Immaculate Conception Parochial School
272 Gen. Luna St. Concepcion, Malabon City
GENERAL CHEMISTRY 1
Midterms
(1st Semester)
Learner’s Activity Module
Photo Credit: https://assets.bwbx.io/images/users/iqjWHBFdfxIU/i55aetWklPc8/v1/-1x-1.jpg
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Chapter 1:
Properties of Matter and
Chemical Substances
Fig. 1. Two types of properties of matter
Remember Madonna singing her favorite lines 'I am a material girl, and I live in a material world'? Well, we
have to agree with Madonna: at least as far as her claim that she lives in a material world. It is a fact. We
live in a material world, that is, everything around us is material or matter in various forms.
Our world is made up of matter: matter being anything which takes up space and has mass. A drinking
glass, pair of eyeglasses, a purse: these all take up space and have mass and therefore are matter. On the
other hand, ideas, spoken words, thoughts are not matter because they do not have mass nor do they
occupy space. The amount of matter that an object contains is referred to as its mass.
II. LESSON OBJECTIVES:
Lesson Standard:
The learners demonstrate understanding of . . .

The properties of matter and its various forms.
Learning Competencies:



Use properties of matter to identify substances and to separate them
Recognize the formulas of common chemical substances.
Describe various simple separation techniques such as distillation, chromatography.
III. LESSON PROPER:
Phases of Matter
Matter can exist in any of the following states: solid, liquid, or gas and few other extreme phases such as
plasma, critical fluids, or degenerate. Under these different phases, matter exhibits characteristics
consistent with the state they are.
Different Phases of Matter
1. Solids are characterized by rigidity, a fixed volume and definite shape. There is very little space
between particles and cannot slide past one another. Molecules are very close together and cannot
move around. Generally, when a solid is heated or as pressure decreases, it will change to a liquid
form, and will eventually become a gas.
2. Liquids flow and take the form of the container they occupy. Molecules are close together but
more loosely packed than in solids and move around slowly. The molecules can slide past one
another and this accounts for its fluidity.
3. Gases have the capacity to diffuse and fill the container where they are placed. This is because its
molecules are widely separated and move around freely at high speeds. Because of the wide
space among the molecules, it is easier to compress gases compared to solids and liquids. It has
also the quality of expandability so that when a gas is transferred to a bigger container, it will
assume the volume and the shape of the new container.
4. Plasma is a special type of gas in which some of the atoms have become free floating ions and
free electrons. It is also referred to as ionized gas. As such, plasma is capable of conducting
electrical currents.
5. Supercritical Fluids are highly compressed gases which combine properties of gases and liquids.
These supercritical fluids have unique characteristics: they possess the density of a liquid and the
mobility of a gas.
6. Degenerate matter is basically a collapsed state of matter. This happens when the usual atomic
structure has broken down because electromagnetic forces are overcome by gravity. When atomic
nuclei collapse, electrons are unbounded from the nuclei and this is a state called Electron
Degeneracy.
Properties of Matter
1. Physical Properties - qualities of matter that are observable and measurable without changing its
composition. Physical properties can either be intensive or extensive.
a) Intensive properties are those properties that do not depend on the amount of the matter
present. Examples are color, odor, luster, malleability, ductility, conductivity, hardness,
melting and freezing point, boiling point and density. When properties are independent of
the quantity of matter present, these are intensive properties of the substance.
b) Extensive properties are those that depend on the amount of matter present such as:
 Mass which is the measurement of the amount of matter in an object.
 Weight which is the measurement of the gravitational force of attraction of the earth
acting on an object.
 Volume which is the measurement of the amount of space a substance occupies.
 Length which is the measurement of the reach or distance from one point to another.
2. Chemical properties – the characteristics of matter that it exhibits when it undergoes a change in
composition. They also describe the behavior of substances in the presence of other substances.
Chemical properties are attributes of matter that are affected when matter undergo chemical
changes and new substances are formed. Example: rusting of an iron nail and burning of paper.
Changes in Matter
Physical changes also occur when a substance is heated or cooled. When the melting point of a
substance is attained, a substance changes from solid phase to liquid phase. When the boiling point is
reached, a liquid change to a gas. These are only changes of state but the substance retained their
chemical properties.
When matter undergoes physical changes, its composition remains the same.
Physical Change
Change in Shape
Change in Size
Change in State
Change in Appearance
Process
Gold bar is hammered/molded in order to create
jewelry.
Cutting rolls of paper into long bond and short
bond sizes.
Boiling water so that it is converted to steam.
Dissolving sugar in water.
Whenever a substance undergoes a change so that one or more new substances with different
characteristics are formed, a chemical change has taken place. Chemical change will alter the original
substances to form new substances whose properties are different from the original.
To recognize whether new substances have formed, these are usually exhibited by the formation of
precipitates, production of a gas, a color change or there is release of energy in the form of heat. These
are just a few of the indicators of chemical changes but are not in any way absolute.
Classification of Matter
The schematic on the classification of matter shows that matter has two broad categories: as a
substance or as a mixture. When matter consists of a particular kind, exhibits uniformity throughout and
has a definite composition it is known as a substance. On the other hand, a mixture has no fixed
composition i.e. you can mix sugar and salt at varying proportions.
1. A substance is simply matter with definite chemical composition and distinct properties.
Substances cannot be separated into components by physical separation techniques. Some
substances, like water, can be broken down into elements by a chemical reaction (to break
chemical bonds). A substance can be solid, liquid, gas.
A substance can be an element or a compound but NOT a mixture. It can also be matter that exists in
its pure form, usually called a pure substance. A few examples of substances include Water (H₂O),
Hydrogen (H₂) and Neon (Ne).
