Scope: From Periodicity to Chemical Reactions Subject: General Chemistry

1. CHEM 101 LEC
I.
INTRODUCTION
SCIENCE – “SCIRE” means to know or knowledge
Divided into 2 major groups:
A. Natural – natural objects and phenomena
B. Social – Deal with the study of human behavior
SOCIAL SCIENCES
 Anthropology – man and his culture
 Economics – study of how man allocates his resources
 Sociology – origin and constitution of society
NATURAL – It is subdivided into two groups:
A. PHYSICAL – deal with the predictable behavior around us and it includes:
1. Astronomy – heavenly bodies
2. Chemistry – structure, properties and composition of substances and its changes
3. Geology – earth
4. Math – relationship of quantities
B. BIOLOGICAL – deal with biotic factors
1. Microbiology – microorganisms
2. Botany – plants
3. Chemistry – organic and biochemistry
4. Medicine – detection and prevention and cure of diseases
5. Zoology – animals
CHEMISTRY – study of matter, its properties and the changes in composition, which matter
undergoes.
2 STEPS IN CHEMISTRY
1. Analysis – taking a part to analyse
EX: Determination of potassium content in pechay leaves
2. Synthesis – combining different parts to produce new and better products
EX: Preparation of furniture polish from spent cooking oil
ADVANTAGES OF CHEMISTRY
a) Improvement of health care
b) Conservation of natural resources
c) Protection of the environment
d) Provision of our everyday needs
e) Clothing and shelter
f) Discovered pharmaceutical needs
DISADVANTAGES OF CHEMISTRY
a) Some chemicals have the potential of harming our health on the environment
CLASSIFICATION OF CHEMISTRY
1. Inorganic Chemistry – study of all ELEMENTS and their COMPOUNDS
2. Organic Chemistry – study of carbon.
THERE ARE 116 ELEMENTS.
FIELDS OF CHEMISTRY
1. Biochemistry – living organisms
2. PhysicalChemistry – structure of matter, energy and its changes
3. AnalyticalChemistry – concerns the identification, separation and quantitative
determination
Divided into 2 subgroups: Quantitative & Qualitative
4. IndustrialChemistry - transformation of raw materials into finished product
OTHER FIELDS OF CHEMISTRY
5. Agriculture – utilized in soil analysis, and in the manufacture of fertilizers
6. Industry – new & better products are being produced from raw materials
7. ChemicalEngineering – a combination of chemistry and engineering, that improves industrial
process
8. ColloidChemistry – study of behavior of particles

2. 9. Electrochemistry – chemical reactions that are produced by electric current
10. Nuclearchemistry – radioactivity
IMPORTANCE OF CHEMISTRY: serves as central science.
II.
MATTER
MATTER – Anything that occupies space and has mass.
Divided into TWO SUBGROUPS:
A. PureSubstances – fixed compositions; cannot be purified. Classified into two:
1. Elements – cannot be subdivided by chemical or physical means. It is combine
chemically to form:
2. Compounds: elements that are united in fixed ratios.
B. Mixtures – combination of two or more pure substances. It is subdivided into two:
1. Homogeneous – uniform composition throughout
EX:
2. Heterogeneous – matter that is non-uniform composition
EX:
CHANGES IN MATTER:
a) Chemical Change (Chemical Reaction) - substances that are used up and others formed in
their place.
Ex: When propane burns in air, propane& oxygen are converted into CO2 and H2O
b) Physical Change – matter doesn’t lose its identity. A common physical change is a change of
state.
Ex: Ice turns into water
CLASSIFICATION OF MATTER
1) Element – a substance that consist of identical atoms
 There are 116 known elements
 88 occur in naturel others have been made by chemist & physicist
 Symbols consist of one or two letters
 Names are derived from a variety of sources:
The English Name of the Element; people important in atomic science, geographical
location, planets and mythological sources
2) Compound – elements united in fixed ratios
FORMULA OF COMPOUND: tells us the ratios of its constituent elements & identifies each
element by its atomic symbol
MIXTURE
 Substances may be present in any mass ratio
 Each substances has a different set of physical properties
 Homogeneous or Heterogeneous
 If we now the physical properties of the individual components of the mixture, we can use
appropriate physical means to separate the mixture into its component parts
INTENSIVE AND EXTENSIVE PROPERTIES
1) Intensive/Intrinsic Properties – property of matter that is independent of the quantity of the
substances
EX: Density and Specific Gravity
2) Extensive/ Extrinsic Properties – depends on the quantity of substance
EX: Mass, Volume
Table Salt is a pure substance when not dissolved in water
Diamond is an element
Bronze is a mixture (mixture of copper & tin)
Brass is a mixture (copper & zinc)
III.
THE ATOM

