Entry Test Preparation 2015
Chemistry
Chapter # 1
Basic
CONCEPTS
Chemistry:
Chemistry is the study of the composition, structure, properties,
changes in matter and energy, and the laws and principles, which govern these
changes.
Substance:
A sample of pure matter is called substance.
Element:
Element is a substance having single type of identical atoms.
Compound:
Compound is substance having two or more type of atoms. These
constituent atoms have definite ratio by weight through out the substance.
Compounds can be decomposed chemically into simpler compounds or
elements.
Physical Properties:
The properties that distinguish one element or compound from other
element or compound on physical basis like colour, boiling point, and odour. A
physical change does not alter the chemical composition of the substance.
Chemical Properties:
Chemical properties are related to the ability of a substance to
undergo chemical changes.
Atom:
Atom is the smallest particle of an element, which may or may not exist
free in nature except those of inert gases. For example, Helium and Neom exist
in atomic state.
Evidence of Atoms:
It is not possible to see the atoms but the nearest possibility to its
direct evidence is by using and electron microscope.
Ordinary optical microscope can measure the size of an object up to or
above
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500 nm
(1nm = 10 -9 m).
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Objects of the size of an atom can be observed in an electron
microscope.For smaller objects wavelength should be smaller from the size of an
object.
- The diameter
of atoms are in the order of
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2 x 10-10 m which is 0.2 nm.
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- Masses of
atoms are found in the order of 10-27 to 10-25 Kg.
- They are
often expressed in atomic mass units where.
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1
a.m.u. = 1.661 x 10-27
kg.
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Molecule:
Molecule is the smallest uncharged individual unit of compound formed
by the combination of two or more atoms.
The number of atoms present in a molecule determines its atomicity.
Important Properties of Molecules:
Some of the important properties of molecules are as follows;
- Molecules of
the same substance are similar in all respects.
- Molecules
have empty spaces (intermolecular spaces) in between them.
- They possess
random movements, which are maximum in gases, less in liquids, and the
least in solids.
- Polar
molecules have attractive forces for each other.
- Molecules
possess definite kinetic (due to random motion) and potential (due to
attraction) energy.
- Their
collisions due to random movements cause chemical reaction.
Ion:
A charge carrying atom is called an ion.
Cation:
A positive ion is called cation.
Anion:
A negative ion is called anion.
Molecular ion:
When a molecule loses or gains electrons, it forms a molecular ion,
e.g., CH4+,CO+, N2+ .
Cationic molecular ions are more abundant than anionic ones. These ions
can be generated by passing high energy electron beam or α-particles through a
gas.
Relative Atomic Mass:
It is the mass of one atom of an element compared to that of the mass
of one atom of C12 taken as 12 amu.
Atomic Number:
It is the number of protons in an atom. It is the identity of atom.
Molecular Mass:
It is the sum of atomic masses of all the atoms present in the simple
formula of the substance.
Example:
Formula weight of Sodium Chloride (NaCl) is 58.5 (Na =23 + Cl =58.5).
Isotopes:
Atoms of an element with different masses are called isotopes. They posses
same chemical properties and same position in the periodic table. Isotopes of
and element have same number of protons and electrons but they differ in number
of neutrons.
Examples:
Four stable isotopes of Carbon are 116C, 126C,
136C, and 146C expressed as C-11, C-12, C-13 and C-14 respectively.
Hydrogen has three stable isotopes 11H, 21H, 31H called protium, Deuterium and Tritium
respectively.
Oxygen has three, nickel has five, calcium has six, palladium has six,
cadmium has nine and tine has eleven isotopes.
At present more than 280 different isotopes occur in nature. The
include 40 radioactive isotopes as well. Besides these, about 300 unstable
radioactive isotopes have been produced through artificial disintegration.
Arsenic, fluorine, iodine and gold, etc have only a single isotope. They are
called mono-isotopic elements.
In general, the elements of odd atomic number almost never possess more
than two stable isotopes.
