9: Charge transfer rxns: acids and bases and oxidation-reduction.

9: Charge transfer rxns: acids and bases and oxidation-reduction

What are acids?

• Acids:

• sharp,sour taste, feel prickly on skin,

• react with metals to release H2(g),

• react with carbonates to release CO2(g),

• turn litmus red,

• neutralize bases.

What are bases?

• Bases:

• bitter taste

• slippery to touch

• turn litmus blue

Arrhenius definitions

• Acid increases __ conc (hydrogen ion, proton) when dissolved in water.

• HCl (aqCl-(aq)

• Acid has to have __ to be acid.

Arrhenius definitions

• Base increases ____ (hydroxide ion) conc when dissolved in water

• NaOH Na+(aq) + OH-(aq)

• Base has to have __

Arrhenius acid-base rxn

• Arrhenius acid-base neutralization rxn yields salt + water.

• Arrhenius limited to aqueous solution.

But NH3 is a base!!

• Arrhenius says bases have to have OH-’s and acids H+’s.

• However NH3 is a base but has no OH- in its formula.

• Now have enlarged defn of acid and base that’s not restricted to aqueous soln

• A Bronsted-Lowry (B-L) acid

• A B-L base

• The acid-base neutralization rxn is the transfer of a H+ from an acid to a base.

• Concept of base in enlarged (H+ acceptor)

• Acid still has to have H+.

Identify acids and bases in

• SO32- + HCl HSO3- + Cl-

• Note that the products of the rxn are acids and bases themselves.

• H3PO4 + H2O H3O+ + H2PO4-

• H3O+: hydronium ion

• HCO3- + H2O H2CO3 + OH-

Water

• Note that water can act both as a base and as an acid.

• H3PO4 + H2O H3O+ + H2PO4-

• HCO3- + H2O H2CO3 + OH-

• Amphiprotic (amphoteric)

Conjugate acid-base pairs

• Conjugate acid-base pairs are species on opposite sides of the rxn that differ by one proton only.

• Are these conjugate acid/base pairs?

• H 2O/OH-

• H2SO4/SO42-

• H2SO4/ HSO3-

• Given base, to get conjugate acid: ___

• What is the conjugate acid of:HS- SO4

2-

NH3 S2-

• Given acid to get conjugate base: _________

• What is the conjugate base of:

• HS- NH3 HSO4- S2-

• 9.36: Identify conjugate acid-base pairs in:

• NH4+(aq) + CN--(aq) NH3(aq) +HCN(aq)

• CO32-(aq) + HCl(aq) HCO3

-(aq) + Cl-(aq)

Strength of acids and bases

• Some species are better proton donors than others; some species are better proton acceptors than others.

• The better the species is at donating a proton to the base, the stronger the acidic properties of the species.

• Strength is a measure of the degree of dissociation of the acid or base in solution. Stronger the acid(base) the larger the value of K for that dissociation.

• The better the species is at accepting the proton from the acid, the stronger the basic properties of the species.

• We are talking about strong and weak here in relative terms.

• Generally talk about aqueous sol’ns.• Some acids and bases are strong acids and

bases meaning that they are

• HCl + H2O Cl- orHCl(aq) H+(aq) + Cl-(aq)

• NaOH(aq) Na+(aq) + OH-(aq)• Ca(OH)2(aq)Ca2+(aq) + 2OH-(aq)

• Some acids and bases are not 100% ionized. These are weak acids (and bases). (weak electrolytes)This means that K < 1 for these rxns. Instead an equilbrium is set up:

• HF + H2O H3O+ + F- Ka= 7.1 x 10-4

• HCOOH + H2O HCOO- + H3O+

Ka = 1.7 x 10-4

• Larger value of Ka stronger the acid

How do you know what acids and bases are strong?

• See fig 9.2 p239Strong acids Strong bases

• HCl OH- (metal )

• H2SO4 (1st ioniz. only) S2-

• HNO3 O2-

• H3O+

• The stronger the acid, the weaker the conj. base of the acid. The weaker the acid, the stronger the conj. base of the acid, etc.

• What does this imply?

• Acid (dec ) Base (inc • HCl Cl-

• H3O+ H2O

• HSO4- SO4

2-

• HF F-

• CH3COOH CH3COO-

• NH4+ NH3

• Which is the stronger acid--HF or HSO4-?

• Which is the stronger base--H2O or F-?

