POST RN BScN SEMESTER II

1. Presented by
Niamat Ullah
Life Saving son

2.  Solutions are homogeneous mixtures of two
or more pure substances.
 In a solution, the solute is dispersed
uniformly throughout the solvent.

3.  The clinical lab almost always uses solutions. A solution
means that something has been dissolved in a liquid. In the
clinical laboratory the solvent we measure most of the time
is human plasma. The solute is whatever the substance is we
want to measure.
 Mixtures of substances – the substances in a
solution are not in chemical combination with
one another.
 Dispersed phase - the substance is dissolved (the
solute)
 The substance in which the solute is dissolved is
the solvent.
 Solute + Solvent = Solution
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4. The intermolecular
forces between solute
and solvent particles
must be strong enough
to compete with those
between solute particles
and those between
solvent particles.

5. As a solution forms, the solvent pulls solute
particles apart and surrounds, or solvates,
them.

6. If an ionic salt is
soluble in water, it is
because the ion-
dipole interactions
are strong enough to
overcome the lattice
energy of the salt
crystal.

7.  Unsaturated
Less than the
maximum amount of
solute for that
temperature is
dissolved in the
solvent.

8.  Saturated
Solvent holds as much
solute as is possible at
that temperature.
Dissolved solute is in
dynamic equilibrium
with solid solute
particles.

9. A supersaturated solution is one in which the solution
has had different forces applied to it, allowing for a
higher concentration of solute in the liquid than would
be possible if no forces were applied. The solubility of a
solution increases as the temperature is increased.
This means that as the temperature rises, more solute
can be dissolved. However, as the temperature of the
supersaturated solution decreases, maximum
saturation decreases again meaning sedimentation
and settling will occur. Increased pressure also
increases maximum possible saturation and allows for
a supersaturated solution.

10.  Chemists use the axiom
“like dissolves like”:
Polar substances tend to
dissolve in polar solvents.
Nonpolar substances tend
to dissolve in nonpolar
solvents.

11. The more similar the
intermolecular
attractions, the
more likely one
substance is to be
soluble in another.

12. Glucose (which has
hydrogen bonding) is
very soluble in
water, while
cyclohexane (which
only has dispersion
forces) is not.

13.  Vitamin A is soluble in nonpolar compounds
(like fats).
 Vitamin C is soluble in water.

14.  In general, the
solubility of gases in
water increases with
increasing mass.
 Larger molecules have
stronger dispersion
forces.

15.  The solubility of
liquids and solids
does not change
appreciably with
pressure.
 The solubility of a
gas in a liquid is
directly
proportional to its
pressure.

16. In chemistry, the equivalent concentration
or normality of a solution is defined as the
molar concentration divided by an equivalence
factor :
mEq

17.  n chemistry, the molar concentration, is defined as
the amount of a constituent (usually measured in
moles – hence the name) divided by the volume of
the mixture : It is also called molarity, amount-of-
substance concentration, amount concentration,
substance concentration, or simply concentration.
Osmotic concentration, formerly known
as osmolarity, is the measure of solute
concentration, defined as the number of
osmoles (Osm) of solute per litre (L)
of solution (osmol/L or Osm/L).

19.  Common prefixes and abbreviations that are added to units of
measure:
 deci (d) 10-1
 centi (c) 10-2
 milli (m) 10-3
 micro ( μ) 10-6
 nano (n) 10-9
 pico (p) 10-12
 femto (f) 10-15
 Example: A common unit of liquid measurement is a deciliter( dl ),
or one – tenth of a liter
 Combine a prefix with a basic unit results in a statement of a specific
length, weight or volume
 Reporting clinical chemistry results may be in units such as :
 mg / dL
 g / dL
 mEq / L
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20.  Atomic mass, Molecular weight, and Formula
Weights
 Moles: 1mole = 6.022 x 1023
(atoms, molecules
or formula units)

