This chemistry blog is aimed mainly at senior high school students or first year university students. It covers general chemistry topics required in Colleges and Universities. However, chemistry topics of general interest are going to be included.
Solutions are homogeneous mixtures. Many chemical reactions are carried out in solutions and solutions are closely related to our every day lives. The air we breathe, the liquids we drink and blood and the fluids in our body are all solutions. The components of a solution are:
The solvent - the major component of a solution
The solute - the minor component of a solution
Solutions especially liquid solutions, usually have different properties than either the pure solvent or the solute. For example, pure water freezes at 0 C, but aqueous solutions freeze at lower temperatures. Some of these properties depend only on the number of dissolved particles and not their identity. Such properties are called colligative properties.
The major colligative properties are the following:
freezing-point lowering
boiling point raising
vapor-pressure lowering
osmotic pressure
References
P. Atkins, J de Paula, “Physical Chemistry: Thermodynamics, Structure and Change”, 10th Edition, W. H. Freeman, 2014
D. A. McQuarrie, J. D. Simon,“Physical Chemistry: A Molecular Approach”, 1st Edition, University Science Books, 1997
K. J. Laidler, J.H. Meiser, B.C. Sanctuary, “Physical Chemistry”, 4th Edition, Brooks Cole, 2002
Boyle’s law: The volume of a fixed amount of gas maintained at a constant temperature is inversely proportional to the gas pressure:
P ∝ 1/V or P = k/V or P*V = k (moles n and temperature T constant) (1)
Equation (1) shows that the product of the pressure and volume of a fixed amount of gas at a constant temperature T is a constant k.
Charles’ law: The volume of a fixed amount of gas at constant pressure is directly proportional to the temperature T (Kelvin)
V ∝ T or V = c * T (where c, pressure P and moles n constant) (2)
Avogadro’s law: Equal volumes of different gases compared at the same temperature and pressure contain equal numbers of molecules.
V ∝ n (P and T constant) (3)
By combining (1), (2) and (3) above into one proportionality:
V ∝ n*T/P (4)
Proportionality (4) can be replaced by an equality if a proportionality constant R would be included:
V = R*n*T/P or P*V = n*R*T (5)
This proportionality constant is known as the gas constant R.
Any gas that obeys (1), (2) and (3) will also obey equation (5) which is called the ideal gas equation (Fig. I1) . All gases that obey this equation are called ideal gases.
Under suitable conditions some real gases do approach the behavior of ideal gases and make equation (5) very useful.
Fig. I1: Interrelationship of the gas laws. Any gas that obeys (1), (2) and (3) will also obey equation (5) which is called the ideal gas equation
The ideal gas equation can be used to establish molecular weights of gases. For this purpose it is helpful to alter the equation slightly by substituting where n (moles of gas) with its equivalent m/MW (where m is the mass of gas and MW its molecular weight) to get the following equation:
P*V = (m/MW)*R*T (5a)
A solved example regarding the determination of the molecular weight of an ideal gas is presented in the following video:
Other Gas Laws
Some other gas laws of note are Raoult’s law, the law of Gaseous Diffusion, Graham’s law and Gay Lussac’s law.
Raoult’s law states: i) the partial pressure of a solute is proportional to the mole fraction of the solute in the solution and ii) the vapor pressure of a solution is directly proportional to the mole fraction of solvent present.
Psoln = xsolvent * Posolvent(6)
Where Psoln the observed vapor pressure of the solution
Posolvent the vapor pressure of the pure solvent
xsolvent is the mole fraction of the solvent in the solution
From equation (6) can be derived that for a solution that contains half solute molecules and half solvent molecules – xsolvent is 0.5 – the vapor pressure of the solution would be half of the vapor pressure of the solvent.
The effect of the solute on the vapor pressure of a solution gives us a convenient way to “count” molecules and thus provides a means for experimentally determining molar masses. Suppose a certain mass of a compound is dissolved in a solvent, and the vapor pressure of the resulting solution is measured. Using Raoult’s law, we can determine the number of moles of solute present. Since the mass of this number of moles is known, we can calculate the molar mass.
