C1501
Introduction
• Chemistry: the branch of science that describes matter; it’s chemical and
physical properties, the chemical and physical changes it undergoes and
the energy changes that accompany these processes.
• Matter: anything that has mass and occupies space.
• Atoms – are the smallest units that we associate with the chemical behavior of matter ; atoms
make up matter.
• Mass – is a measure of the quantity of matter that an object contains; mass does not vary
with location, i.e. moon or earth object has same mass.
• Chemistry can be described as a central science. It relies on the
foundations of mathematics and physics, and in turn underlies the life
sciences i.e. biology and medicine. To understand living systems fully, we
need to first understand the chemical reactions and chemical influences
that operate within them.
Introduction
• In our search for understanding, we eventually have to ask fundamental
questions such as:
1. How do substances combine to form other substances? How are energy
changes involved in chemical and physical changes?
2. How is matter structured? How are atoms and the ways that they
combine related to the properties of the matter that we can measure,
such as colour, hardness, chemical reactivity and electrical conductivity?
3. What fundamental factors influence the stability of a substance? How
can we force a desired but energetically unfavourable change to take
place? What factors control the rate at which a chemical change takes
place?
The Discovery of Cisplatin
The Discovery of Cisplatin
• In 1964 Barnet Rosenberg and his co-workers at Michigan State University
were studying the effects of electricity on bacterial growth. They inserted
platinum wire electrodes into a live bacterial culture and allowed an
electric current to pass.
• After 1 or 2 hours, they noted that the cell division in the bacteria had
stopped. Further experiments showed that cell division was inhibited by a
substance containing platinum, produced from the platinum electrodes by
the electric current. The researchers realised that a substance such as this
one could be useful as an anticancer drug, because cancer involves
uncontrolled cell division.
• Later, research confirmed this view, and today the platinum-containing
substance cisplatin is a leading anticancer drug.
The Discovery of Cisplatin
• This story illustrates 3 significant reasons to study chemistry:
1. Chemistry has important practical applications.
2. Chemistry is an intellectual enterprise, a way of explaining our
material world.
3. Chemistry figures prominently in other fields.
The Scientific Method
• Science is a frame work of gaining and organizing knowledge; it is not
simply a set of facts, but also a plan of action i.e. a procedure for
processing and understanding certain types of information. The
process that lies at the centre of scientific enquiry is called the
scientific method.
• There are many scientific methods depending on the nature of the
specific problem under study and on the particular investigator
involved. However, below is a useful general framework for a generic
scientific method:
The Scientific Method
Steps in the scientific method
1. Making Observations
Observations may be qualitative, e.g the sky is blue, water is a liquid,
etc. Observations may also be quatitative, e.g water boils at 100°C,
10.0mL of water weighs 10.0 g, etc. A qualitative observation does not
involve a number. A quantitative observation(Called a measurement)
involves both a number and a unit.
2. Formulating Hypothesis
A hypothesis is a possible explanation for the observation.
The scientific Method
3. Making Predictions
The hypothesis is then used to make predictions that can be tested by
performing an experiment.
4. Performing Experiments
An experiment is carried out to test the hypothesis. This involves
gathering new information that enables the investigator to decide
whether the hypothesis is correct i.e. whether it is supported by the
new information learned from the experiment. Experiments always
yield new observations, and this brings the process back to the
beginning.
The Scientific Method
• An experiment can be defined as an observation of natural
phenomena carried out In a controlled manner so that results can be
duplicated and rational conclusions obtained.
• As scientist observe nature, they often see that the same observation
applies to many different systems. Such generally observed behaviour
is formulated into a statement called a natural law (or just simply a
law). A law is a concise statement or mathematical equation about a
fundamental relationship or regularity of nature.
• For example, the observation that there is no observable change is
the quantity of matter during a chemical reaction or during a physical
change is known as the law of conservation of matter/mass.
The Scientific Method
• The law tells us what happens, but it offers no explanation for why we
observe a certain phenomenon. To try to explain why, we continue to
make observations, formulate hypothesis and test these against
observations.
• If a hypothesis successfully passes many tests it becomes known as a
theory. A theory is a tested explanation of basic natural
phenomenon.
experiments
results
hypothesis
further exps
negative results
positive results
theory
further exps
12
Law of Conservation of Mass
• The French chemist, Antoine Lavoisier (1743-1794) was one of the
first chemists to insist on the use of the balance in chemical research.
By weighing substances before and after a reaction, he demonstrated
the Law of conservation of Mass. This law states that the total mass
remains constant during a chemical reaction.
