MATTER
Is governed by
Is made up of
atoms
Can react to form
ions
-
Law of conservation of matter
Law of definite proportions
Law of multiple proportions
Are composed of
molecules
-
Electrons
Protons
neutrons
Dalton’s atomic theory
Can either be
constitute
Covalent compounds
anions
cation
Can combine to form
Ionic compounds
I.
DALTON’S ATOMIC THEORY
- John Dalton proposed an atomic theory of matter that can explain chemical observation as
predicted by the fundamental laws:
1. Law of Conservation states that matter is not created nor destroyed.
2. Law of definite proportion states that different samples of the same compound always contain
the constituent elements in the same proportion by mass.
3. Law of multiple proportion states that if two elements form more than one compound, the various
masses of one element combining with a fixed mass of another element are related by small wholenumber ratios.
1.
2.
3.
4.
ATOMIC THEORY
Atoms are the smallest particles of matter. They cannot be divided into smaller particles, they also
cannot be created nor destroyed.
All atoms of an element are identical, but the atoms of one element are different from atoms of other
elements in terms of mass, size, and properties.
Compounds are composed of atoms of more than one element which are combined in fixed ratios.
Atoms retain their identity during chemical reactions, which involve combination, separation, and
rearrangement. They are indestructible.
II.
ATOM
- It is the smallest unit of matter.
- It has subatomic particles, namely proton, neutron, and electron.
- The mass of an atom contains protons and neutrons which are densely concentrated at
the nucleus and the electrons are orbiting around it.
o The sub atomic particles a. Proton
- It is the first subatomic particle discovered by Eugene Goldstein.
- It has positive electrical charge found in the nucleus.
b. Electron
- It is the second subatomic particle discovered by James Thompson.
- It has negative electrical charge found outside the nucleus.
c. Neutron
- It is the third subatomic particle discovered by James Chadwick.
- It has neutral charge found in the nucleus with proton.
The number of neutrons in neutral atom is derived by the mass number minus the atomic number or the
number of proton/ electron.
III.
ATOMIC NUMBER AND MASS NUMBER
•
Atomic number – refers to the number of protons in its nucleus
Atomic number=number of protons=number of electron
•
Mass number – refers to the sum of the numbers of protons and neutrons in its
nucleus.
Mass number=number of protons + number of neutrons Number of
neutron = Mass number – number of protons
•
Isotopes – this are atoms with the same
atomic number but different mass number.
Example:
Atomic
number
Mass number
1
2𝐻
Symbol of the element
Examples on determining the atomic#, Mass #, # of proton, # of electron, # of neutron.
Name
Symbol
Mass#
Atomic #
# of proton
# of
# of
electron
neutron
1𝐻
Protium
1
1
1
1
0
1
Deuterium
2𝐻
1
2
1
1
1
1
Tritium
3𝐻
1
3
1
1
1
2
Oxygen 16
16𝑂 8
16
8
8
8
8
Oxygen 17
Oxygen 18
17𝑂 8
17
18
8
8
8
8
8
8
9
10
18𝑂 8
Isotope samples and their uses
Isotope/s
Uses
Oxygen – 16, Oxygen – 17, Oxygen –
18
Forensics
Determining the origin of a rock or an asteroid
Tritium (hydrogen isotope)
Making glow-in-the-dark objects (such as clockfaces and
wristwatches)
Carbon – 14
Determining the age of organisms
Carbon – 11
Positron emission
Uranium – 238 and Potassium – 40
Determining the age of very old rocks
Cesium – 137
Cancer treatment
Krypton – 85
Fluorescent lamps and flash lamps in high-speed photography.
Measuring the thickness of plastic, rubber, paper, and other
materials.
Iodine – 131
Direct radioisotope therapy to treat hyperthyroidism
Iodine – 123
Diagnostic imaging
Radiation sources in radiation therapy
Monitoring the function of the thyroid gland
Sodium – 24
Tracing gas leaks
Tracing oil leaks from oil pipes
Radiotracer in biological research
Studies of body electrolytes
Cobalt – 60
Inspecting materials to reveal internal structure, flaws, or foreign
objects (in place of x-ray)
Phosphorus – 32
Blood volume determination
Iodine – 131
Measuring of thyroid activity and treatment of thyroid disorders
Gadolinium – 153
Measuring density of bones
Iridium – 192
Industrial tracer
IV.
MOLECULE AND IONS
A. Molecule – refers to the combination of at least two atoms in a definite proportion bound together by
covalent bond.
1. Diatomic molecule – molecules that contains only two atoms. Ex:H2 – it is made
up of two hydrogen atom.
CO – consist of one carbon atom and one
Hydrogen atom
2. Polyatomic molecule – a molecule that contains two
or more atoms.
Ex:H2O – two hydrogen atoms and one oxygen
atom. CO2 – two oxygen atoms and one carbon
atom.
B.
Ion – this refers to atoms that gains positive or negative
charge.
If an atom loses electron, it gains positive (+) charge.
(Cations) If an atom gains electron, it gains negative (-)
charge. (Anions)
Ex: table salt (NaCl) Ionic bond
Na
Cl
Transfer of electron
Na+
Cl-
gained electron
lost one electron
Examples of Atoms and Ions
Na
Atom
Na+
Ion
Mg
Atom
Mg+2
Ion
Cl
Atom
Number of protons
11
11
12
12
Number of electrons
11
10
12
Net charge
0
+1
0
Cl-1
O-2
Ion
O
Atom
17
17
8
8
10
17
18
8
10
+2
0
-1
0
-2
1. Monoatomic Ion – ions that contains only one atom.
Ex: Na+, Mg+2, Cl-, O-2
2. Polyatomic ion – ions that consist of more than one atom.
Ex: OH-, NH4+
V.
CHEMICAL FORMULA
- A chemical formula is written chemical structure of a compound.
1. Molecular formula – indicates the actual number of each element in a compound.
2. Empirical formula – it is the simplest chemical formula.
It shows the relative ratio between the number of atoms of different elements in the
compound.
To further differentiate chemical formula from empirical formula, consider the examples below.
Compound
Chemical formula
Empirical formula
Naphthalene
C10H8
C5H4
Hydrogen Peroxide
H2O2
HO
Benzene
C6H6
CH
Glucose
C6H12O6
CH2O
Most compounds have the same molecular formula and empirical formula. For example, the molecular
and empirical formula of water is H2O; ammonia is NH3; methane is CH4; ethanol is C2H6O; and carbon
tetrachloride is CCl4.
VI.
STRUCTURAL AND MOLECULAR MODELS
- The structural formula shows how atoms are bonded to one another in a molecule.
- A molecular model represents the structural formula using artistic methods like ballandstick model and space filling model.
Molecular and structural formulas and molecular models of some common molecules.
Molecule
Molecular
Structural formula Ball-and-stick model
Space-filling model
formula
Water
H2O
H
O
H
Ammonia
NH3
H
N
H
H
Methane
CH4
H
H
C
H
H
Hydrogen
peroxide
H2O2
Carbon
Tetrachloride
CCl4
H–O–O–H
Cl
Cl
Cl
C
Cl
Oxytocin is a hormone secreted by the posterior
lobe of the pituitary gland, a pea-sized structure
at the base of the brain. It's sometimes known as
the "cuddle hormone" or the "love hormone,"
because it is released when people snuggle up or
bond socially.
This is the structural formula for Oxytocin.
Chemical formula C43H66N12O12S12.
VII.
NAMING OF COMPOUNDS A. Ionic Compounds
Ionic Compounds are made of cations and anions. Hence in naming ionic compounds, you should be
familiar with the name of cations and anions and take note of the following rule.
1. For binary compounds, metal cations take their names from the elements while the anions take
the first part of the name of the element, and add the suffix – ide at the end.
Cation
Anion
Compound
Name of the Compound
Na+
O-2
Na2O
Sodium oxide
Mg+2
N-3
Mg3N2
Magnesium nitride
Al+3
O-2
Al2O3
Aluminum oxide
2. For ternary compounds which contain three elements, the cation goes first in its name before the
polyatomic ion which usually ends with – ite or -ate.
Cation
Polyatomic
Anion
Compound
Name of the Compound
Na+
NO3-1
NaNO3
Sodium nitrate
Na+
NO2-1
NaNO2
Sodium nitrite
Mg+2
PO4-3
Mg3(PO4)2
Magnesium phosphate
Mg+2
PO3-3
Mg3(PO3)2
Magnesium phosphite
Ca+2
CO3-2
CaCO3
Calcium carbonate
Ca+2
BrO4-1
Ca(BrO4)2
Calcium perbromate
NH4+
C2O4-2
(NH4)2C2O4
Ammonium oxalate
3. For compounds containing a metallic ion of variable charge, either the classical method or the
stock method of naming may be used. In the classical method, the name of the metallic ion ends in
-ous (for lower charge) and -ic (for higher charge). In the stock method, the metal is named first
followed by the value of the charge written in Roman numeral (enclosed in parenthesis).
Cation
Anion
Compounds
Classical Name
Stock Name
Fe+2
Cl-1
FeCl2
Ferrous chloride
Iron (II) chloride
Fe+2
Cl-1
FeCl3
Ferric chloride
Iron (III) chloride
Sn+2
OH1
Sn(OH)2
Stannous hydroxide
Tin (II) hydroxide
Sn+4
OH-1
Sn(OH)4
Stannic hydroxide
Tin (IV) hydroxide
Cu+1
SO4-1
Cu2SO4
Cuprous sulfate
Copper (I) sulfate
Cu+2
SO3-2
CuSO3
Cupric sulfite
Copper (II) sulfite
B. Molecular Compounds
1. For one pair of elements that form several different compounds, Greek prefixes are used to determine
the number of each element in the compound. For the first element, the prefix “mono” is omitted.
1 – mono
6 – hexa
2 – di
7 – hepta
3 – tri
8 – octa
4 – tetra
9 – nona
5 – penta
10 –deca
Examples:
CO – carbon monoxide
NO2 – nitrogen dioxide
CO2 – carbon dioxide
N2O4 – dinitrogen tetroxide
2. For binary compounds, place the name of the first element; then, follow it with the second element.
The second element is named by adding -ide to the root of the element name.
Examples:
HCI – hydrogen chloride
HBr – hydrogen bromide
HI – hydrogen iodide
SiC – silicon carbide
3. For binary compounds considered as acids, use the prefix hydro- followed by the stem name of
anion ending -ic, then by the word “acid”.
Examples:
HCI – hydrochloric acid
HBr – hydrobromic acid
HI – hydroiodic acid
HF – hydrofluoric acid
4. Oxy-acids, those that contain hydrogen, oxygen, and another element, is named in two ways:
a. for anions ending with -ate, change -ate to -ic; then, follow it with the word “acid”.
b. for anions ending with -ite, change – ite to -ous; then, follow it with the word “acid”.
Examples:
Oxy – acids
Anion used
H3PO4
Phosphoric acid
PO4-3
Phosphate
H3PO3
Phosphorus acid
PO3-3
Phosphite
HNO3
Nitric acid
NO3-1
Nitrate
HNO2
Nitrous acid
NO2-1
Nitrite
H2C2O4
Oxalic acid
C2O4-2
Oxalate
HC2H3O2
Acetic acid
C2H3O2-1
acetate
WHAT IS CHEMISTRY?

