Download Zn 2+

Chapter 11
EDTA Titrations
11-1 Metal-Chelate Complexes
Metal ion (M)
+
Ligand (L)
Lewis acid
Lewis base
Electron-pair
acceptor
Electron-pair
donor
=
Complexes (ML)
Adduct
coordination
compound
(
Monodentate:
(Chapter 6)
)–
Multidentate:
BOX 6-2 Notation For Formation Constants
- Complex formation reaction is stepwise reaction
M  X  MX
MX  X  MX2
••
•
MXn1  X  MXn
K1  [MX] [M][ X ]
K 2  [MX2 ] [MX][ X ]
••
•
K n  [MXn ] [MXn1 ][ X ]
K : formation constant (Kf, or stability constant)
M  X  MX
M  2 X  MX2
••
•
1  [MX] [M][ X ]  K1
2
 2  [MX 2 ] [M ][ X ]  K 1K 2
••
•
 n  [MX n ] [M ][ X ]  K 1K 2    K n
βn: overall(cumulative) formation constant
M  nX  MXn
n
• Chelating ligand (or Chelates) ; bind to a metal ion through more than
one donor atoms in a single ligand. (multidentate ligand)
Ex) bidentate, tetradentate, pentadentate, hexadentate, …
H2NNH2: Bidentate
ligand
EDTA: Hexadentate
ligand
Mn+
Mn+
Mn+
Macrocycles
(Ionophore)
Chelating
ligand
• Chelate effect: ability of multidentate ligands to form stronger metal
complexes compared to monodentate ligands.
(Bidentate ligand)
(11-1)
(Monodentate ligand)
(11-2)
2 ethylenediamine molecules binds tighter than 4 methylamine molecules
⇒ Larger K value for multidentate ligand (Chelate increase the stability
of complex.)
The octadentate ligand in Figure 12-3 is being evaluated as an anticancer
agent.5
The chelate is covalently attached to a monoclonal antibody, which is a
protein produced by one specific type of cell in response to one specific
foreign substance called an antigen.
11-2 EDTA
Acid-Base Properties
 EDTA (Ethylenediaminetetraacetic acid)
• One of the most common chelating agents as a titrant.
• EDTA has 2N & 4O in its structure giving it 6 free electron pairs.
- High Kf values with metal ions, - polyprotic acid
H6Y2+
- Neutral acid(H4Y)
- Commonly used EDTA reagent : Na2H2Y∙2H2O (Na2EDTA∙2H2O)
(⇒ Reagent grade, dried to the composition Na2H2Y∙2H2O at 80℃)
H6Y2+
Low pH

-H+
+H+
H5Y+
-H+
+H+
H4Y
-H+
+H+
H3Y-
-H+
+H+
H2Y2-
-H+
+H+
HY3-
-H+
+H+
Y4-
High pH
Acid-Base Forms
- EDTA exists in up to 7 different acid-base forms depending on the
solution pH.
- The most basic form (Y4-) is the one which primarily reacts with
metal ions.
Fractional composition diagram for EDTA
Fraction of EDTA in the form Y4- (=αY4-)
☞ Fraction (α) of the most basic form of EDTA(Y4-) is defined
by the H+ concentration and acid-base equilibrium constants.
αY 4  
[H6 Y 2 
[Y4  ]
] [H5 Y  ] [H 4 Y ] [H3 Y  ] [H2 Y 2  ] [HY 3  ] [Y4  ]
αY 4 
[Y4  ]
(11-3) where [EDTA] is the total concentration

of all free EDTA species in solution
[EDTA]
Y
 K 1K 2 K 3 K 4 K 5 K 6 {[ H  ]6  [H  ]5 K 1
4
 [H  ] 4 K 1K 2  [H  ]3 K 1K 2 K 3  [H  ] 2 K 1K 2 K 3 K 4
 [H  ]K 1K 2 K 3 K 4 K 5  K 1K 2 K 3 K 4 K 5 K 6 }
αY 4  
K 1K 2K 3K 4K 5K 6
D
(11-4)
αY4- is depended on the pH of the solution
Ex) αY4- of 0.10 M EDTA at pH 6.00?
Sol) at pH 6.00→[Y4-]=1.0×10-6M
αY 4 
[Y4  ] 1.8 10 -6


