COMPOSITION OF

ATOM
Engr. Charito Ilagan
• The atom is a basic unit of an element that can enter
into chemical combination.
• It has a dense central nucleus surrounded by a cloud
of negatively charged electrons.
The atomic nucleus contains a mix of positively
charged protons and electrically neutral neutrons
(except in the case of hydrogen-1, which is the only
stable nuclide with no neutrons).
• The electrons of an atom are bound to the nucleus by
the electromagnetic force. Likewise, a group of atoms
can remain bound to each other by chemical
bonds based on the same force, forming a molecule.

• An atom containing an equal number of protons and

electrons is electrically neutral, otherwise it is positively
or negatively charged and is known as an ion.
• An atom is classified according to the number of
protons and neutrons in its nucleus:
the number of protons determines the chemical
element and the number of neutrons determines
the isotope of the element.
• Atoms can only be observed individually using special
instruments such as the scanning tunneling microscope.

• Over 99.94% of an atom's mass is concentrated in the

nucleus, with protons and neutrons having roughly equal
mass.
Their masses in kilograms :
Neutron (n0) = 1.6749286 x 10-27 kg
Proton (p+) = 1.6726231 x 10-27 kg
Electron(e-) = 9.1093897 x 10-31 kg
• The electrons determine the chemical properties of an
element, and strongly influence an atom’s magnetic
properties.

• The principles of quantum mechanics have been

successfully used to model the observed properties of
the atom.
➢EARLY ATOMIC THEORY
• Early thoughts concerning atoms were proposed by a
Greek philosopher named Democritus.

He suggested that the world was made of two things:

an empty space and tiny particles called atoms. His
belief , however, was not supported by experimental
evidence.
➢DALTON’S ATOMIC THEORY
John Dalton postulated the following in 1803:
▪ Each element is composed of extremely small
particles called atoms.
All atoms of a given element are identical , having
the same size, mass, and chemical properties.
The atoms of one element are different from the
atoms of all other elements.
▪ Compounds are composed of atoms of more than
one element.
In any compound the ratio of the number of
atoms of any two of the elements present are
either an integer or a simple fraction.

▪ A chemical reaction involves only the separation,

combination or rearrangement of atoms; it does
not result in their creation or destruction.
➢Dalton’s theory explains several simple laws of
chemical combination:

▪The Law of Definite Composition states that

different samples of the same compound
always contains its constituent elements in the
same proportion by mass.
▪ The Law of Multiple Proportion states that if two
elements can combine to form more than one
compound, the masses of one element that combine
with the fixed mass of the other are in the ratios of
small whole numbers.

▪ The Law of Conservation of Mass states that the total

mass of materials present after the reaction is the
same as the total mass before the reaction.
• Substances which chemically react are called reactants and
the newly formed substances are called products.

• A French chemist, Antoine Lavoisier, who is known as the

father of modern chemistry, changed chemistry from a
qualitative to a quantitative science. He proved that the
mass of the products in a chemical reaction is equal to the
mass of the reactants. There are no more atoms at the end
of the chemical reaction than there were at the beginning.
▪ The Law of Conservation of Mass states that
matter can neither be created nor destroyed in
a chemical reaction.
▪ Applying the Law of Definite Composition
1. A 0.001 g sample of magnesium when combined with
oxygen yields 0.166g of magnesium oxide. A second
magnesium sample with a mass of 0.144g is also
combined with oxygen. What mass of magnesium oxide
is produced from this second sample?
Given: 1st sample 2nd sample
m of Mg = 0.001 g m of Mg = 0.144 g
m of MgO = 0.166 g m of MgO = ?
• Required: m of MgO in the second sample
Solution:
1st sample 2nd sample
𝑚𝑎𝑠𝑠 𝑜𝑓 𝑀𝑔 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑀𝑔
=
𝑚𝑎𝑠𝑠 𝑜𝑓 𝑀𝑔𝑂 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑀𝑔𝑂
2) Use the information provided in first sample in
problem no. 1 to determine the mass of magnesium
contained in 0.500 g of MgO.
▪ Law of Conservation of Mass
1) A 0.455g sample of magnesium is allowed to burn in 2.315g
oxygen gas. The sole product is magnesium oxide. After the
reaction, the mass of unreacted oxygen is 2.015 g. What
mass of magnesium oxide was produced?
Given:
Before the reaction After the reaction
mMg = 0.445g munreacted O = 2.015 g
mO = 2.315 g
Required: mass of magnesium oxide
Solution:
Mg + O2 MgO + unreacted O2

