Unit 4 (Chapter 4): Aqueous Reactions & Solution Stoichiometry Chemistry, The Central Science, 10th edition Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten Unit 4 (Chapter 4): Aqueous Reactions & Solution Stoichiometry John D. Bookstaver St. Charles Community College St. Peters, MO  2006, Prentice Hall, Inc.

Salts: Ionic Solids: (metal-nonmetal) dissociate (dissolve) by separation into ions Electrolytes: ions in solution that conduct electricity

Non Weak Strong C11H22O11 CH3OH H2O NaOH HNO3 KCl CH3COOH HNO2 NH3 ions SOME ions ALL ions only molecules partially ionize HW p.159 #33 completely dissociate

Electrolytes: Strong, Weak, or Non? (ions conduct electricity) Compound metal-nonmetal nonmetals (Covalent) Ionic Molecular Acid (H____) Not Acid STRONG KBr CaI2 FeCl3 NaOH Ca(OH)2 (strong bases) STRONG (6) WEAK (& NH3) NON C11H22O11 C2H5OH H2O HCl, HBr, HI HNO3 H2SO4 HClO4 CH3COOH HNO2 HF HW p.157-159 #1,2,4,5,38

Precipitation Reactions Double Replacement: (precipitate) 2 KNO3(aq) + PbI2(s) 2 KI(aq) + Pb(NO3)2(aq)  precipitate: insoluble ionic compound (as predicted by solubility rules) Pb2+ I–

Solubility Rules ALWAYS Soluble Ions: * WS Solubility & NIE’s #1 Na+, K+, etc. group I (alkali metals) NH4+ ammonium NO3– nitrate HCO3– bicarbonate C2H3O2– (CH3COO–) acetate (ethanoate) ClO3–, ClO4– chlorate, perchlorate * WS Solubility & NIE’s #1 Common Precipitates form with: examples Ag+, Pb2+, Hg2+ (halogens) AgCl, PbI2 OH– (hydroxide) Cu(OH)2 CO32– (carbonate) CaCO3

HCl + H2O  H3O+ + Cl– NH3 + H2O  NH4+ + OH– + – + – O O Cl Cl Acid: proton (H+) donor Base: proton (H+) acceptor NH3 + H2O  NH4+ + OH– H + – H N O N O H H H H H H H H

Strength of Acids and Bases STRONG (complete dissociation) HA(aq) + H2O(l) H3O+(aq) + A–(aq)   B(aq) + H2O(l) BH+(aq) + OH–(aq) WEAK (partial dissociation) HA(aq) + H2O(l) H3O+(aq) + A–(aq)   B(aq) + H2O(l) BH+(aq) + OH–(aq) Why is HF considered a weak acid if it etches glass? H bond together and even form complexes like F- - H3O+

Strong Acids: Only 6 strong acids: Nitric (HNO3) Sulfuric (H2SO4) Hydrochloric (HCl) Hydrobromic (HBr) Hydroiodic (HI) Perchloric (HClO4) Strong Acids: HI + H2O  H3O+ + I– proton (H+) donors

ase Strong Bases: The strong bases are soluble hydroxides (OH–) of… Group 1 (Li,Na,K) CBS (Ca, Ba, Sr) Mg(OH)2 not as soluble Strong Bases: OH– + H3O+  H2O + H2O proton (H+) acceptors ase

Electrolytes: Strong, Weak, or Non? QUIZ!!! (at the bell) Electrolytes: Strong, Weak, or Non? Compound metal-nonmetal nonmetals (Covalent) Ionic Molecular Acid (H____) Quiz – Electrolytes (to start class) Not Acid STRONG STRONG (6) WEAK (& NH3) NON

Acid-Base Neutralization Reactions strong acid (H+A–) strong base (M+OH–) ionic compound (M+A–) water H2O (HOH) ACID + BASE SALT + WATER + – O Cl Na Na Cl H O H H H HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l) HW p.158 #40a

AgNO3(aq) + KCl(aq)  AgCl(s) + KNO3(aq) Molecular Equation The molecular equation lists the reactants and products in their molecular form. AgNO3(aq) + KCl(aq)  AgCl(s) + KNO3(aq)

Ionic Equation AgNO3(aq) + KCl(aq)  AgCl(s) + KNO3(aq) In the ionic equation all strong electrolytes (strong acids, strong bases, and soluble ionic salts) are dissociated into their ions. This more accurately reflects the species that are found in the reaction mixture. AgNO3(aq) + KCl(aq)  AgCl(s) + KNO3(aq) Ag+(aq) + NO3–(aq) + K+(aq) + Cl–(aq)  AgCl(s) + K+(aq) + NO3–(aq)