 1.8  10 -5
[EDTA]
0.10
EDTA Complexes
•
The basic form of EDTA (Y4-) reacts with most metal ions to form a 1:1 complex.
(Other forms of EDTA will also chelate with metal ions)
(11-5)
Note: This reaction only involves Y4-, but not the other forms of EDTA
⇒The equilibrium constant for the reaction of a metal with a ligand is called
the formation constant (Kf) or the stability constant:
•
Recall: the concentration of Y4- and the total concentration of
EDTA ([EDTA]) are related as follows:
αY 4 
[Y4  ]

[EDTA ]
[Y4  ]  αY 4  ΕDΤΑ

where αY4-is dependent on pH (From table value)
•
The basic form of EDTA (Y4-) reacts with most metal ions to form a 1:1 complex.
+n ion:
Mn+
+
Y4-
MYn-4
[MY n  4 ]
Kf  n  4 
[M ][Y ]
Conditional formation constant
M
n
 Y 4
MY
n 4
[MY n  4 ]
Kf 
[M n  ][Y4  ]
(Kf : Table 11-2)
[Y4  ]  αY 4  ΕDΤΑ

[MY n  4 ]
Kf 
[M n  ]αY4- [EDTA]
K f  K fαY4-
(αY4- : Table 11-1)
where [EDTA] is the total
[MY n  4 ]

(11-6) concentration of EDTA added to the
[M n  ][EDTA]
solution not bound to metal ions
Kf’: Conditional formation constant (at given pH)
⇒ If pH is fixed by a buffer, then αY4- is a constant (Table 11-1)
that can be combined with Kf (Table 11-2)  evaluate Kf‘
⇒ Kf ‘ is constant for a given pH
Example
Using the Conditional Formation Constant
What is the concentration of free Ca2+ in a solution of 0.10 M CaY2- (Kf=1010.65)
at pH 10.00 and at pH 6.00 ?
Solution
Ca2+ + EDTA
K f  K fαY4-
CaY2-
K f  K fαY4-
[CaY 2- ]

[Ca2  ][EDTA]
at pH 10.00 : Kf’=(1010.65)(0.30)=1.3×1010
at pH 6.00 : Kf’=(1010.65)(1.8×105)=8.0×105
Ca2+ + EDTA
Initial conc.(M)
Final conc.(M)
K f  K fαY4-
0
x
0
x
CaY20.10
0.10-x
[CaY 2- ]
0.10  x


[Ca2  ][EDTA]
x2
∴x=[Ca2+] = 2.7×10-6 M (at pH 10.00 )
= 3.5×10-4 M (at pH 6.00 )
☞ M-EDTA complexes becomes less stable at lower pH
(higher concentration of [Ca2+] at lower pH)
 Complexometric Titrations are based on
the reaction of a metal ion with a chemical agent(ligand)
to form a metal-ligand complex.
- Determination of metal ion concentration
- Standard solution: chelating agent
- Ligand forms strong 1:1 complexes with
most metal ion (The stoichiometry is 1:1
regardless of the charge on the ion)
☞Higher Kf  complete titration
reaction (~99.9%) at the equivalence point
 pH effect on EDTA titration
• Note that the metal–EDTA complex
becomes less stable as pH decreases and
Kf decreases.
[Ca2+]=2.7×10-6 M at pH 10.0
[Ca2+]=3.5×10-4 M at pH 6.0
• In order to get a “complete”(say, 99.9%)
titration, EDTA requires a certain minimum
pH for the titration of each metal ion.
Ex) pH effect for the titration of Ca2+
⇒ Below pH≈8, the end point is not sharp
enough to allow accurate determination.
(∵The K for CaY2- is just too small for
“complete” reaction at low pH.)
Minimum pH for Effective
Titration of Metal Ions
11-3 EDTA Titration Curves
• The titration of a metal ion with EDTA is similar to the titration of a strong
acid (Mn+) with a weak base (EDTA)
Mn+ + EDTA
MYn-4
K f  K fαY4-
[MY n  4 ]
K f  K fαY4- 
[M n  ][EDTA]
• The titration curve has three distinct regions:
Region 1: Before the equivalence point
: excess Mn+ left
: calculation unreacted Mn+ (free Mn+)
Region 2: At the equivalence point
: exactly as much EDTA as metal
: calculation free Mn+ from dissociation of
MYn-4 ([EDTA]=[Mn+])
Region 3: After the equivalence point
: excess EDTA left
: calculation free Mn+ from dissociation of
MYn-4
(11-7)
Titration curve
(VEDTA vs. pM)
Titration Calculations
Ex. Construct the titration curve for 50.0 ml of a 0.0400 M Ca2+ solution
(buffered at pH 10.00) with 0.0800 M EDTA
Titration reaction: Ca2+ + EDTA → CaY2-
(From Table 11-2, 11-3)
Kf(CaY2-)=1010.65 , αY4-=0.30 (at pH 10)
 K f  K fαY4-  (1010.65 )(0.30) 1.34  10 10
⇒ Kf’ is large, the reaction goes to completion
with each addition of titrant.
☞ The equivalence volume (Ve) is,
50.00 mL   0.0400 M   V e   0.0800 M
mmol Ca2+
mmol EDTA
 V e  25.00 mL
Titration curve: pCa2+ vs. VEDTA

