Download Compulsory textbook Recommended textbooks Topics of the first

Compulsory textbook
Fundamentals of Analytical Chemistry, D. A. Skoog,
D. M. West, F. J. Holler and S. R. Crouch,
Brooks/Cole 2004
Brooks/Cole,
2004, 8th Edition
Recommended textbooks
Principles of Instrumental Analysis, D. A. Skoog, F.
J. Holler and T. A. Nieman, Saunders College
Publishing 1999
Publishing,
1999.
Quantitative Analytical Chemsitry, J. S. Fritz and G.
H. Schenk, Allin and Bacon, 1987.
Instrumental Methods of Analysis, H. H. Willard et
al., Wadsworth Publ. Co., 1988.
Topics of the first semester
1.
2
2.
3.
4.
5.
6.
7
7.
8.
General introduction (1)
Fundamental concepts in analytical chemsitry (1)
Gravimetric methods in the analysis (1)
Titrimetry – general principles and concepts (2)
Precipitation titrations (argentometry) (1)
Neutralization titrations (acidi-alkalimetry) (2)
Complexometric titrations (chelatometry) (1)
Redox titrations (oxidi-reductometry) (3)
1
Analytical chemistry
involves
separating,
identifying and
d t
determining
i i the
th relative
l ti amounts
t of
f the
th
components (analytes) of the sample
Qualitative analysis
what is present? – chemical identity of the
species in the sample (preceeds quant. anal.)
Quantitative analysis
how much is present? – percentage or mass of
the analyte in the sample
Separation techniques (chromatographies)
different components may interfere one with
another
Analysis types
1. Complete analysis – each constituent is analysed
y
– each element is
2. Ultimate ((elemental)) analysis
analysed
3. Partial analysis – the amount of selected
compounds/atoms/components
Examples
water analysis
blood sample analysis
N,S,P,C-content in foodstuffs
serial analysis of a pharmaceutical product
household gas analysis
air analysis
etc., etc., etc.
2
Methods of analytical chemistry
Classical
gravimetry
i t
volumetric methods or titrations
Instrumental
electroanalysis
spectrometric analysis
magnetic methods
thermal methods
miscellaneous methods
To be considered
1. accuracy & reliability required
vs. economics
2. no. of samples to be analysed
3. complexity of the samples
3
representative sampling –
when the sample truly
represents the object to
be analysed
grinding (homogeneity)
drying (deliquescence)
homogeneous sample: its
constituents can be
d
distinguished
h d visually
ll or
with the aid of a light
microscope
( heterogeneous sample)
4
replicate samples:
portions of the material
of (approximately) the
same size that are
carried through the
analytical procedure
weighing (by an
analytical balance –
measurement of
f mass))
pipetting (by a pipette –
measurement of volume)
preparing aqueous
solutions
solubilization (digestion)
5
interference:
species other than the
analyte, which
interferes with the
results of the
measurement, i.e.,
causes errors
The measured property, X has
to vary in a known and
reproducible way with the
concenctration of the analyte,
cA
Ideally
cA = k×X
X – the signal
k – characteristic to the method,
usually
s ll unknown
k
(except
( x
t
gravimetry and coulometry)
calibration – the process of
determining k
6
For the calculations
1. experimental data
2. stoichiometry
3. instrumental data
are required
uncertainties
associated with the
measurements
must be known –
errors in the chemical
analysis
7
Chapter 2.
Chemicals, Apparatus and
Unit Operations of Analytical Chemsitry
- dealt with in p
practical (compulsory)
( mp
y)
Chapter 3.
Using Spreadsheets in Analytical Chemistry
- dealt with in practical (optional)
Calculations used in analytical chemistry
Atom –the smallest particle of an element
Molecule - the smallest particle of a compound
Compounds are combination of elements –
molecules are made up of atoms
The important thing for an (analytical) chemist is the number of
atoms reacting (and not the mass)
Atomic mass (Ar): relative masses based on the 12C isotope
Molecular mass (Mr): the sum of the atomic masses of the atoms
that make up the molecule
The chemical mass unit: the mole (1 mole = 6.022×1023 atoms of an
element or molecules of a compound)
Number of moles (n) =
grams of material (m)
_____________________________
formula mass (Ar or Mr)
8
Expressing concentration of solutions 1.