1.1
Elements are pure in nature. They are made up of atoms and cannot be decomposed into
simpler substances by ordinary chemical means.
Elements are categorized as:
a. Metals. Metals are good conductors of heat, have properties of ductility [may be drawn into a
wire], malleability [may be hammered or rolled into shape], exhibit metallic luster, have high
melting points and high densities. Examples of metals are Aluminum, Iron, Gold.
b. Non-metals are dull and poor conductors of heat and electricity. Examples of non-metals are
Carbon, Nitrogen, Phosphorus. The non-metals are to the far right of the Periodic Table.
c. Metalloids bears both characteristics of metals and non-metals. Elements bordering the
stairstep line [an imaginary stairs that descends from the Group IIIA to VIA] resemble metals in
some of their properties and non-metals in others. These are referred to as metalloids. The
metalloids are: Boron, Silicon, Germanium, Arsenic, Antinomy and Tellurium.
1.2
Compounds are pure substances that are made from two or more elements that have reacted
chemically with each other. Compounds may be broken down or decomposed by chemical
change into two or more simpler pure substances.
a. Organic - compounds made up of the carbon element in combination with other elements.
b. Inorganic compounds made up of other elements other than Carbon.
 Salts - formed when an acid is neutralized by a base. It consists of an anion from the
acid and a cation from the base.
 Acids - compounds that produce Hydrogen [H] ion in aqueous solution.
 Bases - compounds that contain hydroxide [OH] ion in aqueous solution.
2. Mixtures are formed when two or more substances are mixed together. Since there are no
chemical bonds that are formed between the two substances, they can be separable by physical
means. Also, the constituents retain their identity since their physical and chemical properties have
not been changed or altered but simply mixed.
A mixture can involve two or more substances of the same phase or different phases. It is possible
to mix a liquid and a solid (like water and sand), or solid and solid (like sugar and salt), or liquid and
liquid (water and oil), or in the case of gas and gas(Nitrogen and Oxygen).
Mixtures can either be homogeneous or heterogeneous.
2.1.
Homogeneous mixtures are perfectly uniform in composition throughout. When two or
more substances are mixed and the final substance formed appears to have the same
chemical composition, it is considered as a homogeneous mixture.
Homogeneous mixtures consist of a solute (dissolved material) in a solvent (dissolving
material). When the mixture is between two liquids, the larger volume is referred to as the
solvent and the smaller quantity is the solute. The solute is scattered throughout the
solvent attributing to the particulate nature of matter, thereby forming a solution that
appears to be uniform. When there is uniformity throughout the solute-solvent mixture, it is
classified as a solution. Solutions are classified as homogeneous materials.
2.2.
Heterogeneous mixtures result when the phases that were mixed are distinguishable
from each other. Such boundary between the phases is usually observable to the naked
eye. A heterogeneous mixture does not have uniform properties and composition.
Examples of heterogeneous mixtures are sand and gravel, kerosene and water.
Methods of Separating Mixtures
Since mixtures are made up of substances that are not bound chemically, it is possible to separate them
into their original components. Some methods that can be used to separate mixtures are the following:
Centrifugation
Centrifugation is the process of separating heterogeneous mixtures by applying centripetal force to a
mixture using a centrifuge machine. A centrifuge machine is a piece of equipment that is equipped with a
fixed axis and a rotating unit. When the rotating unit spins, it provides centripetal acceleration that will
cause heavier substances to move outward and settle at the bottom of a tube. More dense components of
the mixture migrate away from the axis and gather at the bottom of the tube while the less-dense
components migrate towards the axis and can be withdrawn by pouring out or by the use of a pipette.
Chromatography
This method is often used in the food industry. It is used to identify chemicals (coloring agents) in foods or
inks. Separation of the components of inks is done in the following manner: A blob of ink is smeared on a
special paper called filter paper. The paper is placed in a trough of solvent. The solvent used depends on
the chemicals in the ink blob. As the paper gets soaked upwards, it attracts the various chemicals in the ink
blob. Because different chemicals have different rates of attractions to the solvent, the chemicals will travel
upwards in different amounts.
Decantation
Decantation is a process that allows a heterogeneous mixture of solid and liquid to be separated. The
mixture of solid and liquid are allowed to settle and separate by gravity. Then, the liquid is poured off,
leaving the solid behind. Typically, you would expect that a small amount of the lighter liquid will not be
completely withdrawn from the mixture and will be left behind.
Distillation
In distillation, the components of the mixture (which is a solid dissolved in a liquid) are separated. A
distillation set-up is used for the separation process. By application of heat, the liquid is converted to steam;
the vapors pass through a condenser and as w 'er passes through the outer tube of the condenser, the
vapors are cooled and converted back to its liquid state. On the other hand, the solid component of the
mixture is left behind as a residue.
Evaporation
Mixtures of a soluble solid and a solvent can be separated by the process of evaporation. This process
involves the application of heat to the solution to allow the solvent to evaporate leaving behind the solid
component of the mixture as a residue.
Filtration
Filtration can be resorted to for mixtures of solids and liquids where the solids are of smaller size that do
not settle at the bottom of a container despite long periods of standing. This process involves the use of a
filter paper that is lined into a funnel. The end of the funnel empties into a beaker and the mixture of liquid
and solid is poured into the funnel. The liquid drains through the filter paper into the beaker, leaving the
solid particles trapped on the filter. The liquid that is collected is called the filtrate and the solid that
remains on the filter paper is called the residue.
Fractional Distillation
Fractional distillation is a technique that can separate two liquids that. dissolved in each other (miscible
liquids). These liquids, having different boiling points is the basis for this technique. The set-up is like the
simple distillation process. Here, however, the temperature is raised initially to the temperature that
corresponds to that of the lower boiling point of the two liquids. This will allow one liquid to evaporate, and
then condense back to liquid as it passes through the condenser. Once this 'fraction of the mixture is
collected, the temperature is elevated to the boiling point of the other liquid. Then the second 'fraction' is
collected corresponding to the higher boiling component of the two substances.
