The document discusses chemical bonding and provides a pre-assessment to test the reader's understanding. It explains that ionic bonding occurs through the complete transfer of electrons between a metal and non-metal, forming ions. Covalent bonding is formed by the sharing of electron pairs between non-metals. Metals bond through metallic bonding where free-floating electrons are attracted to stationary positive ions. The document aims to help readers understand how different types of bonds are formed and the properties they produce.

1. CHEMICAL
BONDING

2. Objectives:
 Explain the formation of ionic and covalent
bonds.
 Recognize different types of compounds (ionic or
covalent) based on their properties such as
melting point, hardness, polarity and electrical
and thermal conductivity.
 Explain properties of metals in terms of their
structure.
 Explain how ions are formed.

3. Pre-Assessment:
1. Which of the properties of atoms is the most suitable
reference for the kind of bond that will take place
between/among them?
a. atomic size b. electronegativity
b. ionization energy d. electron affinity
2. What kind of particle is produced after covalent
bonding?
a. atom b. molecule
c. ion d. electron

4. 3. How does ionic bonding take place?
a. Two non-metallic elements of different kinds form strong
forces of attraction.
b. Two non-metallic element of the same kind form strong
forces of attraction.
c. A non-metallic element like fluorine is attracted to a
metallic element like sodium.
d. A metallic element like sodium transfers an electron to a
non-metallic element like fluorine.
4. What kind of chemical bond will form between two oxygen
atoms?
a. ionic bond b. metallic bond
c. polar covalent bond d. nonpolar covalent bond

5. 5. Which of the following type of bonds will have the
highest electrical and thermal conductivity?
a. ionic bond b. metallic bond
c. polar covalent bond d. nonpolar covalent bond
6. Why can metals be hammered without breaking?
a. They are ductile.
b. They are not brittle.
c. They are malleable.
d. Its particles are strong

6. 7. When does covalent bonding take place?
a. It takes place when atoms share electrons with
one another.
b. It takes place when the attraction between atoms
is strong.
c. It takes place when atoms collide with one
another.
d. It takes place when atoms attain stability.
8. Nitrogen (N) belongs to family 5A and it is diatomic.
How many nonpolar covalent bonds will there be in N2
molecule?
a. 1 b. 2 c. 3 d. 4

7. 9. Which of the following will have the highest melting
temperature?
a. sodium chloride (salt) b. paraffin wax (candle wax)
c. sucrose (table sugar) d. lead wire
10. Which among the following shows that an atom is stable?
a. having 2 valence electrons b. having 4 valence electrons
c. having 6 valence electrons d. having 8 valence electrons

8. Question:
What do you notice about
the electronic configuration
of the noble gases?

9.  Except for Helium, all of them have eight
electrons at the outermost energy level. The
sharing or the complete transfer of
electron causes an atom to have the same
electronic configuration as that of the
nearest noble or inert gas. The sharing or the
complete transfer indicates that the atom has
attained stability. Either the sharing or the
complete transfer of electrons leads to the
formation of compounds.

10. Valence Electron
 the number of electrons at the
outermost energy level of the atom

11. Perform Activity 1:
Mapping the
Periodic Table

12. Q1. Where can you find metals, non-metals,
and noble gases in the periodic table of
elements?
 Facing the periodic table, metals are located before
the metalloids, on the left side of the periodic table of
element (Groups I -13).
 Facing the periodic table, non-metals are found after
the metalloids, on the right side of the periodic table
(Groups 14-17).
 Noble gases are at the rightmost column of the
periodic table (Group 18).

13. Q2. Which number tells you the number
of valence electrons?
 The representative elements or main group elements are found in
the s- block and p-block of the periodic table.
 The group number of the s-block (Groups 1 and 2) elements tells us
the number of valence electrons. For example, all elements
belonging to Group 2 have two (2) valence electrons. An exception
is helium in Group 18.
 For the p-block elements (Groups 13 to 18), the number of valence
electrons is the group number minus 10. This pattern does not
include the transition elements.

14. Q3. What do you notice in the number of
valence electrons of metals, non-metals, and
noble gases?
Metals have lesser valence electrons
than non-metals. All noble gases have
8 valence electrons except Helium,
which has 2 valence electrons

15. Q4. What kind of element has:
a. Less than 4 valence electrons?
-Metals
b. More than 4 valence electrons?
- Non-metals
c. Low electronegativity?
- Generally, metals
d. High electronegativity?
-Generally, nonmetals

16. e. Low ionization energy
- metals
f. High ionization energy
-Non-metals

17. Electronegativity
 is a measure of the tendency of an atom to attract
electrons, the higher its value, the higher its tendency to
attract electrons.
Ionization energy
 is the energy needed to pull or remove one or more electron/s
from a neutral atom
 The lower the ionization energy the easier it is to remove its
valence electrons.