Molar concentration (molarity) the number of
moles of solute present in 1 L of solution
c=
number of moles of the solute
volume of solution
unit: mole/litre or mole/dm3 or M
unit
(equal to mmol/mL!!!)
Expressing concentration of solutions 2.
Molal concentration (molality or Raoult’s- concentration) –
m=
number of moles of solute
mass of solvent
unit: mole/kg
Advantage of m over c: m is independent of
temperature
9
Expressing concentration of solutions 3.
Mole fraction –
X=
number of moles of solute
the moles of solvent + the moles of solute
unit: Grams per volume – the mass of the solute divided by the
volume of solution
mass of solute
volume of solution
unit: g/L
Expressing concentration of solutions 4.
ppm – the mass of the solute in mg divided by the volume of the
solution in litre
mass of solute in mg
concentration in ppm =
volume of solution in litre
The mass of 1 litre of water equals to 1000g
Unit: ppm – part(s) per million,
ppb – the mass of the solute in µg divided by the volume of the
solution in litre
mass of solute in µg
concentration in ppb =
volume of solution in litre
Unit: ppb, part(s) per billion
10
Expressing concentration of solutions 5.
Mass percent – the mass of solute divided by the mass of solution
mass percent =
mass of solute
mass of solution
x100
Unit : g/100g or m/m%
Volume percent - the volume of solute divided by the volume of
solvent
vol% =
volume
l
of
f solute
l t
volume of solution
x100
Unit : mL/100mL or V/V%
Expressing concentration of solutions 6.
Analytical molarity – the total number of moles of solute present in
a given volume of solution (it says nothing about the actual state of
the solute, whether it ionizes or not, etc.)
S b l c or cT
Symbol:
Equilibrium molarity – the concentration of ions or molecules actually
present in solution, taking into account the possible dissociation of
the solute into ions
Symbol: […]
The analytical concentration is equal to the sum of the equilibrium
concentrations of the various forms of the solute
Example:
HAc
H+ + Ac-
cHAc = [Ac-] + [HAc] – this is a mass balance equation
11
Errors in chemical analysis –
how certain can we be about the results we obtain?
Obtained values (results) for a given quantity from N replicates:
x1, x2, x3, …, xN
Mean (median, arithmetic mean, average), x
x=
1
N
N
∑x
i =1
i
Precision – the reproducibility of the measurements, or the closeness of
results that have been obtained exactly in the same way; can be
obtained
b i db
by repeating
i the
h measurements
(MEMO-technique: pre=rep)
Accuracy – the closeness of our measurements to the true or accepted
values; cannot be obtained by repeating the measurements; expressed in
terms of the absolute error: E = xtrue - xi
12
Types of errors
Random (indeterminate) error – causes data to be scattered
symmetrically
y
y around the mean value; associated with the
precision (or the reproducibility) of the measurement
Systematic (determinate) error – (for example instrumental,
method or personal error); causes the mean of the data set
to differ from the true value; associated with the accuracy
of the measurement
Gross error – they occur occasionally and lead to outliers
Characterization of random errors, i.e., the
precision of the measurement
Gaussian (normal error) curve shows the symmetrical distribution of data
around the mean of an infinite set of data
13
Sample standard deviation, s –
the measure of precision of a measurement
N
s=
∑ (x − x )
2
i
i =1
N −1
di = x – xi – the deviation of the i-th result, xi from the mean;
N - the total number of the measurements;
N-1
N
1 – number of degrees of freedom
s – standard deviation
s2 – sample variance
Significant figures
The significant figures in a number are all the certain digits plus
the first uncertain digit
– most important example: reporting a burette reading
l t us say, that
let
th t th
the b
burette
tt iis of
f 25 mL
L capacity
it
smallest division is 0.1 mL
12.24 mL
the first three digits are certain
the last digit is estimated, i.e., uncertain
if you report 12 mL – rounding off error (ca. 1.8 %)
if you report 12.24478 mL – nobody believes you (rightly so)
if you try to read < 0.1 mL on the same burette –
meaningless result
- another very important example: reading an analytical balance
0,9668 g
the sample weighed must be at least 100 mg (to keep error at
≤1% level)
14
Sampling
The analytical method of choice depends on the sample size and
constituent type
sample size
> 0.1 g
0.01-0.1 g
0.0001 g – 0.01 g
< 0.0001 g
type of analysis
macro
semimicro
micro
ultramicro
analyte
l
level
l
l
1% - 100%
0.01%(100 ppm) -1%
100 ppm – 1 ppb
< 1 ppb
type of
f constituent
major
minor
trace
ultratrace
Minimizing errors in analytical procedures
1.