Magnetism
Mixtures that consist of metallic and non-metallic substances can be separated with the use of a magnet.
Since the metallic substance exhibits properties such as being attracted to the magnet, you are able to
cause the metallic substance to be withdrawn from the mixture leaving behind the non-metallic substance.
Separatory Funnel
To facilitate the separation of liquid-liquid mixtures that did not mix, a separatory funnel can be used. This is
a simple laboratory equipment made mostly of glass that has a tap fitted at the end of its conical section.
When two liquids do not mix and they are placed in a separatory funnel, the high density liquid settles at the
bottom while the lower density liquid stays on top. Allow the liquid mixture to settle in the separatory funnel
after which the tap is opened to allow the high density liquid to pass through and collect in a container
provided. As soon as the high-density liquid has passed through, the tap is closed. What is left in the
separatory funnel is the low-density liquid.
Chemical Symbols to Represent Substances
All matter in our universe is made up of elements or elements in combination with other elements to form
compounds. Elements are represented by chemical symbols which consist of the first one or two letters of
the name of the element. The first letter is capitalized while the second letter is written in lowercase. Since
the name of some elements are derived from Latin names, their chemical symbols are usually a derivative
from these Latin names. Example Sodium (Natrium-latin) is Na and not S or So; Potassium (Kalium - latin)
is K and not P or Po; Iron (Ferum -latin) is Fe and not I or Ir.
These chemical symbols are used in order to write chemical formulas of compounds. The numbers that are
written as subscripts after each element represent the proportion of these various elements in the
compound. This proportion is the same throughout the compound. Example: Table sugar has the chemical
formula C₂H₂O. This means that table sugar is composed of the elements Carbon, Hydrogen and Oxygen
in the proportion of 12 Carbons atoms, 22 Hydrogen atoms and 11 Oxygen atoms to form 1 molecule of
sugar.
Common Name
Alcohol
Baking soda
Simple sugar
Vinegar
Sand
Table salt
Milk of magnesia
Chemical Name
Ethyl alcohol
Sodium bicarbonate
Glucose
Acetic Acid, Ethanoic Acid
Silicon dioxide
Sodium Chloride
Magnesium hydroxide
Formula
C2H5OH
NaHCO3
C6H12O6
CH₃COOH
SiO2
NaCl
Mg(OH)2
IV. ACTIVITIES/EXERCISE:
I.
Identify whether the following processes involve Chemical or Physical Changes:
___________1. Hydrogen and Oxygen combine to form Water
___________2. Fireworks during New Year
___________3. Making of ice cubes for sale during summer
___________4. Digestion of food
___________5. Preparing a cup of coffee with sugar
II.
Classify each of the following as Homogeneous or Heterogeneous:
____________1. Blood
____________2. Liquid Milk
____________3. Brass (alloy of Copper and Zinc)
____________4. Salt water in the ocean
____________5. Iron and aluminum filings
III.
Classify each of the following as Element, Compound or Mixture:
____________1. Helium
____________2. Ethyl Alcohol (C₂H₂OH)
____________3. Mercury
____________4. Air
____________5. Gasoline
V. ASSESSMENT:
Direction: Encircle the letter of the correct answer.
1. A phase of matter characterized by its rigidity, a fixed volume and definite shape.
a. Solid
c. Gas
b. Liquid
d. Plasma
2. This is a process of separating heterogeneous mixtures by applying centripetal force to a mixture
using a centrifuge machine.
a. Decantation
c. Filtration
b. Centrifugation
d. Fractional Distillation
3. You’re working with your teacher to perform a demonstration about different separation technique.
Which among the following would you do if you’re asked by your teacher to show how a powdered
chalk and water could be separated.
a. Chromatography
c. Magnetism
b. Distillation
d. Decantation
4. Which among the following chemical formula has an original name of “vinegar”?
a. NaHCO3
c. CH₃COOH
b. NaCl
d. Mg(OH)2
5. You have a younger sister who’s trolling along the social media. Suddenly, she saw an unfamiliar
chemical formula of “C6H12O6“ and she tried to ask you about its name. Using your knowledge
about common chemical substances, which of the following would be your response to your
younger sister?
a. Milk of magnesia
c. Table salt
b. Ammonia
d. Simple sugar
VII. ADDITIONAL RESOURCES:
The learners could visit the following links for further knowledge for the following lessons.
Properties of Matter and Chemical Substances
-
https://courses.lumenlearning.com/boundless-chemistry/chapter/classification-of-matter/
https://courses.lumenlearning.com/boundless-chemistry/chapter/physical-and-chemicalproperties-of-matter/
VIII. REFERENCES:
 General Chemistry 1 with Laboratory Manual (2016), G. M. Garcia, Unlimited Books Library Services
& Publishing Inc.
Chapter 2:
Atomic Structure and
Nomenclature
Fig. 1. A comparative representation of structures of the inner planets in the Solar system to the atom.
What we know today… what we accept as laws that govern chemical processes… what we hold as sacred
and irrefutable facts are the result of a long history of work, sweat and toil of scientist who attempted to find
explanations for their observations and various phenomena around them.
Thanks to generations who preceded us, a long string of names in the Who's Who of science: Antoine
Lavoisier, Joseph Proust, Lord Ernest Rutherford, Niels Bohr, Marie Curie, John Dalton and other names
whose works provided the foundation for our understanding of matter. This understanding allows us to
harness matter to benefit human kind and to open possibilities yet unknown.
II. LESSON OBJECTIVES:
Lesson Standard:
The learners demonstrate understanding of . . .


