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Chapters 8 and 9
Ionic and Covalent Bonding
Forming Chemical Bonds

Chemical Bond
 Force
that holds 2 atoms together
 Attraction between + nucleus and e- or
between + and – ions
 Only valence e- involved in bonding
 Bonding occurs to have complete outermost
energy levels – to become like noble gases
Forming Ions

Ion
 An
atom or bonded group of atoms that have
a + or – charge because valence e- have
been lost or gained

Cation
+

charged ion; one that has lost e-
Anion
-
charged ion; one that has gained e-
Formation and Nature of Ionic Bonds

Ionic Bond
 Electrostatic
force holds oppositely charged
particles together
 Usually occurs between a metal and nonmetal
 #e- lost by one particle = #e- gained by another
 Overall charge on a compound is always 0

Subscript
 Shows
how many atoms of an element are in
a compound
 Applies to the element to its immediate left
 If there is no subscript, it means there is only
one atom of that element.
 What is the number of atoms in each of these
compounds:
H2O
 CO2
 H2SO4
 CO

Example:

Na +1 and F -1 combine to form NaF (1:1 ratio)
Na will lose 1e- and F will gain 1eOverall charge on NaF is (+1 + -1 = 0)
Swap Charges: Na +1
F -1  NaF
Types

Oxide
 Ionic

compound with metal and oxygen
Salt
 Most
other ionic compounds, usually between
a metal and a halogen

Binary Compounds
 Contains

1 metallic and 1 nonmetallic ion
Ternary Compounds
 Contains
1 polyatomic and 1 monatomic ion
Important Vocab

Formula Unit
 Simplest
ratio of ions in an ionic compound
 The formula of the compound
 Ex: NaCl; KBr

Oxidation Number
 Charge
on an ion
 #e- atom must gain/lose/share to complete its
outermost energy level
 Metal = #valence e Non-metal = #valence e- - 8
Properties of Ionic Compounds

Form crystals: + and – ions are packed into a
regular repeating pattern. Crystal shape
depends on size and # of ions bonded

IB are strong because particles are strongly
attracted to each other
 Bonds
require much E to break apart
 Compounds have high melting and boiling points
 Mostly solids at room temp.
 Hard, rigid, and brittle
Properties of Ionic Compounds

When in aqueous solution (dissolved in water) ions
move about and conduct electricity (electrolyte)

Forming compounds from ions is exothermic
 Gives

off Energy
Breaking down compounds into ions is endothermic
 Requires
or takes in Energy
Writing Formula Units
1.
Identify the cation and write it’s symbol.
1.
2.
Identify the anion and write it’s symbol.
1.
3.
4.
It will be the first element written in it’s
entirety.
It will be the second element written as the
root of the name + “ide”
Find the oxidation number for each
element and write it above each element
without the sign.
Swap the oxidation numbers and write
the formula.
Examples:
Calcium Chloride
 Ca+2 Cl-1 -> Ca2 Cl1
 Ca1Cl2 ---- CaCl2
 Sodium Oxide
 Na+1 O-2 -> Na1 O2
 Na2O1 -> Na2O

Covalent Bonds

Chemical bond resulting from sharing valence e Usually
between elements close on periodic table
(usually nonmetals)
 Not as strong as ionic bonds
 Forms molecules or diatomic molecules
 Atoms are too far away from each other to have a
strong attraction. 1 atoms + nucleus attracts
another’s e- cloud. But both clouds repel each
other. Distance is right then to share e-, and not
transfer them.
Molecules and Diatomic Molecules

Molecule
2

different elements bonded covalently
Diatomic Molecules
2
of the same element bonded covalently
 H2, N2, O2, F2, Cl2, Br2, I2
Representing Molecules

Molecular Formula
 H2

Electron Dot Structure
H

Lewis Structure


H
H–H
Each dot represents an e-. A line represents a
bonding pair of shared e-
Examples of Single Bonds
2 2

H + OH  H – O – H
2 2

2H2 + O2  H2O

H
H + H + NH  H – N – H
2 2

H2
H–H
Properties of Covalent
Compounds
Melting and boiling points are lower than
those of ionic compounds
 Many are gases at room temperature
 If in the solid form, molecules are soft or
brittle

Examples:
 O=O
O
+
O
N
+
N 
N=N