Net Ionic Equation Ag+(aq) + NO3–(aq) + K+(aq) + Cl–(aq)  Cross out Spectator Ions that do not change (same state & same charge) from the left side of the equation to the right. The only species left are those things that react (change) during the course of the reaction. Ag+(aq) + NO3–(aq) + K+(aq) + Cl–(aq)  AgCl(s) + K+(aq) + NO3–(aq) NIE: Ag+(aq) + Cl–(aq)  AgCl(s)

Balanced Net Ionic Equations comp – diss – cross – net – bal Write a complete molecular equation. Dissociate all strong electrolytes (aq) . Cross out spectators (same charge & state) Write the net ionic equation with the species that remain and balance it. (solubility rules)

Balanced Net Ionic Equations comp – diss – cross – net – bal + 2– 2+ – + – (NH4)2SO4 + Ba(NO3)2 → NaOH + MgBr2 → BaSO4 + NH4NO3 Ba2+ + SO42– → BaSO4 + – 2+ – + – NaBr + Mg(OH)2 Mg2+ + 2 OH– → Mg(OH)2(s) HW p.158 #21

Neutralization Reactions When a Strong Acid reacts with a Strong Base, the net ionic equation is… H+ + OH–  H2O HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l) H+ + Cl– + Na+ + OH–  Na+ + Cl– + H2O Salt is (metal/nonmetal) is dissolved (aq)

Neutralization Reactions When a Weak reacts with a Strong, the net ionic equation is… HX + OH–  X– + H2O HF(aq) + KOH(aq)  KF(aq) + H2O(l) HF + Na+ + OH–  Na+ + F– + H2O Salt is (metal/nonmetal) is dissolved (aq) HW p.159 #40 (finish)

Balanced Net Ionic Equations comp – diss – cross – net – bal + 2– 2+ – + – (NH4)2SO4 + Ba(NO3)2 → BaSO4 + NH4NO3 Ba2+ + SO42– → BaSO4(s) + – + – HF(aq) + KOH(aq)  KF(aq) + H2O(l) HF + OH–  F– + H2O WS Solubility & NIE’s #2

Gas-Forming Reactions H2 Demo (M0) (H+) (M+) (gas) Single Rep: Metal + Acid Metal Ion + H2 Ex: Zn(s) + H2SO4(aq) ZnSO4(aq) + H2(g) NIE: Zn(s) + 2 H+(aq) Zn2+(aq) + H2(g) + 2– 2+ 2– (gas) H2O(l) + CO2(g) CO2 Demo (H+) (CO32–) Double Rep: Acid + Carbonate Salt + Ex: HCl(aq) + CaCO3(s) CaCl2(aq) + H2O(l) + CO2(g) NIE: 2 H+(aq) + CaCO3(s) Ca2+(aq) + H2O(l) + CO2(g) H2CO3(aq) (or Bicarbonate) (HCO3–) (decomposes immediately) H2 Demo – small bit of Zn in large test tube, then add but of 6 M H2SO4, hold burning splint for test. CO2 - Demo HW p. 159 #43 CH3COOH + NaHCO3  CH3COONa + H2O + CO2

Oxidation-Reduction Reactions (REDOX) video clip (One cannot occur without the other) LEO says GER lose e-, charge up….gain e-, charge down (Video – OxRed I)

Oxidation Numbers Is it a redox reaction? To find out… assign oxidation numbers* (or oxidation states) to each element in a reaction. check if any oxidation states changed (↓ reduced , ↑ oxidized) *oxidation numbers of elements describe electrons that would be lost or gained IF the compound was 100% ionic. of ions show electrons transferred IN an ionic compound *charges

Assigning Oxidation Numbers All elements are 0. (all compounds are 0) Monatomic ion is its charge. (Ex. Na+ ion) Most nonmetals tend to be negative, but some are positive in certain compounds or ions. (Ex. SO3) O is −2 always, but in peroxide ion is −1 (O2–2). H is +1 with nonmetals, −1 with a metals. F is always −1. other halogens are −1, but can be positive, like in oxyanions. Ex. ClO3– or NO3– or SO42–

Oxidation Numbers The sum of the ox. #’s in a neutral compound is 0.ex: NaCl = 0 The sum of the ox. #’s in a polyatomic ion is the charge on the ion. Ex: PO4=-3 Calculate the oxidation number of each: Sulfur in… SO3 Chromium in… K2Cr2O7 Nitrogen in… NH4+ Cobalt in… [CoCl6]3–