Atomic structure
Formulas and names of compounds
Learning Competencies:



Recognize common isotopes and their uses.
Represent compounds using chemical formulas, structural formulas and models.
Name compounds given their formula and write formula given the name of the compound.
III.
LESSON PROPER:
The idea of atoms was first suggested by Democritus, an ancient Greek who lived in the 4th B.C. However,
his idea of the atom could not support chemical phenomena. He called these small pieces of matter
"atomos," the Greek word for indivisible. Democritus, theorized that atoms were specific to the material
which they composed.
John Dalton (1766-1844), an English chemist and physicist was able to relate chemical changes to the
level of individual atoms and state his atomic theory with the following ideas:
DALTON'S ATOMIC THEORY
1. All elements are composed of tiny indivisible particles called atoms.
2. Atoms of the same element are identical. The atoms of any one element are different from those of
another element.
3. Atoms of different elements can combine with one another in simple whole number ratios to form
compounds.
4. Chemical reactions occur when atoms are separated, joined or rearranged. However, atoms of one
element are not changed into atoms of another by a chemical reaction.
Dalton’s idea gave birth for the following laws:
1. Law of Conservation of Mass. This was postulated by French chemist Lavoisier in 1785. It states
that when a chemical reaction takes place, there will be no detectable change in masses of the
substance. Expressed in equation form: Mass of reactants = Mass of the products
Ex. 5 grams of Magnesium Chloride completely reacted with 20 grams of Ammonium Oxalate to
produce Magnesium Oxalate and Ammonium Chloride. What could be the mass of Ammonium
Chloride if the mass of Magnesium Oxalate is 15 grams?
REACTANTS
Magnesium Chloride = 5 grams
Ammonium Oxalate = 20 grams
Total = 25 grams
PRODUCTS
Magnesium Oxalate = 15 grams
Ammonium Chloride =?
Total =?
The Law of Conservation of Mass would explain that if the original reactants totaled 25 grams, then
the products would be unchanged. That means that the final products also weigh 25 grams. Since
we are given the weight of Magnesium Oxalate as 15 grams, then we can determine that the
weight of Ammonium Chloride is 25 grams less 15 grams or 10 grams.
2. Law of Definite Proportion. This Law states that when elements combine to form a given
compound, they do so in a fixed and invariable ratio by weight.
Ex. Calcium Chloride has the Molecular Formula of CaCl₂. This means that the Elements: Calcium
and Chlorine make up the compound in a 1:2 ratio.
Under the right conditions, for as long as there are Calcium and Chloride in a quantity for a 1:2
ratio, Calcium Chloride will form in a definite proportion of 1:2.
3. Law of Multiple Proportion. When atoms combine to form a compound, they always combine in
definite ratio and proportion expressed in small whole numbers. These same elements may also
combine in a different proportion to yield a different compound.
Ex.
a. Nitrogen and Oxygen can combine in a variety of proportions: 1:1 ratio or 1:2 ratio
NO (Nitric Oxide)
NO2 (Nitric Dioxide)
b. Tin and Oxygen can combine in different proportions to form different compounds: 1:1 ratio or
1:2 ratio
SnO [Tin (II) Oxide]
SnO2 [Tin (IV) Oxide]
Atomic Structure
In the decades following the discovery of the neutron, scientists discovered more subatomic particles.
Physicists refer to two families of particles: the leptons, the electron, the mu-meson, taumesons, and
neutrinos; the second family is called the hadrons that include the proton and neutron, and quarks.
Most chemical reactions can be explained by the atomic structure advanced by John Dalton and for
purposes of this study and discussion, the structure of the atom shall consider the three subatomic
particles: proton, electron, and neutron.
1) Electrons (e) - are the negatively charged sub-atomic particles of an atom. They are located
outside the nucleus and occupy electron orbitals. (Discovered by Sir Joseph J. Thompson)
2) Protons (p+) - the positively charged sub-atomic particle of an atom. They are located in the
nucleus. (Discovered by Baron Ernest Rutherford)
3) Neutrons (nº) - are the neutral sub-atomic particles of an atom and are also located in the nucleus.
(Discovered by Sir James Chadwick)
The number above the Element's symbol is the Atomic Number. The atomic
number is the whole number that increases as you read across each row of the
periodic table from left to right. This also corresponds to the number of protons.
Since the element is a neutral entity, the positive charge of the protons must be
cancelled out by a similar number of electrically charged particles. This would
mean that the number of protons should also be the number of electrons in an
element's atom. The number below the element's symbol corresponds to the
Atomic Mass expressed as atomic mass unit (amu). By international agreement,
Fig. 2. The element the atomic mass standard is the pure isotope Carbon-12, which is assigned a
“Gold”
mass of exactly 12 atomic mass units (12u). Based on this standard, an atomic
mass unit (amu) is exactly one-twelfth the mass of a carbon-12 atom. ¹
The term atomic mass replaced the older term atomic weight which is still used by the International Union
of Pure and Applied Chemistry (IUPAC). These two terms are used interchangeably. Since the mass of the
atom is concentrated in the nucleus (which is made up of protons and neutrons), given that protons must
equal the number of electrons, the number of neutrons can thus be mathematically determined by
subtracting from the Atomic Mass (protons + neutron) the Atomic Number (protons).
IONS
In their elemental state, elements have the same number
of protons and electrons. Thus, the positive charges
cancel out the negative charges and elements are neutral.
Sometimes atoms gain or lose electrons. When they do so
they either gain a 'negative charge or a 'positive charge
because the number of electrons does not equal the
Fig. 3. Formation of sodium cation from
number of protons in the atom or molecule. When atoms
element “sodium”.