3. ATOM – is the smallest unit of matter that can retain the properties of element
EARLY ATOMIC THEORIES
 5TH CENTURY BC – Democritus of Abdera conceived the existences of “atomos”, which
means indivisible
 4TH CENTURY BC - Plato conceptualized the existent of elements
 3RD CENTURY BC – Aristotle developed the theories made by Plato and stated that “all
substances were made up of matter on which different forms could be impressd”
MODERN ATOMIC THEORIES
 18TH CENTURY – Dalton introduced the modern atomic theory which states that:
1) Element are composed of minute indivisible particles called atoms which
preserve the individuality in all chemical changes
2) All atoms in the same element are identical in all respects, particularly weight
3) Chemical combination between two or more element consist of the union of
atoms in simple numerical ratios to form the smallest unit of compound called
MOLECULE
4) Atoms of the same elements can unite in more than one ratio to form more than
one compound
EVIDENCE OF DALTON’S ATOMIC THEORY
a) Antoine Laurent Lavoisier (1743-1794) – LAW OF CONSERVATION OF MASS,
states that matter can be neither destroyed or created
b) Joseph Proust – French chemist, LAW OF CONSTANT COMPOSITION, states that
any compound is always made up of elements I the same proportion of mass
 1886 – E. Goldstein discovered the existence of proton
 1897 – Sir JJ Thomson discovered that are negatively charged subatomic particles called
electron. Thomson formulated the first atomic model called “Plum Pudding Model”
 1911 – Ernest Rutherford observed that the weight of an atom is concentrated at the center.
From the observation, he suggested the existence of nucleus & formulated the “Nuclear
Model of an Atom”
 1913 – Neil Bohr formulated the “Planetary Model of an Atom”, stated that electrons found
outside the nucleus travel around the nucleus in definite orbits
 1926 – Erwin Schrodinger used quantum theory to come up with modern atomic model
which is QUANTUM MECHANICAL MODEL
 1932 – James Chadwick discovered neutrally charged subatomic particle NEUTRON, which is
also located at the nucleus of atom. # of neutrons is equal to number of protons in the
nucleus.
Milikan De Broglie Heisenberg –
ATOMICNOTATION – is a system of writing an atom showing its atomic number, atomic symbol and
mass number.
MASS NO.
ATOM. NO.
A
Z
X
ATOMIC SYMBOL
ATOMICSYMBOL – short hand notation of an element
Au – Gold (Aurum)
Sn – Tin (Stannum)
H – Hydrogen
ATOMICNUMBER – number of protons present in an atom. Atom is neutral, meaning number of
protons = number of electrons
To get the # OF NEUTRONS: Atomic Mass – Atomic Number
MASS NUMBER – sum of protons & neutrons in an element
Ex: He (Helium) has two protons and two neutrons; therefore the mass number is 4.
To get the MASS NUMBER: # of PROTONS + Neutrons
PROPERTIES AND LOCATION WITHIN ATOMS OF PROTONS, ELECTRONS AND NEUTRONS

