Chapter 4. Aqueous Solutions

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1 Chapter 4 Reactions in Aqueous Solutions Aqueous Solutions Solvent: the dissolving medium Solute: the substance dissolved Solution: the homogeneous mixture of the two Saltwater: Solvent = water Solute = salt 1

2 Aqueous Solutions Saturated solution: maximum solute dissolved at equilibrium Supersaturated solution: more solute dissolved than in a saturated solution How can this be? Ionic Compounds Forces within ionic compounds are very strong MP NaCl = 801 o C Dissolving ionic compounds requires breaking these forces Why would any ionic compound be soluble in water? 2

3 Ionic Compounds Water is polar Oxygen relatively negative Hydrogen relatively positive Water molecules gang up on ions Precipitation Reactions Not all ionic compounds dissolve in water. Many are insoluble 2 NaI (aq) + Hg(NO 3 ) 2 (aq) HgI 2 (s) + 2 NaNO 3 (aq) 3

4 Net Ionic Equation 2 NaI (aq) + Hg(NO 3 ) 2 (aq) --> HgI 2 (s) + 2 NaNO 3 (aq) 2Na + (aq) + 2 I - (aq) + Hg 2+ (aq) + 2 NO 3- (aq) --> HgI 2 (s) + 2Na + (aq) + 2 NO 3- (aq) Hg 2+ (aq) + 2 I - (aq) --> HgI 2 (s) Net ionic equation for reaction Na + and NO 3 - are spectator ions What Ionic Compounds Dissolve? Some ionic compounds dissolve in water and some do not. We need to be able to predict which compounds are soluble. Use a set of rules that help us make these predictions solubility rules. 4

5 Solubility Rules Use to predict whether a precipitation rxn will occur Compounds containing Exceptions alkali metals, NH 4 + are soluble NO 3 -, HCO 3-, ClO 3 - are soluble Halides usually are soluble Ag +, Hg 2 2+, Pb 2+ SO 4 2- usually are soluble Ag +, Hg 2 2+, Pb 2+, Sr 2+, Ba 2+, Ca 2+ (slightly) OH - seldom are soluble Group I, Ba 2+ S 2-, CO 3 2-, PO 4 3-, CrO 4 - seldom are soluble Group I, NH 4 + Solubility Rules Question: Would you expect Fe(OH) 3 to be water-soluble? Ba(OH) 2? Question: If you wanted a solution containing CrO 4 2- ion, what solid(s) would you get from the stockroom? Question: What would be the net ionic equations for the reactions (if any) between aqueous solutions of Na 3 PO 4 and NiCl 2 Pb(NO 3 ) 2 and KCl Ba(OH) 2 and Fe 2 (SO 4 ) 3 5

6 Acids and Bases Acids and bases: classification scheme for chemicals Taste sour Turn litmus red React with active metals React with carbonates Acids Taste bitter Turn litmus blue Feel slippery Bases Brönsted-Lowry Definition Acid: proton (H + ) donor Base: proton (H + ) acceptor Question: In the following reactions HCl(aq) + H 2 O ----> H 3 O + (aq) + Cl - (aq) CH 3 COOH (aq) + NH 3 (aq) > NH 4+ (aq)+ C 2 H 3 O 3 - (aq) What are the acids? The bases? 6

7 Brönsted-Lowry Definition H 3 O + = hydronium ion HCl(aq) + H 2 O ----> H 3 O + (aq) + Cl - (aq) frequently written as HCl(aq) ----> H + (aq) + Cl - (aq) Inorganic acids Recognizing Acids Generally have acidic protons first in formula HCl, H 2 SO 4, H 3 PO 4 Organic acids COOH group is organic acid functional group CH 3 COOH, C 6 H 5 COOH 7

8 Strong Acids Some acids are strong electrolytes: completely dissociated in solution Common strong acids HCl hydrochloric acid (also HI, HBr but not HF) HNO 3 nitric acid H 2 SO 4 sulfuric acid HClO 4 perchloric acid Know these: assume anything not strong is weak HCl, HNO 3, HClO 4 are monoprotic, H 2 SO 4 is diprotic First proton is strong Strong Bases Common strong bases Group IA and barium hydroxides LiOH lithium hydroxide NaOH sodium hydroxide KOH potassium hydroxide RbOH rubidium hydroxide CsOH cesium hydroxide Ba(OH) 2 barium hydroxide Again, know these 8

9 Weak Acids and Bases Weak acids: all acids that are not strong HC 2 H 3 O 2, HF Weak bases Many common weak bases are related to ammonia, NH 3 Example: Methylamine CH 3 NH 2 Acid-Base Reactions Neutralization reactions HNO 3 (aq) + KOH (aq) ----> H 2 O + KNO 3 (aq) HC 2 H 3 O 2 (aq) + NaOH (aq) > H 2 O + NaC 2 H 3 O 2 (aq) These are proton transfer reactions. 9

10 Acid-Base Reactions HNO 3 (aq) + KOH (aq) ----> H 2 O + KNO 3 (aq) HC 2 H 3 O 2 (aq) + NaOH (aq) > H 2 O + NaC 2 H 3 O 2 (aq) The net ionic equations for these reactions are H + (aq) + OH - (aq) > H 2 O HC 2 H 3 O 2 (aq) + OH - (aq) > H 2 O + C 2 H 3 O 2 - (aq) Question: Why don t we separate HC 2 H 3 O 2? Oxidation-Reduction 2 Na (s) + Cl 2 (g) ----> 2 NaCl (s) Na, Cl 2 are uncharged In NaCl Na + Cl - Electron-transfer reaction 10