Chemistry is the study of matter, its chemical and physical properties, the chemical and physical
changes it undergoes, and the energy changes that accompany those processes.
BRANCHES OF CHEMISTRY
1. ANALYTICAL CHEMISTRY
Analytical chemistry is the study involving how we analyze the chemical components of samples. For
example; how much caffeine is really in a cup of coffee? Are there drugs found in athlete’s urine samples?
What is the pH level of my swimming pool? Examples of areas using analytical chemistry include forensic
science, environmental science, and drug testing.
2. BIOCHEMISTRY
Biochemistry is the study of life or more aptly put, of chemical processes in living organisms. Biochemists
research includes cancer and stem cell biology, infectious disease as well as membrane and structural
biology and spans molecular biology, genetics, mechanistic biochemistry, genomics, evolution and
systems biology.
3. INORGANIC CHEMISTRY
Inorganic chemistry is the field that focuses on the study of elements and compounds other than carbon
or hydrocarbons. Simply put, inorganic chemistry covers all materials that are not organic and are termed
as non-living substances – those compounds that do not contain a carbon hydrogen (C-H) bond.
Compounds studied by inorganic chemists include crystal structures, minerals, metals, catalysts, and most
elements on the periodic table. An example is the strength of a power beam used to carry a specific
weight or investigating how gold is formed in the earth.
4. ORGANIC CHEMISTRY
The study of carbon compounds such as fuels, plastics, food additives, and drugs. An opposite of inorganic
chemistry that focuses on non-living matter and non-carbon-based substances, organic chemistry deals
with the study of carbon and the chemicals in living organisms. An example is the process of
photosynthesis in a leaf because there is a change in the chemical composition of the living plant.
5. PHYSICAL CHEMISTRY
The study of the physical properties of molecules, and their relation to the ways in which molecules and
atoms are put together. Physical chemistry deals with the principles and methodologies of both chemistry
and physics and is the study of how chemical structure impacts physical properties of a substance. An
example is baking brownies, as you’re mixing materials and using heat and energy to get the final product.
MATTER: IT`S PROPERTIES AND MEASUREMENT
Matter is anything that has mass and occupies space. The changes that matter undergoes always involve
either gain or loss or energy.
A. PROPERTIES OF MATTER (Properties are characteristics of matter that enable us to identify them.)
1. PHYSICAL PROPERTIES are those properties that can be observed or measured without changing the
fixed composition of the substance
Examples:
Density- amount of matter in a given volume of material
Melting point- the temperature at which a solid turn into a liquid
Boiling point- temperature at which a liquid change into gas
Thermal and electrical conductivity
Malleability- can be pounded and shaped into very thin sheets without breaking. Examples: platinum,
aluminum, and tin
Ductility- can be stretched into wires or threads; changes shape when stretched. Examples: silver, gold,
platinum, jewelries, and tire wire.
Solubility
Specific gravity
Color, taste, odor, hardness
2. CHEMICAL PROPERTIES of matter are those associated with a change in the fixed composition of a
substance. Examples are rusting metal, burning fuel, milk turning sour, etc.
3. PHYSIOLOGICAL PROPERTIES refer to the action or effect of substances on the body. For example;
sodium bromide is a sedative because it soothes the nerves and induces sleep.
2. CLASSIFICATION OF MATTER
A. PURE SUBSTANCE
- its percentage composition is always the same regardless of its amount and origin.
- its component particles are identical
1. ELEMENTS- a pure substance whose atoms have the same atomic number; each has unique
characteristics
a. Natural elements- occur in nature (ex. Au, Sn, H, O)
b. Synthetic elements- artificially produced in the laboratory (ex. Astatine, Plutonium, etc.)
2. COMPOUNDS- pure substances that are made of combinations of elements
GROUPS OF COMPOUNDS
1.
2.
3.
4.
5.
Organic compounds- contain carbon
Inorganic compounds- all other compounds that are not considered organic
Acids- tastes sour, corrosive
Bases- bitter taste and slippery feel, corrosive like acid
Salt-acid plus base; formed by metal and nonmetal
B. MIXTURES



