Preparation of solutions.pptx

1. Weighing & Preparation of
Solutions of Different
strengths & their Dilution
AND
Handling Techniques of
Solutions
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2. 2
1. Solutions
2. Measuring Chemicals
3. Different chemical concentrations
4. Dilution of stock solutions
5. Labelling
6. Lab safety
7. Conclusions

3. • Solution uniform homogenous mixture of two
or more substances i.e, solute and solvent.
Solution= solute + solvent
• Standard solution: very precise solution, usually to 3-4 significant
figures, used in quantitative analysis or an analytical procedure.
• Saturated solution: a solution that contains the maximum
amount of a particular a solute that will dissolve at that
temperature
• Supersaturated solution: a solution that contains more solute
than equilibrium condition allow; it is unstable & the solute may
precipitate upon slight agitation or addition of single crystal
3

4. Preparing Solutions
Solutions of known concentration can be prepared in a number of
different ways depending on the nature of the analyte and/or the
concentration required:
• Weighing out a solid material of known purity, dissolving it in a
suitable solvent and diluting to the required volume
• Weighing out a liquid of known purity, dissolving it in a suitable
solvent and diluting to the required volume
• Diluting a solution previously prepared in the laboratory
• Diluting a solution from a chemical supplier.
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5. Weight Measurements
Basic protocol 1: measuring mass using a top-loading Balance
1. Turn on balance and wait for display to read 0.0 g.
2. Place weighing vessel on the balance pan (e.g., creased weighing
paper, weigh boat)
3. Press tare button so that display reads 0.0g.
4. Gently add the substance being weighed to the weighing sample.
5. Record mass.
6. Remove weighed sample.
7. Clean spills off balance with brush or absorbent laboratory tissue.
Discard any disposable weighing vessel. 5

6. Contd...
• Basic protocol 2: measuring Mass using an analytical
Balance
1. Turn on balance and wait for display to read 0.0000 g.
2. Check the level indicator & do not lean on table while
weighing.
3. Place weighing vessel on the balance pan (e.g., creased
weighing paper, weigh boat)
4. Close the sliding doors & wait for stability light
indicator, indicating that the weight is stable.
5. Press tare button so that display reads 0.0g.
6. Gently add the substance being weighed to the
weighing sample.
7. Record mass.
8. Remove weighed sample.
9. Clean spills off balance with brush or absorbent
laboratory tissue. Discard any disposable weighing
vessel
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7. Volume Measurements
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Micropipettes
Volumetric or transfer
pipettes

8. Volumetric Containers
1.Beakers & Erlenmeyer flasks
2. Volumetric flasks
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3. Graduatedcylinders

9. Procedure for preparing a solution of known concentration from a
known amount of a solid material
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10. Procedure for preparing a solution of known
concentration by dilution
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11. Common P r a c t i c a l Units f o r
Reporting Concentration
Name Units Symbol
Molarity Moles of solute / litres of solution M
Normality Number of EWs solute / Litre of solution N
molality Moles of solute / Kg of solvent m
Weight % g of solute / 100 g of solution % w/w
Volume % mL of solute / 100 mL of solution % v/v
Weight-to-Volume % g of solute / 100 mL of solution % w/v
•Weight per unit volume e.g., g/L, mg/ml
•Parts per million(ppm) or ppb
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12. 12
1. Molar solutions
• Molarity is number of moles of a solute that are dissolved per liter of
total solution.
• A 1 M solution contains 1 mole of solute per liter total volume.
Example:
A 1M solution of H2SO4 contains 98.06 g of sulfuric acid in 1 liter of
total solution.
"mole" is an expression of amount
"molarity" is an expression of concentration.