2.
3.
4.
5.
6.
7.
Choosing the correct blank solution
Application of separation techniques – elimination of
i t f
interferences
s
Saturation – deliberate addition of large amount of
interfering components to all samples and standards (this
may degrade sensitivity and detectability)
Matrix modification – a non-interfering component is
added to modify the response, to make it independent of
the presence of the interfering species
Adding
g of a masking
g agent
g
– it selectively
y reacts with the
interfering component and makes it „invisible”
Dilution method
Matrix matching method – for example, synthetic
seawater
15
Classical methods of chemical analysis
includes
gravimetry
titrimetry
argentometry
acidi-alkalimetry
complexometry
redox titrations
Gravimetric methods
are quantitative methods that are based on determining the mass of
the pure compound to which the analyte is chemically related
The mass is always measured on an (accurate) analytical balance
Types
T
p of
f gravimetric
im t i m
methods
th ds
precipitation gravimetry
volatilazation gravimetry
electrogravimetry
thermogravimetry
gravimetric titrimetry
Steps of precipitation gravimetry:
1 an excess precipitating reagent added to the sample,
1.
sample thus the
analyte converted into sparingly soluble product (precipitate)
2. precipitate is filtered
3. precipitate is washed from impurities
4. precipitate is dried or ignited (to convert it to a product with
known composition)
5. precipitate is weighed
16
Precipitation gravimetry
A successfull gravimetric determination meets the following criteria
1
1.
The
h analyite
l
must be
b completely
l
l (quantitatively)
(
l ) precipitated
d
2. The precipitating agent reacts selectively or, at least, specifically with
the analyte
•
Selective reagent reacts only with a single chemical species (rare)
•
Specific reagent reacts with several, but limited number of
chemical species (more common)
3 The precipitate must easily filtered and washed free from
3.
contaminants
4. Must be of sufficiently low solubility (to avoid loss of the analyte)
5. Must be unreactive with constituents of the atmosphere
6. Its weighed form must be of known composition (gravimetric factors)
Solubility of precipitates
Precipitate – it is formed from a solution which is
supersaturated with respect to the solute; when no more
precipitate is ible to form, the remaining solution is
called saturated solution
Types of electrolyte solutions:
1. Non-saturated (or undersaturated)
2. Saturated
3. Supersaturated
1
2
3
17
Solubility of precipitates
Solubility (or equilibrium solubility,S): the concentration of
a saturated solution in molarity at a given temperature;
characteristic to the given salt (depends on solvent and
temperature)
MA
M+ + A-
MxAy
xMy+ + yAx-
S = [My+ ]/x = [Ax-]/y
Solubility
y product
p
(L, Ksp) : the equilibrium
q
constant for the
components of the precipitate in a saturated solution
(i.e., in a solution, which contains some precipitate)
Ksp = [M+ ][A-]
Ksp = [My+ ]x[yAx-]y
Solubility and solubility product
[
K sp = M y +
] [A ]
x
S = x+ y
x− y
= ( Sx) x ( Sy ) y
K sp
xx y y
E
Examples:
l
calculate
l l t solubility
l bilit f
for
AgCl in water, at 25 oC Ksp = 1.0.10-10
Ag2CrO4 in water, at 25 oC Ksp = 1.1.10-12
Bi2S3 in water, at 25 oC Ksp = 1.0.10-72
18
Factors influencing the solubility of a precipitate
1. Common ion effect – common ion will reduce the
concentration of the other ion (and therefore the
solubility) of ppt
(unless the common ion forms complex compound with
the ppt)
2. Effect of pH –
• if the anion gets protonated, decrease of pH
increases solubility
• if the cation hydrolyses, increase of pH increases
solubility
3. Effect of complexation – complexation always
increases solubility
4. Effect of foreign ions – foreign ions in small quantities
increase, while in large quantities decrease solubility
(latter is called salting-out)
…now back to gravimetry…
Steps of precipitation gravimetry:
1. Precipitation: an excess precipitating reagent
added to the sample,
sample thus the analyte converted
into precipitate
2. Filtration: precipitate is separated from the
solution via filtration
3. Washing: precipitate is washed from impurities
4. Drying: precipitate is dried or ignited (to
convert it to a product with known composition)
5. Weighing: precipitate is weighed
19
1. Precipitation