Before the Equivalence Point
Ex. 50.0 ml of a 0.0400 M Ca2+ (at pH 10.00) with 5.00 mL of 0.0800 M EDTA
mmoles of Ca2+=original mmoles of Ca2+ – mmoles of EDTA added
 50.00 mL   0.0400 M   5.00 mL   0.0800 M   1.60mmol
Volume is 55.00 mL (=50.00 mL + 5.00 mL)
1.60 mmol
 0.0291 M
55.00 mL
 pCa  log[Ca2  ]  -log(0.0291)  1.54
[Ca2  ] 
At the Equivalence Point
Ex. 50.0 ml of a 0.0400 M Ca2+ (at pH 10.00) with 25.00 mL of 0.0800 M EDTA
- Virtually all of the metal ion is now in the form CaY2- (∵Kf≫1)
- Just enough EDTA has been added to consume Ca2+
- pCa determined by dissociation of CaY2[CaY 2  ]  (0.0400M)
Initial conc.(M)
Final conc.(M)
(50.0mL)
 0.0276M
(50.0mL  25.0mL)
Ca2+ + EDTA
0
x
K f' K fαY4-  1.34  10
0
x
10
CaY2-
0.0276
0.0276-x
[CaY 2  ]
0.02760  x


[M Ca2  ][EDTA]
x2
x=[Ca2+]=1.4×10-6 M ⇒ ∴pCa2+ = -log(1.4×10-6) = 5.85
After the Equivalence Point
Ex. 50.0 ml of a 0.0400 M Ca2+ (at pH 10.00) with 26.00 mL of 0.0800 M EDTA
☞ Virtually all of the metal ion is now in the form CaY2- and there is
excess, unreacted EDTA.  A small amount of free Ca2+ exists in
equilibrium with CaY2- and EDTA.
- Calculate excess, unreacted moles of EDTA:
mmols of total EDTA – mmoles of Ca2+
=(26.00mL)(0.080) – (50.0mL)(0.040) = 0.08mmol
- Calculate excess, unreacted [EDTA]:
[EDTA] 
0.08mmol
 1.05  10 3 M
(50.0mL  26.0mL)
- Calculate [CaY2-]:
[CaY 2  ]  (0.0400M)
Initial conc.(M)
Final conc.(M)
K f' K fαY4-  1.34  10
10
(50.0mL)
 2.63  10 2 M
(50.0mL  26.0mL)
Ca2+ + EDTA
0
x
0.00105
0.00105+x
CaY2-
0.0263
0.0263-x
[CaY 2- ]
0.0263 - x
0.0263