4. SUBATOMIC
CHARGE
MASS
MASS
MASS (amu
Location in the
PARTICLE
(g)
(amu)
to 1 S.F)
atom
PROTON
+1
1.6726 x 10-24
1.0073
1
in the nucleus
ELECTRON
-1
9.1094 x 10-28
5.4859 x 10-4
0.0005
outside nucleus
-24
NEUTRON
0
1.6740 x 10
1.0078
1
in the nucleus
The unit of mass is the ATOMIC MASS UNIT (amu). By definition, 1 amu = 1/12 the mass of
carbon atom with 6 protons and 6 neutrons
1 amu = 1.6605 x 10-24
ISOTOPES- in some instances, the number of neutrons of atom may be altered by radioactive
means.
- an atom with the same atomic number but with different atomic mass
EX:
1
2
3
1H – protium
1H – deuterium
1H – tritium
 Oxygen has 3 isotopes, namely 16, 17 and 18. Calculate the number of neutrons of each
isotope. Remember: To get the neutron – Atomic mass minus the Atomic Number
Calculating Atomic Weight &Percent Abundance:
AtomicWeight = (massax %abua) + (massbx %abub)
AverageAt. Mass = (Mass of 1stiso) (x) + (Mass of 2ndIso) (1-x)
Missing Isotopes = __________________________________________________________
ELECTRONICCONFIGURATION – is a designation of flow of electrons are distributed among various
orbitals in principal shell & subshells
In Modern Atomic Theory, electrons travel around the nucleus in a quantized manner.
Fixed amount – Quantum Number
Main Energy level – Principal Quantum Number, orbits & electron shell (n)
n = 1 is the electron shell NEAREST to the nucleus and is the first atomic orbit
n = 2 is the second electron shell/ second atomic orbit
n = 7 is the farthest atomic orbit
ENERGYSUBLEVELS – atomic orbitals or electron subshell. They are the s, p, d and f.
# OF ORBITAL (rep. in each
PRINCIPAL
SUBSHELL
# OF E’S
MAX. OF E’S
subshell
1
s
2
2
s=1
2
s,p
2,6
8
s = 1; p = 3
3
s, p, d
2,6,10
18
s = 1; p = 3; d =5
4
s, p, d, f
2, 6, 10, 14
32
s =1; p=3; d=5; f=7
5th
50
THE AuBau Principle

5. -
In distributing the electrons, the lowest energy levels are the filled to the capacity before
going to the next level.
HUND’S RULE
- Electrons must fill up orbitals of the same energy level before pairing with another
electron
Pauli’s Exclusion Principle
- States that no more than 2 can occupy a single orbital
Orbital – region in space around the nucleus where the probability of finding the electron is greatest
Valence shell & Valence Electron
1) Valence – in atoms, means OUTERMOST
2) Valenceshell – outermost shell of an atom. It is the orbital that has the highest ENERGY
LEVEL
3) ValenceElectrons – outermost shell of an atom, where electrons are found
ELECTRON DOT/ LEWIS STRUCTURE of ATOMS:*
- shows the # of valence electrons of an element by using dots
5 METHODS OF REPRESENTING ELECTRONIC DISTRIBUTION
1) Spectronic /s, p, d, f notation
2) Half shell method
3) Orbital Diagram
4) Electron dot formula/ Lewis dot formula
5) Nobel gas notation
QUANTUM NUMBERS
 Electrons of an element may be described by 4 quantum numbers
1) Principal Quantum Numbers (n) – described as #1-7
2) Secondary/ Azimuthal Quantum Numbers (l) the shape of the orbital. May range
from 0 to n-1. Electrons in an s-orbital have a value of 0.
s=0
REMEMBER: If the value of l is 0, the value
p= 1
of m is 0.
d= 2
f= 3
3) Magnetic Quantum Number (m) – describes the orientation of electrons in the
orbital. Value of m can range from -1, 0 or +1. If n =1 the only possible value of m is
0.
4) Magnetic spin Quantum Number (s) – describe the spin of the electrons. Value of s
can only be +1/2 or -1/2 which describes clockwise &counter clockwise spin
respectively
IV.
PERIODICITY – THE PERIODIC TABLE
PERIODIC CLASSIFICATION OF THE ELEMENTS
 The long form periodic table or periodic chart is a list of all the known elements (116)
arranged in order of increasing atomic number in horizontal rows of such a length that
element which is chemically alike fall directly beneath one another.
 It serves as a guide in predicting the electronic arrangement of atoms. The physical and
chemical behavior of the elements can also be related to their position in the periodic
table.
CLASSIFICATION OF ELEMENTS
NOBLE GASES
 Colorless monatomic gases, which are chemically unreactive and diamagnetic (with the
exception of the helium, which has the configurations of 1s2)
 All the noble gases have outer configurations of ns2np6, a very stable arrangement
 Group 8 or group zero in old, group 18 in new
REPRESENTATIVE ELEMENTS