Combination of pure substances
Components are combined physically
The proportion of components may vary
1. HOMOGENOUS MIXTURES- has only one distinct phase
a. Solution- component particles are too small to be seen. Solute particles are evenly spread throughout
the solvent. For example, sugar solution, salt solution, and others.
· Solvent-one that dissolves (Maybe solid, liquid, or gas in a greater amount. A common solvent is water
which is known as the universal solvent.)
· Solute- one that is dissolved (may be solid, liquid, or gas in a lesser amount)
· Concentration- the proportion of the components in a solution
b. Colloids- have component particles that are bigger than those in a solution but they are still too small
to be seen with an ordinary microscope. Examples are milk, mayonnaise, toothpaste, fog, and others.
2. HETEROGENOUS MIXTURE- component particles can be clearly distinguished from each other and are
large enough to be seen with a microscope.
a. Suspension- a solid in a liquid mixture (ex. Muddy water, a mixture of the dust of chalk and water)
3. THE SCIENTIFIC METHOD
Chemistry deals with chemical situations that are problematic- hence good chemists must be good
problem solvers.
Scientists search for answers to questions and solutions to problems by using a procedure called
the scientific method. This procedure consists of making observations, formulating hypotheses, and
designing experiments; which leads to additional observations, hypotheses, and experiments in repeated
cycles
4. MEASUREMENT OF MATTER
Taking measurements is vital to scientific observations, particularly to physical science like Chemistry and
Physics. It will be useful to remember that a quantitative observation consists of two parts- a number and
a unit.
Table 1. Fundamental Units of Measurement (SI)
Physical quantity
length
mass (not weight)
time
temperature
amount of substance
electric current
luminosity
Unit used
meter
kilogram
second
Kelvin
mole
ampere
candela
Abbreviation
m
kg
s
K
mol
A
Cd
UNITS OF MEASUREMENT IN CHEMISTRY
1. MASS AND WEIGHT- Mass is the measure of the quantity of a body contains. On the other hand, the
weight of a body is a measure of the gravitational attraction of the earth (or any heavenly body) for the
body. The basic unit of mass in the SI system is kilogram and gram for the metric system.
2. LENGTH- the meter is the standard unit of length in both SI and metric systems. The meter is defined
as the distance light travels in a vacuum in 1/299,792,468 seconds.
3. VOLUME- volumes are often measured in liters or milliliters in the metric system.
Conversion of Units
Sample Problems: (cancel like terms)
1. A length of glass tubing is 0.525 m. How many inches long is the tubing?
Given: length- 0.525 m x = inches
Solution: 0.525 m x 100 cm x 1 inch = 20.66 inch
1m
2.54 cm
2. Mark woke up late for his first class. He rushed to get ready and rode his bicycle. The speed of his
bicycle is 90 km/hr, what is its speed in m/sec?
Given: speed= 90 km/hr x= speed in m/sec
Solution: 90 km x 1000 m x 1 hr x 1 min = 90,000 m = 25 m/sec
hr
1 km
60 min 60 sec
3600 sec
THE PERIODIC TABLE
ATOMIC THEORY
THE BEGINNING





