This document provides an overview of atomic theory and the laws of chemical combination. It discusses the early Greek philosophers' debates on the nature of matter and whether it is continuous or made of discrete particles. John Dalton developed the modern atomic theory in the early 19th century, which included five main points. The document outlines the contributions of scientists like Thomson, Rutherford, and Bohr to models of atomic structure. It describes the three states of matter and defines the fundamental laws of conservation of mass, definite proportions, and multiple proportions discovered by scientists like Lavoisier, Proust, and Dalton. Examples are provided to illustrate applications of these laws.

1. CHAPTER ONE
Atomic Theory and
Law of Chemical
Combination

2. Introduction
• The concept of an atom can be traced to debates among Greek
philosophers that took place around the sixth century B.C.
• One of the questions that interested these thinkers was the nature of
matter.
• they asked, Is matter, continuous or discontinuous?
 That is, if you could break apart a piece of chalk as long as you
wanted, would you ever reach some ultimate particle beyond which
further division was impossible? Or
 Could you keep up that process of division forever?

3. Cont.
• Although, the debate over ultimate particles was never resolved,
the proponent (supporter) of the ultimate particle concept-
Democritus (470–380 B.C), named those particles atoms. In Greek,
atoms means "indivisible."
• However, Greek philosophers had no interest in testing their ideas
with experiments. They preferred to choose those concepts that
were most sound logically.

4. Cont’d
• Atomic theory is the idea that says matter is made up of little
units called atoms.
• As of 1897, the British scientist J.J. Thomson discovered that
atoms are in fact made up of smaller particles.
• English chemist John Dalton revived the old idea of atomic
theory , and used it to solve various problems that chemists
were grappling with at the time.
• His atomic theory was popularized and confirmed
experimentally over the course of the early 19th century.

5. Dalton's atomic theory: DAT
• However, John Dalton was the first to develop the modern atomic
theory, that there exists an ultimately small particle which
cannot be further divided. Dalton called this the atom.
 Dalton's atomic theory had five main points
1) All elements consist of tiny particles called atoms.
2) All atoms of a given element are identical to each other. (Atoms
of the same element are identical in all respects, having the
same size, shape and structure, and especially mass. )

6. Cont’d
3) All atoms of a given element are different than those of other
elements.(Atoms of different elements have different properties
and different masses)
4) Atoms of one element combine with other elements to create
compounds. They always combine in equal amounts.
5) Atoms cannot be created, divided, nor destroyed.

7. Drawbacks of Dalton's atomic theory of matter
• The indivisibility of an atom was proved wrong: an atom can be
further subdivided into protons, neutrons and electrons. However
an atom is the smallest particle that takes part in chemical
reactions.
• According to Dalton, the atoms of same element are similar in
all respects. However, atoms of some elements vary in their
masses and densities. These atoms of different masses are called
isotopes. For example, chlorine has two isotopes with mass
numbers 35 and 37.
• Dalton also claimed that atoms of different elements are
different in all respects. This has been proven wrong in certain
cases: argon and calcium atoms each have an atomic mass of 40
amu. These atoms are known as isobars.

8. Cont’d
• According to Dalton, atoms of different elements combine in
simple whole number ratios to form compounds. This is not
observed in complex organic compounds like sugar (C12H22O11).
• The theory fails to explain the existence of allotropes; it does not
account for differences in properties of charcoal, graphite,
diamond
 Merits of Dalton's atomic theory
• The atomic theory explains the laws of chemical combination.
• Dalton was the first person to recognize a workable distinction
between the fundamental particle of an element (atom) and that of
a compound (molecule).

9. Objectives
• Students will explain that atoms are the smallest
unit of an element and are composed of
subatomic particles.
• Students will analyze models of the scientific
theory of atoms.
• Students will analyze models and describe the
motion of particles in solids, liquids, and/or
gasses.

