This document discusses chemical bonds and how they form compounds. It explains that there are three main types of bonds: ionic, covalent, and metallic. Ionic bonds form when electrons are transferred between atoms. Covalent bonds form when atoms share pairs of electrons to achieve stable electron configurations. Metallic bonds involve delocalized electrons that are shared throughout the metal. The document also discusses naming and writing formulas for ionic and covalent compounds according to systematic rules based on the elements present and their bonding structure.

1. PowerPoint Lectures
to accompany
Physical Science, 10e
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Chapter 9
Chemical Bonds
An Introduction to Chemistry

2. Core Concept
Electron structure will explain how and
why atoms join together in certain
numbers.

3. Compounds and Chemical
Change
• Atom - smallest elemental
unit
• Molecule
– smallest particle still
retaining the characteristic
chemical properties of a
substance
– Examples:
• oxygen, hydrogen gas -
diatomic molecules
• Ozone - triatomic oxygen
molecule
• Noble gases: helium, neon
- “monatomic” molecules

4. Chemical Reactions
• Formation and/or
breaking of chemical
bonds to form new
molecules (products)
from old ones
(reactants)
• Chemical energy -
internal bonding
potential energy
• Chemical equation -
symbolic summary of
chemical reaction

5. Valence Electrons and Ions
• Outer electrons
determining the
chemical properties of
an atom
• Octet rule
– Atoms attempt to acquire
an outer shell of eight
electrons
– Electrons can be
gained/lost/shared in the
process
• Example: sodium (Na)

6. Chemical Bonds
• Attractive forces holding
atoms together in
compounds
• Can be described in terms of
molecular (delocalized) or
atomic (localized) orbitals
Three types:
• Ionic
– Electrons transferred
between atoms
– Electrostatic force =
binding force
• Covalent
– Octets achieved through
sharing electrons
– Typically between
nonmetallic elements,
r.h.s of periodic table
• Metallic bonds
– Outer electrons move
freely throughout metal
– “Electron gas” within rigid
lattice of metal atoms
– Conduct heat and
electricity well

7. Ionic Bonds
• Chemical bond of
electrostatic attraction
• Form crystalline solids
with orderly geometric
structure
• Example: NaCl
• Na loses; Cl gains
• No single NaCl
molecule, per se

8. Energy and Electrons in
Ionic Bonding
• Reaction energy
released = heat of
formation
• Divided conceptually
into half-reactions
Electron transfer rules
• Electrons lost/gained to
form closed octets
• Number gained =
number lost

9. Ionic Compounds and
Formulas
Formulas
• List elements in compound and
their proportions
• Proportions decided by electron
gain/loss
Ionic compounds
• Characterized by ionic bonds
• White, crystalline solids soluble
in water
• Families IA and IIA lose
electrons and form positive ions
• Families VIA and VIIA gain
electrons to form negative ions

10. Covalent Bonds
• Chemical bonds formed by sharing
pairs of electrons
• Electrons shared to form octets, ideally
• Overlap of shared electron clouds
between nuclei yields net attraction
• Atoms within covalent compounds are
electrically neutral, or nearly so

11. Covalent Compounds and
Formulas
• Covalent compound -
held together by
covalent bonds
• Electrons shared in
covalent bonds
• Electron dot
representation
– Bonding pairs shared
– Lone (non-bonding) pairs
not shared

12. Multiple Bonds
• Sharing of more
than one electron
pair
• Examples
– Ethylene - double
bond
– Acetylene - triple
bond

13. Bond Polarity
• Result of unequal
sharing of electrons
• Electronegativity
– Measure of an atom’s
ability to attract electrons
– Differences:
• 1.7 or greater - ionic
• 0.5-1.7 - polar covalent
• Less than 0.5 - covalent

14. Electronegativities

15. Composition of Compounds
• Millions of different combinations of over 90
elements
• Common names
– Often related to historical usage (baking soda,
washing soda,…)
– Difficult to relate to actual molecular composition
• Modern approach - systematic sets of rules
– Different for ionic and covalent compounds
– One common rule - “-ide” means compound
contains only two different elements

16. Ionic Compound Names
• Name of metal (positive) ion
first; then nonmetal
(negative) ion
• Many elements have
variable charges
• Historical suffix usage
– “-ic” for higher of two;
– “-ous” for lower
• Modern approach
– English name of metal
followed by Roman
numeral indicating
charge

17. Ionic Compound Formulas
• Two rules
– Write symbol for
positive ion first
followed by negative
ion symbol
– Assign subscripts to
assure compound is
electrically neutral
• Example: Calcium
chloride

18. Covalent Compound Names
• Molecular -
composed of two or
more nonmetals
• Same elements can
combine to form a
number of different
compounds
Two rules
• First element in formula
named first with number
indicated by Greek
prefix
• Stem name of second
element next; Greek
prefix for number;
ending in “-ide” (for two
elements)

19. Covalent Compound Formulas
• Examples: carbon dioxide,
carbon tetrachloride
• Valence
– Number of covalent
bonds an atom can form
– Hydrogen valence = 1
– Oxygen = 2; single and
double bonds
– Nitrogen = 3; single,
double and triple bonds
– Carbon = 4 - single,
double and triple bonds