This document discusses chemical bonding, including ionic and covalent bonding. It explains the octet rule and how atoms gain, lose, or share electrons to achieve stable electron configurations. Ionic bonding occurs through the transfer of electrons to form cations and anions, while covalent bonding involves the sharing of electron pairs between nonmetals. The properties and naming conventions of ionic and covalent compounds are also outlined.

1. CHEMICAL
BONDING
Chapter 7

2. THE OCTET RULE
• Atoms tend to gain, lose, or share electrons
in order to get a full set of valence electrons.
• “octet” – most atoms need 8 valence
electrons for a full set
• Gaining or losing → ions = ionic bonding
• Sharing = covalent bonding
• “Dogs Teaching Chemistry”
● https://www.youtube.com/watch?v=_M9khs87x
Q8

3. IONIC
BONDING

4. PROPERTIES OF IONIC COMPOUNDS
• High melting points
• Brittle
• Usually salts
• Many dissolve in water
● Can conduct electricity because ions separate
and are charged in the solution

5. IONIC BONDS
• Electrons are transferred from one atom to
another creating ions
• Cations are attracted to anions (positives and
negatives attract)
• Metal + nonmetal
● Metals form cations
● Nonmetals form anions

6. TYPES OF IONS
• Monatomic = “one-atom”
● H+, Ca2+, Br-, N3-
• Polyatomic = “many-atoms”
● NH4
+, OH-, SO4
2-,

7. LEWIS DOT STRUCTURES
• Developed by American chemist Gilbert
Lewis (1875-1946)
• Valence electrons represented by dots
around the element symbol
● No more than two dots per side
• Can be used to show rearrangement of
electrons during chemical reactions

9. BINARY IONIC COMPOUNDS
• Contain ions of only two elements
• Formula: Cation written first, then anion
● Charges of ions written as superscripts, # of
atoms in a compound written as subscripts
• Ratio written in lowest terms = empirical
formula
● (REMEMBER THIS!)

10. BINARY IONIC COMPOUNDS
• Draw the Lewis Dot Structures for sodium
and chlorine
• Using an arrow, identify how the transfer of
1 electron can create 2 new ions

11. • Sodium
transfers an
electron to
chlorine.
• Sodium
becomes a
positive ion
with a +1
charge.
• Chlorine
becomes a
negative ion
with a -1
charge.

12. BINARY IONIC COMPOUNDS
Na+ + Cl- → NaCl
• The total (net) charge on the compound
should be zero.
• You must determine how many of each ion
will need to be in the compound to balance
out the charges.

13. COMPOUND FORMULA
PRACTICE
magnesium ion + oxide ion
Mg2+ + O2- →
Mg2+ + O2- → MgO
calcium ion and bromide ion
Ca2+ + Br- →
strontium ion and nitride ion
Sr2+ + N3- →
Mg2O2
CaBr
2
Sr3N
2

14. THE CRISSCROSS METHOD
FOR WRITING COMPOUND
FORMULAS
• Write the ion symbols (with their charges as
superscripts) for the cation and anion
• Criss-cross the two charges, moving them
diagonally from one ion’s superscript to the
other ion’s subscript
● Drop the sign!

15. CRISSCROSS METHOD
PRACTICE
magnesium ion and chloride ion
Mg2+ Cl-1
Mg with Cl =
MgCl2

16. NAMING IONIC COMPOUNDS
• Name the cation using its element name.
• Name the anion by dropping the ending of the
element name and adding –ide.
Ca3P2
calcium phosphide
• If the anion is polyatomic, simply name it using
the ion’s name
Mg3(PO4)2
magnesium phosphate

17. NAMING IONIC COMPOUNDS
• If the cation has more than one valence (it
can have different charges), indicate the
charge using roman numerals in parenthesis
after the cation name.
FeO = iron (II) oxide
Fe2O3 = iron (III) oxide

18. COVALENT
BONDING

19. COVALENT BONDS
• Formed by a shared pair of electrons
between two atoms
• Make up molecules (which make up
molecular substances)
• Between nonmetals

20. FORMULAS
• Empirical formula gives the lowest ratio of
types of atoms in a compound
• Molecular formula gives the exact number
of atoms of each element in a single
molecule of a compound
• Structural formula shows how atoms are
bonded together

21. FORMULA EXAMPLE: GLUCOSE
molecular formula
C6H12O6
empirical formula
CH2O
structural formula

22. LEWIS DOT STRUCTURES
• For molecules:
● Show pairs of electrons that are shared between
atoms using 2 dots or 1 dash.
● Leave electrons not involved in bonds as dots.

23. LEWIS DOT STRUCTURES
Draw the Lewis dot structures for:
F2
NH3
H2O
H2CO
C2H2

24. EXCEPTIONS TO
THE OCTET RULE
• Less than an octet
● BF3
• More than an octet
● SF4
• Odd number of electrons
● NO

25. PROPERTIES OF COVALENT
BONDS
• Polar covalent bonds = Unequal sharing
● Due to electronegativity difference
● More electronegative atom gets slightly negative
charge (higher electron density)
● Less electronegative atom gets slightly positive
charge (lower electron density)
• Nonpolar covalent bonds = equal sharing
● No electronegativity difference

26. PROPERTIES OF COVALENT
BONDS
• Low melting points
• Soft, flexible
• Many won’t dissolve in water
● Cannot conduct electricity even if they do
dissolve (due to no charges being present)

28. NAMING COVALENT
COMPOUNDS (MOLECULES)
• Similar to naming
ionic compounds, but
prefixes must be
added to tell the ratio
of atoms in the
compound.
mono- 1
di- 2
tri- 3
tetra- 4
penta- 5
hexa- 6
hepta- 7
octa- 8
nona- 9
deca- 10

29. NAMING COVALENT
COMPOUNDS
(MOLECULES)
• Most electronegative element written last in
formula and name.
● Drop ending of this element’s name and add
–ide.
Si2Br6
disilicon hexabromide
• Don’t include mono- prefix for 1st element listed.
CF4
carbon tetrafluoride

30. NAMING COVALENT
COMPOUNDS
(MOLECULES)
• Shorten prefixes to make names easier to say.
H2O
dihydrogen monoxide
not dihydrogen monooxide
• Sometimes common names are used.
O2 = oxygen
NH3 = ammonia

31. HYDRATES

32. NAMING HYDRATES
• Hydrates are ionic compounds that absorb water
into their solid structures.
● Anhydrous substances are water-free
● Example:
MgSO4 • 7 H2O