This document discusses covalent bonding and molecular compounds. It defines a chemical bond as a force that holds atoms together, and describes covalent bonding as atoms sharing electrons. As two atoms approach each other to form a bond, their potential energy decreases to a minimum at the bond length. Bond length and bond energy vary between different bonded atoms. The octet rule states atoms want 8 electrons in their valence shell. Practice problems classify bonds and identify valence electrons.

1. Lesson 1: Introduction to Chemical
Bonding
Covalent Bonding Unit

2. Targets
I can define chemical bond.
I can describe covalent bonding.
I can classify bonding type according to
electronegativity differences.

3. Definitions
Chemical Bond – mutual electrical attraction between the
nuclei and valence electrons of different atoms that bind
the atoms together
Valence electrons – outermost electrons that are available
to be lost, gained, or shared to form a chemical bond

4. Chemical Bond
A force that holds groups of 2 or more atoms together and
makes them function as a unit
Atom – smallest unit of an element
Molecule – Group of covalently bonded atoms
Atoms
Molecule

5. Types of Chemical Bonds
Ionic Bonding – (covered in next chapter) a type of bond
in which a metal and a nonmetal transfer electrons
Covalent Bonding – type of bond in which 2 or more
nonmetal atoms share electrons

6. IONIC – Metal + nonmetal
Periodic Table COVALENT – 2 nonmetals

7. Types of Covalent Bonds
Nonpolar covalent bond – electrons are shared equally

8. Types of Chemical Bonds
Polar covalent – electrons are not shared equally because
one atom attracts the shared electrons more than the other
atom

9. Bond Types
Video

10. Electronegativity
Electronegativity - measure of an atom’s ability to attract
electrons.
Electronegativities tend to increase across a period and
decrease down a group

11. Classifying Chemical Bonds
The polarity of a bond depends on the difference between
the electronegativity values of the atoms forming the
bonds.
Nonpolar covalent – 0 to 0.3
Polar covalent – 0.4 to 1.7
Ionic – greater than 1.8

12. Electronegativity Values
Increases from left to right across a period
Decreases down a group of representative elements

13. Practice
Use electronegativity values to classify the following bonds:
a. Sulfur and Hydrogen
b. Lithium and Fluorine
c. Potassium and Chlorine
d. Iodine and Bromine
e. Carbon and Hydrogen

14. Practice
Use electronegativity values to classify the following bonds:
a. Sulfur and Hydrogen 2.5 – 2.1 = 0.4; polar covalent
b. Lithium and Fluorine 4.0 – 1.0 = 3.0; Ionic
c. Potassium and Chlorine 3.0 – 0.8 = 2.2; Ionic
d. Iodine and Bromine 2.8 – 2.5 = 0.3; Nonpolar
covalent
e. Carbon and Hydrogen 2.5 – 2.1 = 0.4 ; polar covalent

15. Covalent Bonding
Covalent Bonding and Molecular Compounds

16. Targets
I can explain why most atoms form chemical bonds.
I can explain the relationships among potential energy,
distance between approaching atoms, bond length and bond
energy.
I can state the octet rule.
I can determine the number of valence electrons for a given
atom.

17. Formation of a Covalent Bond
Nature favors chemical bonding because most atoms have
lower potential energy when they are bonded to other
atoms.

18. Formation of a Covalent Bond
Each atom has a positive nucleus in the center and negative
electrons surrounding the nucleus in a spherical pattern.
The positively charged nuclei are attracted to the negatively
charged electrons.

19. Formation of a Covalent Bond
As the atoms approach each other, the charged particles
interact: nucleus on one atom attracts electrons on the other
atom.

20. Formation of a Covalent Bond
As the atoms approach one another, the potential energy
decreases.
A bond forms when the potential energy is at a minimum.

21. Formation of a Covalent Bond
If the atoms continue to approach one another once the bond
forms, the nuclei will begin to repel one another and the
potential energy will start to increase.

22. Characteristics of the Covalent Bond
Bond length – distance between two bonded atoms at
their minimum potential energy or the average distance
between two bonded atoms
Bond energy – energy required to break a chemical bond
and form neutral isolated atoms
- kilojoules per mole (kJ/mol)
Bond lengths and bond energies vary with the types of atoms
that have combined

23. The Octet Rule
The octet rule states that atoms tend to lose, gain or
share electrons until they are surrounded by 8
electrons in their valence shell.
The number of valence electrons is equal to the group
number. (Groups 13-18; Group # -10)
LABEL YOUR PERIODIC TABLE
1A 8A
2A 3A 4A 5A 6A 7A

24. Practice
What is the relationship between bond energy and bond
length?

25. Practice
What is the relationship between bond energy and bond
length?
The bond length decreases as the strength of the
bond increases.

26. Practice
Arrange the following in order of increasing bond strength:
C–Cl, C–I, H–F, and I–I
SKIP

27. Practice
Arrange the following in order of increasing bond strength:
C–Cl, C–I, H–F, and I–I
I-I, C-I, C-Cl, H-F

28. Practice Problems
Which pair of bonded atoms has the strongest bond?

29. Practice Problems
Which pair of bonded atoms has the strongest bond?
H – F

30. Practice Problems
Which pair of bonded atoms has the weakest bond?

31. Practice Problems
Which pair of bonded atoms has the weakest bond?
I – I

32. Practice Problems
Arrange the following bond lengths in order of increasing bond
strength: 72 pm, 149 pm, 53 pm, and 398 pm
SKIP

33. Practice Problems
Arrange the following bond lengths in order of increasing bond
strength: 72 pm, 149 pm, 53 pm, and 398 pm
398 pm, 149 pm, 72 pm, 53 pm

34. Practice Problems
Determine the number of valence electrons in each of the
following atoms.
Lithium
Sulfur
Carbon
Neon

35. Practice Problems
Determine the number of valence electrons in each of the
following atoms.
Lithium - 1
Sulfur - 6
Carbon -4
Neon - 8