Basic concepts
Boyle's law
Boyle's law (sometimes referred to as the Boyle–Mariotte law) is an experimental gas
law that describes how the pressure of a gas tends to increase as the volume of a gas
decreases. The absolute pressure exerted by a given mass of an ideal gas is inversely
proportional to the volume it occupies if the temperature and amount of gas remain
unchanged within a closed system.
Mathematically, Boyle's law can be stated as:
PV  k
(1)
where P is the pressure of the gas, V is the volume of the gas, and k is a constant.
Ideal gas law
The “ideal-gas” is the equation of state of a theoretical ideal gas. It is a good
approximation of the behavior of many gases under many conditions, although it has
several limitations. It was first stated by Émile Clapeyron in 1834 as a combination of
the empirical Boyle's law, Charles' law and Avogadro's Law. The ideal gas law is often
written as:
PV  nRT
(2)
where, P is pressure, V is volume, T is absolute temperature, R is universal gas
constant and n is number of gas moles.
The mole volume
Equation 2 shows that a mole of ideal gas under definite conditions of temperature and
pressure always occupies a definite volume in spite of the nature of the gas. This
volume is called the mole volume. A gram mole of ideal gas at 0 ºC and pressure of 760
mmHg occupies 22.4 liters. The mole volume under any conditions of temperature and
pressure is easily calculated by means of equation 2.
V 2 V 1  P1 / P2 T 2 / T1  (3)
Dalton’s law
In Chemical engineering, Dalton's law (also called Dalton's law of partial pressures)
states that in a mixture of non-reacting gases, the total pressure exerted is equal to the
sum of the partial pressures of the individual gases. This empirical law was observed by
John Dalton in 1801 and is related to the ideal gas laws.
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Mathematically, the pressure of a mixture of non-reactive gases can be defined as the
summation:
n
Ptotal   Pi
i 1
(4)
where Pi represents the partial pressure of each components.
Pi  Ptotal y i
(5)
where yi is the mole fraction of the ith component in the total mixture of n components .
Energy and mass conservation law
Energy conservation law is a fundamental concept of physics, states that the total
amount of energy remains constant in an isolated system. It implies that energy can
neither be created nor destroyed, but can be change from one state to another. The
mass conservation law demonstrates that the total amount of mass remains constant in
and isolated system in spite of physical or chemical changes that may take place.
The conservation law can be formulated practically as below:
Q (input) – Q (output) + Q (generation) – Q (consumption) = Q (accumulation)
where Q can be expressed as energy of mass flow.
Chemical reaction
A chemical reaction is a process that leads to the transformation of one set of chemical
substances to another. Classically, chemical reactions include changes that only involve
the positions of electrons in the forming and breaking of chemical bonds between
atoms, with no change to the nuclei (no change to the elements present), and can often
be described by a chemical equation. Nuclear chemistry is a sub-discipline of chemistry
that involves the chemical reactions of unstable and radioactive elements where both
electronic and nuclear changes may occur.
Chemical equations are used to graphically illustrate chemical reactions. They consist of
chemical or structural formulas of the reactants on the left and those of the products on
the right. They are separated by an arrow (→) which indicates the direction and type of
the reaction; the arrow is read as the word "yields". The tip of the arrow points in the
direction in which the reaction proceeds. A double arrow (⇌) pointing in opposite
directions is used for equilibrium reactions. Equations should be balanced according to
the stoichiometry, the number of atoms of each species should be the same on both
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sides of the equation. This is achieved by scaling the number of involved molecules (A,
B, C and D in a schematic example below) by the appropriate integers a, b, c and d.
aA  bB  cC  dD
Four basic types of reaction
1) Synthesis
In a synthesis reaction, two or more simple substances combine to form a more
complex substance. These reactions are in the general form:
A  B  AB
Two or more reactants yielding one product is another way to identify a synthesis
reaction. One example of a synthesis reaction is the combination of iron and sulfur to
form iron (II) sulfide:
8Fe  S 8  8FeS
Another example is simple hydrogen gas combined with simple oxygen gas to produce
a more complex substance, such as water.[18]
2) Decomposition
A decomposition reaction is when a more complex substance breaks down into its more
simple parts. It is thus the opposite of a synthesis reaction, and can be written as:
AB  A  B
One example of a decomposition reaction is the electrolysis of water to make oxygen
and hydrogen gas:
2H 2O  2H 2  O 2
3) Single replacement
In a single replacement reaction, a single uncombined element replaces another in a
compound; in other words, one element trades places with another element in a
compound. These reactions come in the general form of:
A  BC  AC  B
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One example of a single displacement reaction is when magnesium replaces hydrogen
in water to make magnesium hydroxide and hydrogen gas:
Mg  2H 2O  Mg (OH ) 2  H 2 
4) Double replacement
In a double replacement reaction, the anions and cations of two compounds switch
places and form two entirely different compounds. These reactions are in the general
form:
AB  CD  AD  CB
For example, when barium chloride (BaCl2) and magnesium sulfate (MgSO4) react, the
SO42− anion switches places with the 2Cl− anion, giving the compounds BaSO4 and
MgCl2.
Another example of a double displacement reaction is the reaction of lead (II) nitrate
with potassium iodide to form lead (II) iodide and potassium nitrate:
Pb (NO 3 ) 2  2KI  PbI 2  2KNO 3
Representation of four basic chemical reactions types: synthesis, decomposition, single
replacement and double replacement.
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