GAIN electrons, the result is that there are more electrons
than protons. This would result to a net 'negative charge’. When atoms LOSE electrons, the result is that
there are less electrons than protons and a net 'positive charge of the ion’. A positively charged ion is called
a CATION while a negative charged ion is called an ANION.
ISOTOPES
Isotopes have different atomic masses (mass number). Isotopes of an element have nuclei with the same
number of protons (the same atomic number) but different numbers of neutrons. Isotopes have different
mass numbers, which give the total number of nucleons (protons + neutrons).
Example, the most common isotope of Hydrogen
has no neutrons at all but there is also a Hydrogen
isotope called Deuterium which has one neutron;
and Tritium which has two neutrons. The nucleus of
the above isotopes is represented as follows:
To symbolize the composition of an isotope, two
numbers are written to the left of the chemical
symbol. The mass number is written as a
superscript (above) and the atomic number is
written as a subscript (below).
Fig. 4. Different isotopes of hydrogen element.
Example: Three isotopes of Carbon are Carbon-12, Carbon13 and Carbon-14. Write the Chemical symbol for each.
Answer: Since Carbon has an atomic number of 6, all Carbon
atoms have 6 protons
NUCLIDE
Fig. 5. Isotopes of Carbon
Nuclide is any particular atomic nucleus with a specific atomic number “Z” and mass number “A”. It is
equivalently an atomic nucleus with a specific number of proton and neutrons. Collectively, all the isotopes
of all the elements form the set of nuclides. The terms isotope and nuclide are often used interchangeably.
Isotope is best used when referring to several different nuclides of the same element; nuclide is more
generic and is used when referring one to one nucleus or several nuclei of different elements. For example,
it is more correct to say that an element such as Fluorine consists of one stable nuclide rather than that it
has one stable isotope.
The Periodic Table
A table is a way by which
information can be efficiently
organized. A simple glance at
tables can already tell a volume of
information one has to know.
Having studied properties of
elements and having found that
certain elements showed
similarities in chemical and
physical properties, scientists
Fig. 6. The Periodic Table of Elements
sought ways to classify the elements.
Early attempts at classification were rudimentary. In 1817, John Dobereiner, a German chemist grouped
the metals Calcium, Barium and Strontium because they showed very similar chemical properties. He
referred to these 3 elements as the triad.
Doing X-ray experiments, Henry Moseley (1887-1915), a British scientist found the reason for the
exceptions to Mendeleev's period law. This led to the revision of the periodic law that arranged elements
according to their atomic numbers instead of their atomic masses. The periodic law thus states: The
properties of elements are a periodic function of their atomic numbers.
The Modern Periodic Table that we know today is loaded with information. Other than the atomic number,
atomic mass, and symbol of the element it also indicates the number of electrons surrounding the nucleus
and the order with which they fill the sublevels or orbitals (location of electrons).
At this point, we will use the Periodic Table, only to be familiarized with their GROUP and PERIOD to help
us in writing chemical formulas. A more in-depth discussion of the periodic table will be taken up in future
discussions as we expand on concepts that are fundamental to chemistry. A column of elements is known
as a group which carries the label IA, IIA, IIIB, etc. A period is a horizontal arrangement of elements and
are labeled as 1,2,3, etc.
CHEMICAL FORMULAS
From individual elements in the periodic table, chemical compounds are formed. Millions of chemical
compounds have been identified and today we have not yet run out of possibilities. These compounds are
represented by chemical formulas to show the kinds and numbers of atoms that are present in the
smallest representative unit of the substance.
The chemical formula of a molecular compound is called a molecular formula. The molecular formula
contains the symbol of the elements that have combined. If there are any subscripts (numbers written
below the element), these indicate the number of atoms of said element that are present in a molecule of a
compound.
In certain instances, ionic compounds are formed instead of molecular compounds. In molecular
compounds, the elements are chemically bonded together. For ionic compounds, however, the ions are
bonded together by electrostatic forces. When dissolved in water, the ions dissociate and freely move
around as cations (+) and anions (-). The chemical formula of an ionic compound is called a formula unit.
The formula unit is the lowest whole number of ions in an ionic compound.
In order to write chemical formulas, we need to know the types of ions that atoms tend to form. We also
need to know the ionic charges of the elements. These can be determined by using the periodic table.
For Group IA, IIA and IIIA, the ionic charge is positive and is numerically equal to the group number. Group
IA forms cations and have a +1 charge. Group IIA form cations with a +2 charge. Group IIIA forms a +3
cation. The elements in Group IVA and Group O do not commonly form ions. Group IVA ordinarily form
molecular compounds. Group O seldom form compounds and for that are referred to as Noble Gases.
Unlike the Group A elements, the transitions metals that make up Group B of the Periodic Table have more
than one common ionic charge. A few transition metals have only one ionic charge. These are initial
information that are helpful in writing chemical formulas.
The periodic table here is presented in such a way that
the valence of the elements used to combine with
other elements or ions is shown. The valence figures
significantly in the writing of chemical formulas of
compounds both molecular and ionic.
One other group of ions, the polyatomic ions, has to
be explained. Polyatomic Ions are tightly bound
groups of atoms that behave as a unit and carry a
Fig. 7. Periodicity of Elements by Oxidation Number
charge. An example of a polyatomic ion is the carbonate
ion represented as (CO). A carbonate polyatomic ion is composed of one carbon atom and three oxygen
atoms that form a unit and have a charge of -2.
Fig. 8. Common Polyatomic Ions
Writing Formula
In writing chemical formulas, the general rule is to write the symbol of the positive ion first followed by the
negative ion. Writing the subscript becomes only troublesome when polyatomic ions are involved. This is
remedied by enclosing the polyatomic ion in open and close parenthesis and then followed by the subscript
written outside the close parenthesis.
