John Dalton proposed the first scientific atomic theory, which stated that each chemical element is composed of atoms of a single, unique type that can combine to form chemical compounds. An atom is the fundamental unit of matter and consists of a nucleus containing protons and neutrons surrounded by electrons. Atoms of the same element contain the same number of protons but can vary in the number of neutrons, resulting in different isotopes of that element. Molecules are the smallest fundamental units of compounds made of two or more atoms bonded together.

1. Basic Concepts in
Chemistry
BASIC DEFINITIONS AND CONCEPTUAL QUESTIONS
BY KHUSHBAKHT RIAZ

2. Atom
• A fundamental piece of matter ,can take part in any chemical reaction ,having as same properties as its
respective element has.
• Made up of three tiny kinds of particles called subatomic particles: protons, neutrons, and electrons.
• The protons and the neutrons make up the center of the atom called the nucleus and the electrons fly around
above the nucleus in a small cloud, their pathways are known as shells.
History
• The word atom coined by the ancient Greek philosophers Leucippus and his pupil Democritus, from 'atomos'
meaning "indivisible".
Pupil
means
Student

3. Atomic Theory
John Dalton proposed that each ‘’Chemical element is composed of atoms of a
single, unique type, and though they cannot be altered or destroyed by chemical
means, they can combine to form more complex structures (chemical compounds)’’.
This marked the first truly scientific theory of the atom, known as Atomic theory.
Background
• Near the end of the 18th century, two laws about chemical reactions emerged
without referring to the notion of an atomic theory.
• The first was the law of conservation of mass, formulated by Antoine Lavoisier in
1789, which states that the total mass in a chemical reaction remains constant.
• The second was the law of definite proportions. First proven by the French
chemist Joseph Louis Proust in 1799, this law states that if a compound is broken
down into its constituent elements, then the masses of the constituents will
always have the same proportions, regardless of the quantity or source of the
original substance.

4. Molecule
A group of atoms bonded together, representing the
smallest fundamental unit of a chemical compound that can
take part in a chemical reaction.
Ion
An atom or molecule with a net electric charge due to the
loss or gain of one or more electrons.Cation is an ion with +ive
charge due to loss of e- and anion with -ive charge by gaining
e- .
Compound
• A compound is a substance formed when two or more
chemical elements are chemically bonded together.
• E.g., Water: two hydrogen atoms bonded to an oxygen
atom. It lies in the category of compound because of
chemical reaction in between its components.
&
-
&
+
&
+
Water-A Compound
Molecule-A part of compound
Ions

5. Anions are larger than the corresponding
neutral atom, since adding electrons increases
the number of electron-electron repulsion
interactions that take place. Cations are
smaller than the corresponding neutral atoms,
since the valence electrons, which are furthest
away from the nucleus, are lost.
I got a larger
size than you
I am
small
Interesting
to know
Feeling
sad for
cation.

6. Interesting to know that-
• A zwitterion is a neutral molecule with positive and
negative charges at different locations within that
molecule.
• An amino acid contains both acidic (carboxylic acid fragment)
and basic (amine fragment) centres. The isomer on the right is a
zwitterion.

7. Mixture & Types
• A mixture is made from different substances that are not chemically joined.
• For example, powdered iron and powdered sulphur mixed together makes a mixture of iron and sulphur.
• A heterogeneous mixture is simply any mixture that is not uniform in composition. Using various means,
the parts in the mixture can be separated from one another.
• Pizza is a heterogeneous mixture of dough, sauce, cheese, and other toppings.
• Homogeneous mixture is a solid, liquid or gaseous mixture that has the same proportions of its
components throughout a given sample.Water is an example. It often contains dissolved minerals and
gases, but these are dissolved throughout the water. Tap water and rain water are both homogeneous.

8. Difference in between Mixture & Compound
Mixture Compound
1. Impure matter 1. Pure matter
2. Constituents combine in any ratio to
form mixture.
2. Constituents combine in fixed ratio to
form compound.
3. Constituents retain their properties. 3. Constituents don’t retain their
properties because a new substance has
formed.
4. Constituents can be separated by
physical processes.
4. Constituents can’t be separated by
physical processes.
5. Iron and sulfur powder mixed with
another.
5. Water is an example.

9. Atomic no. & Atomic mass
Example- Helium atom
• p stands for proton.
• n for neutron.
 The atomic
mass (ma) is
the mass of
an atom (Proton
+ Neutrons). Its
unit is the
unified atomic
mass units (amu)
 Where 1
unified atomic
mass unit is
defined as ​1⁄12 of
the mass of a
single carbon-
12 atom, at rest.
 The number of
protons in the
nucleus of an
atom us called
atomic no.