Thales of Miletus said that all things come from water.
Leucippus conceived the idea of indivisible units called atoms.
Democritus believed that matter is consisted of tiny particles called atoms and that the infinite
variety of observable things could be explained by the combinations of different sizes and shapes
of these particles.
Empedocles proposed that matter is made up of four elements – earth, water, air and fire.
Aristotle believed that there is no limit to subdividing matter.
JOHN DALTON’S ATOMIC THEORY

-
John Dalton – one of the fathers of modern physical science
Pictured an atom as a tiny indestructible sphere with mass


Atom – smallest particle of an element which can enter into a chemical combination
Postulates of the theory
1. Matter is made up of extremely small indestructible particles called atoms.
2. All atoms of a given element are alike.
3. Atoms enter into a combination with other atoms to form compound but remained unchanged
during ordinary reactions.
4. Atoms can combine in simple numerical ratios such as 1:1, 1:2 2:3 and so on.
THOMSON’S ATOMIC MODEL
-
By Joseph John Thomson
The atom was composed of a positively charged sphere in which the electrons are loosely
embedded on the surface
NUCLEAR MODEL OF AN ATOM
-
By Ernest Rutherford
-
Electrons in the atom revolve around the nucleus

Nucleus – tiny, positive, central core where the mass of the atom is concentrated
BOHR’S ATOMIC MODEL
-
By Neils Bohr
-
The electrons may exist in any allowed circular orbit around the nucleus

This circular orbit was referred to as the principal quantum numbers specified by whole number
integer n with values 1, 2, 3, and so on.
BOHR-SOMMERFELD MODEL OF AN ATOM
lines
Arnold Sommerfeld introduced the concept of elliptical orbits to explain the splitting of the spectral

This elliptical orbit was introduced as the second quantum number for the sublevel within an orbit.
ATOMIC STRUCTURE
Atoms are the smallest particles of matter. They are extremely small, electrically neutral particles that have
tiny massive core or nucleus and one or more electrons relatively far outside the
nucleus
STRUCTURE OF THE NUCLEUS