13. Contd..
13
• "Millimolar", mM, millimole/L.
– A millimole is 1/1000 of a mole.
• "Micromolar", µM, µmole/L.
– A µmole is 1/1,000,000 of amole.
HOW MUCH SOLUTE IS NEEDED FOR A SOLUTION OF APARTICULAR MOLARITY
AND VOLUME?
(g solute ) X (mole) X (L) = g solute needed
1 mole L
or
FW X molarity x volume = g solute needed

14. TO MAKE SOLUTION OF GIVEN
MOLARITY AND VOLUME
1. Find the FW of the solute, usually from label.
2. Determine the molarity desired.
3. Determine the volume desired.
4. Determine how much solute is necessary by
using the formula.
5. Weigh out the amount of solute.
6. Dissolve the solute in less than the desired
final volume of solvent.
7. Place the solution in a volumetric flask or
graduated cylinder. Add solvent until exactly
the required volume is reached, Bring To
Volume, BTV. 14

15. 15
2. Normal Solutions
• Normality is defined as the gram Eq.Wt. of the solute
per L of the solvent.
1N sol. = 1 EW solute / 1L of sol.
• Conc. Of acids and alkalis are usually expressed in
this unit.
• gram Eq.Wt. is the M.W divided by the no. of H+ or
OH- ions released from 1 molecule of the acid or
base, respectively in solutions.
Eq. Wt. = MW of the substance / replaceable no. of H+ or OH-

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Example:
1N Sulphuric Acid
M.W of H2SO4 = 98 g
Each molecule of acid releases 2 H+ ions in solutions.
Eq. Wt. = 98/2
= 49
So, 1L of 1N H2SO4 solution contains 49 g ofH2SO4
Chemical M.W Eq. Wt. 1N of solution contains
NaOH 40 1 40 g
KOH 56 1 56g
Na2CO3 106 2 53g
HCl 36.45 1 45g

17. 3. Molal solutions
• Molality expresses the no. of moles per 1000 g or
1 Kg of solvent.
• It is dependent on the density of solvent.
• It is different from Molarity as the later refers to
volume of the solution, which is temperature
dependent.
• Molal solutions are not usually used in
biochemical exp.
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18. 19
4. Percent solution
• Mass percent solutions are defined based on the grams of solute
per 100 grams of solution.
Example: 20 g of sodium chloride in 100 g of solution is a 20% by
mass solution.
• Volume percent solutions are defined as ml of solute per 100 mL
of solution.
Example: 10 mL of ethyl alcohol + 90 ml of H2O (making approx.
100 mL of solution) is a 10% by volume solution.
• Mass-volume percent solutions are also very common. These
solutions are indicated by w/v % & are defined as the grams of
solute per 100 mL of solution.
Example: 1 g of phenolphthalein in 100 mL of 95% ethyl alcohol
is a 1 w/v % solution.

19. 20

20. 5. PPM and PPB
ppm: The number
of parts of solute
per 1 million parts
of total solution.
ppb: The number
of parts of solute
per billion parts of
solution.
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Example
5 ppm chlorine = 5 g of chlorine
in 1 million g of
solution,
Or
5 mg chlorine in 1 million mg of
solution,
Or
5 pounds of chlorine in 1 million
pounds of solution

21. CONVERSIONS
To convert ppm or ppb to simple weight per
volume expressions:
5 ppm chlorine = 5 g chlorine =
106 g water
= 5 mg/1 L water
5 g chlorine
106 mLwater
= 5 X 10-6 g chlorine/ 1 mLwater
= 5 micrograms/mL

22. A COMPARISON OF METHODS OF EXPRESSING THE CONCENTRATION OF A
SOLUTE
CONCENTRATION OF SOLUTE AMOUNT OF SOLUTE AMOUNT OF WATER
(Na2SO4)
1 M 142.04 g Na2SO4 BTV 1 L with water
1 m 142.04 g Na2SO4 Add 1.00 kg of water
1 N 71.02 g Na2SO4 BTV 1 L with water
1 % 10 gNa2SO4 BTV 1 L with water
1 ppm 1 mg BTV 1L
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23. 24
PREPARING DILUTE SOLUTIONS
FROM CONCENTRATED ONES
• Concentrated solution = stock solution
• Use this equation to decide how much stock
solution you will need:
C1V1=C2V2
Where, C1 = concentration of stocksolution
C2 = concentration you want your dilute solution to be
V1 = how much stock solution you willneed
V2 = how much of the dilute solution you want tomake

