Chem 28 – analytical chemistry
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Chem-28-Finals-Reviewer
- Bu sahifa navigatsiya:
- COMPLEX FORMATION TITRATION/ COMPLEXOMETRIC TITRATION
- Water Hardness Category CaCO 3 /Li Soft <17.1 Moderately Soft
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PRECIPITATION TITRATION Effect of Concentration on Titration Curves – increase in concentration enhances ÄpX at equivalence point region (same trend as in acid-base titration) Effect of Reaction Completion on Titration Curves – greatest ÄpX I titration of ion which forms the least soluble Ag+ salt (represents most nearly complete reaction) 𝐴𝑔 m
n ↔ 𝐴𝑔𝑋
𝐾 = 1 [𝐴𝑔 m ][𝑋
n ] = 1 𝐾 (X The smaller the K sp , the larger the K for titration Endpoints of Argentometric Titrations: Chemical - indicators Potentiometric – pH meter; voltage converted to pH (Potential) Amperometric – current measured by ammeter Chemical Indicators for Precipitation Titrations: Analyte A, Titrant B and Indicator In Titration Reaction: 𝐴 + 𝐵 → 𝐴𝐵 Endpoint Reaction: 𝐼𝑛 + 𝐵 → 𝐼𝑛𝐵 In B – with significant difference in appearance in solution than In at end point; color change Types of Argentometric titrations: 1. The Mohr Method – formation of a 2 nd precipitate) – widely applied for the titration of Cl, Br, or Cn with standard AgNO 3
2 CrO
4
Titration Reaction: (White Precipitate) 𝐴𝑔 m + 𝐶𝑙 n ↔ 𝐴𝑔𝐶𝑙
(() Endpoint Reaction: (Red Precipitate) 2𝐴𝑔 m
ƒ }n ↔ 𝐴𝑔2𝐶𝑟𝑂 ƒ(() K sp AgCl = 1.8 x 10 -10
K sp Ag 2 CrO
4 = 1.2 x 10 -12
4 2- ] protonates; pH >10 Ag 2 O precipitates Solubility – Ag 2 CrO 4 is more soluble than AgCl so all x- precipitate as Ag 2 CrO 4 with correct pH 2. Volhard Method – formation of a colored complex; Ag + titrated with standard KSCN; titration in acidic solution to prevent precipitation of Fe 3+ as hydrated oxide; most important application – indirect determination of X - ( add excess Ag + and back titration with standard SCN - ) Titration Reaction: 𝐴𝑔 m + 𝑆𝐶𝑁 n ↔ 𝐴𝑔𝑆𝐶𝑁 (() Endpoint Reaction: (blood red complex wherein Fe 2+ is the
indicator) 𝐹𝑒 }m + 𝑆𝐶𝑁 n ↔ 𝐹𝑒𝑆𝐶𝑁 }m K f = 1.05 x 10 5
3. Fajan’s Method – make use of adsorption indicators Adsorption Indicators – an organic compound that tends to be absorbed unto the surface of the solid in precipitation titration e.g fluorescein e.g Titration of Cl - with AgNO 3 which will form a colloidal precipitate 𝐴𝑔 m + 𝐶𝑙 n ↔ 𝐴𝑔𝐶𝑙 (() Before Equivalence Point (with excess Cl - ):
(AgCl) | Cl - | M +
After Equivalence Point (with excess Ag + ):
(AgCl) | Ag + | x - / FI
- (pinkish red) (can be NO 3 -
