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Analytical Chemistry > Exam 2 > Flashcards

Exam 2 Flashcards

1
Q

Which ions are usually soluble when apart of a molecule?

A

Cl-, Br-, I-

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2
Q

Which ions are an exception to the solubility rules?

A

Ag+, Pb2+, Hg22+

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3
Q

Ksp

A

Solubility product. The product of the concentration of the ions of a substance in a saturated solution of the substance, the concentrations being expressed as molarity.

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4
Q

When does the reaction form a precipitate?

A

-If two solutions, each containing one of the ions of the sparingly soluble salt, are mixed, no precipitation will occur unless the reaction quotient exceeds the solubility product, namely, unless QSP > KSP

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5
Q

Ksp depends on what

A

Depends on temperature

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6
Q

How to set up an ICE table?

A

See picture

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7
Q

What do you not include in equilibrium expression?

A

Pure solids and pure liquids

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8
Q

Common ion effect (product)

A

-shifts eq to the left

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9
Q

Approximation for equilibrium

A

When the concentration is very small (10^-3)

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10
Q

Solubility of PbI2 when common ion is present

A

-Lead(II) “flows” from Pb2+ into all other soluble forms and is itself continuously replenished as more PbI2(s) dissolves
-Holes in the sand fill with an ocean wave: Lead(II) “flows” from Pb2+ into all other soluble forms and is itself continuously replenished as more PbI2(s) dissolves
-Because of all these equilibria, PbI2(s) is more soluble in solutions of I- than in pure water.

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11
Q

Do precipitation reactions ever go to completion? Why is this important?

A

-No
-In the gravimetric analysis of silver as AgCl(s)

Ag+(aq) + Cl-(aq) → AgCl(s)

there will always be a small amount of unprecipitated Ag+ in equilibrium with AgCl(s)
-This may be ignored for a small ksp

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12
Q

When would you use AAS technique?

A

-When Ksp is not too small.

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13
Q

Complex Ion Formation equilibria

A

-Analytes that form complex ions require attention
-Need to minimize their formation if the aim is to precipitate the metal

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14
Q

Metal ions exists as what in equilibria?

A

As complex ion: Ag+, AgCl2-(aq)

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15
Q

Weak acids in equilibria

A

-Have different protonation state: H2SO4, HSO4-, SO4^2-
-Predominant species depends on pH in comparison to pKa value(s)

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16
Q

Solubility depends on what?

A

Temperature

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17
Q

Metal ions exists as what?

A

Exist in different oxidation state

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18
Q

Analyte-Reagent Equilibria (excess reagent)

A

Generally, reactions are going to completion, but some require excess reagent, because they are equilibria, or they could have slow kinetics.

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19
Q

Examples of analyte-reagent equilibria

A
  • Endpoint of titration of weak acid HA with NaOH: A- + H2O ⇄ HA + OH-
  • Endpoint of titration of weak base B with HCl: BH+ ⇄ H+ + B
  • REDOX titration follows Nernst equation: equilibrium between two oxidation states
  • Acid-Base indicator is an equilibrium strategically selected to be at pH of equivalence point
  • In EDTA titration, M-EDTA and M-Indicator are in equilibrium with free Mn+
  • During liquid-liquid extraction into organic layer, control solution pH to have all HA or all B
  • Acid-Base buffer keeps analyte in proper speciation: HPLC; Selective Electrode; absorbance
    measurement of acid-base indicator; EDTA and indicators need just the right pH
  • AAS: include easily ionizable element
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20
Q

Techniques that rely on equilibria for their operation

A

-Potentiometric Sensor
-Ion Chromatography
-GC with headspace sampling
-Liquid Chromatography
-mass spec

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21
Q

Potentiometric Sensor

A

pH electrode, ion selective electrode – electrochemical potential

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22
Q

Ion Chromatography

A

ion-exchange equilibria between mobile and stationary phase

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23
Q

GC with headspace sampling

A

partition equilibrium between sample and gas phase

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24
Q

Liquid chromatography

A

retention depends on equilibrium partitioning between stationary and mobile phases