21.  Equivalents Weights / Liter
 Equivalent weight is equal to the gram molecular
weight of a substance divided by its valence
 Valence = the electrical charge of an ion, or the
number of moles that react with 1 Mole H+
 Example
 The MW of calcium = 40 grams
 Calcium ions carry a +2 electrical charge ( valence = 2 )
 Equivalent Weight of calcium = 40 / 2 = 20 gram
equivalent weight
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22. N= number of grams of solute
Gram equivalent weight of solute
1.00 L of solution
 Normality (N)
 N = Molarity (M) x valence
 Molarity = N / valence
 M is always < N
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23. % w/w – percentage weight per weight
 Most accurate method of expressing concentration
 % w/w = gram of solute OR gram of solute per 100.0 g of
solution
100.0 g of solution
 How many grams of NaOH are needed to make a 25.0%
w/w solution using deionized water as the solvent?
25.0% w/w = X g of solute in 100 g of solution
X= 25.0 g NaOH
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24. % w/v – percentage weight per volume
 Easiest & most commonly used, very accurate if temperature controlled.
 %w/v= g of solute OR g of solute per 100.0 mL of solution
100 mL of solution
What is the %w/v of a solution that has 15.0 g of NaCl dissolved into a
total volume of 100 mL deionized water?
X% w/v = 15.0 g NaCl
100 mL of solution
X= 15.0 %
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25.  % v/v –percentage volume per volume
 Least accurate, but used when both substances are
liquids
 Note: volumes of liquids are not necessarily
additive
 %v/v= mL of solute OR milliliter of solute per 100
mL of solution
100 mL of solution
 How many milliliters of ethanol are needed to make a
75.0% v/v solution using deionized water as the
solvent?
75.0% v/v EtOH = X mL EtOH in 100 mL of solution
= 75.0 mL EtOH
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26.  Molarity (M)= moles/liter or mmoles/mL
 Normality(N) = equivalence/liter or meq/mL
Molality(m) = moles/1000g solvent

27.  Density = mass solute /unit volume
 Specific Gravity = Dsolute/DH20
 DH2O = 1.00000 g/mL @ 4o
C
 DH2O = 0.99821 g/mL @ 20o
C

28. Mass % of A =
mass of A in solution
total mass of solution
× 100

29. moles of A
total moles in solution
XA =
 In some applications, one needs the mole
fraction of solvent, not solute—make sure you
find the quantity you need!

30. mol of solute
L of solution
M =
 volume is temperature dependent, molarity can
change with temperature.

31. mol of solute
kg of solvent
m =
Because both moles and mass do not change
with temperature, molality (unlike molarity) is
not temperature dependent.

32.  Some substances form semi permeable
membranes, allowing some smaller particles to
pass through, but blocking other larger particles.
 In biological systems, most semi permeable
membranes allow water to pass through, but
solutes are not free to do so.

33. In osmosis, there is net movement of solvent from
the area of higher solvent concentration (lower
solute concentration) to the are of lower solvent
concentration (higher solute concentration).

34.  If the solute
concentration outside
the cell is greater than
that inside the cell,
the solution is
hypertonic.
 Water will flow out of
the cell.

35.  If the solute
concentration outside
the cell is less than
that inside the cell,
the solution is
hypotonic.
 Water will flow into
the cell, and hemolysis
results.

36. Suspensions of particles larger than individual
ions or molecules, but too small to be settled
out by gravity.

37.  Colloidal suspensions
can scatter rays of light.
 This phenomenon is
known as the Tyndall
effect.

38. Some molecules have a
polar, hydrophilic
(water-loving) end and
a nonpolar,
hydrophobic (water-
hating) end.

39.  A suspension of liquid droplets or fine solid particles in a
gas is called an aerosol. In the atmosphere these consist of
fine dust and soot particles, and cloud droplets.
 suspension: system does not stays stable and settle
 Examples of Suspensions
 Mud or muddy water, is where soil, clay, or silt particles are
suspended in water.
 Paint
 Chalk powder suspended in water.
 Dust particles suspended in air.
 Algae in water
 Milk of Magnesia
 Flour suspended in water, as pictured to the right.

40. MATTER
Heterogeneous
mixture
Is it uniform
throughout?
No
Homogeneous
Yes
Can it be separated
by physical means?
Pure Substance Homogeneous
Mixture (solution)
Can it be decomposed
into other substance by
a chemical process?
Element Compound
No yes
No yes