A systematic way of writing formulas has been devised using the 'crisscross' method. Illustrations on how
this is done:
1. Rule 1: Criss-Cross Rule - Chemical compounds are electrically neutral. The total number of
positive charges is equal to the number of negative charges hence valence of positive entity
[whether ion or radical] equals the subscript of the negative entity.
However, when the valence number is one, the convention is to eliminate the subscript; then the
symbol itself stands for one unit.
Example: for the chemical compound Sodium Chloride, the Sodium is the positive unit and the Chloride
is the negative unit.
Step 1 - the symbol and the valence of sodium is Na' and that
of Chlorine is Cl-.
Step 2 - each of these units has the same number of charges
namely +1 and -1.
Only one unit of each is required in the chemical formula of
NaCl.
Fig. 9. Forming chemical formula of
table salt (NaCl)
For the compound Calcium Phosphate, the Calcium is the
positive unit and the Phosphate is the negative unit.
Step 1- the symbol and the valence of Calcium is Ca2 and the
formula and valence number of Phosphate is PO³
Step 2 - by criss-crossing, the -3 valence of Phosphate
becomes the subscript of Calcium and the +2 valence of
Calcium becomes the subscript of Phosphate, thus the formula
is: Ca3(PO)₂
Fig. 9. Forming chemical formula of
Calcium Phosphate.
2. Rule 2. All radicals taken more than once must be enclosed in parentheses () or brackets [].
Example: Al₂(SO4)3 or (NH4)2S
3. Rule 3. All subscripts must be reduced to lowest terms.
Chemical Nomenclature
During the earlier days of chemistry, the discoverer of a new compound gave it a name. The name usually
described some physical or chemical property of the substance or its source. As the science of chemistry
developed, it became apparent that some systematic method of naming chemical compounds was needed.
Naming of compounds then came under a set of rules that defined the norm:
1. BINARY COMPOUNDS- combination between two (2) elements:
Ionic Binary Compounds - a combination between a metal and a nonmetal.
I.
Metals with only ONE oxidation number
Rule: Name the metal followed by the root name of the non-metal + the suffix IDE.
Example: Silver = Ag+1 + Iodine = I-1 turns into “Silver Iod-IDE or Silver Iodide = AgI”
II.
Metals with two (2) oxidation numbers
Rule: Give the root name of the metal + the suffix-OUS for the lower oxidation number; the suffix-IC for
the higher oxidation number + the root name of the non-metal + the suffix-IDE.
Example:
1. Cobalt=Co+2 and Phosphorus = P-3
Since Cobalt has 2 oxidation numbers namely +2 and +3
Lower oxidation number is used, the suffix -OUS applies
Cobalt Cobaltous
Phosphorus = Phosp-IDE
Cobalt = Co+2 and Phosphorus = P-3
Formula = Co3P2
2. Cobalt = Co3 and Phosphorus = P-3
Cobalt has 2 oxidation numbers namely +2 and +3
Higher oxidation number, use suffix -IC
Cobalt = Cobaltic
Phosphorus = Phosp-IDE
Cobaltic Phosphide = Co3(P)3
Apply Rule of reducing to lowest terms: Co3(P)3 = CoP
(Just made the other pages in picture form)
IV. ACTIVITIES/EXERCISE:
I. Write the correct formulas for the following compounds:
a.
b.
c.
d.
e.
Sodium Sulfide
Lead Iodide
Copper (II) Nitrate or Cupric Nitrate
Iron (III) Permanganate or Ferric Permanganate
Ammonium Arsenate
__________________________________
__________________________________
__________________________________
__________________________________
__________________________________
II. Name the following compounds according to the Rules on Chemical Nomenclature:
a)
b)
c)
d)
e)
CCl4
HClO
NaMnO4.5H2O
HgCl2
Al2S3
____________________________
____________________________
____________________________
____________________________
____________________________
V. ASSESSMENT:
Direction: Choose the letter of the correct answer by changing its font color into “RED”,
1. He is an English chemist and physicist was able to relate chemical changes to the level of
individual atoms and state his atomic theory
a. John Dalton
C. Henry Moseley
b. Antoine Von Lavoiser
D. Henri Becquerel
2. Idea in chemistry which states that when a chemical reaction takes place, there will be no
detectable change in masses of the substance.
a. Law of Definite Proportion
C. Law of Multiple Proportion
b. Law of Conservation of Mass
D. Law of Partial Pressure
3. Zhayie is skimming along the periodic table of elements. While skimming each of the chemical
symbols of elements, she asked herself about how many electrons does an element “Germanium”
have if its atomic number is 32.
a. 31 eC. 32 eb. 35 eD. 34 e4. Zhairuz is playing an educational game about chemical bonding. He observed that the element
Magnesium (Mg) tends to donate its electrons to element Oxygen (O) when they bond to each
other. Which among the statement will be agree about the mentioned situation above?
a. The Magnesium and Oxygen atoms will turn to be ANIONS.
b. The Magnesium and Oxygen atoms will turn to be CATIONS.
c. The Oxygen turns to be the CATION while the Magnesium will be an ANION.
d. The Magnesium turns to be the CATION while the Oxygen will be an ANION.
5. Zhaifa was daydreaming inside her chemistry class that’s why she was being called by her teacher
and asked her about the description of an isotopes. Which of the following would she answer to her
chemistry teacher?
a. Isotopes of an element have nuclei with the same number of protons but different numbers of
neutrons.
b. Isotopes of an element have nuclei with the different number of protons but have sane
numbers of neutrons.
c. Isotopes of an element have nuclei with the same number of protons but different numbers of
electrons.
d. Isotopes of an element have nuclei with the same number of neutron but different numbers of
protons.
VII. ADDITIONAL RESOURCES:
The learners could visit the following links for further knowledge for the following lessons.