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25
Mass spectrometry
Requires charged particles i.e. M/z * Electrospray ionization equilibria (protonation/deprotonation) affects ion formation and sensitivity * Chemical ionization – equilibrium between reagent gas and analyte molecules for ion formation
26
Does weak acid equilibria have a bottomless supply of H2SO3?
No
27
Charge balance
Statement of electrical neutrality in soln
28
Mass balance
Statement of the law of conservation of mass
29
How do you calculate mass balance?
F is the formal concentration (Concentration that you weighed) FNaA=[Na+]
30
What are some assumptions that are made when using the HH equation
-OH- is neligable is solns more acidic than pH of 6 -Addition and substration of H+ can be neglected if FNaA and FHA are not too small or Ka is not too large.
31
Buffer
-Connectes to analyte-reagent equilibria -Solution containing weak acid HA and its conjugate base salt NaA; acid dissoc Ka -Within buffer capacity, strong base consumes HA and converts it to NaA
32
When do we use buffers?
Buffers are employed in pH-sensitive measurements – * EDTA titration of Ca2+ and Mg2+ with Eriochrome Black T at pH 10 * Fluoride ion selective electrode (ISE) measured at pH 5 with TISAB
33
Why is buffer capacity important?
Buffer capacity (defined below) is important * in instrumental methods (eg Ion Chromatography) * and biological systems (pH 7.2 to 7.4) with HCO3- + CO32- buffer * with buffer capacity of 0.01 (low end) to 0.05 (high end) 𝑀 𝑝𝐻
34
When are buffers more effective?
pH range of pka +- 1
35
What is buffer capacity?
Is the amount of H+ or OH- that would consume ≈90% of the acid or base form of the buffer system
36
How do you calculate buffer capacity
see pic
37
What is buffer capacity affected by?
Dilution, temperature, and ionic strength
38
Why do the solubility increase when salts are added to the soln?
-Ionic atmosphere
39
Do ions other than the ions present in soln affected eq?
Yes
40
Ionic atmosphere
* Region of solution around an ion or charged particle * Contains an excess of the oppositely charged ions The greater the ionic strength of solution, the higher the charge on each ionic atmosphere. * Cation and anion attraction is reduced (relative to attraction in distilled water) * Tendency of ions to come together is reduced * Solubility of salt is increased
41
How to calculate ionic strength
42
How to account for ionic strength?
Concentrations are replaced by activities (activity and activity coefficient)
43
Activity
Value that replaces concentration in a thermodynamically correct equilibrium expression Activities ( represent the "effective" concentration of an ion in solution, influenced by interactions with other ions.
44
Activity coefficient
* Measures the deviation of behavior from ideality * Numeric value used to adjust concentration for ionic strength * If gamma= 1, behavior would be ideal
45
What does the correct form of the equilibrium constant account for?
Ionic strength
46
How to determine activity coefficients
Look at pic
47
Metal ions in metal chelate complexes
-Metal ions are Lewis acids, accepting electron pairs from electron-donating ligands that are Lewis bases.
48
Monodentate ligands
bind to a metal ion through only one atom
49
Multidentate (or chelating) ligands
Bind to a metal ion through more than one ligand atom.
50
Chelating Effect
The chelate effect is the ability of multidentate ligands to form more stable metal complexes than those formed by similar monodentate
51
Complexometric titrations
* are based on metal-ligand complex formation. * Ligands form strong 1:1 complexes with all metal ions. * Except univalent ions such as Li + , Na+ , and K+ * Stoichiometry is 1:1 regardless of the charge on the ion. -EDTA is used
52
What's the most widely used chelator?
EDTA
53
Every element can be measured using...
Direct or indirect titration with EDTA
54
EDTA properties
-Forms strong 1:1 complexes with most metal ions -Employed in quantitative analysis
55
pH change for EDTA
See pic
56
Conditional Formation constant (Kf)