Proton – positively charged fundamental particle found in the nucleus
The atomic number Z of the element is the number of protons in the nucleus. Since an atom is
electrically neutral, the number of protons must be equal to the number of electrons.
Neutron – has a mass roughly equal to that of proton but has no electric charge. Discovered by
James Chadwick in 1932.
FUNDAMENTAL PARTICLES OF AN ATOM
Name of particle
Electron
Proton
Neutron

Symbol
ep+
n0
Mass (g)
1.1096 x 10-28
1.6726 x 10-24
1.6749 x 10-24
Mass (amu)
5.4859 x 10-4
1.0073
1.0087
Charge
-1
+1
0
ISOTOPES – Atoms of the same element with different masses. They differ in the number of
neutrons but this does not change the identity of the element.
ARRANGEMENT OF ELECTRONS IN ATOMS
Electrons of atoms are arranged around their nuclei in positions known as energy level. Outer electrons
are said to be in higher energy levels and inner electrons are in lower energy levels. As atoms become
more complicated, energy levels become filled. Energy levels proceeding outward from the nucleus are
numbered 1, 2, 3 and 4.
Number of energy level
1
2
3
4
ENERGY SUBLEVELS
Number of electrons
2
8
18
32
Within the main energy levels, there are energy sublevels. The energy sublevels were given names
suggested by the appearance or position of 18 lines in the spectra of excited elements.
Energy sublevels
s
p
d
f
sharp
principal
diffuse
fundamental
Number of electrons
2
6
10
14
Within a given energy level, electrons in an s sublevel are closer to the nucleus than the p sublevel,
those in the p sublevel are closer to the nucleus than the d sublevel and so on.
The energy sublevels of atoms are divided further into regions of space called orbitals. Each orbital
can accommodate a maximum of two electrons.
Energy sublevel
s
p
d
f
Number of orbitals
1
3
5
7
ORDER OF FILLING UP ENERGY SUBLEVELS
The electronic configuration describes the manner in which energy sublevels are filled with electrons.
The filling of energy sublevels in the order of increasing energy is called Aufbau principle (building up).
ORDER OF FILLING ORBITALS
The order of filling orbitals follow the Hund’s rule of multiplicity. This rule states that when electrons
enter a sublevel containing more than one orbital they will spread out over the available orbitals with their
spins in the same direction before they pair up with the opposite spins. Electrons occupy orbitals singly
before pairing. Proposed by Friedrich Hund.
QUANTUM NUMBERS
Schroedinger in 1926 developed an equation that related the wave properties associated with electrons to
their energies, the wave function describes the region within which an electron is most likely to be found,
the orbital. The energy states that an electron can occupy in an atom can be described by 4 quantum
numbers:
n is the principal quantum number – gives the size and energy of an orbital. n = 1, 2, 3, 4
l is the azimuthal or subsidiary quantum number – tells what kind of energy sublevel the electron is in
and describes the shape of the orbital. The allowed values of l depend upon the value of n and ranges
from 0 to n – 1.
n
l
Energy sublevel
1
0
1s
2
0
2s
2
1
2p
3
0
3s
3
1
3p
3
2
3d
4
0
4s
4
1
4p
4
2
4d
4
3
4f
m is the magnetic quantum number – denotes the particular orientation of the orbital which the
electron is occupying within an energy sublevel. The allowed values of m are –l to +l.
Energy sublevel
s
p
d
f
l
0
1
2
3
m
0
+1, 0 -1
+2, +1, 0, -1, -2
+3, +2, +1, 0, -1, -2, -3
s is the spin quantum number – describes the two ways in which an electron may be aligned in a
magnetic field.
s = + ½ unpaired electron
s = - ½ paired electron
Example: B = 5;
1s2
2s2
n
1
1
2
2
2

2p1
l
0
0
0
0
1
m
0
0
0
0
+1
s
+1/2
-1/2
+1/2
-1/2
+1/2
The Pauli Exclusion Principle
States that no two electrons within the same atom have the same quantum number. If two electrons
occupy the same orbital they must have different values on s.
THE PERIODIC TABLE
HISTORY OF THE PERIODIC TABLE

1871 – Johann W. Dobereiner
Triads of elements e.g. Cl, I, Br and Li, Na, K

-
1864 – John Newlands
Elements are arranged in the order of increasing atomic masses (law of octaves)

1869 – Lothar Meyer
-
Devised a classification of the elements that accounted for the periodic variations in properties
-
Based primarily on physical properties of elements
-
Dmitri Mendeleev – arranged the elements in the order of increasing atomic weights
-
Based primarily on chemical properties of the elements

-
1914 – Henry G. Moseley
Arranged the elements according to increasing atomic numbers
PERIODS AND GROUPS