24. EXAMPLE
• How would you prepare 1000
mL of a 1 M solution of Tris
buffer from a 3 M stock of
Tris buffer?
– The concentrated solution is 3
M, and is C1.
– The volume of stock needed is
unknown, ?, and is V1.
– The final concentration required
is
1 M, and is C2.
– The final volume required is
1000 mL and is V2.
SUBSTITUTING INTO THE
EQUATION:
C1 V1 = C2 V2
3 M (?) 1 M (1000 mL)
? = 333.33 mL
So, take 333.33 mL of the
concentrated stock solution
and BTV 1 L.
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25. 26
Preparation of exact 1N HCL
• Dilute 100 ml of HCl with water to 1 L. Mix well.
• Prepare exact 1N sol. Of Na2CO3 by dissolving 5.30g
anhydrous Na2CO3 in 100 ml H2O.
• Phenolphthalein indicator
– Dissolve 250 mg indicator in 50 ml of 50% alcohol.
• Titration
– Take 10ml of acid sol & 10ml H2O in a small beaker
– Add 2-3 drops of indicator
– Titrate with Na2CO3 sol from a 25ml burette till a faint red
colour is obtained
– Note the vol.(x ml) of base consumed at the end point

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Contd..
• Calculate the exact normality of the acid by
formula
– Normality of base X vol. of base = normality of
acid X vol. of acid
– So, the normality of HCl = 1x X/10
– Normality of base is 1 vol of base is x ml.
• After calculating the exact normality of the
acid, it is proportionately diluted with water to
obtain to exact 1 normal sol.

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Exact 1N NaOH solution
• Eq. Wt. Of NaOH is 40g
So 40g dissolved in 1L of H2O fo approx. 1N sol. & used.
• But for exact normality it is titrated against Oxalic acid
sol.(6.3g in 100ml water)
• Take 10ml oxalic acid + 10ml water in a beaker, add 2-3
drops of phenolpthalein indicator. Titrate against the NaOH
from a burette till a faint red colour is obtained.
• Calculate exact normality of sodium hydroxide sol as in
case of acid & dilute proportionately with water to obtain
exact 1N sol.

28. • Do not use chemicals
from unlabeled
containers
• Do not place labels on
top of one another.
• Label chemicals clearly
and permanently.
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29. You make it- you label it
1. identity of contents
2. concentration
3. your name
4.date of preparation
5.Hazard alert (if applicable)
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An unla bele d container w i l l become
t o m o r r o w ’s

30. Do NOT
× eat, drink or smoke in the
laboratory .
× pipette by mouth
× leave equipment using
water, gas or electricity on
overnight
× Never add water to conc.
Sulphuric acid
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31. ALWAYS
 Keep your working area clean
and tidy.
 Open bottles near window
where ventilation is available.
 Handle conc. Acids & liquor
Ammonia with care.
 label containers & solutions.
 secure the tops of reagent
bottles immediately after use
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 wear a lab coat & appropriate
eye protection
 wash hands after using any
substances hazardous to health,
on leaving the laboratory.
 keep broken glassware & sharps
separate from other waste &
dispose of in the appropriate
containers

32. Know the solutions & different conc. to
represent them.
Documentation, labeling & recording what
was done
Traceability
SOPs & SPs
Maintenance and calibration of instruments
Stability and expiration date recorded
Proper storage
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References
Gallagher Sean R.; A.Wiley Emily; Current
protocols, essential laboratory techniques; 2nd
Ed. ; Wiley-Blackwell a John Wiley & sons, Inc.