Classifying REDOX Reactions All rxns (but…NOT double replacement) Synthesis A + B → AB 2 → 1 (0 0 → +/–) Decomposition AB → A + B 1 → 2 (+/– → 0 0) Single Replacement AB + C → A + CB (+/– 0 → 0 +/–) Combustion CxHy + O2 → CO2 + H2O (–/+ 0 → +/– +/–) Zn in acid releases H2 gas 26

Single Replacement (REDOX) silver ions oxidize copper metal What is reduced? Cu(s) + 2 Ag+(aq)  Cu2+(aq) + 2 Ag(s) X Cu2+(aq) + 2 Ag(s)  Cu(s) + 2 Ag+(aq)

Activity Series of Metals increasing ease of oxidation Cannot displace H+ from acid to make H2(g)

Writing REDOX Reactions Write an overall equation for the synthesis of calcium sulfide from its elements. Then write the RED and OX half-equations to identify the REDOX process. +2 –2 Ca + S  CaS OX: Ca0  Ca+2 + 2 e– RED: 2 e– + S0  S–2 Ca + S  CaS

Writing REDOX Reactions Write an overall equation for the decomposition of aluminum oxide into its elements. Then write the RED and OX half-equations to identify the REDOX process. +3 –2 Al2O3  Al + O2 3 ( ) OX: 2 O–2  O20 + 4 e– 6 O–2  3 O20 + 12 e– 4 ( ) RED: 3 e– + Al+3  Al0 12 e– + 4 Al+3  4 Al0 2 Al2O3  4 Al + 3 O2

Writing REDOX Reactions Write the net ionic equation for the reaction of solid zinc in a solution of hydrochloric acid. comp – diss – cross – net – bal +1 –1 +2 –1 Mg(s) + HCl(aq)  MgCl2(aq) + H2(g) Mg + H+  Mg+2 + H2 2 Classify the reaction in two ways. Single-Replacement and Redox

Mg + 2 H+  Mg2+ + H2(g) ox red What is red & what is ox? WS 5c #1-2 Title: Reaction of magnesium with acid. Caption: The bubbles are due to hydrogen gas.

+ Solutions: homogeneous mixtures. ( _____ throughout) same solvent is present in greatest abundance. solute dissolved in/by solvent + same

Molarity units: mol/L or mol∙L–1 moles of solute (mol) Molarity (M) = Molarity (M) is a measure of the concentration of a solution. moles of solute (mol) liters of solution (L) Molarity (M) = units: mol/L or mol∙L–1 What is the molarity of a solution with 29.2 g of sodium chloride in 250. mL of water? mol/L like g/mol 1 mol NaCl 58.44 g NaCl 2.00 M NaCl 0.500 mol NaCl 29.2 g NaCl x = = 0.250 L 35

Preparing a Solution 1-mass solute 2-add solvent, swirl to dissolve WS #1-2 Conc. Calc’s 1-mass solute 2-add solvent, swirl to dissolve 3-add solvent to mark HW p.160 #60, 67 36

WS Concentration & Dilutions #1 1 mol NaHCO3 84.01 g NaHCO3 1 L NaHCO3 5.00 g NaHCO3 x x = 0.100 mol NaHCO3 0.595 L NaHCO3 #2 1.20 mol CuSO4 1 L CuSO4 159.62 g CuSO4 0.275 L CuSO4 x x = 1 mol CuSO4 52.7 g CuSO4

Dilution M1V1 = M2V2 1-calc M1V1=M2V2 2-pipet V1 from concentrated 3-fill to mark w DI 38

Dilution WS #3-4 Dilutions M1V1 = M2V2 What volume of water must be added to prepare 2.0 L of 3.0 M CuSO4 from a stock solution of concentration 8.0 M ? 39

WS Concentration & Dilutions M1V1 = M2V2 M1V1 = M2V2 #3 (12.0 M)V1 = (1.25 M)(500 mL) V1 = 52.1 mL (0.0521 L) #4 M1V1 = M2V2 (2.50 M)V1 = (0.200 M)(250 mL) V1 = 20.0 mL (0.0200 L)

Solution Stoichiometry Rxn: A(aq) + 2 B(aq)  C + 2 D molar mass A molarity A (M) g A mol A L of A g A 1 mol A mol A 1 L mol-to-mol ratio HW p. 161 #81 g B 1 mol B mol B 1 L molar mass B molarity B (M) g B mol B L of B 41