PERIOD or SERIES – single horizontal row in the periodic table

GROUP or FAMILY – a vertical column of elements that have similar physical and chemical
properties
CLASSIFICATION OF GROUPS
1. Group 1A or 1 – alkali metals
Soft shiny metals with high ductility, relatively low melting point and are excellent conductors of
heat and electricity
-
React vigorously with water, releasing H and forming strong caustic solutions
1.
2.
3.
4.
5.
6.
-
Group 2A or 2 – alkaline earth metals
Group 3A or 13 – Boron family
Group 4A or 14 – Carbon family
Group 5A or 15 – Nitrogen family
Group 6A or 16 – Oxygen family
Group 7A or 17 – Halogen family
Exists as combinations of two atoms forming diatomic molecules
1. Group 8A or 18 – Noble gases
Are quite unreactive elements, are seldom found in combination with other elements and have very
stable configuration
CLASSIFICATION OF ELEMENTS
1. METALS – high electrical conductivity that decreases with increasing temperature
-
High thermal conductivity
-
Metallic gray or silver luster
-
Almost all are solid
-
Malleable and ductile
-
Outer shells contain few electrons (usually 3 or fewer)
-
Form cations by losing electrons
-
Form ionic compounds with nonmetals
-
Solid state characterized by metallic bonding
2. NONMETALS – poor electrical conductivity (except carbon in the form of graphite)
-
Good heat insulators (except carbon in the form of diamond)
-
No metallic luster
-
Solids, liquids, gases
-
Brittle in solid state
-
Nonductile and not malleable
-
Outer shells contain four or more electrons
-
Form anions by gaining electrons
-
Form ionic compounds with metals and covalent compounds with nonmetals
-
Covalently bonded molecules; noble gases are monoatomic
3. METALLOIDS – elements adjacent to the ladder-like line
-
They are metallic (or nonmetallic) only to a limited degree
PERIODIC LAW: The properties of the elements are periodic functions of their atomic numbers.
PERIODIC PROPERTIES
METALLIC PROPERTY
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
Metallic character increases from top to bottom and decreases from left to right with respect to
position in the periodic table.
Nonmetallic character decreases from top to bottom and increases from left to right in the periodic
table.
ATOMIC RADIUS – determined by measuring the distance between the nuclei of the bonded atoms
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


Interatomic bonded distance – distance between the nuclei of two atoms bonded together
Angstroms – unit that expresses atomic radii (1 angstroms = 10-10 m, 1 angstrom = 0.1nm)
Going across any period in the periodic table, the atomic radius decreases due to increasing
nuclear charge.
Within a group, atomic radius increases from top to bottom as electrons are added to shells further
from the nucleus.
IONIZATION ENERGY – measure how tightly electrons are bound; maybe expressed in electron volt
per atom(eV/atom), kilocalories per mole ( kcal/mol) or kilojoules per mole (kJ/mol)

First ionization energy, IE1 – also called first ionization potential
Minimum amount of energy required to move the most loosely bound electron from an isolated
gaseous atom to form an ion with +1 charge

Second ionization energy, IE2 – amount of energy required to remove the second electron from the
gaseous atom.
o The second ionization energy is always higher than the first because it is more difficult to
remove a negatively charged electron from a positively charged ion than from the
corresponding neutral atom.
o Ion – an atom or group of atoms that has a net positive or negative charge. This charge
results from an unequal number of protons and electrons.
o The first ionization energy for the atoms going across a period from left to right generally
increases. This is due to the increasing nuclear charge, although the atoms belong to the
same energy level.
o The first ionization energy decreases from one atom to the next as one goes down a family.
This is due to the fact that valence electrons are in a higher energy level or at a farther
distance from the nucleus.
ELECTRON AFFINITY – the amount of energy released when a gaseous atom gains an electron
The amount of energy absorbed when an electron is added to an isolated gaseous atom to form an
ion with a -1 charge



Anions (negative ions) – formed when atoms acquire an electron or electrons from other atoms
Electron affinity increases within a period from left to right.
As one goes down a group, electron affinity decreases.
o F is an exception to the general rule due to its small size which causes more repulsion
among the valence electrons.
IONIC RADII





Cations – positively charged ions
Isoelectronic – ions/elements with the same number of electrons thus the same electronic
configuration
The greater the positive charge, the smaller is the ion because the electrons are more tightly held.
The greater the negative charge, the larger is the size of the ion.
Some guidelines:
1. Simple positively charged ions (cations) are always smaller than the neutral atom from which they
are formed.
2. Simple negatively charged ions (anions) are always larger than the neutral atom from which they
are formed.
3. The sizes of cations decrease from left to right across a period.
4. The sizes of anions decrease from left to right across a period.
5. Within an isoelectronic series, radii decrease with increasing atomic number because of increasing
nuclear charge.
6. Both cation and anion sizes increase going down a group.
ELECTRONEGATIVITY – ability of an atom in a compound to attract additional electron toward itself. The
greater the electronegativity, the greater the attraction for electrons.


Electronegativity tends to increase from left to right across each period.
Electronegativity tends to decrease from top to bottom in each family.
resources:
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DELA CRUZ, M.P.(2009).General Chemistry.Valenzuela City, Philippines:Mutya publishing house, Inc.
Extracted from BSU chemistry lecture notes
Lecture notes prepared by Marvileen B. Samidan
CHEMICAL COMPOUND
CHEMICAL COMPOUNDS
CHEMICAL COMPOUNDS are substances which are made of more than one atom or element. What holds
these atoms together are strong forces of attraction known as chemical bonds.
VALENCE ELECTRONS – number of electrons in the outermost energy level; responsible for the
combining capacity of an element
LEWIS ELECTRON DOT STRUCTURE (LEDS)
-
Developed by Gilbert Lewis
This structure consists of the chemical symbol of an element surrounded by a number of dots. The
chemical symbol represents the nucleus of the atom and the inner electrons. The dots represent the
valence electrons.