6.  The term representative element is related to the stepwise addition of electrons to the s and
p sublevels of the atoms.
 Found in A families
 They exhibit a wide range of chemical behavior and physical characteristics
 Some of the elements are diamagnetic and some are paramagnetic
TRANSITION ELEMENTS
 The column 1B t0 8B contain the transition elements or the B family/group.
 Group starts with 3B up to 8B which has 8 columns and then end with 1B and 2B. These
sequences which contain 10 elements each, are related to the stepwise addition of the 10
electrons to the d sublevel of the atoms.
TRANSITION ELEMENTS
 These elements are all metallic, dense, lustrous, good conductor of heat and electricity and
in most cases, hard.
INNER-TRANSITION ELEMENTS
 The additional horizontal rows comprise two groups of elements which were discovered to
have similar characteristics as Lanthanum in the 6th period called Lanthanoids (rare-earth
elements) and Actinium in the 7th period called Actinoids (heavy rare elements). These two
series are called INNER-TRANSITION ELEMENT
DEVELOPMENT OF THE PERIODIC TABLE
1. John Dalton’s atomic theory paved the way to the establishment of the periodic table. Dalton
assigned weights and combining capacities to the atoms that by the early 19th century, the
approximate atomic weights for more than 20 elements have been determined.
2. DOBEREINER’S TRIADS (1829) – Johann WolfgangDobereiner grouped elements, which exhibit
very similar characteristics in three or triads. The atomic weight of the second element was
found to be the average of the first and third elements.
3. NEWLAND’S TABLE ( 1864) – John Newland developed a periodic table in which the elements
were arranged in order of increasing atomic weights.
- He proposed the any given element is similar to the eight elements following it.
- He arranged the elements at intervals of eight, similar to the octave of the musical scale.
4. DMITRI MENDELEEV (1834-1907) - arranged the known elements in order of increasing atomic
weight beginning with hydrogen.
- He observed that when elements are arranged in this manner, certain sets of properties
recur periodically.
- He then arranged elements with recurring sets of properties in the same column
(vertical row); Li, Na, and K, for example, fall in the same column and start new periods
(horizontal rows).
- Periodic Table: Fluorine, chlorine, bromine, and iodine fall in the same column
5. MEYER’S TABLE (1870) – unaware of Mendeleev’s study, Julius Meyer had been working on his
periodic table which consisted of 56 elements. He also maintained that the properties of the
elements were periodic functions of their atomic mass.
6. MOSELEY’S TABLE (1914) – the early periodic tables were arranged according to increasing
atomic weights and this misplaced several elements, such as Ar, and K, Co and Ni, and Te and I,
in the periodic table
PERIODIC LAW - states that the physical and chemical properties of the elements are a periodic
function of their atomic number.
PERIODS, GROUPS AND FAMILIES
The seven horizontal rows in the periodic table are called periods.
 Period 1 and 2 has 2 elements corresponding to 2 electrons in the s sublevel.
 Period 2 and 3 have 8 elements corresponding to 8 electrons in the s and p sublevels.
 Period 4 and 5 have 18 elements corresponding to 18 electrons in the s, p and d sublevels.
 Period 6 has 32 elements corresponding to 32 electrons in the s, p, d, and f sublevels.
 Period 7 is still incomplete but elements fill up s, p, d and f sublevels.