any anion) Fluorescein - adsorption indicator (HFI) Color Change – yellow green to pinkish red; due to adsorption process and not precipitation Ion Product Q > K eq , precipitation
COMPLEX FORMATION TITRATION/ COMPLEXOMETRIC TITRATION: Coordination Compounds/ Complexes – formed by the reaction of a metal ion with ligands via a Lewis Acid and Base Reaction wherein the metal acts as the Lewis Acid (e- pair acceptor) and the ligands act as the Lewis Base (e- pair donor) Chelate – complex produced when metal ions coordinates with 2 or more donor groups of a single ligand to form a heterocyclic ring EDTA – hexadentate ligand with 6 donor atoms (2N and 4O) with 4 acidic Hs with 4 K a s
𝐶 ž = [𝑌 ƒn ] + [𝐻𝑌
•n ] + [𝐻
} 𝑌 }n ] + [𝐻 • 𝑌 n ] + [𝐻
ƒ 𝑌]
𝛼
= [𝐻 ƒ 𝑌] 𝐶 ž 𝛼 ‚ =
[𝐻 • 𝑌 n ] 𝐶 ž 𝛼 } = [𝐻 } 𝑌 }n ] 𝐶 ž 𝛼 • =
[𝐻𝑌 •n ] 𝐶 ž 𝛼 ƒ =
[𝑌 ƒn ] 𝐶 ž Plot of á values vs. pH; K’ - conditional & effective formation constant (pH dependent K eq
equilibrium concentration of a metal ion and complex at any point in the titration curve 𝑀 2m
ƒn ↔ 𝑀𝑌
(2nƒ) 𝐾 ¢Ž = 𝐾 'W(
= [𝑀𝑌
(2nƒ) ] [𝑀 2m ][𝑌
ƒn ] 𝐾′ ¢Ž =
[𝑀𝑌 (2nƒ)
] [𝑀 2m ]𝐶 ž = 𝛼 ƒ 𝐾 ¢Ž EDTA Titration: pM vs V EDTA 1. Initial pM (V EDTA = 0ml) : pM = -log [M n+ ] 2. Preequivalence Point (V EDTA = xml) 𝑝𝑀 = − 𝑙𝑜𝑔 𝑙𝑜𝑔 𝑚𝑜𝑙𝑠 ¢
£Šže 𝑉 ž1;'9 3. Equivalence Point (V EDTA = yml) ¤𝑀𝑌 (2nƒ)
¥ = 𝑚𝑜𝑙𝑠
¢ 𝑉 +f:-9. [𝑀 2m ] = 𝐶 ž 𝐾′ ¢Ž = [𝑀𝑌
(2nƒ) ] [𝑀 2m ]𝐶 ž [𝑀 2m ] } =
[𝑀𝑌 (2nƒ)
] 𝐾′ ¢Ž 4. Post Equivalence Point (with excess EDTA) ¤𝑀𝑌
(2nƒ) ¥ =
𝑚𝑜𝑙𝑠 ¢ 𝑉 ;1;'9 [𝑌 ƒn ] = 𝐶 ž = 𝑚𝑜𝑙𝑠 £Šže
− 𝑚𝑜𝑙𝑠 ¢ 𝑉 ž1;'9 𝐾′ ¢Ž = [𝑀𝑌
(2nƒ) ] [𝑀 2m ]𝐶 ž [𝑀 2m ] } =
[𝑀𝑌 (2nƒ)
] 𝐾′ ¢Ž Cations with larger K f provide good end points even in acidic solutions Fe 3+ and In 3+ can be ti trated even at strongly acidic solutions (pH = 2) Ca 2+ , Zn 2+ , Al 3+ . Fe
2+ in moderately acidic pH only Auxillary Complexing Agent – sometimes added to keep metal ion in the solution at pH required for the titration; decreases the sharpness of the endpoints e.g. Zn 2+ titration in NH 3 -NH
4 Cl
buffer – NH 3 prevents formation of Zn(OH) 2 by forming NH 3 complexes with the Zn 2+
Titration Reaction: 𝑍𝑛(𝑁𝐻 • ) ƒ }m + 𝐻𝑌 •n → 𝑍𝑛𝑌
}n + 3𝑁𝐻
• + 𝑁𝐻
ƒ m Metallochromic Indicators – colored organic compounds (dyes) that form colored chelates with metal ions; chelates of different color for free indicator; widely used Eriochrome Black T (EBT) 𝐻 }
n (𝑟𝑒𝑑) + 𝐻 } 𝑂 ↔ 𝐻𝐼𝑛
}n (𝑏𝑙𝑢𝑒) + 𝐻 • 𝑂
𝐻𝐼𝑛 }n (𝑏𝑙𝑢𝑒) + 𝐻 } 𝑂 ↔ 𝐼𝑛 •n (𝑜𝑟𝑎𝑛𝑔𝑒) + 𝐻 • 𝑂
At pH >7, blue HIn 2- predominate if there are no metal ions Endpoint Reaction: 𝑀𝐼𝑛
n (𝑟𝑒𝑑) + 𝐻𝑌 •n ↔ 𝐻𝐼𝑛
}n (𝑏𝑙𝑢𝑒) + 𝑀𝑌 }n K
