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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

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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 = 801oC

• Dissolving ionic compounds requires breaking these forces

• Why would any ionic compound be soluble in water?

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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(NO3)2 (aq) HgI2 (s) + 2 NaNO3 (aq)

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Net Ionic Equation

• 2 NaI (aq) + Hg(NO3)2 (aq) --> HgI2 (s) + 2 NaNO3 (aq)

• 2Na+(aq) + 2 I-(aq) + Hg2+(aq) + 2 NO3-(aq) --> HgI2(s) + 2Na+(aq) + 2 NO3

-(aq)

• Hg2+(aq) + 2 I- (aq) --> HgI2 (s)

• Net ionic equation for reaction

• Na+ and NO3- 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.

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Solubility Rules

• Use to predict whether a precipitation rxn will occur

Compounds containing Exceptions

alkali metals, NH4+ are soluble

NO3- , HCO3

-, ClO3- are soluble

Halides usually are soluble Ag+, Hg22+, Pb2+

SO42- usually are soluble Ag+, Hg2

2+, Pb2+, Sr2+, Ba2+, Ca2+(slightly)

OH- seldom are soluble Group I, Ba2+

S2-, CO32-, PO4

3- , CrO4- seldom

are soluble

Group I, NH4+

Solubility Rules

• Question: Would you expect Fe(OH)3 to be water-soluble? Ba(OH)2?

• Question: If you wanted a solution containing CrO42- 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

– Na3PO4 and NiCl2

– Pb(NO3)2 and KCl

– Ba(OH)2 and Fe2(SO4)3

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Acids and Bases

• Acids and bases: classification scheme for chemicals

Taste sour Taste bitter Turn litmus red Turn litmus blue React with active metals Feel slippery React with carbonates

Acids Bases

Brönsted-Lowry Definition

• Acid: proton (H+) donor

• Base: proton (H+) acceptor

• Question: In the following reactions

– HCl(aq) + H2O ----> H3O+(aq) + Cl-(aq)

– CH3COOH (aq) + NH3 (aq) —> NH4+

(aq)+ C2H3O3- (aq)

– What are the acids? The bases?

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Brönsted-Lowry Definition

• H3O+ = hydronium ion

• HCl(aq) + H2O ----> H3O+(aq) + Cl-(aq) frequently written as

HCl(aq) ----> H+(aq) + Cl-(aq)

Recognizing Acids

• Inorganic acids

– Generally have acidic protons first in formula

– HCl, H2SO4, H3PO4

• Organic acids

– COOH group is organic acid functional group

– CH3COOH, C6H5COOH

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Strong Acids

• Some acids are strong electrolytes: completely dissociated in solution

• Common strong acids

– HCl hydrochloric acid (also HI, HBr but not HF)

– HNO3 nitric acid

– H2SO4 sulfuric acid

– HClO4 perchloric acid • Know these: assume anything not strong is weak

• HCl, HNO3, HClO4 are monoprotic,

• H2SO4 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

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Weak Acids and Bases

• Weak acids: all acids that are not strong

– HC2H3O2, HF

• Weak bases

– Many common weak bases are related to ammonia, NH3

– Example: Methylamine CH3NH2

Acid-Base Reactions

• Neutralization reactions

• HNO3 (aq) + KOH (aq) ----> H2O + KNO3 (aq)

• HC2H3O2 (aq) + NaOH (aq) —> H2O + NaC2H3O2 (aq)

• These are proton transfer reactions.

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Acid-Base Reactions

• HNO3 (aq) + KOH (aq) ----> H2O + KNO3 (aq)

• HC2H3O2 (aq) + NaOH (aq) —> H2O + NaC2H3O2 (aq)

The net ionic equations for these reactions are

• H+ (aq) + OH- (aq) —> H2O

• HC2H3O2 (aq) + OH- (aq) —> H2O + C2H3O2- (aq)

Question: Why don’t we separate HC2H3O2?