• Lavoisier conducted a series of combustion reactions and showed
that when a material burns, a component of air (which he called
oxygen) combines chemically with the material. E.g. when mercury is
heated in air, it burns or combines with oxygen to give a red-orange
substance whose modern name is mercury(II) oxide. We can
represent the chemical reactions as follows:
Law of Conservation of Mass
Mercury + oxygen  mercury (II) oxide
The arrow in the above representation means “changes to.”
• By strongly heating the red-orange substance, Lavoisier was able to
decompose it to yield the original mercury and oxygen gas.
• Therefore the law of conservation of mass can be represented as
follows:
MASSReactants = MASSProducts
Law of Conservation of Mass
Example 1
a)
You heat 2.53 grams of metallic mercury in air, which produces
2.73 grams of a red-orange residue. Assume that the chemical
change is the reaction of the metal with oxygen in the air.
Mercury + oxygen  red-orange residue
What is the mass of oxygen that reacts?
b)
When you strongly heat the red-orange residue, it decomposed
to give back the mercury and releases the oxygen, which you
collect. What is the mass of oxygen you collect?
Law of Conservation of Mass
Solution 1
Strategy: Apply the law of conservation of mass to the reactions.
MASSReactants = MASSProducts
a) From the law of conservation of mass,
Mass of mercury + mass of oxygen = mass of red-orange residue
2.53 g + mass of oxygen = 2.73 g
Mass of oxygen = 2.73g – 2.53g
= 0.20 grams
Therefore the mass of oxygen that reacts is 0.20 grams
Law of Conservation of Mass
b) When the red-orange residue is strongly heated, the mass of oxygen
collected is 0.20 grams.
Law of Conservation of Mass
• Definition of Terms
Mass: the quantity of matter in a material. It remains the same
regardless of where you are.
Weight: The force exerted on a body by gravity. It changes as
gravitational acceleration changes e.g. on earth vs on the moon.
Fg = mg
Fg – gravitational force
m – mass
Law of Conservation of Mass
Example 2
Methane gas burns in oxygen according to the reaction:
CH4(g) + 2O2(g)  CO2(g) + 2H2O(l)
Reactants : Methane gas (CH4)
Oxygen gas (O2)
Products: Carbon dioxide gas(CO2)
Water (H2O)
Note: subscripts tell how many of a particular type of atom are inside a molecule.
E.g. there are 4 hydrogen atoms in a molecule of methane and 2 hydrogen atoms
in a molecule of water.
Co-efficients tell how many of each particle is involved in the reaction. E.g. 2
molecules of oxygen react with 1 molecule of methane.
Law of Conservation of Mass
a) How would you draw this reaction as particles and show
conservation of mass?
Solution:
Key:
=C
=H
=O
+

+
Law of Conservation of Mass
Note about showing “conservation“ in particle diagram:
If you have the reaction:
A2
+
B2

A3B
You should show conservation by drawing:
+

i.e. 3A2
+
B2 
2A3B
Do not simply add stray “atoms” to molecules. It changes them into different
substances.
Law of Conservation of Mass
b) If 16 grams of CH4 reacts completely with 64 grams of O2, what mass of
products should form?
Solution:
CH4(g) +
2O2(g) 
CO2(g) + 2H2O(l)
16 g
64 g
Xg
MASSReactants = MASSProducts
∴ (mass of methane + mass of oxygen) = mass of carbon dioxide + mass of
water)
Substituting:
Mass products = 16g + 64g
= 80g
Law of Conservation of Mass
c) If 32g of CH4 reacts completely with 128g of O2, and 88g of CO2
forms, how many grams of H2O form?
CH4(g) + 2O2(g)  CO2(g) + 2H2O(l)
mass of methane + mass of oxygen = mass of carbon dioxide + mass of
water
Substituting:
32g + 128g = 88g + mass of water
Mass of water = (32g + 128g) – 88g
= 72g
Law of Conservation of Mass
d) If 8g of CH4 reacts completely with oxygen, and 22g of CO2 and 9g H2O
form, how much oxygen was consumed?
Solution:
CH4(g)
+
2O2(g)

CO2(g)
+
2H2O(l)
(mass of methane + mass of oxygen) = (mass of carbon dioxide + mass of
water)
Mass of oxygen= (mass of carbon dioxide + mass of water) – mass of
methane
= (22g + 9g) – 8g
= 31g – 8g
= 23g
Law of Conservation of Mass
e) 4 g of CH4 reacts with 20 g O2. The CH4 is used up completely, but
there is some O2 left over. Given that 20g of products was formed, how
much oxygen was used up?