·
The elements of even atomic number usually have larger number of
isotopes. Out of 280 isotopes that occur in nature, 154 have even mass number
and even atomic number.
·
Isotopes whose mass number are multiples of four are particularly
abundant. For example, 16O, 24Mg, 28Si, 40Ca
and 56Fe form nearly
50% of the earth’s crust.
Mass Spectrometer:
Mass spectrometer is an instrument which used to measure the exact
masses of different isotopes of an element.
Dempster’s Mass Spectrometer:
Dempster’s mass spectrometer is used for the identification of elements
in solid state.
Aston’s Mass Spectrograph:
Aston’s mass spectrograph is used to identify the isotopes of an
element on the basis of their atomic masses.
Objectives:
1-no. of isotopes of an element
2-reletive abundance of isotopes
3-relative atomic masses of isotopes
working:
Evaporation and Ionization:
Vapour state with 10-6 to 10-7 torr→ ionization
using fast moving electrons→ atoms with positive charge
Acceleration and Detection:
5oo to 2000 volts→ ions are accelerated
m/e
= H2r2/2E
If E is increased by keeping H constant then radius will increase and
positive ions of particular e/m will fall at different places. This can also be
done by changing the magnetic field.
Electrometer: develops current→ comparison with C-12→
comparison of current strength of both give exact mass no. of each isotope.
Relative Abundance of Isotopes:
The isotopes of all the elements have their own natural abundance. The
properties of a particular element, which are mentioned in the books generally,
correspond to the most abundant isotope of that element.
Some of the most Abundant
Element:
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Element
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Abundance %
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Hydrogen
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99.985 1H
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0.0152H
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Carbon
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98.893 12C
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1.10713C
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Nitrogen
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99.634 14N
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0.36615N
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|
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Chlorine
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75.53 16Cl
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24.4737Cl
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|
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Oxygen
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99.759 16O
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0.03717O,
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0.20418O
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95.032S
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0.7533S,
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4.2234S,
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0.01736S
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Uranium
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99.27%.U238
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0.72%,U235
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Empirical Formula:
The simplest whole number ratio of atoms in the compound represented by
the symbols of the elements is called empirical formula. It represents the
chemical composition of the compound in terms of smallest possible number of
atoms of each element present in the formula. Used for both molecular as well
as ionic compounds.
Determination of empirical and
molecular formula:
Chemical analysis:
1- Qualitative Analysis, type of elements.
2- Quantitative Analysis, mass of an element, % of an element.
3- no.of moles (divide the % by atomic mass)
4- mole ratio of elements
5- convert into whole no.
Combustion analysis:
Compounds containing C, H and O can be analyzed very easily by
combustion analysis.
C → CO2 absorbed
in 50%KOH
H → H2O absorbed
in Mg(ClO4)2
% H = 2.016/18.0 x mass of H2O/mass
of org.c x100
%
C = 12.0/44.0 x mass of CO2/mass of orgC x 100
%
O = 100 - ( % C + % H )
Molecular formula:
A chemical formula based on an actual molecule of a substance is called
the molecular formula of the substance. It represents the actual number of
atoms of each element that are combined in each molecule of the substance.
Molecular
formula = n x empirical formula
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Avogadro’s Hypothesis
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Under the same conditions of pressure and temperature, equal volumes of
gases contain equal number of molecules.
Avogadro’s
Number:
Avogadro’s number is the number of atoms,
molecules or ions in one gram atoms one gram molecule or one gram ion of a
substance, respectively.
Its value is 6.022 x 1023, and is
represented by NA
Applications
of Avogadro’s Law:
- In
calculating the atomicity of elementary gases, For example, the atomicity
of oxygen is 2.
- To
find the relationship between molecular weight and vapour density of a gas
which is given as:
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Molecular weigh = 2 x Vapour density
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·
To find the relationship between weight and volume of a gas.
22.4 litre of any gas at STP is equal to the
molecular weight of the gas expressed in gram.