7.  The vertical columns are called group or families, which are divided into A and B subgroups.
Some of the A families are designated by names.
 Group 1A – Alkali Metals
 Group 2A – Alkaline Earth Metals
 Group 7A – Halogens
 Group 8A – Noble Gases
 The other subgroups are classified according to the first element in the column.
 Group 3A – Boron Family
 Group 4A – Carbon Family
 Group 5A – Nitrogen Family
 Group 6A – Oxygen Family
CLASSIFICATION OF ELEMENTS
Metals
 are solids at room temperature (except for Hg, which is a liquid), shiny, conduct electricity,
and are ductile and malleable.
 form alloys (solutions of one metal dissolved in another); brass, for example, is an alloy of
copper and zinc.
 In chemical reactions, they tend to give up electrons.
Nonmetals
 Except for hydrogen (H), they lie on the right side of the Periodic Table.
 Except for graphite, do not conduct electricity.
 In chemical reactions, they tend to accept electrons.
Metalloids
 They have some of the properties of metals and some of nonmetals; for example, they are
shiny like metals but do not conduct electricity.
 Six elements are classified as metalloids: boron, silicon, germanium, arsenic, antimony, and
tellurium.
 One of the metalloids, silicon, is a semiconductor; it does not conduct electricity under
certain applied voltages, but becomes a conductor at higher applied voltages.
The halogens (Group 7A elements)
The alkali metals (Group 1A elements)
The noble gases, Group 8A elements
Periodic Property
 The Periodic Table was constructed on the basis of trends (periodicity) in chemical
properties. With an understanding of electron configuration, chemists realized that the
periodicity of chemical properties could be understood in terms of periodicity in electron
configuration.
 The Periodic Table works because elements in the same column (group) have the same
configuration in their outer shells.
Atomic Size
The size of an atom is determined by the size of its outermost occupied orbital.
Example: the size of a chlorine atom is determined by the size of its three 3p orbitals, the
size of a carbon atom is determined by the size of if its three 2p orbitals.
1. Metallic Property
The lesser the number of valence electrons and the farther the valence electrons from
the nucleus, the greater is the metallic property.
Metallic property - is the ability of the atom to donate electrons.
Metallic property thus increases from top to bottom within the same group and
decreases from left to right within the same period.
2. Ionization Energy - the energy required to remove the most loosely held electron from an
atom in the gaseous state.
Example: when lithium loses one electron, it becomes a lithium ion; it still has three
protons in its nucleus, but now only two electrons outside the nucleus.

8. In general, it increases across a row; valence electrons are in the same shell and subject
to increasing attraction as the number of protons in the nucleus increases. It increases
going up a column; the valence electrons are in lower principle energy levels, which are
closer to the nucleus and feel the nuclear charge more strongly.
3. Electron Affinity
Non-metals, in contrast with metals, tend to gain electrons to form negative ions. The
energy in this process is termed as electron affinity.
Electron affinity increases from left to right within the same period and decreases from
top to bottom within the same group.
4. Electronegativity - a measure an atom’s attraction for the electrons it shares in a chemical
bond with another atom.
5. Ionic Size
When an atom loses or gains an electron it becomes a positively or negatively charged
particle ion.
Within the group, the ionic size increases from top to bottom, the ionic size decreases across
a period for ions possessing the same electronic configuration
V.
CHEMICAL BONDS, LEWIS STRUCTURE, RESONANCE STRUCTURE, FORMAL CHARGE
 Forming Chemical Bonds
According to the Lewis model:
An atom may lose or gain enough electrons to acquire a filled valence shell and
become an ion. An ionic bond is the result of the force of attraction between a
cation and an anion.
An atom may share electrons with one or more other atoms to acquire a filled
valence shell. A covalent bond is the result of the force of attraction between
two atoms that share one or more pairs of electrons.
 Ionic Compounds
According to the Lewis model, an ionic bond is formed by the transfer of one or
more valence electrons from an atom of lower electronegativity to an atom of
higher electronegativity;
The more electronegative atom gains one or more valence electrons and
becomes an anion.
The less electronegative atom loses electrons and becomes a cation.
The compound formed by the combination of an anion and a cation is called an
ionic compound.
 Forming an Ionic Bond
In forming sodium chloride, NaCl, one electron is transferred from a sodium atom to
a chlorine atom:
N a(1s22s 22p6 3s1) + Cl(1s22s 22p6 3s23p 5)
S od iu m atom
Chlorine atom
N a+ (1s 22s2 2p6 ) + Cl -(1s 22s2 2p6 3s23p 6)
Sodium ion
Chloride ion
We use a single-headed curved arrow to show this transfer of one electron:

Na +
Forming a Covalent Bond
Cl
Na+
Cl
-
A covalent bond is formed by sharing one or more pairs of electrons.
The pair of electrons is shared by both atoms and, at the same time, fills the
valence shell of each atom.