For the Group A elements, the number of valence electrons (dots in LEDS) for the neutral atom is
equal to the group number except H and He.
OCTET RULE – refers to the rare gas configuration where there are exactly eight electrons in the valence
shell
IONIC BOND
-
Characterized by the transfer of electrons from one atom to another
It results from the electrostatic attraction between two oppositely charged ions, the cation and the
anion
Ionic bonds are normally formed between the elements in group 1A, 2A or 3A and the elements in
group 6A or 7A
-
Examples:
1. Na and Cl
1. Mg and F
2. Al and O
3. Li and N
PROPERTIES OF IONIC COMPOUNDS
1.
2.
3.
4.
5.
6.
They are solids with melting points (typically > 400 degrees Celsius).
Many are soluble in polar solvents such as water.
Most are insoluble in nonpolar solvents, such as hexane, C6H14, and carbon tetrachloride, CCl4.
Molten compounds conduct electricity well because they contain mobile charged particles (ions).
Aqueous solutions conduct electricity well because they contain mobile charged particles.
They are often formed between two elements with quite different electronegativities, usually a
metal and a nonmetal.
COVALENT BOND
-
Characterized by sharing of electrons in the valence shells
-
Normally formed between H and elements in groups 4A, 5A, 6A and 7A
In general, when a nonmetallic element combines with another nonmetallic element, electrons are
neither gained nor lost by the atoms, but are shared.
-
Represented by dashes between the atoms


Triple bond > double bond > single bond
The stronger the chemical bond, the shorter the bond distance
LEWIS STRUCTURE FOR COVALENT COMPOUNDS
Tips in Writing the LEDS for Covalent Compounds
1. Add up the total valence electrons.
2. The atom with the highest covalency number (number of covalent bonds formed by the atom is
considered as the central atom.
3. Bond the other atoms to the central atom by a single bond. Take into account the number of
electrons used in bonding.
4. Distribute the remaining valence electrons to the attached atoms first and then to the central atom
last.
5. If there is a deficiency in the octet rule form a multiple bond.
watch the video and try to answer the example.
Play Video
Examples: Show the acceptable Lewis structures for:
1. CO2
2. SO2
3. SCl

RESONANCE – a situation where there are more than one probable structures for a species
PROPERTIES OF COVALENT COMPOUNDS
1.
2.
3.
4.
5.
They are gases, liquids or solids with low melting point (typically < 300 degrees Celsius)
Many are insoluble in polar solvents.
Most are soluble in nonpolar solvents.
Liquid and molten compounds do not conduct electricity.
Aqueous solutions are usually poor conductors of electricity because most do not contain charged
particles.
6. They are often formed between two elements with similar electronegativities, usually nonmetals.
CHEMICAL FORMULAS
The formula for a substance shows its chemical composition. By composition, we mean the elements
present as well as the ratio in which atoms occur in the compound.
The formula for a single atom is the same as the symbol for the element. Thus, Na represents a
single sodium atom. A subscript following the symbol of an element indicates the number of atoms of that
element in a molecule. A chemical formula then is a combination of symbols for atoms or ions that are
held together chemically.
OXIDATION NUMBER / OXIDATION STATE
It is the number of electrons an element is capable of losing or acquiring in a chemical reaction to become
chemically stable.
In ionic compound, the oxidation number of an ion is the same as the charge of the ion. The atom that
loses electrons has the positive (+) oxidation number and the atom that gains electrons has the negative () oxidation number.
In covalent compounds, the oxidation number of an atom does not necessarily correspond with the
number of covalent bonds joining the atom to other atoms. Although oxidation numbers may not indicate
the complete loss or gain of electrons, they are essentially a way of counting electrons. The atom that is
more electropositive has the positive oxidation number and the atom that is more electronegative has the
negative oxidation number.
WRITING AND NAMING COMPOUNDS
BINARY COMPOUNDS CONTAINING A METAL AND A NONMETAL
Binary compounds consist of two elements. The binary compounds formed by metals and
nonmetals are usually ionic in nature.

Steps in Writing the Chemical Formula
1. Write each ion with its charge starting with the positive ion.
2. Crisscross the numbers (but not the plus or minus sign) and write them as subscripts.
Notes:
a. a subscript of one is not written anymore
b. subscripts are always reduced to lowest terms
Example: Write the formula of aluminum chloride.
Al+3
Cl-
Al1Cl3
AlCl3
Naming Binary Compounds
1. The metal (with fixed oxidation number) is named first followed by the nonmetal with the ending –
ide.
Example: CaBr2
calcium bromide
1. In naming compounds with metals which has variable oxidation numbers, the oxidation number
must be specified.
1. Classical system – the suffixes –ous and –ic are used to denote the lower and higher
oxidation state respectively
2. Stock system – the oxidation number of the metal is indicated by a Roman numeral in
parenthesis
Examples: 1. Cu+ + O-2
Cu2O
2. Fe+3 + O-2
Cuprous oxide
ferric oxide
Copper (I) oxide
iron (III) oxide
for more information watch the video below:
Play Video
Exercises: 1
Fe2O3
a. Write the formula for the following:
1.
2.
3.
4.
5.
Strontium fluoride
Calcium sulfide
Calcium phosphide
Barium oxide
Sodium oxide
1.
2.
3.
4.
5.
6.
What is the name for the following compounds:
CaBr2
Mg3N2
Al2S3
KCl
Na2O
1. Write the chemical formula and give the name of the compound using the classical and stock
systems formed by each of the following pairs of ions:
2. Fe+2 and Cl3. Sn+4 and F4. Cu+2 and O-2
5. Pb+2 and I6. Hg+2 and BrBINARY COMPOUNDS CONTAINING TWO NONMETALS