Atomic Structure and Nomenclature
-
https://courses.lumenlearning.com/introchem/chapter/isotopes/
https://courses.lumenlearning.com/boundless-chemistry/chapter/namingcompounds/
VIII. REFERENCES:
 General Chemistry 1 with Laboratory Manual (2016), pp. 56 – 74, G. M. Garcia, Unlimited Books
Library Services & Publishing Inc.
Chapter 3:
Molar Mass and Percentage
Composition
Fig. 1. Shows the formula in getting the percentage composition of each molecule and compound.
Qualitative understanding of matter limits our knowledge of the world we live in. We need to have a good
grasp of matter's quantitative concept too. In previous lessons, you have dug deeper into the
submicroscopic world of atoms and molecules, the fundamental building blocks of matter. At this point in
our study, you will learn how matter is measured by a quantity unit called mole.
As well as the principles of stoichiometry are applied wherein, the quantitative relationships between and
among substances are derived from the chemical formulas. After you learned of mole quantity in terms of
particles and mass, you are now ready to do some mole calculations.
You need to be reminded that to change from one unit to another, you use the mole quantity and its
numerical values as an intermediate step. The factor label method is advised to be used because it can
help one to follow changes in the units during calculations. In using the factor label method, you use the
mole conversion factors in converting among mass, moles and number of particles.
II. LESSON OBJECTIVES:
Lesson Standard:
The learners demonstrate understanding of . . .


The mole concept in relation to Avogadro’s number and mass.
The relationship of percent composition and chemical formula
Learning Competencies:


Calculate the empirical formula from the percent composition of a compound
Calculate molecular formula given molar mass
III.
LESSON PROPER:
You must remember that molecules and atoms are extremely small objects both in size and mass. Because
of this, they are invisible to our naked eyes. Counting them individually would be very difficult or even
impossible. However, chemists are able to keep track of atoms and molecules contained in a given sample
of materials. How do they do that? Chemist’s "count" particles of matter by using a standard called mole.
Mole as a Counting Unit
Mole is a counting unit for chemist, the same way that vendors in the market sell eggs by the dozen.
Counting units are used to speed up counting because you count a collection of materials and not
individual or by pieces. A counting units represents a specified number of particles.
Consider the analogy:
1 dozen eggs = 12 pieces of eggs
1 pair of shoes = 2 pieces of shoes
1 pack of cigarettes = 20 pieces of cigarettes
1 ream of bond paper = 500 pieces of bond papers
All the counting units cited are analogous to a mole. If all these counting units represent specific number of
objects, a mole is also equivalent to a specific number. One (1) mole is equal to
602,214,199,000,000,000,000,000 simply 6.02 x 1023. Therefore,
1 mole of a substance = 6.02 x 1023 particles of the substance
The exact number of particles contained in a mole is a very large number. This number, 6.02 x 1023 is often
referred to as Avogadro’s number represented by the symbol N. Thus, another way of referring to a mole is
that:
1 mole = 6.02 x 1023 or Avogadro’s Number
Mole as a Measurement Unit
Mole is one of the seven fundamental units in the International System (SI) of Units. In the SI system, the mole is the
amount of a substance that contains as many elementary entities (atoms, molecules, ions and other particles) as
there are atoms in exactly 12 grams of the carbon -12 isotope.
Mole is used to count representative particles. Representative particle refers to the species or entity
present in substance. The representative particle of an ionic compound is a formula unit. Examples of a
formula unit are NaCl, MgBr2, AI2O3. The representative particle of a covalent compound is a molecule
such as CO2, N2, NH3, SO3, and many others. The representative particle of an element is the atom. One
mole of any substance contains the same number of particles equal to Avogrado’s number; 6.02 x 1023.
Molar Mass
Molar mass of a substance is the mass in grams that is numerically equal to the atomic or formula mass of
the substance.
“Molar mass of an element = atomic mass of the element expressed in grams/mole”
To find the mass of a compound or its molar mass, you need to determine the formula mass or formula
weight of the compound and express the value as grams/mole.
“Molar mass of a compound or molecule = formula mass of the compound or molecule expressed
in grams/mole”
The formula of a compound is a shorthand notation that tells you the number of atoms of each element in a
representative particle of a compound. For instance, the formula of sodium chloride is NaCl. It is composed
of one atom of sodium (Na) and one atom of chlorine (CI). Similarly, formula of aluminum oxide, Al₂O, has
2 atoms of Al and 3 atoms of O.
Therefore, to determine the formula mass or formula weight of the compound, you get the total atomic
masses of all the atoms making up the formula unit. Refer to the periodic table for the atomic masses of the
elements.
For instance, the formula mass of Al2O3 is calculated this way:
Al = 2 x 27 g/mol = 54 a.m.u
O = 3 x 16 g/mol = 48 a.m.u
102 a.m.u
Since the formula mass is numerically equal to its molar mass, the molar mass of Al2O3 is 102g/mole.
Percentage Composition
Percentage composition can be defined as the proportions of the constituent elements in a compound
expressed as the number of grams of each element per 100 grams of the compound.
Formula: P.C = _____molar mass (element)____
molecular mass of the compound
X 100
Example:
Given the compound K2CrO4, determine its percentage composition.
Given: Atomic masses: K = 39 g
Cr = 52 g
O = 16 g
Molecular mass:(2 x K) + (1 x Cr) + (4 x O)
(2 x 39 g) + (1 x 52 g) + (4 x 16 g)
78 g + 52 g + 64 g
= 194 g/mol of K2CrO4
Required: P.CK = ? ; P.CCr = ? ; P.CO = ?