Before the reaction = After the reaction

2) A 7.12 g sample of magnesium is heated with 1.8 g
bromine. All the bromine were used up and 2.07 g
of magnesium bromide is the only product.
▪ Law of Multiple Proportion
1) A student finds that 15.2 g of nitrogen will react with 17.37g,
34.74g or 43.43g of oxygen to form three different
compounds. What is the ratio of oxygen per gram of
nitrogen in each compound?
Given :
mN = 15.2 g m1 O = 17.37 g
m2 O = 34.74 g
m3 O = 43.43g
Required: Ratio of oxygen per g of nitrogen
Solution:
17.37 𝑔
Compound 1 = =
15.2 𝑔
34.74𝑔
Compound 2 = =
15.2 𝑔
43.43𝑔
Compound 3 = =
15.2𝑔

Ratio =
2) A chemist prepares three compounds containing
9.8g fluorine in 23.2g, 16.6g and 19.3g of sulfur.
Determine the ratio of the amount of sulfur per
gram of fluorine.
Early Research on Atomic
Principles

General Chemistry
▪ The first cathode ray tube (CRT) was made by Michael
Faraday in 1830.
▪ In passing electricity through the evacuated glass
tubes, he discovered cathode rays (a type of radiation
emitted by the negative terminal of the tube or
cathode)
▪ All subsequent experiments using the cathode ray lead
to the discovery of the following:
1) Protons - in 1896 by Eugene Goldstein
2) Electrons – in 1897 by LJ Thomson
3) Neutons in 1932 by James Chadwick
4) X-rays in 1895 by Wilhelm Roentgen
5) Radioactivity – in 1896 by Antoine Henri
Becquerel
Early Atomic Model
General Chemistry
 PLUM PUDDING MODEL

In 1897, J. J. Thomson
• imagined an atomic model with a
large and massive positively charged
entity, with tiny, far less massive
electrons embedded therein.
NUCLEUS MODEL
• In 1909, Ernest Rutherford discover that the positive charged
in an atom was concentrated in a very small volume called
the "atomic nucleus," rather than being spread out.
BOHR MODEL
In 1913, Niels Bohr
• depicts the atom as small, positively charged nucleus surrounded by
electrons that travel in circular orbits around the nucleus—similar in
structure to the solar system, but with attraction provided by
electrostatic forces rather than gravity.
SOMMERFELD’S MODEL
• In 1916, Sommerfield refined the assumptions of Bohr. He
believed that lower energy electrons revolved in a circular
path around the nucleus and the higher energy electrons
flow an elliptical path.
• WAVE MECHANICAL MODEL

• Physicists W. Heisenberg, L. de Broglie and E. Schrodinger

developed a model of an atom, which was based on
quantum mechanics advanced by Max Planck and Albert
Einstein.

• Quantum mechanics provides a mathematical description of

much of the dual particle-like and wave-like behavior and
interactions of energy and matter
• In 1924, L. de Broglie originated the idea that an
electron has dual properties – a particle and a wave.
• Schrodinger provided mathematical calculations of
the energy of each electron in an atom.

• This model shows that electrons occupy various

energy levels and sublevels.
• QUANTUM MECHANICAL MODEL
Schrodinger’s mathematical model for the atom led to the
present quantum mechanical model. The probability of
finding an electron is in a given region of space called
orbital.
• In 1926, Heisenberg proposed the Uncertainty Principle which
states that it is impossible to determine simultaneously the
exact position and exact momentum of a body as small as an
electron.
• Thus the precise description of the path of an electron in an
atom is impossible using Bohr’s approach.
• Instead Schrodinger used de Broglie’s relation to develop an
equation that describes the electron in terms of its wave
character.
➢Atomic Number and Mass Number

Each atom is characterized by its atomic number, Z,

and its mass number, A.

It is represented by a chemical symbol for the element

with the atomic number of the element placed at the
lower left and the mass number placed at the upper
left.
• The atomic number gives the number of protons; it
also refers to the number electrons in an electrically
neutral atom.

• The mass number refers to the total number of

protons and neutrons.
• Each element has at least one isotope with an
unstable nucleus that can undergo radioactive
decay.

• This can result in a transmutation that changes the

number of protons or neutrons in a nucleus.
• Isotopes are atoms of the same element with the
same atomic number but differ in their atomic
masses.

• Example: 35Cl 37Cl

• The average atomic mass is found by multiplying the

exact atomic mass of each isotope by its percent
abundance, then take the sum of their products.
Example:
1. Solve for the average atomic mass of chlorine given its 2
isotopes and their corresponding masses and the %
abundance:
35Cl – isotopic mass = 34.969 amu % abundance = 75.53%

37Cl – isotopic mass = 36.966 amu % abundance = 24.47%

Solution:
35Cl = 34.969 ( 0.7553) = 26.410

37Cl = 36.966( 0.2447) = 9.046

35. 458amu
2) Given:
12Carbon: isotopic mass = 12 amu % abundance = 98.892%
13Carbon: isotopic mass = 13.00335 % abundance = 1.108%
Solve for the average atomic mass.