• Particle size and filterability/washability – the larger the
better
• The factor determining the particle size:
relative supersaturation = (Q-S)/S (where Q is the
concentration of the supersaturated solution)
• Nucleation • Particle growth • At large relative supersaturation the rate of nucleation is
large – large naumber of small particles are formed
• At small relative supersaturation the particle growth
dominates, large particles are formed ☺
• In practice: elevate temperature to increase solubility, use
dilute solution (to decrease Q) and add the precipitating
agent slowly and under vigorous stirring
Filtration and washing
• Filtration may happen on paper filter or on glass filter
• Mother liquor – is the liquid from which the precipitate is
formed
• Washing liquids – distilled water or water saturated with the
precipitate
• Peptization – is a process by which the precipitate returns to
it dispersed state (behaves as a solution again)
• Coprecipitation – soluble components other than the analyte
are removed from the solution together
g
with the precipitate
p
p
– surface adsorption
– mixed crystal formation
– occlusion and mechanical entrapment
• Precipitation from homogeneous solution
20
Drying and weighing
• Drying/ignition is necessary to obtain constant mass for the
precipitate
• Drying/ignition leads to the weighing form – the form of the
analyte with accurately known composition (or stoichiometry)
• Drying: t < 200 oC
• Ignition: t = 6-800 oC
• If the filtration is done with filter paper, ignition can be done,
if glass filter is used, only drying is allowed
• Weighing is always done by using an analytical balance
• The weighed mass must always be larger than 100 mg (to have
accuracy better than 1%)
• Examples: SO42- ions in the form of BaSO4
Ca2+ ions in the form of Ca(COO)2.H2O
Fundamentals of titrimetry
A chemical reaction between the titrant solution and
the analyte is suitable for titrimetry,
titrimetry if
1. it takes place according to one kind of known
stoichiometry
2. it is quantitative (conversion is > 99.9%, no
excess of reactant is needed)
3. it isreasonably fast
4. completion of the reaction can be indicated
21
Terms used in titrimetry
Standard solution – is a reagent of exactly known
concentration that is used in the titrimetric analysis
Titration – is a process in which the standard solution is added
to the analyte until the reaction between the analyte and
the reagent is complete
Equivalence point – the point in the titration, when the amount
of reagennt added to thge solution is exactly equivalent
to the amount of the analyte (theoretical value)
End point – the point in the titration, when a physical change
occurs that is associated with the chemical equivalence
q
(practical value, this is what we obsrerve)
Titration error – Et = Vep – Veq, where Vep is the actual volume
of reagent required to reach the end point and Veq is the
theoretical volume to reach the equivalence point
Perfect titration - Vep = Veq,
Terms used in titrimetry
Titration curves – plot the reagent volume on the horizontal axis
and some function of the analyte on the vertical axis; the
equivalence point can be read off the titration curve; it
can either be sigmoidal or linear segment curve.
Indicators – they are added to the analyte solution to produce a
visually observable physical change (usually colour change)
at or very near to the equivalence point
Primary standard – is an ultrapure compound that serves as a
reference material for titrimetric method of analysis
high purity
atmospheric stability
f hydrate
y
water
absence of
reasonable cost
reasonable solubility
large molar mass
Secondary standard – a compound, whose purity has been
established by chemical analysis and serves as a reference
material for titrimetric method of analysis
22
Terms used in titrimetry
Preparation of standard solutions:
1. Direct method – includes (i) accurate weighing of a
1
primary standard which is then (ii) dissolved in a suitable
solvent and (iii) diluted to exactly known volume in a
volumetric flask
2. Indirect method – includes preparation of the tatrant
solution by approximate weighing and dilution to an
approximately known volume, followed by standardization,
which means (i) titrating a weighed quantity of a primary
or a secondary standard or (ii) titrating a known volume of
a standard solution
23