9. Example: In forming H2, each hydrogen contributes one electron to the single
bond.
the single line represents
a shared pair of electrons
H. + .H
H H
 Polarity of Covalent Bonds
All covalent bonds involve sharing of electron pairs, they differ in the degree of
sharing:
Nonpolar covalent bond: electrons are shared equally
Polar covalent bond: electron sharing is not equal
The degree of sharing depends on the relative electronegativity of the bonded
atoms.
Electronegativity
Difference Between
Bonde d Atoms
Les s than 0.5
0.5 to 1.9
Gre ater than 1.9
Most Lik ely to
Form Betwe en
Type of Bond
N onpolar covalent
Polar covalen t
Ionic
Two nonmetals or a
nonmetal and a metalloid
A metal and a nonmetal
 Concept of Bond Polarity
Concept of bond polarity is useful in describing the sharing of electrons between
atoms.
“The greater the difference in electonegativity between two atoms, the more polar the
bond.”
 Polarity of Covalent Bonds
Bo nd
H -Cl
O -H
N -H
N a-F
C-M g
C-S
D i f fe re nce i n
El e c tr ne g ati vi ty
o
3.0
3.5
3.0
4.0
2.5
2.5
-
2.1
2.1
2.1
0.9
1.2
2.5
=
=
=
=
=
=
0.9
1.4
0.9
3.1
1.3
0.0
Ty pe o f Bo nd
po l ar co val e nt
po l ar co val e nt
po l ar co val e nt
i o ni c
po l ar co val e nt
no np ol ar co val e n t
In a polar covalent bond,
The more electronegative atom gains a greater fraction of the shared electrons
and acquires a partial negative charge; indicated by d- or the head of a crossed
arrow
The less electronegative atom has a lesser fraction of the shared electrons and
acquires a partial positive charge; indicated by d+ or the tail of a crossed arrow
 Lewis Structure
Lewis Structure can predict the molecular structure of a molecule that is the way
atoms are positioned in space relative to each other. The following are guidelines in
illustrating the Lewis structure.
1.
2.
3.
4.
5.
To predict the arrangement of atoms within the molecule.
A. Hydrogen is always a terminal atom and is connected to only one atom.
B. The atoms with lowest affinity for electrons in the molecule or ion is the central atom.
Find the total valence electrons (TVE)
Place one pair of electrons (a sigma bond) between each pair of bonded atoms.
Subtract from the total number of valence pairs the number of bonds drawn from step 3.
Place lone pairs about each terminal atom, except hydrogen to satisfy the octet rule. If pairs
are still at this point, assign it to the central atom. If the central atom from the third or higher
period it can accommodate more than four electron pairs.

10.  Lewis Structures for Several Small Molecules. (The number of valence electrons is given in
parentheses after the molecular formula)
O
H
H
H
H 2O (8)
Water
H
N
H
H
N H 3 (8)
Am m on ia
H
C
C
H
C H
H
CH 4 (8)
Meth ane
H
H
HCl (8)
Hyd ro g en ch lo ride
O
H
H
C
C
H
C
l
C
O
H
H
C2 H 4 (12)
C2 H 2 (10)
H
CH 2O (12)
Eth ylene
Acetylene
Fo rmald ehy de
H
C
H
O
O
H 2CO 3 (24)
Carbo nic acid
 Exceptions to the Octet Rule
Atoms of period 2 elements use 2s and 2p orbitals for bonding:
These four orbitals can contain a maximum of 8 electrons; hence the octet rule.
Atoms of period 3 elements have one 3s orbital, three 3p orbitals, and five 3d
orbitals:
These nine orbitals can accommodate more than eight electrons, by using 3d
orbitals; period 3 atoms can have more than eight electrons in their valence
shells.
10 electrons in
8 electrons in
10 electrons in
the valence
the valence
th e valen ce
shell of P
sh ell of P
sh ell of P
Cl
O
Cl
Cl
P
H-O-P-O-H
H-P-H
Cl
Cl
O-H
H
Phosph in e
Phosph orus pentachloride Phosph oric acid
8 electron s in
th e valence
sh ell o f sulf ur
10 electron s in
the v alence
shell o f sulf ur
: :
:
: :
12 electro ns in
the v alen ce
s hell of su lfu r
O
O S O
H- O-S-O-H
H- S- H
O
Hy drog en su lfid e S ulfu r d ioxid e
S ulf uric acid
 Resonance Structure
Resonance structure is two or more of the Lewis Structure.
Examples:
1. NO3
2. O3
3. HNO3
 Formal Charge & Electronegativity
Formal charge decides which resonance is most stable. The atoms in molecules
should have a formal charge as small as possible and the structure becomes most stable
when any negative (-) formal charge resides on the most electronegative.
FC= valence electron – ½ (no. of shared e’s)- no of unshared e’s
 Formal Charge
Examples:
1. BrO3-1
6. H3PO4
-3
2. PO4
7. SO4-2
3. NH4+1
8. SF6
4. CH3CH3
9. PCl5
+
5. H3O
10. ICl4-