10. Atoms
• Matter is anything that takes up space and has
mass. All matter is made of atoms.
• Atoms are the basic building blocks of matter.
They make up everything around us; Your desk,
the board, your body, everything is made of
atoms!
• Atoms are too small to see without powerful
microscopes.
• a molecule is an independent structural unit
consisting of two or more atoms chemically
bound together, Elemental oxygen, for example,
occurs in air as diatomic (two-atom) molecules.

11. Atomic Structure
There are two basic components in every atom:
Electron Cloud
Nucleus

12. Three subatomic particles make up every atom:
Subatomic Particles
Subatomic Particle Charge Location
Proton Positive (+) Nucleus or “Core”
Neutron No Charge (0) Nucleus or “Core”
Electron Negative (-) Electron Cloud

13. Subatomic Particles
Electron Cloud:
• Electrons orbit the
nucleus.
Nucleus or “Core”:
• Protons and Neutrons
are found in the
nucleus.

14. Atomic Theory
Changes over time…

15. Atomic Theory
• Because we can not see atoms, we use models to teach and learn
about atoms.
• The atomic theory has changed over time as new technologies
have become available.
– Remember: Scientific knowledge builds on past research and
experimentation.
Scientist Information Model
John
Dalton
All matter is made of atoms.
Atoms are too small to see,
indivisible and indestructible.
All atoms of a given element
are identical.

16. Scientist Information Model
J.J Thomson
Discovered the negative
electron, and predicted that
there also must be a positive
particle to hold the electrons
in place.
Atomic Theory
Timeline
He is credited for the discovery of the electron and of isotopes,
and the invention of the mass spectrometer.
Thomson was awarded the 1906 Nobel Prize in Physics for the
discovery of the electron and for his work on the conduction of
electricity in gases

17. Scientist Information Model
Ernest
Rutherford
Discovered the nucleus of an atom and
named the positive particles in the
nucleus “protons”. Concluded that
electrons are scattered in empty space
around the nucleus.
Atomic Theory
Timeline
James
Chadwick
Discovered that neutrons were also
located in the nucleus of an atoms
and that they contain no charge.
Neutrons

18. Scientist Information Model
Neils Bohr Concluded that electrons are
located in planet-like orbits
around the nucleus in certain
energy levels.
Niels Bohr (1885 –1962) was a Danish physicist who made
fundamental contributions to understanding atomic structure and
quantum mechanics, for which he received the Nobel Prize in
Physics in 1922.
Scientist Information Model
(Many Scientists!)
The Modern
Atomic Theory
Electrons do not orbit the nucleus
in neat planet-like orbits but
move at high speeds in an
electron cloud around the
nucleus.

19. Three states of matter
• Solid:- The particles in a solid are very tightly packed and vibrate
in place.
• Solids have a definite volume and shape
Liquid :- The particles in a liquid are close together but can move
and flow past one another.
• Liquids have a definite volume but they do not have a definite
shape. This is why liquids like water take the shape of the
container they are in.
At room temperature most substances exist in one of three
physical states.

20. Gases
• Particles in a gas have higher amounts of energy than
those in a solid or liquid.
• Gases do not have a definite shape or volume. When
placed in a container, it fills up the entire container and
spreads out as far as possible.

21. 1.2. The Fundamental Laws of Chemical combination
• 1.2.1. Law of Conservation of Mass
• The total mass of material present after a chemical reaction
is the same as before the reaction.
• This Law was discovered by Antoine Lavoisier in about
1789.
• In a turnabout of the Scientific Method, Lavoisier had always
assumed this Law was true, and sought out experiments
which would verify his assumptions. As a result of numerous
combustion experiments conducted on systems in closed
containers, so as to retain any gases present,
• Lavoisier was able to unambiguously verify his assumptions and
formally state the Law of Conservation of Mass.