Dissociation of water• HOH(l) + HOH(l) H3O+ + OH-• Acid1 base2 acid2 base1 • Or H2O(l) H+(aq) + OH-(aq)• Find from experiment that for pure water at 25oC

that [H3O+] = [OH-] = • Can write the ion product for the dissociation of

water as ion product = [H3O+][OH-] = ____________

at 25oC• H3O+ and H+ stand for the same thing.

• Can write the ion product for the dissociation of water as

ion product = [H3O+][OH-] = 1 x 10-14 at 25oC

• This is also called Kw.

• This holds for all solutions not just pure HOH

• Kw = [H3O+][OH-] = 1 x10-14

• as [H3O+]

• If [H3O+] = 4.6 10-3M, [OH-] =?

• as [OH-]

• If [OH-] = 1.22 x 10-4M, [H3O+]=?

pH or how to avoid the exponents

• define pH = -log10[H3O+]

• Neutral sol’n: [H3O+] =[OH-] = 1 x10-7MpH=_

• Acidic sol’n; [H3O+] >OH-] ; [H3O+] > 1x10-7M pH<_

• Basic sol’n: [OH-]> [H3O+]; [OH-]>1x10-7MpH >_

• [H+]M pH • 10 -1• 1 0• 0. 1 1 • 0.001 3• 1x10-7 7 • 1x10-10 10• 1x10-14 14 • 1x10-15 15 • As pH [H+]

• What’s pH for a sol’n whose [H3O+] = 4.46 x10-3M?

• If pH = 9.72, what’s [H3O+]?• If pH = 4.45, what’s [H3O+] ?• What is [OH-] in the above solutions?

• Calc pH of 0.0158M HCl

• Calc pH of 1 x 10-4M KOH and of 0.036M KOH

• 9.53 and 54. How many times more basic is a soln at pH 6 relative to pH 4?

• How many times more acidic is a soln at pH 7 relative to pH 11?

9.3 Rxns btn acids and bases• Neutralization rxn

HCl(aq) + NaOH(aq) NaCl(aq) + HOH(l)• acid + base salt + water• In the dissociated form

H+(aq) +Cl-(aq) +Na+(aq) + OH-(aq) Na+(aq) + Cl-(aq) + HOH(l)

• H+(aq) + OH-(aq) H2O(l)• This is the same for any strong acid-strong

base neutralization

Titration

• Used to determine the conc of unknown acid or base solution.

• 6.00mL of a 0.500M HCl soln is needed to neutralize 10.00mL of a NaOH soln. What is the concentration (molarity) of the NaOH soln?

• 15.00mL of a 0.200M NaOH soln is needed to neutralize 25.00mL of a HCl soln. What is the concentration (molarity) of the HCl soln?

• Equivalence pt and end point, indicators

Polyprotic acids

• Some acids have more than one acidic hydrogen as: H2SO4, H2CO3, H3PO4

• H2CO3 +2NaOH Na2CO3 + 2H2O

• Undergo stepwise dissociation as:

• H2CO3 + H2O HCO3- + H3O+

• HCO3- + H2O CO3

2- + H3O+

9.4 Acid-base buffers

• Solutions that contain a

• Buffers have the ability of withstanding small additions of added acid or base without large fluctuations in pH.

How do buffers work?

• How do buffers work?

• HA H+ + A-

• add acid A- + H+ HA

• add base HA + OH- HOH + A-

• Have 100mL of 3.6x10-5M HCl; pH = 4.44

• Add 1mL of 0.10M HCl to the above soln, the pH changes to 2.99; add 1mL of 0.10M NaOH to the above soln, pH changes to 11.99.

• Have 100mL of a buffer that is 0.050M Na acetate and0.10M acetic acid, pH = 4.44

• Add 1mL of 0.10M HCl to the buffer, pH changes to 4.43; add 1mL of 0.10M NaOH to the buffer, pH changes to 4.46.

• Are these buffers:

• NaOH/NaCl

• NaOH/HCl

• NH3/NH4Cl

• NaF/HF

• CH3COOH and CH3COONa

• HNO3 and KNO3

• HBr and MgCl2

• H2CO3 and NaHCO3

Calculating the pH of a buffer

• HA(aq) + H2O(l) H3O+(aq) + A-(aq)

• Ka = [H3O+][A-] [HA]

• Or [H3O+] = Ka[HA][A-]

• [H3O+] = Ka[acid] [salt]

• 9.65 and 66 What is [H3O+] for a buffer soln that is 0.200M in acid and 0.500M in the corresponding salt if the weak acid has Ka= 5.80 x 10-7? What is the pH of the solution?

• What is the pH of a solution that is 0.1M in CH3COOH and 0.1M in CH3COONa? Ka for acetic acid is 1.8 x 10-5.