[Ca2  ][EDTA] [Ca2  ](0.00105 x] [Ca2 ](0.00105)
 [Ca 2  ]  1.9 10 -9M  pCa
2
 8.73
The Titration Curve
: Ca2+ and Sr2+ show a distinct break
at the equivalence point, where
the slope is greatest.
Kf’ & pH effects on titration
- The equivalence point is sharper for
Ca2+ than Sr2+. This is due to Ca2+
having a larger Kf’.
- If the pH is lowered, the Kf’ decrease
(because αY4- decrease), and the end
point becomes less distinct.
⇒ The completeness of these reactions is
dependent on αY4- and correspondingly pH.
⇒ The pH cannot be raised arbitrarily high,
because metal hydroxide precipitate.
The pH is an important factor in setting the
completeness and selectivity of an EDTA titration.
11-5 Auxiliary Complexing Agents
 Metal Hydroxide (M(OH)x)
: In general, as pH increases a titration of a metal ion with EDTA will have a
higher Kf.
⇒ Larger change at the equivalence point as pH increases.
⇒ Exception: If Mn+ reacts with OH- to form an insoluble metal hydroxide
 Auxiliary Complexing Agents: a ligand can be added that complexes with
Mn+ strong enough to prevent hydroxide formation.
- Binds metal weaker than EDTA(⇒Auxiliary complexing
agents are displaced by EDTA during the titration).
- Ammonia, tartrate, citrate or triethanolamine, …..
Ex) Zn2+ in ammonia buffer (pH 10.00) to prevent Zn(OH)2(s)
- At pH=10.00 ([OH-]=104 M),
Ksp(Zn(OH)2)=3.0×1016=[Zn2+](104)2 ⇒[Zn2+]=3.0×108 M
→[Zn2+] should be less than 3.0×108 M to prevent Zn(OH)2(s)
☞ Fix the pH 10.00 for Zn2+ solution,
i) by OH-: [Zn2+]>3.0×108 M→Zn(OH)2 precipitation→titration (X)
ii) By ammonia buffer: soluble Zn-NH3 complex ion→titration (O)
⇒ Ammonia complexes the metal ion to keep it in solution at pH 10.
Metal-Ligand Equilibria
- Consider a metal ion that form two complexes with
the auxiliary complexing ligand L:
M  L  ML
M  2L  ML
[ML]
(11-13)
[M][L]
[ML 2 ]
β2 ( K 1K 2 ) 
(11-14)
[M][L]2
β1 ( K 1 ) 
2
βn: overall(cumulative) formation constant
• Fraction of free metal ion(αM): the fraction of metal ion in the
uncomplxed state
[M]
αM 
M tot
M
tot
(11-15)
[M]: conc. of metal ion in the uncomplxed state
Mtot: total conc. of all forms M (=M, ML, ML2)
 [M]  [ML]  [ML 2]
[ML]  β1 [M][L]
[ML 2 ]  β2 [M][L]2
M
tot
 [M]  β1 [M][L] β2 [M][L]2
 [M]{1  β1 [L] β2 [L]2 }
 αM 
[M]
[M]
1


CM
[M]{1  β1 [L] β2 [L]2 } 1  β1 [L] β2 [L]2
⇒ depends on the equilibrium constants or cumulative
formation constants
(11-16)
Ex)
Ammonia Complexes of Zinc
Example
Calculate αZn2+ in ammonia buffer (NH3=0.10M) solution.
⇒ All Zinc species: Zn2+, Zn(NH3)2+, Zn(NH3)22+, Zn(NH3)32+, Zn(NH3)42+
From Appendix I
Solution
Zn2+
+ NH3 →
Zn(NH3)2+
β1=
Zn2+ + 2NH3 → Zn(NH3)22+
β2=
Zn2+ + 3NH3 → Zn(NH3)32+
β2=
Zn2+ + 4NH3 → Zn(NH3)42+
β2=
[Zn(NH3)2+]
=102.18
2+
[Zn ][NH3]
[Zn(NH3)22+]
=104.43
[Zn2+][NH3]2
[Zn(NH3)22+]
=106.74
2+
2
[Zn ][NH3]
[Zn(NH3)22+]
=108.70
2+
2
[Zn ][NH3]
Zntot = [Zn2+]+[Zn(NH3)2+]+[Zn(NH3)22+]+[Zn(NH3)32+]+[Zn(NH3)42+]
= [Zn2+]/(1+β1[NH3]+β2[NH3]2 +β3[NH3]3 +β4[NH3]4)
 αZn2 
[Zn 2  ]
1


Zn tot
1  β1 [NH 3 ] β2 [NH 3 ]2  β3 [NH 3 ]3  β4 [NH 3 ]4
(11-17)
L(=[NH3])=0.10M
 αZn2 
[Zn 2  ]
1
5



