Chapter 4 Type of Chemical Reactions and Solution Stoichiometric
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1 Chapter 4 Type of Chemical Reactions and Solution Stoichiometric Water, Nature of aqueous solutions, types of electrolytes, dilution.Types of chemical reactions: precipitation, acid-base and oxidation reactions.Stoichiometry of reactions and balancing the chemical equations.
2 Water is the dissolving medium, or solvent. Aqueous SolutionsWater is the dissolving medium, or solvent.
3 Figure 4. 1: (Left) The water molecule is polar Figure 4.1: (Left) The water molecule is polar. (Right) A space-filling model of the water molecule.
4 Figure 4.2: Polar water molecules interact with the positive and negative ions of a salt assisting in the dissolving process.
5 Some Properties of Water Water is “bent” or V-shaped.The O-H bonds are covalent.Water is a polar molecule.Hydration occurs when salts dissolve in water.
6 Figure 4.3: (a) The ethanol molecule contains a polar O—H bond similar to those in the water molecule. (b) The polar water molecule interacts strongly with the polar O—H bond in ethanol. This is a case of "like dissolving like."
7 A Solute dissolves in water (or other “solvent”) changes phase (if different from the solvent)is present in lesser amount (if the same phase as the solvent)
8 A Solvent retains its phase (if different from the solute) is present in greater amount (if the same phase as the solute)
9 General Rule for dissolution Like dissolve likePolar dissolve polar (water dissolve ethanol)Non-polar dissolve nonpolar (benzene dissolve fat)
10 Figure 4.5: When solid NaCl dissolves, the Na+ and Cl- ions are randomly dispersed in the water.
11 pure water, sugar solution ElectrolytesStrong - conduct current efficientlyNaCl, HNO3Weak - conduct only a small currentvinegar, tap waterNon - no current flowspure water, sugar solution
12 Figure 4.4: Electrical conductivity of aqueous solutions.
13 hydrochloric and sulfuric acid AcidsStrong acids - dissociate completely to produce H+ in solutionhydrochloric and sulfuric acidHCl , H2SO4Weak acids - dissociate to a slight extent to give H+ in solutionacetic and formic acidCH3COOH, CH2O
14 Bases Strong bases - react completely with water to give OH ions. sodium hydroxideWeak bases - react only slightly with water to give OH ions.ammonia
16 Figure 4.7: An aqueous solution of sodium hydroxide.
17 Figure 4.8: Acetic acid (HC2H3O2) exists in water mostly as undissociated molecules. Only a small percentage of the molecules are ionized.
18 MolarityMolarity (M) = moles of solute per volume of solution in liters:
19 Common Terms of Solution Concentration Stock - routinely used solutions prepared in concentrated form.Concentrated - relatively large ratio of solute to solvent. (5.0 M NaCl)Dilute - relatively small ratio of solute to solvent. (0.01 M NaCl): (MV)initial=(MV)Final
20 Figure 4.10: Steps involved in the preparation of a standard aqueous solution.
21 Figure 4.12: Dilution Procedure (a) A measuring pipet is used to transfer 28.7mL of 17.4 M acetic acid solution to a volumetric flask. (b) Water is added to the flask to the calibration mark. (c) The resulting solution is 1.00 M acetic acid.
22 Practice ExampleHow many moles are in 18.2 g of CO2?
23 Practice Example Consider the reaction N2 + 3H2 = 2NH3 How many moles of H2 are needed to completely react 56 g of N2?
24 Practice ExampleHow many grams are in mole of caffeine C8H10N4O2
25 Practice ExampleA solution containing Ni2+ is prepared by dissolving g of pure nickel in nitric acid and diluting to 1.00 L. A mL aliquot is then diluted to mL. What is the molarity of the final solution? (Atomic weight: Ni = 58.70).
26 Practice ExampleCalculate the number of molecules of vitamin A, C20H30O in 1.5 mg of this compound.
27 Practice ExampleWhat is the mass percent of hydrogen in acetic acid HC2H3O2
29 Simple Rules for Solubility 1. Most nitrate (NO3) salts are soluble.2. Most alkali (group 1A) salts and NH4+ are soluble.3. Most Cl, Br, and I salts are soluble (NOT Ag+, Pb2+, Hg22+)4. Most sulfate salts are soluble (NOT BaSO4, PbSO4, HgSO4, CaSO4)5. Most OH salts are only slightly soluble (NaOH, KOH are soluble, Ba(OH)2, Ca(OH)2 are marginally soluble)6. Most S2, CO32, CrO42, PO43 salts are only slightly soluble.
30 Figure 4.13: When yellow aqueous potassium chromate is added to a colorless barium nitrate solution, yellow barium chromate precipitates.
32 Describing Reactions in Solution (continued) 3. Net ionic equation (show only components that actually react)Ag+(aq) + Cl(aq) AgCl(s)Na+ and NO3 are spectator ions.