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Generally covalent in nature
Naming covalent compounds containing only two elements is similar to naming ionic compounds
except that prefixes are used
Greek prefixes used in naming covalent compounds:
Prefix
Mono
Di
Tri
Tetra

Number
1
2
3
4
Prefix
Penta
Hexa
hepta
Number
5
6
7
The first element named is the one to the left in the periodic table.
Examples:
1. CO – carbon monoxide
2. CO2 – carbon dioxide
for more information watch using this link:
Play Video
Exercises: 2
a. Name the following compounds:
1. NO
Prefix
Octa
Nona
deca
Number
8
9
10
2. SO2
3. CCl4
4. P2O5
5. PCl5
b. Write the chemical formula of the ff:
1. nitrogen dioxide
2. dinitrogen trioxide
3. carbon disulfide
4. sulfur trioxide
5. diphosphoruspentoxide
METALS WITH POLYATOMIC IONS


Polyatomic ion – a stable group of atoms that carries an overall charge
Parentheses are placed around the polyatomic ion and the subscript is written just after the close
parenthesis whenever a multiple of the polyatomic ion is necessary. The parentheses are not used
when a single polyatomic ion is present.
Examples:
a. magnesium hydroxide
b. potassium chlorate
c. calcium carbonate
Exercises: 3
a. Write the correct formula for:
1. strontium bicarbonate
2. ammonium nitrate
3. iron (III) sulfate
4. magnesiumphosphate
5. copper (II) carbonate
b. Give the name of each of the following:
1. K2CrO4
2. Zn(NO3)2
3. CuClO3
4. FeSO4
Mg(OH)2
KClO3
CaCO3
5. Sn(OH)4
D.
HYDROGEN AND A NONMETAL

A binary compound which is composed of hydrogen and a more electronegative element is named
like any other binary compound of nonmetals.
Examples:
a. HCl hydrogen chloride
b. HF hydrogen fluoride
c. HBr hydrogen bromide
d. HI Hydrogen iodide

When these substances are dissolved in water, they become aqueous acids. The prefix hydro- is
attached to the root word of the nonmetal and the suffix –ic is added. The word acid becomes the
last term.
Examples:
a. H2S (aq)
b. HCl (aq)
hydrosulfuric acid
hydrochloric acid
1. ACIDS
The term acid is generally taken to mean a water solution of acidic substances. Acids are named by
modifying the name of the anion to end in –ic.
Examples:
Formula of acid
HBRO3
HIO3
HClO3
HMnO3
H2SO4
H3PO4
Notes:
Name as ionic compound
Hydrogen bromate
Hydrogen iodate
Hydrogen chlorate
Hydrogen manganate
Hydrogen sulfate
Hydrogen phosphate
Anion
Bromate
Iodate
Chlorate
Manganate
Sulfate
phosphate
Common name of acid
Bromic acid
Iodic acid
Chloric acid
Manganic acid
Sulfuric acid
Phosphoric acid
a. the suffix –ic denotes the common acid.
b. The suffix –ous denotes an acid containing one less oxygen atom than the common acid.
c. The prefix hypo- with the suffix –ous denotes an acid with less two oxygen atoms than the common acid.
d. The preficper- with the suffix –ic, denotes an acid containing one more oxygen atom than the common
acid.
e. The prefix hydro- with the suffix –ic denotes an acid with no oxygen.
Examples:
HClO4
perchloric acid
HClO3
chloric acid
HClO2
chlorous acid
HClO
hypochlorous acid
HCl
hydrochloric acid
F. BASES
These are compounds composed of a cation plus the hydroxyl ion, OH-. To name bases, name the cation
first followed by the word hydroxide.
Examples:
G.
NaOH
sodium hydroxide
KOH
potassium hydroxide
Ca(OH)2
calcium hydroxide
Ba(OH)2
barium hydroxide
SALTS
Salts are ionic compounds, therefore they are named by combining the names of the cation followed by
the name of the anion.
Examples:
NaClO4
sodium perchlorate
NaClO3
sodium chlorate
NaClO2
sodium chlorite
NaClO
sodium hypochlorite
Hydrated salts refer to the crystalline form of salts which contain water molecules in a definite ratio. The
water present is known as water of hydration. A dot is used in the formula to separate the salt from the
water of hydration. The complete name of a hydrate must include both the name of the salt and the
number of water of hydration, indicated by a number or a numerical prefix.
Examples:
MgSO4 • 7H2O
magnesium sulfate 7-hydrate
or magnesium sufateheptahydrate
Na2CO3 •10H2O
or sodium carbonate decahydrate
sodium carbonate 10-hydrate