1.80

10
2
3
4
Zn tot
1  β1 (0.10) β2 (0.10)
 β3 (0.10)
 β4 (0.10)
(⇒ Very little zinc is in the form Zn2+ in the presence of 0.10 M NH3)
EDTA Titration with an Auxiliary Complexing Agents
• In the presence of auxiliary complexing agents, use a new conditional
formation constant that incorporates the fraction of free metal at a fixed pH.
[MY n  4 ]
: at a fixed pH (consider αY4-)
K f  K fαY4- 
n
[M ][EDTA]
[M n  ]
αM 
 [M n  ]  αM C tot
C tot
[MY n  4 ]
: at a fixed conc. of auxiliary complexing
K   K fαY4- 
agent (consider αM)
αM C tot[EDTA]
[MY n  4 ]
K f  αM αY4-K f 
C tot[EDTA]
(11-18)
K”f : Effective(or conditional)
formation constant
(☞ at a fixed pH and fixed conc. of auxiliary complexing agents)
EDTA Titration in the Presence of Ammonia
Example
Consider the titration of 50.0 mL of 1.00×103M Zn2+ with 1.00×103M EDTA at
pH 10.00 in the presence of 0.10 M NH3. Find pZn2+ after addition of 20.0, 50.0,
and 60.0 mL of EDTA.
Solution - Y4 = 0.30 (pH=10, from Table 11-1)
- Zn2+ = 1.8×105 (from Eq. 11-17)
⇒ Conditional formation constant (K”f)
= αZn2+αY4-Kf = (1.8×105)(0.3)(1.00×1016.5) = 1.7×1011
(a) Before the equivalence point (20.0 mL of EDTA)
- Zinc not bound to EDTA(CZn2+) is bound to ammonia: calculation CZn2+
mmoles of Zn2+ = original mmoles of Zn2+ - mmoles of EDTA added
 50.00 mL   0.0010 M   20.00 mL   0.0010 M   0.030mmol
Volume is 70.00 mL (=50.00 mL + 20.00 mL)
 C Zn2  
0.030mmol
 4.3  10  4 M
70.00mL
- The concentration of free Zn2+ ([Zn2+])
αZn 2 
[Zn2  ]

 [Zn 2  ]  αZn 2  C Zn 2  = (1.8×105)(4.3×10-4) = 7.7×10-9 M
C Zn 2 
∴ pZn2+ = -log[Zn2+] = 8.11
Check reality!: Zn(OH)2 precipitation at pH 10 in the presence of 0.10 M NH3.?
(Ksp of Zn(OH)2) =10-15.52 )
Q=[Zn2+][OH-]2 = (10-8.11)(10-4.00)2 = 10-16.11 <10-15.52
⇒ Do not precipitate Zn(OH) (s).
2
(b) At the equivalence point (50.0 mL of EDTA)
Zn2+ + Y4- = ZnY2
Kf” = (Zn2+)(Y4-)(Kf ) = 1.7×1011
- Virtually all of the zinc ion is now in the form ZnY2-
(∵Kf (ZnY2-) ≫ Kf Zn(NH3)x2+)
- Just enough EDTA has been added to consume Zn2+
- pZn determined by dissociation of ZnY2[ZnY 2  ]  (0.0010M)
Initial conc.(M)
Final conc.(M)
K f" 1.7  10
11
(50.0mL)
 5.0  10 -4 M
(50.0mL  50.0mL)
CZn2+ + EDTA = ZnY20
x
0
x
5.00×10-4
5.00×10-4 – x
[ZnY 2- ]
5.00  10 -4 - x


[Czn2  ][EDTA]
x2
⇒ x = CZn2+ = 5.4×10-8 M
- The concentration of free Zn2+ ([Zn2+])
αZn 2 
[Zn2  ]

 [Zn2  ]  αZn 2  C Zn 2 
C Zn 2 
⇒[Zn2+] = Zn2+CZn2+ = (1.8×105)(5.4×10-8)
= 9.7×10-13 M
∴pZn = -log[Zn2+] = 12.01
(c) After the equivalence point (60.0 mL of EDTA)
- Virtually all of the zinc ion is now in the form MgY2- and there is
excess, unreacted EDTA  A small amount of free Zn2+ exists in
equilibrium with ZnY2- and EDTA.
- Calculate excess, unreacted [EDTA]:
[EDTA] 
(0.0010M)(10.0mL)
 9.1  10 5 M
(50.0mL  60.0mL)
- Calculate [ZnY2-]:
[ZnY 2  ]  (0.0010M)
(50.0mL)
(50.0mL  60.0mL)
4
 4.5  10 M
Zn2+ + EDTA = ZnY2-
Initial conc.(M)
0
Final conc.(M)
x
9.1×10-5
4.5×10-4
9.1×10-5+x 4.5×10-4-x
10.65
K f' K fαY4-  (0.30)(10
)  9.5 10 15 (Not Kf”)
[ZnY 2- ]
4.5 10 -4  x


