reaction occurs. Step 2: write the balanced net ionic equation. Step 3: calculate the moles of reactants. Step 4: determine which reactant is limiting. Step 5: calculate the moles of product or product as required. Step 6: convert to grams or other units, as requires. Example 4.11 When aqueous solutions of Na 2 SO 4 and Pb(NO 3 ) 2 are mixed, PbSO 4 precipitates. Calculate the mass of PbSO 4 formed when 1.25L of 0.0500M Pb(NO 3 ) 2 and 2.00L of 0.0250M Na 2 SO 4 are mixed.

17. 17 4.8 Acid – Base Reactions: Acids Have a sour taste. Taste of vinegar is due to acetic acid. Citrus fruits contain citric acid. React with certain metals to produce hydrogen gas. React with carbonates and bicarbonates to produce carbon dioxide gas Have a bitter taste. Feel slippery. Many soaps contain bases. Bases

18. 18 4.8 Acid – Base Reactions: Arrhenius acid is a substance that produces H + (H 3 O + ) in water Arrhenius base is a substance that produces OH - in water

19. 19 4.8 Acid – Base Reactions: A Brønsted acid is a proton donor A Brønsted base is a proton acceptor acidbaseacidbase A Brønsted acid must contain at least one ionizable proton!

20. 20 4.8 Acid – Base Reactions: acid + base salt + water HCl (aq) + NaOH (aq) NaCl (aq) + H 2 O H + + Cl - + Na + + OH - Na + + Cl - + H 2 O H + + OH - H 2 O Describing Reactions in Solution They are also called neutralization reaction

21. 21 4.8 Acid – Base Reactions: x acid + y base salt + water They are also called neutralization reaction Main reaction is titration, the key terms are: Titrant - solution of known concentration used in titration Analyte - substance being analyzed Equivalence point - enough titrant added to react exactly with the analyte Endpoint - the indicator changes color so you can tell the equivalence point has been reached. a. Write the correct balanced acid–base reaction. b. Use the following equation: y. (M.V) acid = x. (M.V) base The neutralization reaction calculation:

22. 22 4.8 Acid – Base Reactions: In a titration a solution of accurately known concentration is added gradually added to another solution of unknown concentration until the chemical reaction between the two solutions is complete. Equivalence point – the point at which the reaction is complete Indicator – substance that changes color at (or near) the equivalence point Slowly add base to unknown acid UNTIL the indicator changes color 4.7

23. 23 4.9 Oxidation-Reduction Reactions: (electron transfer reactions) 2Mg (s) + O 2 (g) 2MgO (s) 2Mg 2Mg 2+ + 4e - O 2 + 4e - 2O 2- Oxidation half-reaction (lose e - ) Reduction half-reaction (gain e - ) 2Mg + O 2 + 4e - 2Mg 2+ + 2O 2- + 4e - 2Mg + O 2 2MgO

24. 24 4.9 Oxidation-Reduction Reactions:

25. 25 4.9 Oxidation-Reduction Reactions: Rules for Assigning Oxidation States 1. Oxidation state of an atom in an element = 0 2. Oxidation state of monatomic element ions = charge 3. Oxygen =-2 in covalent compounds (except in peroxides where it = -1) 4. H = +1 in covalent compounds 5. Fluorine = -1 in compounds 6.Sum of oxidation states = 0 in compounds Sum of oxidation states = charge of the ion IF 7 F = -1 7x(-1) + ? = 0 I = +7 Oxidation numbers of the elements in the following ?

26. 26 4.10 Balancing Oxidation – Reduction Equations Balancing by Half-Reaction Method (Acidic) 1.Write separate reduction, oxidation reactions 2.For each half-reaction: Balance elements (except H, O) Balance O using H2O Balance H using H+ Balance charge using electrons 3.If necessary, multiply by integer to equalize electron count 4.Add half-reactions 5.Check that elements and charges are balanced

27. 27 4.10 Balancing Oxidation – Reduction Equations Half-Reaction Method - Balancing in Base 1.Balance as in acid. 2.Add OH - that equals H+ ions (both sides!) 3.Form water by combining H +, OH - 4.Check elements and charges for balance

28. 28 4.10 Balancing Oxidation – Reduction Equations MnO 4 - + ClO 2 - → MnO 2 + ClO 4 -