11. VI.
CHEMICAL REACTIONS
 Chemical Reactions and Chemical Equations
A chemical reaction transforms one or more substances into a set of different
substances. The substances that enter into a chemical reaction are called reactants
and the substances formed are called products.
Chemical equations are representations of chemical reactions in terms of symbols of
elements and formulas of compounds involved in the reaction.
 Symbols commonly used in chemical equations
Symbols
Use / s
→
The arrow separates the reactants from
the products. It is read as yields, produces,
or forms.
+
The plus sign separates the formulas of the
reactants/products if there are more than
one reactants/products.
↑
An arrow pointing upward written after
the formula indicates gaseous product.
↓
(s) , (l), (g)
(aq)
An arrow pointing downward written after
the formula indicates a precipitate
These symbols written after the formula indicate
that the substance is a solid, liquid, or a gas
This symbol written after the formula indicates
that the substance is a solution in water, aq for
aqueous
∆
Means undergo heat
Ni
→
Condition of chemical reaction such as
temperature and catalyst used in the reaction are
indicated below or above the arrow
 Types of Chemical Reactions
Composition reaction – is one in which two or more substances (either elements or
compounds) react to form one compound. This reaction is also known as combination,
direct union or synthesis.
General form: A + B → AB
Examples:
1. 2 Na + Cl2 →
2 NaCl
2. CaO + H2O →
Ca(OH)2

12. Decomposition reaction – is one in which one compound decomposes to form two or
more new substances. Usually heat is necessary to cause this reaction to take place. This
is also known as analysis.
General form: AB → A + B
Examples:
1. 2 KClO3
→
2 KCl + 3 O2
2. 2 HgO
→
2 Hg + O2
Single Replacement reaction – is one in which a metal replaces another metal ion from a
solution or a non-metal replaces a less active nonmetal in a compound. This reaction is
also called displacement or substitution.
General form: AX + B → BX + A
(where A and B are metals)
or AX + Y → AY + X
(where X and Y are nonmetals)
Examples:
1. Fe + CuSO4 → FeSO4 + Cu
2. Cu + FeSO4 → no reaction
3. Zn + HCl
→ ZnCl2 + H2
VII.
Balancing an equation
An equation has to be consistent with the fundamental law of nature – The Law of
Conservation of Mass. That is, atoms are neither lost nor gained during chemical
reactions.
To balance a chemical equation
 Begin with atoms that appear only in one compound on the left and one on the
right; in this case, begin with carbon (C) which occurs in C3H8 and CO2.
C3 H8 (g) + O2 (g)
3CO2 ( g) + H2 O(g)
Now balance hydrogens, which occur in C3H8 and H2O:
C3 H8 (g) + O2 (g)
3CO2 ( g) + 4H2 O(g)
If an atom occurs as a free element, as for example O2, balance it last:
C3 H8 (g) + 5O2 ( g)
3CO2 (g) + 4H2 O( g)
 Practice problems: Balance these equations
Ca( OH) 2 ( s) + HCl( g)
Calcium
hy dro xide
CO2 ( g) + H2 O(l)
ph otosynthes is
CaCl2 ( s) + H2 O( l)
Calcium
chlo rid e
C6 H1 2 O6 (aq) + O2 (g)
Glucose
C4 H1 0 ( g) + O2 (g)
Bu tane
CO2 (g) + H2 O(g)