Law of Conservation of Mass
Solution:
CH4(g)
+
2O2(g)

CO2(g)
+
2H2O(l)
4g
Xg
20g
X= 16g was used up.
Since 20g of oxygen was available, the left over oxygen would be;
= 20 g available O2 – 16 g used up O2
= 4g left over
Phases and Classification of Matter
Learning objectives/skills to develop
1. Describe the basic properties of each physical state of matter: solid,
liquid and gas.
2. Define and give examples of atoms and molecules.
3. Classify matter as an element, compound, homogeneous mixture or
heterogeneous mixture with regards to its physical state and
composition.
4. Explain the difference between pure substances and mixtures.
Phases and Classification of Matter
• There are two principal ways of classifying matter:
– By its physical state as a solid, liquid, or gas which is dependent on
the conditions, i.e. temperature, pressure.
– By its chemical constitution as an element, compound, or mixture.
Phases/States of Matter: Solids, Liquids, and Gases
• Solid: the form of matter characterized by
rigidity; a solid is relatively incompressible and
has a fixed shape and volume.
• Liquid: the form of matter that is a relatively
incompressible fluid; liquid has a fixed volume
but no fixed shape.
• Gas: the form of matter that is an easily
compressible fluid; a given quantity of gas will
fit into a container of almost any size and
shape.
Phases and Classification of Matter: Elements,
Compounds and Mixtures
• Matter can be classified into two categories: Pure substances and
Mixtures.
• Pure substances that can not be broken down into simpler substances by
chemical changes are called elements. Iron, silver, gold, aluminium, oxygen,
hydrogen, carbon, copper and sulphur are examples of common elements we
encounter. There are 113 elements in the periodic table of the elements;
about 90 occur naturally and 2 dozen or so have been created in laboratories.
The smallest unit of an element is the atom.
• Pure substances that can be broken down into simpler substances by
chemical changes are called compounds. A compound is, therefore, a
substance composed of two or more elements chemically combined (e.g.
NaCl, H2O). This breakdown may produce either elements or other
compounds or both. E.g. mercury (II) oxide can be decomposed to yield
mercury and oxygen. The smallest unit of a compound is the molecule.
Phases and Classification of Matter: Elements,
Compounds and Mixtures
• The properties of combined elements are different from those of free
elements. For example, white crystalline sugar(sucrose) is a compound
resulting from the chemical combination of the element carbon, which is a
black solid in one of its uncombined forms, and two elements oxygen and
hydrogen, which are colourless gases when uncombined.
• A mixture is a material that is composed of two or more types of matter
that can be separated by physical changes/means such as distillation.
Unlike a pure compound, a mixture has variable compositions. When you
dissolve sodium chloride in water, you obtain a mixture, its composition
depends on the relative amount of sodium chloride dissolved. You can
separate this mixture into solid sodium chloride and water by distillation.
Phases and Classification of Matter: Elements,
Compounds and Mixtures
Mixtures can be classified into two types:
1. Homogeneous mixtures
2. Heterogeneous mixtures
A homogeneous mixture, also known as a solution is mixture that is
uniform in its properties throughout given samples. When sodium
chloride is dissolved in water, you obtain a homogeneous mixture. Air is
a homogeneous gaseous mixture of primarily nitrogen and oxygen
gases. The gases are physically combined but not chemically combined.
Phases and Classification of Matter: Elements,
Compounds and Mixtures
• A heterogeneous mixture is a mixture that consists of physically distinct
parts, each with different properties. E.g. a mixture of sand and water, a
mixture of oil and water, fruit salad, a mixture of potassium dichromate
power and iron fillings, a mixture of salt and sugar mixed together.
• A phase is one of several different homogeneous materials present in a
portion of matter under study.
E.g. a mixture of ice cubes in water is said to be composed of two phases:
one phase is ice(solid) and the other is liquid water.
Ice floating in a solution sodium chloride in water consist of two phases: ice
and the liquid solution.
Note that a phase can be a pure substance in a particular state or a solution
in a particular state (solid, liquid or gaseous).
Phases and Classification of Matter
Law of Definite Proportions
• Antoine Lavoisier’s quantitative experiments showed that combustion
involved oxygen. He also discovered that life was supported by a process
that also involved oxygen and was similar in many ways to combustion.