1 K otherwise premature endpoint is observed Titration Methods Employing EDTA: 1. Direct 2. Back 3. Displacement Determination of Water Hardness – determination of water quality through the titration of EDTA Hardness – total concentration of alkaline earth ions in water [Ca 2+
2+ ] > concentration of other alkaline earth ions Hardness = [Ca 2+ ] + [Mg 2+ ] expressed in CaCO 3 /Li
Water Hardness Category CaCO 3 /Li Soft <17.1 Moderately Soft 17.1-60
Moderately Hard 60-120
Hard 120-180
Very Hard >180
Liebig Titration – titration involving unidentate ligand; endpoint is marked by appearance of turbidity (precipitation of AgCN) 𝐴𝑔 m
n ↔ 𝐴𝑔(𝐶𝑁)
} n 𝐴𝑔 m + 𝐴𝑔(𝐶𝑁)
} n ↔ 𝐴𝑔 m [𝐴𝑔(𝐶𝑁)
} ] 𝑜𝑟 𝐴𝑔𝐶𝑁 Deniyes Modification: 𝐾𝐼 + 𝑁𝐻
• → 𝐴𝑔𝐼 (𝑝𝑎𝑙𝑒 𝑦𝑒𝑙𝑙𝑜𝑤) Generalized Equation for Redox Reaction: 𝐴 *+/ + 𝐵 13 ↔ 𝐴 13 + 𝐵
*+/ Red:
𝐵 13 + 𝑛𝑒 n ↔ 𝐵
*+/
Ox: 𝐴 *+/ ↔ 𝐴 13 + 𝑛𝑒 n
Redox Reactions are typically carried out in an electrochemical cell in which the reactants (OA and RA) are not in direct contact with one another; connected via salt bridge Electrochemical Cell – 2 conductors (electrodes) each of which is immersed in an electrolyte solution Cathode – electrode where reduction occurs Anode – electrode where oxidation occurs Electrodes – can be metal electrodes directly involved in the reaction or an inert electrode like Pt Cathodic Reaction when there is no easily reduced species: 2𝐻 m + 2𝑒 n ↔ 𝐻 }(K) Anodic Reaction when there is no easily oxidized species: 2𝐻 }
}(K) + 4𝐻
m + 4𝑒
n Types of Electrochemical Cells: Galvanic or Voltaic Cells – with spontaneous redox reaction Electrolytic Cell – with nonspontaneous redox reaction and requires an external source of electrical energy for operation Schematic Representation of Cells: Anode | Electrolyte Solution of Anode || Electrolyte Solution of Cathode | Cathode Electrons from the anode to the cathode If some cell is connected to a battery, the electron flow is reversed thus producing an electrolytic cell Reversible Cell – in which the direction of the electrochemical reaction is reversed when the direction of the current (electron flow is reversed) Irreversible Cell – cell wherein the reversing of the current causes an entirely different half-reaction to occur at either one or both the electrodes Electrode potential – measure of the tendency for the reaction to proceed from a nonequilibrium state to equilibrium. ∆𝐺 = −𝑛𝐹𝐸° ,+99 = −𝑅𝑇 𝑙𝑛 𝑙𝑛 𝐾 +f
Cell Potential: Saturated Calomel Electrode – E SCE
= 0.244V 𝐻𝑔 } 𝐶𝑙 } + 2𝑒 n ↔ 2𝐻𝑔
(9) + 2𝐶𝑙
('f) n Standard Hydrogen Electrode (SHE) – basis of oxidation- reduction potentials E ° = 0.00V 2𝐻 m
n ↔ 𝐻