22. Cont’d
• For example, consider combustion reactions of elemental
Carbon. If the mass of the gasses are accounted for, it is found:

23. Example 2
• An 28.4g sample of sodium bicarbonate is
added to a solution of acetic acid weighing
15.0g. The two substances react, releasing
carbon dioxide to the atmosphere. After
reaction, the contents of the reaction vessel
weigh 24.0g. What is the mass of the carbon
dioxide given off during the reaction?

24. 1.2.2. Law of Definite Proportions
A chemical compound, no matter what its origin or its method of
preparation, always has the same composition; i.e., the same
proportions by mass of constituent elements.
• This Law, sometimes known as the Law of Definite Composition, was
first enunciated by Joseph Proust in 1799.
• Proust discovered this law while analyzing samples of Cupric
Carbonate.
• He found two samples, one prepared via synthetic methods, and the
other mined naturally (Malachite Green),
• possessed the same composition of elemental Carbon, Oxygen and
Copper:

25. Cont’d
• So, for example, if we decompose water by electrolysis and we
recover the elemental gases hydrogen and oxygen (not a
difficult task experimentally), and subsequently measure the
masses of each gas respectively, we can determine the composition
of this compound:

26. Cont’d

27. Cont’d
• For example, if Carbon monoxide is 27.29 % carbon and 72.71
% oxygen. So, how much carbon monoxide can be produced from
5.0g of carbon?

28. • Further, this result can be used to determine how much
oxygen would be consumed in the
• reaction forming this compound:
• mass Oxygen = 18.32g - 5.00g = 13.32g
 Home Work Exercise
 0.7 gm of iron unites directly with 0.4 gm of sulphur to form
ferrous sulphide.
 If 2.8 gm of iron are dissolved in dilute HCl and excess of sodium
sulphide solution is added, 4.4 g. of iron sulphide is precipitated.
Prove the law of constant composition.

29. 1.2.3. Law of Multiple Proportions
• The Law of Multiple Proportions was enunciated by John Dalton at
about the same time he postulated his Atomic Theory of Matter in
~1803.
• It was experimental results in the form which suggested the
validity of the Law of Multiple Proportions which provided Dalton
with the data needed to formulate the Atomic Theory.
• This Law, therefore, is a central key player in the development of
modern chemistry

30. Cont’d
• If two elements form more than a single compound, the
masses of one element combined with a fixed mass of the
second are in the ratio of small whole numbers.

31. Cont’d
• states that when two elements A and B combine to form more than
one compound, then the masses of B that combine with a fixed
mass of A are in simple ratio to one another.
• For example, carbon forms two oxides.
• In one, 12 g of carbon is combined with 16 g of oxygen (CO);
• In the other, 12 g of carbon is combined with 32 g of oxygen
(CO2).
• The masses oxygen combining with a fixed mass of carbon is in
the ratio 16:32, i.e. 1:2.

32. Cont’d
• Example
• Elements X and Y form two different compounds. In the first,
0.324 grams X is combined with0.471 gm of Y. in the second,
0.117 gram of X is combined with 0.509 gram of Y. show that
these data illustrate the law of multiple proportions.

33. Cont’d
• Answer
• In the first compound: 0.324g of X combines with 0.471 g of Y
• In the second compound: 0.117 g of X combines with 0.509 g of Y
• Therefore 0.324 g of X will combine with the wt. of Y
• Now, the Wts. of Y that combine with the same wt. of X, i.e., 0.324
g of it, are in the ratio of
• 0.471:1.4095 or 1:3. The ratio, being simple, illustrates the law of
multiple Proportions.

34. Exercise
• An element forms two oxides containing, respectively, 50%
and 40% by weight of the element.
• Show that these oxides illustrate the law of multiple proportions.

35. Example 2
• The mass of oxygen that combines with 1g of
Nitrogen to form two different compounds are
2.284g and 2.855g respectively. Show how this
data illustrate the law of multiple proportions.

37. Solution
• Fe (0.7 g) + S(0.4 g) FeS
• Fe (2.8 g) + 2HCl + Na2S FeS (4.4g) + 2NaCl + H2
• %e =mass of element/mass of compound