The equilibrium constant for a metal-ligand reaction -K f for EDTA is defined in terms of the species Y 4− reacting with metal ion -Does not mean that only Y 4− reacts with metal ions -K f for most EDTA complexes large and tend to be larger for more positively charge ions
57
Conditional Formation Constant (Kf')
The equilibrium constant at a specific pH
58
Metal EDTA complexes become
Less stable at lower pH
59
What can be used in order to select which metal will be titrated by EDTA
Metals with higher formation constants can be titrated at lower pH, so pH can be used to select which metal will be titrated by EDTA
60
Regions of an EDTA titration curve
Three regions of an EDTA titration curve: * Before Ve: Excess unreacted M n+ is present. [Metal] is determined by excess M n+ in solution. * At Ve: There is exactly as much EDTA as metal in the solution. [Metal] is determined by dissociation of MY𝑛−4 according to Kf'. * After V e : Excess unreacted EDTA is in the solution. Use excess [EDTA] in Kf' expression to calculate [metal].
61
Auxiliary Complexing Agents
-An auxiliary complexing agent is a ligand that binds to the metal strongly enough to prevent metal hydroxides from precipitating because at high pH, many metals are prevented from titration with EDTA due to the formation of metal hydroxides. -The ligand binding must be weak enough to give up metal to EDTA during titration. -Common examples are ammonia, tartrate, citrate, and triethanolamine
62
Cumulative formation constants (beta)
-can be rearranged and combined with the mass balance equation to form an equation for the fraction of free metal ion.
63
Effective formation constant (Kf'')
the effective formation constant at a fixed pH and fixed concentration of auxiliary complexing agent.
64
EDTA titration with auxiliary complexing agents
See pic
65
Metal indicators
-If a metal does not freely dissociate from an indicator, the metal is said to block the indicator. -Eriochrome black T is blocked by Cu2+ , Ni2+ Co2+ , Cr3+ , Fe2+ , and Al 3+ . -It can only be used for back titrations of these metals (not direct titrations). - Tiron is an indicator for EDTA titration of Fe(III) at pH 2–3 at 40°C. Color change is from blue to pale yellow.
66
EDTA titration techniques
Indirect, Direct, back, and displacement titration
67
Direct titration
-analyte titrated with standard EDTA How it works: * Analyte is buffered to ensure large formation constant and distinct color changes. * Auxiliary complexing agents may be used to prevent metal ion precipitation
68
Back titration
known excess of EDTA added to analyte and excess EDTA titrated with a standard solution of a second metal; necessary if: * Analyte precipitates in the absence of EDTA * Analyte reacts too slowly with EDTA * Analyte blocks the indicator Note: The metal ion for the back titration must not displace analyte from EDTA.
69
Displacement titration
metal does not have a satisfactory indicator but it can be displaced by a metal that does
70
Indirect titration
anions that precipitate with certain metal ions can still be analyzed with EDTA
71
Masking
prevents one species from interfering in the analysis of another (not restricted to EDTA titrations)
72
Masking Agent
-protects some component of the analyte from reaction with EDTA. * Al3+ can be masked by F− . * CN − masks Cd2+ , Zn2+ , Hg2+ , Co2+ , Cu+ , Ag+ , Ni2+ , Pd2+ , Pt2+ , Fe2+ , or Fe 3+ but not Mg2+ , Ca2+ , Mn2+ , or Pb2+
73
Demasking
releases a metal ion from a masking agent
74
Example for demasking
Cyanide complexes can be demasked with formaldehyde.
75
Hardness
is the total concentration of alkaline earth ions in water, predominantly Ca2+ and Mg2+. Hard water reacts with soap to form an insoluble product.
76
How do you measure the hardness of water?
-To measure hardness, sample is treated with ascorbic acid to reduce Fe3+ to Fe2+ and Cu2+ to Cu+ . -Cyanide is added to mask Fe2+ , Cu+ , and several other minor metal ions. -Titration with EDTA at pH = 10 in NH 3 buffer determines total concentration of Ca2+ and Mg+ . -Separate titration at pH = 13 will cause Mg(OH)2 to precipitate leaving Ca2+ to react with EDTA. -Interference from many metals can be reduced by the right choice of indicators.

Analytical Chemistry (1 decks)

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