33 Performing Calculations for Acid-Base Reactions 1. List initial species and predict reaction.2. Write balanced net ionic reaction.3. Calculate moles of reactants.4. Determine limiting reactant.5. Calculate moles of required reactant/product.6. Convert to grams or volume, as required.Remember: n H+ = n OH-(MV) H+ = (MV) OH-
35 Key Titration TermsTitrant - solution of known concentration used in titrationAnalyte - substance being analyzedEquivalence point - enough titrant added to react exactly with the analyteEndpoint - the indicator changes color so you can tell the equivalence point has been reached.movie
36 Oxidation-Reduction Reactions (electron transfer reactions)2Mg (s) + O2 (g) MgO (s)2Mg Mg2+ + 4e-Oxidation half-reaction (lose e-)Reduction half-reaction (gain e-)O2 + 4e O2-2Mg + O2 + 4e Mg2+ + 2O2- + 4e-2Mg + O MgO
38 Redox ReactionsMany practical or everyday examples of redox reactions:Corrosion of iron (rust formation)Forest fireCharcoal grillNatural gas burningBatteriesProduction of Al metal from Al2O3 (alumina)Metabolic processescombustion
39 Rules for Assigning Oxidation States 1. Oxidation state of an atom in an element = 02. Oxidation state of monatomic element = charge3. Oxygen = 2 in covalent compounds (except in peroxides where it = 1)4. H = +1 in covalent compounds5. Fluorine = 1 in compounds6. Sum of oxidation states = 0 in compounds Sum of oxidation states = charge of the ion
41 Zn (s) + CuSO4 (aq) ZnSO4 (aq) + Cu (s) Zn Zn2+ + 2e-Zn is oxidizedZn is the reducing agentCu2+ + 2e CuCu2+ is reducedCu2+ is the oxidizing agentCopper wire reacts with silver nitrate to form silver metal.What is the oxidizing agent in the reaction?Cu (s) + 2AgNO3 (aq) Cu(NO3)2 (aq) + 2Ag (s)Cu Cu2+ + 2e-Ag+ + 1e AgAg+ is reducedAg+ is the oxidizing agent
42 IF7 F = -1 7x(-1) + ? = 0 I = +7 K2Cr2O7 NaIO3 Na = +1 O = -2 O = -2 Oxidation numbers of all the elements in the following ?F = -17x(-1) + ? = 0I = +7K2Cr2O7NaIO3Na = +1O = -2O = -2K = +13x(-2) ? = 07x(-2) + 2x(+1) + 2x(?) = 0I = +5Cr = +6
43 Balancing by Half-Reaction Method 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
44 Balancing by Half-Reaction Method (continued) 3. If necessary, multiply by integer to equalize electron count.4. Add half-reactions.5. Check that elements and charges are balanced.
45 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.
46 Balancing Redox Equations Example: Balance the following redox reaction:Cr2O Fe Cr3+ + Fe3+ (acidic soln)1) Break into half reactions:Cr2O Cr3+Fe Fe3+
47 Balancing Redox Equations 2) Balance each half reaction:Cr2O Cr3+Cr2O Cr3+Cr2O Cr H2OCr2O H Cr H2O6 e- + Cr2O H Cr H2O
48 Balancing Redox Equations 2) Balance each half reaction (cont)Fe Fe3+Fe Fe e-
49 Balancing Redox Reactions 3) Multiply by integer so e- lost = e- gained6 e- + Cr2O H Cr H2Ox 6Fe Fe e-
50 Balancing Redox Reactions 3) Multiply by integer so e- lost = e- gained6 e- + Cr2O H Cr H2O6 Fe Fe e-4) Add both half reactionsCr2O Fe H Cr Fe H2O
51 Balancing Redox Reactions 5) Check the equationCr2O Fe H Cr Fe H2O2 Cr 2 Cr7 O 7 O6 Fe 6 Fe14 H 14 H
52 Balancing Redox Reactions Procedure for Basic Solutions:Divide the equation into 2 incomplete half reactionsone for oxidationone for reduction
53 Balancing Redox Reactions Balance each half-reaction:balance elements except H and Obalance O atoms by adding H2Obalance H atoms by adding H+add 1 OH- to both sides for every H+ addedcombine H+ and OH- on same side to make H2Ocancel the same # of H2O from each sidebalance charge by adding e- to side with greater overall + chargedifferent
54 Balancing Redox Equations Multiply each half reaction by an integer so that# e- lost = # e- gainedAdd the half reactions together.Simply where possible by canceling species appearing on both sides of equationCheck the equation# of atomstotal charge on each side
55 Balancing Redox Reactions Example: Balance the following redox reaction.NH3 + ClO Cl2 + N2H4 (basic soln)1) Break into half reactions:NH N2H4ClO Cl2