18. Perform Act 2: Lewis Symbol
Lewis Symbol- This symbol is composed of
the chemical symbol of the element and
dots that represent the number of valence
electrons.

19. Octet Rule
 Atoms always strive to attain the most stable
arrangement of electrons. Atoms are stable if their
electrons have the same kind of arrangement as that of
noble gases, where the s and p orbitals are filled with
electrons except for helium, where only the s orbitals
are filled up. All the noble gases except for helium have
8 valence electrons.
 The Octet Rule tells you that elements gain or lose or
share electrons to achieve the electronic configuration
of the nearest noble gas. Thus after chemical bonding,
elements become isoelectronic with the nearest
noble gas in the periodic table.

20. Metals have low electronegativity and
ionization energy, thus they tend to
transfer or loose electrons.
 Non-metals have high
electronegativity and ionization
energy. They have a greater tendency
to attract electrons towards themselves.
Thus non-metals tend to gain electrons.

21. Perform Act 3: Bonding by Transfer of
Electrons
 Thus, 3.0-0.90= 2.1
 If the electronegativity difference, is greater than 1.9,
complete transfer of electron/s is possible.

22. After ionic bonding, sodium (Na) became
isoelectronic with neon (Ne) while chlorine
became isoelectronic with argon (Ar), thus
both sodium and chlorine attained stability.
 Isoelectronic means sodium (Na) attain the
same electronic configuration with neon
(Ne) and in the case of chlorine it acquired
the same configuration with that of argon
(Ar). Thus, both of them become stable.

23. Guide Questions:
 Q1: What kind of element forms cation after ionic
bonding? Why?
-Metals form cations because they completely transfer or
give away electrons.
 Q2: What kind of element forms anion after ionic
bonding? Why?
-Non-metals form anions because they attract electrons
toward themselves.
 Q3: Why do ions form after ionic bonding?
- Ions form after ionic because this type of bond involves
complete transfer of electrons.

24.  Q4: Did the atoms attain stability after ionic bonding? Explain
your answer.
-Yes, after ionic bonding the participating atoms attain the
stable electronic configuration of the nearest noble gas.
 Q5: How can you tell that ionic bonding will take place
between metals and non-metals?
- Ionic bonding will take place between metal and non-metal
with electronegativity difference of greater than 1.9.
 Q6: Will all combinations of metals and non-metals form ionic
bond? Why?
- Not all metal-nonmetal combinations will result in an ionic
bond. Only those with electronegativity difference of more than
1.9 will result to an ionic bond.

25.  The best example of a metal-non-metal
combination which did not result to ionic bond
is AlCl3 because the electronegativity
difference is only 1.5. Thus, instead of ionic
bond the chemical combination formed a polar
covalent bond.

26. Bonding by Sharing Electrons
 Covalent bond- bonding by sharing electrons
 2 Types of Covalent Bond:
1. Polar Covalent Bond- electronegativity
difference is less than 1.9 and more than 0.4,
polar covalent bond is formed.
2. Non-polar Covalent Bond- electronegativity
difference is equal to 0.4 or less

27. You can detect the number of bonding and
nonbonding electrons through a computation
based on octet rule.
a. Get the Total Available Valence Electron (TAVE) in a Compound:
For CO2
carbon atom has 4 valence electrons
oxygen atom has 6 valence electrons
Total Available Valence Electrons (TAVE) = (1 C atom x 4) + (2 O atoms x 6)
= 4 + 12
= 16

28. b. Compute for the Octet Rule requirement that each atom should have 8
valence electrons to become stable.
Number of Electrons based on Octet Rule = (1 C atom x 8) + (2 O atoms x 8)
= 8 + 16
= 24
 c. Subtract a from b, then divide the difference by 2 because a pair of
shared electron is equal to 1 bond. The quotient will give you the number
of bonds around the central atom.

30. In covalent bonding, a pair of shared
electrons is equal to one (1) bond. Notice
that after the sharing of electrons, each of
the atoms in the compound attains a stable
configuration and a covalent compound is
formed.
Such compound could exist as independent
units called molecules. As a whole, the
molecule does not carry a charge.

31. Perform Act 5:
BONDING AMONG METALS
 Metallic Bond- a bond formed by the attraction of cations and the
electrons around them
 Metals have low ionization energy so they easily lose their
outermost electrons
 Electrons are not moving around one nucleus, the positive atomic
nuclei of the metal are surrounded by electrons moving freely
throughout the piece of metal.
 These moving electrons in metals are called a “sea of electrons.”
 So, what holds the metal together are the strong forces of
attraction between the positive nuclei and the freely moving
electrons.

32. Metallic Properties

33. Uses of Metals

34. Summary

36. Glossary