Ralph H. Petrucci, “General Chemistry”, 3rd Edition, Macmillan Publishing Co., 1982
Key Terms
gas ideal gas law, P.V = nRT, the gas laws, Boyle's gas law,Raoult's law, Graham's law, ideal gas law constant, gas law practice problems, Avogadro's, Graham's, Gay Lussac's, gas constant R, chemistry net, ideal gas law equation
Solutions and Concentration - Solution Composition
Chemical reactions often take
place in aqueous solutions. To perform stoichiometric
calculations in such cases the amounts of chemicals present in solution
– the concentration of solution -
must be known.
Concentrationof a solution is a
measurement stating the amount of a solute present in a known amount of
solution:
Concentration = amount of solute / amount of
solution
The terms solute and solution are
usually used for liquid samples but they can be extended to gaseous and solid
samples.
The most common units of
concentration are given in Table I.1:
Common Units of
Concentration
Name
Symbol
Units
molarity
moles solute / liters of
solution
M
molality
moles solute / kg
solvent
m
normality
number of equivalent
weights of solute / liters of solution
N
formality
number of formal weights
of solute / liters of solution
F
weight %
g solute / 100 g of
solution
%
w/w
volume %
ml solute / 100 ml
solution
%
v/v
weight-to-volume %
g solute / 100 ml
solution
%
w/v
parts per million
g solute / 106
g solution
ppm
parts per billion
g solute / 109
g solution
ppb
Note: Another way of describing
solution concentration is the mole
fraction (xi)
Molarity (M)is defined as the number of moles of solute per liter of
solution.
i.e by dissolving 0.1 mol NaOH in
1 l of H2O gives a solution that contains 0.1 mol Na+ and
0.1 mol of OH- in 1 l. The concentration of the solution is
[Na+] = 0.1 M and [OH-] = 0.1 M.
Since molarity depends on the
volume of the solution it changes slightly with temperature.
Another way of describing solution
concentration is molality
(m) which is the number of moles of solute per kilogram of
solvent.
Molality is independent of temperature since it depends on
mass.
In very dilute aqueous solutions
the molarity (M) and molality (m) are nearly the same.
Example
#1
A solution of 1M
H2SO4 has density 1.04 g/cm3. Calculate the
(%w/w) concentration of the solution.
Given
[H2SO4] = 1M
d = 1.04
g/cm3
PH2O =
41 mmHg
MW
H2SO4 = 98 g/mole
Asked
for
(%w/w) = ?
From the definition of (%w/w) =
g solute / 100 g of solution
(1)
The mass of solute (g solute) is
unknown but it can be calculated.
Since [H2SO4] = 1M ⇒1000 cm3 of H2SO4
solution contain 1 mole “pure”
H2SO4 (2)
The mass of 1 mole
“pure” H2SO4 can be calculated as shown
below:
mass (g) = mole* MW = 1 mole * 98 g/mole = 98
g (3)
From (2) and d =
m/V = 1.04 g/cm3 the mass of 1000 cm3 of
H2SO4can be calculated:
m = d* V = 1.04g/cm3* 1000cm3 = 1040 g of H2SO4
solution (4)
From (2), (3) and (4):
Mass of 1040 g of
H2SO4 solution contain 98 g of“pure”H2SO4
Mass of 100 gof H2SO4 solution
contain x = ? g of“pure”H2SO4
x = 98 g“pure”H2SO4 * (100gof H2SO4 / 1040gof H2SO4) = 9.42 g of“pure”H2SO4
Therefore,(%w/w) =
9.42
Provided that the theory and the
definitions of solution concentration units is understood a % solution calculator
can be used.
Meant to be used in both the
teaching and research laboratory, a %
solution calculatorcan be utilized to perform
a number of different calculations for preparing percent (%) solutions when
starting with the solid or liquid material.
An additional solved example
regarding solutions and concentration is shown in the following video