}(K) Strongest Reducing Element – Li metal Strongest Oxidizing Element – Fluorine IUPAC Convention – electronic potential is reversed exclusively for half-reactions written as reductions Reverse reactions are spontaneous when standard state conditions apply Sign of the electrode potential indicates whether the reaction is spontaneous Nernst Equation: 𝐸 = 𝐸° − 𝑅𝑇 𝑛𝐹 𝑙𝑛 𝑙𝑛 𝑎 ‰ , 𝑎 Š / … 𝑎 e ' 𝑎 € W … 𝐸 = 𝐸° − 0.0592 𝑛
𝑎 ‰ , 𝑎 Š / … 𝑎 e ' 𝑎 € W …
If x = solute, a x ~ [x] If x = gas, a x ~ Px, atm If x = pure liquid, solvent or solid, a x ~ 1 Formal Potential – electrode potential when analytical concentrations are used in the place of molar concentrations Thermodynamic Potential of an Electrochemical Cell (E cell
): 𝐸 ,+99 = 𝐸 ,';k1/+
− 𝐸 '21/+
Calculation of Redox Equilibrium Constants from Standard Potentials: 𝐵 13
n ↔ 𝐵
*+/ 𝐴 13 + 𝑎𝑒 n ↔ 𝐴 *+/ Cathode: 𝑎𝐵 13
n ↔ 𝑎𝐵
*+/
Anode: 𝑏𝐴 *+/
↔ 𝑏𝐴 13 + 𝑏𝑒 n
Balanced Reaction: 𝑏𝐴 *+/
+ 𝑎𝐵 13 ↔ 𝑏𝐴 13 + 𝑎𝐵
*+/
At equilibrium ÄG = 0, ÄG = -nFE 𝐸 ,+99 = 0 = 𝐸 ,';k1/+
− 𝐸 '21/+
𝐸 ,';k1/+
= 𝐸 '21/+
𝐸° € − 0.0592 𝑎𝑏 = 𝐸° e − 0.0592 𝑎𝑏 𝑙𝑜𝑔 𝑙𝑜𝑔
[𝐴 *+/
] W [𝐴 13 ] W 𝐸° € −𝐸° e = 0.0592 𝑎𝑏 = 0.0592 𝑎𝑏 𝑙𝑜𝑔 𝑙𝑜𝑔 𝐾 +f 𝑙𝑜𝑔 𝑙𝑜𝑔 𝐾 +f =
𝑎𝑏(𝐸° € −𝐸° e ) 0.0592 Redox Titration Curve: electrode potential for the redox system vs. V
Titrant Indicators Used: OA or RA that responds to the change in potential of the system rather than changes in concentration of any particular product or reactant Electrode Potential System – equilibrium attained after each addition of titrant; system at equilibrium at all times throughout the titration Example: 𝐶𝑒 ƒm
}m ↔ 𝐶𝑒
•m + 𝐹𝑒
•m 𝐸 (y(;+= = 𝐸 ‰+ ¬• − 𝐸 j+ -• If indicator is present: (concentrations vary as titration proceeds; E system varies as well) 𝐸 š2 = 𝐸 (y(;+= = 𝐸
‰+ ¬• − 𝐸 j+ -• Hypothetical Cell of the Titration Mixture: SHE||Ce 4+ ,Ce 3+ ,Fe
3+ , Fe
2+ |Pt
Apply Nernst Equation: 1. Before equivalence point: 𝐸 ‰+
2. Equivalence Point: 𝐸 ‰+
& 𝐸 j+ -• 3. After Equivalence Point: 𝐸 j+
Effect of Completeness of Reaction in Redox Titration Curves: The larger the Keq, the larger the ÄE at the equivalence point region Effect of Concentration of redox Titration Curves: Titration Curves are independent of concentration of reactants and are independent of dilution over a considerable range Types of Redox Indicators: 1. General Redox Indicators – respond to the potential of a system; substance that changes color upon being oxidized or reduced; change in color depends only upon potential of the system 𝐼𝑛 13 ↔ 𝐼𝑛 *+/
𝐸 = 𝐸° š2 − 0.0592 𝑛 𝑙𝑜𝑔 𝑙𝑜𝑔 [𝐼𝑛 *+/
] [𝐼𝑛
13 ]
Color change occurs when there is a change in the ratio of the reactants of about 100 [𝐼𝑛 *+/
] [𝐼𝑛
13 ]
1 10
[𝐼𝑛 *+/
] [𝐼𝑛