• After 1800, chemistry was dominated by scientists who, following
Lavoisier’s example, performed careful weighing experiments to study the
course of chemical reactions and to determine the composition of various
chemical compounds.
• One of these chemists, a French, Joseph Proust (1754-1826) showed that a
given compound always contains exactly the same proportion of elements
by mass. This principle of the constant composition of compounds,
originally called Proust’s Law, is now known as the Law of Definite
Proportions.
– The law of definite proportions states that a pure compound, whatever its source,
always contains definite or constant proportions of the elements by mass; unique
composition for each compound but always the same for a particular compound.
• E.g. copper carbonate is always 5.3 parts copper to 4 parts oxygen to 1 part
carbon (by mass); and sodium chloride always has 1 part sodium & 1.54
parts chlorine.
Law of Definite Proportions
• One of these chemists, a French, Joseph Proust (1754-1826) showed that a
given compound always contains exactly the same proportion of elements
by mass. This principle of the constant composition of compounds,
originally called Proust’s Law, is now known as the Law of Definite
Proportions.
– The law of definite proportions states that a pure compound, whatever its source,
always contains definite or constant proportions of the elements by mass; unique
composition for each compound but always the same for a particular compound.
• E.g. copper carbonate is always 5.3 parts copper to 4 parts oxygen to 1 part
carbon (by mass); and sodium chloride always has 1 part sodium & 1.54
parts chlorine.
Implications of the law of Definite Proportions
• Constant composition implies constant properties (e.g.water always boils
at 100°C and freezes at 0°C).
Law of Multiple Proportions
• Proust’s discovery stimulated John Dalton (1766-1844), an English
school teacher, to think about atoms. Dalton reasoned that if
elements were composed of tiny individual particles, a given
compound should always contain the same combination of these
atoms. This concept explained why the same relative masses of
elements were always found in a given compound.
• But Dalton discovered another principle that convinced him even
more of the existence of atoms. He noted, for example, that carbon
and oxygen form two different compounds that contain different
relative amounts of carbon and oxygen:
Law of Multiple Proportions
Mass of oxygen that combines with 1g carbon
Compound I
1.33 g
Compound II
2.66 g
Dalton noted that compound II contains twice as much oxygen per
gram of carbon as compound I, a fact that could be easily explained in
terms atoms. Compound I might be CO and compound II CO2. This
principle, which was found to apply to compounds of other elements as
well, became known as the Law of Multiple proportions:
When two elements form a series of compounds, the ratios of
the masses of the second element can always be reduced to
small whole numbers.
Law of Multiple Proportions
• Compounds of differing mass ratios of the same elements are found,
but they will have different properties.
E.g. CO, CO2
H2O, H2O2
NO, NO2
Example
1. Given the following date for 3 compounds composed of nitrogen
and oxygen, propose possible formulas for compounds I, II and III:
Law of Multiple Proportions
Compound I
Compound II
Compound III
Solution:
𝐼
1.750 2
=
=
𝐼𝐼 0.8750 1
𝐼𝐼 0.8750 2
=
=
𝐼𝐼𝐼 0.4375 1
𝐼
1.750 4
=
=
𝐼𝐼𝐼 0.4375 1
Mass of Nitrogen that combines with 1g of Oxygen
1.750 g
0.8750 g
0.4375 g
Law of Multiple Proportions
• Compound I contains twice as much nitrogen (N) per gram of oxygen
(O) as does compound II.
• Compound II contains twice as much nitrogen per gram of oxygen as
does compound III.
These data can be summarised by the following sets of formulas:
Compound I
N2O
NO
N 4O2
Compound II NO
or
NO 2
N2O2
Compound III NO2
NO4
N204
Law of Multiple Proportions
2. Hydrogen and Oxygen are known to form form 2 compounds. The
hydrogen content in one is 5.93% and that of the other is 11.2%. Show
that this data illustrates the law of multiple proportions.
Solution:
In the first compound:
Hydrogen = 5.93%
Oxygen = (100-5.93)% = 94.07%
In the first compound the number of parts of oxygen that combine with
94.07
one part by mass of hydrogen =
= 15.86 parts
5.73
Law of Multiple Proportions
In the second compound:
Hydrogen = 11.2%
Oxygen = (100-11.2)%=88.88%
In the second compound the number of parts by mass of oxygen that
88.88
combine with one part by mass of hydrogen =
=7.9
11.2
Ratio of the masses of oxygen that combine with fixed mass of
hydrogen is 15.86 : 7.9 or 2: 1. This is consistent with the law of
multiple proportions.