Oxidation-Reduction

• 2 Na (s) + Cl2 (g) ----> 2 NaCl (s)

• Na, Cl2 are uncharged

• In NaCl

– Na+

– Cl-

• Electron-transfer reaction

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Oxidation-Reduction

• 2 Na (s) + Cl2 (g) ----> 2 NaCl (s)

• Express electron transfer in half-reactions

• Na (s) ----> Na+ + e-

– Oxidation

• Cl2 (g) + 2 e- ------> 2 Cl-

– Reduction

Oxidation-Reduction

• Na (s) ----> Na+ + e- oxidation

• Cl2 (g) + 2 e- ------> 2 Cl- reduction

• Na is oxidized (charge becomes more positive)

– Na is reducing agent

• Cl2 is reduced (charge becomes more negative)

– Cl2 is oxidizing agent

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Oxidation-Reduction

• Redox reactions don’t need to involve ionic cmpds

• CH3OH (l) + 2 O2 (g) ----> CO2 (g) + 2 H2O (l)

• Use oxidation numbers to tally exchange of electrons

– ON are imaginary charges

Oxidation Numbers

• Imaginary charges – Fill in table

below

Species Oxidation Number

Free elements 0

Ions charge of ion

Group IA Group IIA Aluminum Fluorine

Oxygen (usually) Hydrogen (usually)

Sum of oxidation nos = charge

When in doubt, use Periodic Table

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Oxidation Numbers

+1 +2 +3 -1 -2 (except peroxide O2

2- )

+1 (except hydride H- )

Oxidation Numbers • Question: what are oxidation numbers of elements in

compounds below. – Start with the oxidation numbers you know

• CH3OH CO2 KMnO4

• Cr2O72- PCl5

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Redox Reactions

• CH3OH (l) + 2 O2 (g) ----> CO2 (g) + 2 H2O (l) • -2 0 +4 -2

• CH3OH is oxidized (ON becomes more positive)

– Reducing agent

• O2 is reduced (ON becomes more negative)

– Oxidizing agent

Note

• Not doing types of redox reactions

• Pages 139 - 145

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Molarity

• The molarity, (M), is the most common way that chemists express the concentration of solutions.

• Example: How would you prepare a 500.0 mL of a 0.1000 M solution of NiCl26H2O? (MM = 237.7 g/mole) – M = moles/V

– # moles = MV

Molarity

• Example: How many mL of a 0.5000 M solution of NiCl26H2O will contain 0.0250 mole?

– M = moles/V

– V = moles/M

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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.

• M1V1 = M2V2

– M1 and V1 are the molarity and volume of the concentrated solution.

M2 and V2 are the molarity and volume of the dilute solution.

Dilution

• Example: How would prepare 500.0 mL of a 0.0250 M solution from a 0.5000 M solution of CrCl3 ?

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Gravimetric Analysis

• Method of quantitative analysis

• Example: Have 1.894 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(NO3)2, BaCl2, HCl, or Cu(NO3)2

– Write net ionic equation for the reaction

• Weight of solid product is 1.745 g.

• What is % sulfate in original sample?

Gravimetric Analysis

• Have 1.894 g sample.

• Ba2+ + SO42- BaSO4(s)

• Mass BaSO4 = 1.745 g

• What is %SO42- ?

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Gravimetric Analysis

• Have 1500 ml of solution containing Fe+3

• What would you add to precipitate the iron?

– Na2SO4, Na2S, or NaCl

– Write net ionic equation for reaction

• Obtain 0.619 g product

• What is molarity of iron solution?

Titrations

• Standard solution

– Add until reacts with all of unknown

– Equivalence point

– #moles = MSSVSS

• Reaction relates #moles unknown and #moles standard

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Titrations

• Example: Strong acid-strong base titration

• H+ + OH- H2O

– Start with 100.0 mL of HCl in flask

– Add 0.1250 M NaOH from pipet

– First reading: 43.21 mL

– Second reading: 15.42 mL

• What is molarity of HCl solution?

Titrations

• Reaction: #moles H+ = #moles OH-

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Titrations

• Titrate solution of Fe2+ with KMnO4

• 5Fe2+ + MnO4- + 8H+ Mn2+ + 5Fe3+ + 4H2O

• 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 2.100 g iron sample in 250.00 mL

• Add 35.29 mL of .1000M MnO4- solution to reach equivalence

point

• What is percent iron in sample?

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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

• KMnO4 Mn ON = +7

– “Usual” ON of Mn = +2

– +7 Mn tends to be reduced to +2

– KMnO4 is good oxidizing agent

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Oxidation Nos and Reactivity

• Chemicals with unusual oxidation numbers tend to react to produce usual ones

• LiAlH4 H ON = -1

– +1 is “usual” ON for H

– -1 H tends to be oxidized to +1

– LiAlH4 is good reducing agent

Oxidation Nos and Reactivity

• Which of the following are potential oxidizing or reducing agents?

– K2Cr2O7

– Zn

– H2O

– Cl2

– LiAlH4

– KCl

– ClF5

– KH

– PtF6 (Used to produce first noble gas compound)