13 ] ≥ 10 Full color change when 𝐸 = 𝐸°
š2 ±
0.0592 𝑛 Choose indicator which would change color near the equivalence point e.g Ferroin (Phen) 3 Fe
𝑃ℎ • 𝐹𝑒 •m + 𝑒 n ↔ 𝑃ℎ
• 𝐹𝑒 }m 𝐹𝑒𝑟𝑟𝑖𝑖𝑛 (𝑏𝑙𝑢𝑒)(𝑜𝑥) ↔ 𝐹𝑒𝑟𝑟𝑜𝑖𝑛(𝑟𝑒𝑑)(𝑟𝑒𝑑) 2. Specific Indicator – reacts in a specific manner with one of the reactants in titration to produce a color e.g. Starch forms a dark blue complex with I 3 -
end point of titrations with I 2 or KSCN (blood red complex) Iodometry – indirect; I 3 -
2 in
the solution Iodimetry – direct Redox Titration: Redox Titration – widely used in titrimetric analysis; analyte can be present in the sample in more than 1 oxidation state and must be converted to a single oxidation state prior to its titration; auxillary OA (e.g. H 2 O 2 ) or RA (Zn, Cd, Hg) may be added to accomplish this Standard Oxidants: Reagent Reduction Product Standard Potential (V) 1 ° Standard Used KMnO
4
Mn 2+
1.51 Na 2 C 2 O 4 , Fe, As 2 O 3
KBrO 3
Br -
1.44 KBrO
3
Ce 4+
Ce 3-
1.44 Na 2 C 2 O 4 , Fe, As 2 O 3 , K 2 Cr 2 O 7 Cr 3+ 1.33
K 2 Cr 2 O 7 , Fe I 2 I - 0.536
BaS 2 O 3 ∙H 2 O, Na 2 S 2 O 3
KMnO 4 – most widely used of all standard oxidizing agents; readily available; inexpensive, requires no indicator 𝑀𝑛𝑂
ƒ n + 8𝐻 m + 5𝑒
n ↔ 𝑀𝑛
}m + 4𝐻
} 𝑂 E° = 1.51V In Acidic Solution: 𝐻 } 𝐶 } 𝑂 ƒ ↔ 2𝐶𝑂
} + 2𝐻
m + 2𝑒
n E ° = 0.48V Reaction with KMnO4:
2𝑀𝑛𝑂
ƒ n + 5𝐻 } 𝐶 } 𝑂 ƒ + 6𝐻 m + 5𝑒
n ↔ 2𝑀𝑛
}m + 10𝐶𝑂 } + 8𝐻
} 𝑂 Ce 4+ in H
2 SO 4 – powerful oxidizing agent; yellow-orange; stoichiometry of reaction is simple; indefinitely stable but relatively high cost of Ce 4+ compounds; indicator ferroin I 2 – solutions are weak oxidizing agents; used for the determination of strong reductants; lack stability; must be restandardized regularly; I 2 is not very soluble in water; dissolved moderately in concentrated KI 𝐼 • n + 2𝑒
n ↔ 3𝐼
n Iodimetry – I 2 used as OA Iodometry – I 2 used as RA; standard Na 2 S 2 O 3 used to titrate I 2
liberated by reaction of analyte in measured excess KI in slightly acidic solution 𝐼 } + 2𝑆 } 𝑂 • }n ↔ 2𝐼 n + 𝑆
ƒ 𝑂 Y }n 𝐼𝑂 • n + 5𝐼
n + 6𝐻
m ↔ 3𝐼
} + 3𝐻
} 𝑂 3𝐼 } + 6𝑒
n ↔ 6𝐼
n 1 𝑚𝑜𝑙 𝐼𝑂
• n ≡ 3 𝑚𝑜𝑙𝑠 𝐼 } ≡ 6 𝑚𝑜𝑙𝑠 𝑆 } 𝑂
}n 𝐸𝑊 µš— - = 𝐹𝑊 µš— - 6 Indicator: Starch (specific indicator) – deep blue complex with I 2 Standard Reductants – standard solutions of reducing agents tend to react with atomic O2; titrations carried out in reagents under an inert atom; Indirect method used – aliquot containing excess reductant added to the sample and the excess is quickly back titrated with standard oxidant Download 383.93 Kb. Do'stlaringiz bilan baham: 
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