2
[Zn ][EDTA] [Zn2  ](9.1 10 -5  x)
4.5 10 -4

[Zn2  ](9.1 10 -5 )
 [Zn 2  ]  5.3 10 -16 M  pZn
2
 15.28
Titration curves and effect of
auxiliary complexing agents
11-6 Metal Ion Indicators
 Determination of EDTA Titration End Point
- Four Methods:
1. Metal ion indicator (This chapter)
2. Mercury electrode
Potential Measurements
3. pH electrode
(Potential(V)=logM=pM, Ch. 14~16)
4. Ion-selective electrode
Metal Ion Indicator (In): a compound that changes color when
it binds to a metal ion
MIn
In + M
■
(blue)
(red)
: Similar to pH indicator, which changes color with pH
(or as the compound binds H+)
⇒ For an EDTA titration, the indicator must bind the metal ion
less strongly than EDTA
: Similar in concept to Auxiliary Complexing Agents
: Needs to release metal ion to EDTA
Ex) In EDTA titration:
Mg-In + EDTA
Mg-EDTA + In
(11-19)
(red)
(colorless)
Before eq. point
(colorless) (blue)
At eq. point
⇒ End Point indicated by a color change from red to blue
: Most are pH indicators and can only be used over a given pH range. ⇒ Most
indicators can be used only in certain pH ranges.
Ex) Calmagite with metal ion (at pH 10):
- If a metal(M) does not freely dissociate from an indicator(In), the metal is
said to block by the indicator (∵Kf(M-In)>Kf(M-EDTA) or slow reaction).
Ex) For Eriochrome black T(EBT):
1) Direct titration of Cu2+, Ni2+, Co2+, Cr3+, Fe2+, Al3+ ⇒ impossible
(∵blocking of EBT by stable M-In complex)
2) Back titration of Cu2+ ⇒ possible
i) Add excess standard EDTA to Cu2+  Cu-EDTA+free EDTA
ii) Add In; In cannot take Cu from already formed Cu-EDTA
 Cu-EDTA+free EDTA+In
iii) Titration of excess EDTA with standard Mg2+; Mg2+ can only take
free EDTA (∵Kf (Mg2+-EDTA)<Kf (Cu2+-EDTA)
 Cu-EDTA + Mg-EDTA + Mg-In
☞ Color change of back titration at end point?
:Blue(In) to Red(Mg-In)
(∵Kf (Mg2+-In)<Kf (Mg2+-EDTA)
Guide to EDTA titrations of some
common metals; pH ranges,
auxiliary complex agents, indicators
Ex) Pb-EDTA titration;
- Possible pH range for EDTA titration:
pH 3 to 12
- Auxiliary complex agents are need to
pH 9-12
11-7 EDTA Titration Techniques
 Almost all elements can be determined by EDTA titration.
 Some Common Techniques used in these titrations include:
- Direct Titrations
- Back Titrations
+ pH control
- Displacement Titrations
with buffer solution
- Indirect Titrations
- Masking Agents
Direct Titrations
•
•
•
•
Analyte (metal ion) is buffered to appropriate pH and is titrated directly
with standard EDTA.
Kf large
Metal ion indicator does not block the metal.
An auxiliary complexing agent may be required to prevent precipitation
of metal hydroxide.
Back Titrations
•
•
Approach necessary if analyte:
- precipitates in the absence of EDTA (Ex. Al3+ at pH 7→ Al(OH)3(s))
- Reacts slowly with EDTA
- Blocks the indicator
Second metal ion must not displace analyte from EDTA
Step 1) A known excess of standard EDTA is added to analyte.
⇒Free EDTA left over after all metal ion is bound with EDTA
Step 2) The remaining excess of EDTA is then titrated with a standard
solution of a second metal ion.
Example
A Back Titration
Back titration of 25.00 mL of Ni2+ in dilute HCl with standard Zn2+ at pH 5.5
using xylenol orange indicator. ⇒ [Ni2+] = ?