Equation: P.C = _____molar mass (element)____
molecular mass of the compound
X 100
78 𝑔/𝑚𝑜𝑙
Solution and Answer: P.C. for K = 194 𝑔/𝑚𝑜𝑙
X 100 = 40.2%
52 𝑔/𝑚𝑜𝑙
P.C. for Cr = 194 𝑔/𝑚𝑜𝑙 X 100 = 26.8%
64 𝑔/𝑚𝑜𝑙
P.C for O = 194 𝑔/𝑚𝑜𝑙
X 100 = 33.0%
Total = 100%
Determining Empirical Formula using Percentage Composition
If given the percentage composition of a compound, one can compute for its Molecular Formula.
Step 1. Determine the mass of each of the elements in 100 g of the substance.
Step 2. Determine how many moles there are of each element in 100 g of the substance by using the molar
mass of each element.
Step 3. Divide each one by the smallest number of moles
Step 4. Write the empirical formula based on the resulted number of mole in each element of the
compound.
Example:

Empirical formula is the formula containing the smallest whole number ratio of the atoms present
in a molecule. This is the simplest formula representing a compound.

Molecular formula is the formula that contains the actual number of atoms of each element present in one
molecule of the compound.
Determining Molecular Formula using Molar Masses
Sample Problem:
An oxide of Nitrogen gave the following analyses: 3.04 g of Nitrogen, 16.95 g of Oxygen. The molecular
mass of the compound was experimentally found to be 91.0 amu. Determine its molecular formula.
Given: N = 3.04 g
O = 16.95 g
Molecular mass (compound) = 91.0 amu
Required: Molecular formula = ?
Equation: Empirical formula x Multiplier
Solution:
Step 1. Determine the atomic mass of each element in a compound.
Atomic Wt. of N = 14 g/mol
Atomic Wt. of O = 16 g/mol
Step 2. Compute for the number of moles of each element by dividing its molar mass to the given mass in
the problem.
𝐺𝑖𝑣𝑒𝑛 𝑚𝑎𝑠𝑠 (𝑁)
3.04 𝑔
𝐺𝑖𝑣𝑒𝑛 𝑚𝑎𝑠𝑠 (𝑂)
16.95 𝑔
#. of moles (N) = 𝑀𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 (𝑁) = 14 𝑔/𝑚𝑜𝑙 = 0.22
#. of moles (O) = 𝑀𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 (𝑂) = 16 𝑔/𝑚𝑜𝑙 = 1.05
Step 3. Divide each number of moles to the lowest number of moles calculated from step 2.
#. of moles (N) = 0.22 / 0.22 = 1
# of moles (O) = 1.05 / 0.22 = 5
Step 4. Write the empirical formula of the compound and calculate its mass.
NO5 = M.M
= (N x 1) + (O x 5)
= (14 g/mol x 1) + (16 g/mol x 5)
= 14 g/mol + 80 g/mol
= 94 g/mol
Step 5. Determine the multiplier by dividing the empirical formula mass to the given molecular mass.
𝐺𝑖𝑣𝑒𝑛 𝑀𝑜𝑙𝑒𝑐𝑢𝑙𝑎𝑟 𝑚𝑎𝑠𝑠
Determine Multiplier  Multiplier = 𝐸𝑚𝑝𝑖𝑟𝑖𝑐𝑎𝑙 𝑓𝑜𝑟𝑚𝑢𝑙𝑎 𝑚𝑎𝑠𝑠 =
91.0 𝑎𝑚𝑢
94 𝑎𝑚𝑢
= 0.96 or 1
Step 6. Multiply each subscript in the empirical formula by the answer in step 5.
Molecular formula
= Empirical formula x Multiplier
= (NO5) x 1
= NO5
IV. ACTIVITIES/EXERCISE:
Direction: Calculate the following problems about percentage composition and molar masses using the
G.R.E.S.A Method. (5 pts/each)
1. Analysis of a sample of a pure compound reveals that it contains 50.1% of Sulfur and 49.9%
Oxygen by weight. What is the simplest formula to represent the compound?
2. A compound has a percent composition of 58.8% C, 9.8% H, and 31.4% O. If its molecular weight
is 102 g/mol, what is its molecular formula?
V. ASSESSMENT:
Direction: Choose the letter of the correct answer by changing its font color into “RED”,
1. The amount of substance that contains the same number of representative particles as there are
atoms in exactly 12 g of carbon-12 isotope.
a. Mole
C. Molality
b. Molarity
D. Normality
2. Avogadro’s number is used to represent the number of representative particles in a single unit of
mole. Which of the following is the correct value of Avogadro’s number?
a. 6.02 x 1022
C. 6.03 x 1022
b. 6.02 x 1023
D. 6.02 x 1022
3. The proportions of the constituent elements in a compound expressed as the number of grams of
each element per 100 grams of the compound.
a. Avogadro’s number
C. Molality
b. Molarity
D. Percentage Composition
4. The formula containing the smallest whole number ratio of the atoms present in a molecule. This is
the simplest formula representing a compound.
a. Chemical formula
C. Empirical formula
b. Molecular formula
D. Mole ratio
5. The formula that contains the actual number of atoms of each element present in one molecule of
the compound.
a. Chemical formula
C. Empirical formula
b. Molecular formula
D. Mole ratio
VII. ADDITIONAL RESOURCES:
The learners could visit the following links for further knowledge for the following lessons.
Percentage Composition and Molar Mass


https://courses.lumenlearning.com/introchem/chapter/percent-composition-of-compounds/
https://www.thoughtco.com/mass-percent-composition-example-609567
VIII. REFERENCES:
 General Chemistry 1 with Laboratory Manual (2016), pp. 86 – 92, G. M. Garcia, Unlimited Books
Library Services & Publishing Inc.
 Science @ Work 9 (2014), pp. 194-196, M.G. Geroy et. al., Neo Asia Publishing Inc.