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It is known as “Gram Molecular Volume (GMV)
law”
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Mole:
Atomic mass, molecular mass, or ionic mass of a substance expressed in
grams in one gram mole of that substance.
Example:
Molecular mass of water is 18.
18 gm of water = 1 gram mole of water
Gram atom of an element is equal to its atomic weight in grams.
23 gram of Na is 1 gram atom of Na
Gram molecule of a compound is equal to its molecular weight in grams.
18 gm of water = 1 gram molecule of water
Mole and Number of Particles:
1 mole of a substance is the quantity of a substance which contains
6.023 x 1023 (Avogardo number, NA) particles (atoms, molecules, or ions). Mole
is a latin word means huge mass.
Example:
18 gms of water is one mole (gram) of water.
18 gm of water contain 6.023 x 1023 molecules of water.
Mole and Volume of a gas at S.T.P.
1 mole of a gas always occupies 22.4 liters at S.T.P. stands for
standard temperature (273 K or 0o C. and pressure (1 atm) and RTP
stands for room temperature (25oC) and pressure (1 atm).
2 gm (1 mole) of hydrogen gas at S.T.P. = 22.4 liters of Hydrogen gas.
Gram formula:
The formula mass of an ionic compound expressed in grams is called gram
formula of the compound.
71 grmas of Chlorine gas is one gram formula of Chlorine gas.
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Stoichiometry
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Stoichiometry is a branch of chemistry which deals with the quantative
relationship between reactants and products in a balanced chemical equation.
For Stoichiometric calculations we have to assume the following
conditions:
All reactants converted into products
No side rections occur
In calculations law of conservation of mass and law of definite
proportions are obeyed
Limiting Reactant:
The reactant which gives minimum amount of products and is completely
consumed in the reaction is called limiting reactant.
Example:
8 grams of O2 react with 2 grams of H2 to form
water.
The balance equation for this reaction is
2H2 + O2 → 2H2O
Apply CAT’s method to find the limiting reactant
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Quantity of Reactant 1
Quantity of Reactant 2
──────────── : ─────────────
mole in same units mole
in same units
8 2 1
→ ─ : ─ : ─
: 1
32 2 4
Since 1 is
lesser, Oxygen is the limiting reactant
4
In the reaction with the given quantities.
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Theoretical yield:
The theoretical yield of a reaction is the quantity of the product
calculated with the help of a balanced chemical equation.
Actual Yield:
The actual amount of the products(s) obtained in a chemical reaction is
(are) called the actual yield of that reaction.
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Practically, in most of the chemical
reactions the actual quantity of the product obtained is less than the
theoretical yield.
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Efficiency of a reaction or
Percentage yield:
The efficiency of a reaction is expressed by comparing the actual and
theoretical yields in the form of percentage (%) yield.
Actual Yield
Percentage yield = ──────────
× 100
Theoretical Yield
………………………….
Macromolecules:
Macromolecules are a large
molecule with a large molecular mass bonded covalently, but generally the use
of the term is restricted to polymers and molecules which structurally include
polymers. Some of these may include lipids, proteins, mono-, di- and
polysaccharides. DNA is a type of a macromolecule.
Significant Figures:
These are the reliable digits in
a number that are known wit certainty.
Scientific Notation:
In this system, numbers are
written as the product of two factors. The first is a decimal number that
usually ranges between 1 and 0 and the second is 10 raised to an appropriate
power.
Atomicity:
It is the number of atoms present
in a molecule of a substance.
Chemical Formula:
The formula of a compound tells
us which elements it is composed of and how many atoms of each element are
present in formula unit.
Chemical Equation:
It is a shorthand expression for
chemical reaction.
Molar Volume:
One mole of any gas at standard
temperature and pressure occupies a volume of 22.4 dm3 is known as
molar volume.
STP:
Standard temperature is 0oC
or 273 Ao or 273 K. standard pressure is 1 atmosphere or 76cm or 760
mm.
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