11 Oxidation-Reduction 2 Na (s) + Cl 2 (g) ----> 2 NaCl (s) Express electron transfer in half-reactions Na (s) ----> Na + + e - Oxidation Cl 2 (g) + 2 e > 2 Cl - Reduction Oxidation-Reduction Na (s) ----> Na + + e - oxidation Cl 2 (g) + 2 e > 2 Cl - reduction Na is oxidized (charge becomes more positive) Na is reducing agent Cl 2 is reduced (charge becomes more negative) Cl 2 is oxidizing agent 11

12 Oxidation-Reduction Redox reactions don t need to involve ionic cmpds CH 3 OH (l) + 2 O 2 (g) ----> CO 2 (g) + 2 H 2 O (l) Use oxidation numbers to tally exchange of electrons ON are imaginary charges Imaginary charges Fill in table below Oxidation Numbers Species Free elements 0 Ions Group IA Group IIA Aluminum Fluorine Oxidation Number charge of ion Oxygen (usually) Hydrogen (usually) Sum of oxidation nos = charge When in doubt, use Periodic Table 12

13 Oxidation Numbers (except peroxide O 2-2 ) +1 (except hydride H - ) Oxidation Numbers Question: what are oxidation numbers of elements in compounds below. Start with the oxidation numbers you know CH 3 OH CO 2 KMnO 4 Cr 2 O 7 2- PCl 5 13

14 Redox Reactions CH 3 OH (l) + 2 O 2 (g) ----> CO 2 (g) + 2 H 2 O (l) CH 3 OH is oxidized Reducing agent O 2 is reduced Oxidizing agent (ON becomes more positive) (ON becomes more negative) Note Not doing types of redox reactions Pages

15 Molarity The molarity, (M), is the most common way that chemists express the concentration of solutions. Example: How would you prepare a ml of a M solution of NiCl 2 6H 2 O? (MM = g/mole) M = moles/v # moles = MV Molarity Example: How many ml of a M solution of NiCl 2 6H 2 O will contain mole? M = moles/v V = moles/m 15

16 Dilution Another way of preparing a solution is to dilute a more concentrated solution. The key to a dilution problem is the "magic formula" below. M 1 V 1 = M 2 V 2 M 1 and V 1 are the molarity and volume of the concentrated solution. M 2 and V 2 are the molarity and volume of the dilute solution. Dilution Example: How would prepare ml of a M solution from a M solution of CrCl 3? 16

17 Gravimetric Analysis Method of quantitative analysis Example: Have g of a solid that contains sulfate. Want to know the %sulfate in this solid. Start by precipitating the sulfate. What would you add? KCl, Fe(NO 3 ) 2, BaCl 2, HCl, or Cu(NO 3 ) 2 Write net ionic equation for the reaction Weight of solid product is g. What is % sulfate in original sample? Have g sample. Gravimetric Analysis Ba 2+ + SO 4 2- BaSO 4 (s) Mass BaSO 4 = g What is %SO 4 2-? 17

18 Gravimetric Analysis Have 1500 ml of solution containing Fe +3 What would you add to precipitate the iron? Na 2 SO 4, Na 2 S, or NaCl Write net ionic equation for reaction Obtain g product What is molarity of iron solution? Standard solution Add until reacts with all of unknown Equivalence point #moles = M SS V SS Titrations Reaction relates #moles unknown and #moles standard 18

19 Titrations Example: Strong acid-strong base titration H + + OH - H 2 O Start with ml of HCl in flask Add M NaOH from pipet First reading: ml Second reading: ml What is molarity of HCl solution? Titrations Reaction: #moles H + = #moles OH - 19

20 Titrations Titrate solution of Fe 2+ with KMnO 4 5Fe 2+ + MnO H + Mn Fe H 2 O What kind of reaction is this: acid-base, precipitation, or redox? Acid-base: exchange of protons, no change in oxid. nos Precipitation: solid formed Redox: change in oxidation numbers Titrations Dissolve g iron sample in ml Add ml of.1000m MnO 4 - solution to reach equivalence point What is percent iron in sample? 20

21 Oxidation Nos and Reactivity We can use oxidation numbers to predict reactivity of compounds Usual oxidation numbers Group 1A. +1 Group 2A. +2 Al.. +3 Halogens. -1 Oxygen. -2 Hydrogen.. +1 Transition metals. +2 (generally) Oxidation Nos and Reactivity Chemicals with unusual oxidation numbers tend to react to produce usual ones KMnO 4 Mn ON = +7 Usual ON of Mn = Mn tends to be reduced to +2 KMnO 4 is good oxidizing agent 21

22 Oxidation Nos and Reactivity Chemicals with unusual oxidation numbers tend to react to produce usual ones LiAlH 4 H ON = is usual ON for H -1 H tends to be oxidized to +1 LiAlH 4 is good reducing agent Oxidation Nos and Reactivity Which of the following are potential oxidizing or reducing agents? K 2 Cr 2 O 7 Zn H 2 O Cl 2 LiAlH 4 KCl ClF 5 KH PtF 6 (Used to produce first noble gas compound) 22

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