10. Why was carbon-12 selected as the standard
element for atomic mass?
Before 1961, there actually were two sets of atomic masses (and everybody called them atomic weights
then). One scale was used by physicists; the other by chemists. Both were based upon weights compared
to Oxygen, rather than Hydrogen. Oxygen was used because it combines with a lot of things to form oxides.
This made it a better choice as a standard because of the ease of chemical analysis. Oxygen was set to
have an atomic mass of 16, which was just about 16 times as heavy as Hydrogen being 1. Unfortunately,
Chemists picked naturally occurring Oxygen, which is a mixed form of isotopes of Oxygen-16,17, and 18.
After all when one made an oxide of an element he would do so in naturally occurring oxygen. Physicists
picked the pure isotope Oxygen-16, because they tended to make their measurements on the basis of mass
spectrometry.
Though the ratio of any two atom’s masses was the same on either scale, it was horribly confusing, so in
1961, a compromise was reached. Instead of using either Hydrogen, or Oxygen as the standard, the isotope
of Carbon with 6 protons and 6 neutrons in its nucleus (Carbon-12) was given a mass of exactly 12. It was a
good choice, since it was in between the two previously used standards, and meant that nothing had to
change too much. Additionally, Carbon-12’s atomic mass could be measured particularly accurately
compared to the other elements on the periodic table.
So the Atomic Mass of Carbon-12 is defined to be 12 exactly and all other atomic, molecular and formula
masses are referred to this standard. That is why Carbon or C-12 particularly is used as the benchmark for
all atomic masses to be worked out on the Periodic Table ultimately.

12. Molar volume
• The molar volume is the volume occupied by one mole of ideal gas at STP. Its value is: 22.414 L mol¯1.
• One mole of any gas will have same value for this.
• One mole of O2 and H2 having different masses have the same molar volume at STP (standard temperature and
pressure).
Molar Mass (M)
• A physical property defined as mass of a given substance (chemical element or chemical compound) divided
by the amount of substance.
• Basic SI unit for molar mass is kg/mol.
Molecular mass
• Molecular mass or molecular weight is the mass of a molecule.
Formula Mass
• It is the sum of atomic masses of all atoms in a formula, regardless of whether or not the compound is
molecular.
• Unit is Kg/mol.

13. The key difference between
formula mass and molar mass is
that, the formula mass of a
molecule or a compound is the sum
of the atomic weights of the atoms
in its empirical formula while molar
mass is the mass in grams of 1 mol
of substance.
Either formula mass and molecular mass are same or not?

14. . Molecular Formula
A chemical formula that shows the total number and kinds of atoms in a molecule, but not
their structural arrangement. For example, the molecular formula of ascorbic acid is C6H8O6.
 Steps for Determining Molecular Formula
1. Calculate the empirical formula mass.
2. Divide the gram molecular mass by the empirical formula mass.
3. Multiply each of the subscripts within the empirical formula by the number calculated in
Step 2.
Empirical Formula
A formula giving the simpler proportions of the elements present in a compound but not
the actual numbers or arrangement of atoms. Formula for ascorbic acid is C3H4O3.
 Steps for Determining an Empirical Formula
1. Start with the number of grams of each element, given in the problem.
2. Convert the mass of each element to moles using the molar mass from the periodic
table.
3. Divide each mole value by the smallest number of moles calculated.
4. Round to the nearest whole number.

15. Can a compound
has the same
molecular and
empirical
formula?
Yeah!! I have same
molecular and
empirical
formula.It’s C2H6O.
Reason behind is the no.
of oxygen because its 1
and can’t be simplified
more.So,in case, this
compound has the
simplest ratio between
constiuents that is
1(integer).

16. Here‘s an example: What is the molecular formula of a compound that has a gram molecular mass of 34
g/mol and the empirical formula HO?
Solution:
Calculate the empirical formula mass.
You determine this number by finding the mass of HO (1 hydrogen atom and 1 oxygen atom).
So, the empirical formula mass is 17.01 g/mol.
Divide the gram molecular mass by the empirical formula mass.
Dividing the gram molecular mass by this value yields the following:
Multiply each of the subscripts within the empirical formula by the number calculated in Step 2.
Multiplying the subscripts within the empirical formula by this number gives you the molecular
formula H2O2. This formula corresponds to the compound hydrogen peroxide.
 To find molecular formula when the other is given.

17.  Isotopes
• Variants of a particular chemical element differ in neutron
number. But have the same number of protons.And
phenomenon is called isotopy.
 History
• The existence was first suggested in 1913 by the
radiochemist Frederick Soddy, based on studies of
radioactive decay chains that indicated about 40 different
species referred to as radioelements (i.e. radioactive
elements) between uranium and lead, although the periodic
table only allowed for 11 elements from uranium to lead.
Hydrogen isotopes
-atom with
1 p and 0-n
is called
protium.
-atom with
1 p and 0-n
is called
protium.
-atom with
1 p and 0-n
is called
protium.
H-atom
with 1-p
and 0-n is
called as
protium.

18. Relative abundance is the percent composition of
an isotope of a particular kind relative to the total number of
isotopes.
• Its used to find the most abundant isotopic atom of an element in
nature.
• It is measured by using mass spectrometry.
• E.g., H-atom with mass and atomic no. 1 is most abundant than other
types.
Mass spectrometry
• A process for measuring the masses of isotopes, molecules, and
molecular fragments by ionizing them and determining their
trajectories in electric and magnetic fields.
• Apparatus is called spectrometer.

19. Limiting Reactant
The limiting reagent (or limiting reactant, LR)
in a chemical reaction is the substance that is
totally consumed when the chemical reaction is
complete.
Car bodies are limiting
reagent.

20. Percent yield
It is the percent ratio of
actual yield to the theoretical yield.
It is calculated to be the
experimental yield divided
by theoretical yield multiplied by
100%.

21. i hope you
will find this
helpful

22. Thank you