• 9.49 and 50 ForCH3COOH(aq)+H2O(l) CH3COO-(aq)

+H3O+(aq)

• Write the K expression• Explain what happens if • Add strong acid

• Dilute with water

• Add strong base

• Add more acetic acid

Henderson-Hasselbalch equation

• pH = pKa + log

•

• Solve: What is [H3O+] for a buffer soln that is 0.200M in acid and 0.500M in the corresponding salt if the weak acid has Ka= 5.80 x 10-7? What is the pH of the solution? Usong Henderson-Hasselbalch

9.5 Oxidation-reduction processes

• Oxidation (oxd’n) ; originally rxn in which O2 combined with another substance as

C(s) + O2 CO2(g)

• Also defined as

• Reduction (red’n) : originally rxn in which O2 was removed from a substance as

Fe2O3(s) + 3C(s) 3CO(g) + 2 Fe(s)

• Also defined as

Let’s look at gain and loss of electrons more closely

• Look at the reaction2NaCl(s) 2Na(s) + Cl2(g)

• We can divide this rxn into half rxns as

•

And we identify oxidation and reduction as

Na+ + e- Na Cl- Cl2 + 2e-

• The species (reactant, Cl-) that is oxidized is the _________ agent (reducing agent donates e-’s to the species that is reduced)

• The species (reactant, Na+) that is reduced is the ___________ agent (oxidizing agent causes oxd’n of another species by accepting e-’s from that species)

Identify species oxidized, reduced, oxidizing agent,

reducing agent

• Cu2+ + Mg Mg2+ + Cu

• Br2 + 2KI 2KBr + I2

• 4Fe + 3O2 2Fe2O3

• Zn + Cu 2+ Zn2+ + Cu

Corrosion

• Corrosion: eating away of metals by oxidation-reduction processes.

• Rusting of iron4Fe(s) + 3O2(g) 2Fe2O3(s)

Voltaic cells

• Voltaic cells: spontaneous chemical rxn is used to generate electricity

• Cu2+(aq) +Zn(s) Zn2+(aq) + Cu(s)

• Oxidation occurs at the______ and reduction at the _________.

• Half cell rxns: Cu2+ + 2e- Cu Zn Zn2+ + 2e-

• Electrochemical cells: generate electricity (e-’s moving through a wire) by spont. chem. rxn --voltaic cells

• Right now we have a direct transfer of e-’s from the anode to the cathode: short circuit

• Cu2+ + 2e- Cu GER; cathode Zn Zn2+ + 2e- LEO; anode

• Need to design a system such that keeps Cu2+ and Zn from being in direct contact with each other.

Eocell = 1.10V

Zn --> Zn2+ + 2e- Cu2+ + 2e- -->Cu

Batteries: portable source of electricity

• Flashlight battery• anode: Zn(s) Zn2+(aq) + 2e-

• cathode: 2NH4+(aq) + 2MnO2(s) + 2e-

Mn2O3(s) + 2NH3(aq) + H2O(l)

• overall: Zn(s) + 2NH4+ (aq) +MnO2(s) Zn2+(aq)

+ Mn2O3(s) + 2NH3(aq) + H2O(l)

• supplies 1.5V new; not rechargeable; voltage decreases on use (why?)

(inert

electrolyte

Why?

Alkaline version of flashlight battery

• anode: Zn(s) + 2OH-(aq) ZnO(s) + H2O(l) + 2e-

• cathode: 2MnO2(s) + H2O(l) + 2e- Mn2O3(s) + 2OH-

• overall: Zn(s) + 2MnO2(s) Mn2O3(s) + ZnO(s)

• advantage: supplies constant voltage over time period

Mercury battery

• anode: Zn(Hg) + 2OH-(aq) ZnO(s) + H2O(l) + 2e-

• cathode: HgO(s) + H2O(l) +2e- Hg(l) + 2OH-(aq)

• overall: Zn(Hg)+ HgO(s) ZnO(s) + Hg(l)

• Eo constant(1.35V)--why?; long life: good for pacemakers, watches, hearing aids

•Zn(Hg)+ HgO(s) ZnO(s) + Hg(l)

Electrolysis

• Electricity used to cause a nonspontaneous chemical reaction to occur as

• Definitions for oxd’n, etc still hold

Electrolysis of water

• Electrolysis of pure water slow. Why?

• Instead use 0.1M H2SO4 soln

• cathode: 2H+ + 2e- H2

• anode: 2H2O O2(g) + 4H+ + 4e-

• Overall 2H2O 2H2(g) + O2(g)

• cathode anode