analytical chemistry exam 3

titration of strong base with strong acid

• H+ + OH- → H2O

• high starting pH

• before titration, only SB, pOH = -log[starting OH-]

• before Ve - excess unreacted OH

  • (unreacted OH/needed0 x initial concentration x (inital vol/new volume) = OH M

  • use M OH to get pOH, 14-pOH = pH

• at Ve - moles OH = moles H+

  • pH = 7

• After Ve - excess H+

  • initial concentration of H+ x excess/new volume = [H+] → pH

titrating strong acid with strong base

• pH starts low

• before start, ph=log[H+]

• Ve = excess H+

  • fraction of H left x concentration of H x inital vol/new vol

• Ve, pH = 7

• after Ve, excess OH-

  • excess OH/new vol x initial concentration → pOH → pH

titration of weak acid with strong base

• pH starts low

• HA + OH- → A- + H2O

• before start, only weak acid

  • only HA ←→ A- + H+

  • use ka and ice to find x 

  • x = [H+]

• before Ve, excess WA

  • buffer

  • HH equation

  • at ½ Ve, pH = pka

• at Ve, only weak base

  • use Kb = kw/ka

• after Ve, excess strong base

  • excess/new vol x initial concentration → pOH → pH

titration of weak base with strong acid

• pH starts low

• B + H2O → BH+ + OH-

• before start

  • only WB, kb

• before Ve

  • buffer B and BH+

  • HH

• at Ve, only WA, use Ka

• after Ve, excess H+

  • excess/total vol x conc = H+ → pH

diprotic A-B systems

• two Ves

• two buffer areas

• two equations

• before start - only weak k

henderson haselbach equation 

• pH = pka + log (conj base/ WA)

• pH = pkb + log (conj acid/WB)

Kw

kw = ka x kb

ka = kw/kb

kb = kw/ka

kw = 1.0×10^-14

indicators

• A-B indicators are weak acids or bases that have different colors at different protonated forms

• want an indicator with a color change close to the pH at Ve

Metal-chelate complex

• metal ions accept electron pairs

• ligands donate electron pairs

• chelating ligand binds metal ion through more than one ligand atom, form stable complexes

EDTA titrations

• based on metal-ligand complex formation

• binds in 1:1 ratio

• Y4- = fully deprotonated form

αY4-

fraction of EDTA in the fully deprotonated form (Y4-) (not complexed into metal ions)

formation constant, Kf

• equilibrium constant for metal-ligand reaction

• in terms of Y4-

• usually large and tend to be even larger for positive charge ions

• Kf = [MYn-4] / [Mn+] [Y4-]

  • for reaction Mn+ + Y4- ←→ MYn-4

Conditional formation constant

• K’f, at specific pH

• metal complex / metal x EDTA

metal-EDTA complex

• become less stable at lower pH

• metals with higher formation constant can be titrated at lower pH

• need pH low enough so metal hydroxide does not form, but high enough that Ve will be distinct

EDTA titration curve

• before Ve - unreacted metal

  • metal concentration determined by excess

  • unreacted/total x concnetration x initial vol/ new vol

• at Ve

  • metal concentration determined by dissociation of metal complex, according to Kf

  • concentration x inital vol/new vol

• after Ve - unreacted EDTA

  • use excess EDTA in K’f to calculate metal

  • concentration x excess/new vol

auxiliary complexing agent

• ligand that will bind to metal strongly enough to prevent metal hydroxide from precipitating

• but loose enough to give the metal up to EDTA

metal ion indicator

• color change when bound to metal

• must bind metal less strongly than EDTA

• can only be used in certain pH ranges

oxidation

• loss of electrons

• becomes more positive

• reducing agent

reduction

• gain of electrons

• gets more negatove

• oxidizing agent

electrical charge, Q

• quantity of electricity

• measure in coulombs

• charge = n mol F

electric current

• quantity of charge flowing each second through a circuit

• ampere, A = C/sec

electric potential difference, E

• created by electric charge 

• between two points

• work required to move from one point to the other

• measured in volts

work

I work I = E x Q (electrical potential difference x charge)

joules

gibbs free energy

DG = -n F E

battery

• gives energy as heat or work (q or w)

relation between charge and moles

Q = n mol F

relation between work and voltage 

work = E Q

realtion between free energy difference and electrical potential difference

DG = - n F E

galvanic cell

• spontaneous chemical reaction to generat energy

• battery

• oxidizing and reducing agents separated, electrons flow through a wire

anode

• oxidation occurs

• E - 

cathode

• reduction occurs

• E +

E

E = E+ - E-

salt bridge

• negative electrons want to move toward positive electrical potential

• connects two half cells

• allows ions to diffuse and minimize the mixing of ions

• prevents charge build up in half cell

standard potential 

Eo 

found int able 

E, (nernst)

E+ = E-

Eo+= Eo - 0.05916/n log 1/[ ]

• write net cell reaction

• write both half reactions in reduction form

• find Eo on table

• do nernst equaiton for both sides

• E = E+ = E- = V

• net cell reaction

formal potential

Eo’

different conditions

relevant pH

potentiometry

• use of electrodes to measure voltage that provides chemical information

• measures electrical potential difference

• electroactive specied interacts with indicator electrode. \\

  • salt bridhe connects to referemce electrode

reference electrode 

• stable voltage 

• constant 

• don’t want voltage to change during measuerment 

• vcreated different refernce scales with difference reference methods

indicator electrode

• metal electrode or ISE

junction potential

• electric potential that exists at interface of solutions

• different diffusion coefficients

• want oppositely charged ions with very close diffusion coefficients

• minimized mV difference

ka

10 ^ - pka

pka

-logka

strong acids

• HCl

• HBr

• HI

• HClO4

• HNO3

• H2SO4

strogn bases

• NaOH

• LiOH

• KOH

• Ca(OH)2

• Si(OH)2

• Ba(OH)2