(∵ Nickel reacts too slowly with EDTA)
Adding excess 25.00 mL Na2EDTA (0.05283M); [Ni-EDTA + free EDTA]
Neutralized with NaOH  then, pH adjusted to 5.5 with acetate buffer;
[Ni-EDTA + free EDTA]
Indicator (xylenol orange, In) added: [Ni-EDTA + free EDTA + In]
Titration of free EDTA with Zn2+(0.02299 M, 17.61 mL)
at end point ; [Ni-EDTA + Zn-EDTA + Zn-In ]
Solution
mmol of EDTA added = (25.00 mL)(0.05283 M) = 1.3208 mmol
mmol of free(unreacted) EDTA = (17.61 mL)(0.02299 M) = 0.4909 mmol
⇒ mmol of Ni2+ = mmol of EDTA added - mmol of free EDTA
= 1.3208 - 0.4909 = 0.9151 mmol
∴ [Ni2+] = 0.9151 mmol/42.61 mL = 0.03664 M
Ex) Back titration of Al3+: EDTA prevent precipitation of Al(OH)3 at pH 7
(formed stable Al3+–EDTA complex at pH 7)
Displacement Titration
☞ Used for some analytes that don’t have satisfactory metal ion indicators.
Step1) Analyte (Mn+) is treated with excess Mg(EDTA)2-, causes release of
Mg2+. (∵Kf(MgY2-)<Kf(MY2-) ⇒ MY2- + MgY2- + Mg2+)
Step2) Amount of Mg2+ released is then determined by titration with
a standard EDTA solution (∵Kf(MgY2-)<Kf(MY2-)
 Concentration of released Mg2+ equals [Mn+]
Step1: Mn+ + MgY2- ⇌ MYn-4 + Mg2+ (11-20)
Analyte excess
↳ Step2: titrate with standard EDTA
Ex) Hg2+ titration
Hg2+ + MgEDTA2- ⇌ HgEDTA2- + Mg2+
Then, Mg2+ is titrated with standard EDTA
Ex) Ag+ titration
2Ag+ + Ni(CN)42- ⇌ 2Ag(CN)22- + Ni2+
Then, Ni2+ is titrated with standard EDTA
Indirect Titration
☞ Used to determine anions that precipitate with metal ion.
Ex) CO32-, CrO42-, S2-, SO42Step1) Anion is precipitated from solution by addition of excess metal ion
- Ex) SO42- + excess Ba2+  BaSO4(s) (pH 1)
- Precipitate(BaSO4(s)) is filtered & washed
Step2) Precipitate(BaSO4) is then reacted with excess EDTA to bring the
metal ion back into solution (boiled at pH 10)  BaEDTA2- + EDTA
Step3) The excess EDTA is titrated with Mg2+ standard solution.
Alternatively,
Step1) Anion is precipitated from solution by addition of excess standard
metal ion
Step2) Excess standard metal ion in the filtrate is titrated with EDTA.
BOX 11-3 Water Hardness
• Hardness: the total concentration of alkaline earth ion(mainly Ca2+ and Mg2+)
- unit: mg/L as CaCO3
:[Ca2+]+[Mg2+]=1mM→CaCO3=1mM→100mg CaCO3→hardness=100mg/L
- Soft water: 0 to 60 mg/L as CaCO3, Hard water: ~270 mg/L as CaCO3
- Determination of hardness by EDTA titration
[Ca2+]+[Mg2+]: pH=10 (ammonia buffer)
[Ca2+]: pH=13 (without ammonia)(at pH 13, Mg(OH)2(s))
Masking
• Masking Agents
: A reagent added to prevent reaction of some metal ion with EDTA
(⇒remove interferences of specific metal ion)
Al3+ + 6F- → AlF63-
⇒Al3+ is not available to bind EDTA because of the complex with FRequires: K f(AlF6 )  K f(Al(EDTA))
3
Ex) CN- masking: Cd2+ , Cu2+ , Ag2+ , Bi2+, …..
(CAUTION: CN- formed toxic HCN gas below pH 11)
F- masking: Al3+, Fe3+, Ti+, Be2+
(CAUTION: HF formed by F- in acidic solution)
Triethanolamine masking: Al3+, Fe3+, Mn2+
2,3-Dimercaptopropanol masking: Bi3+, Cd2+, Cu2+, Hg2+, Pb2+
• Demasking: refers to the release of a metal ion from a masking agent
Ex) Cyanide demasking with formaldehyde
Masking, demasking, pH control ⇒ selective titration of individual metal
ion from complex mixtures of metal ions