staff.tuhsd.k12.az.us/ldompier/ChmPhyPP%5CCh.%206.ppt
TEKS
 7 (A) name ionic compounds containing main group or transition
metals, covalent compounds, acids, and bases, using International
Union of Pure and Applied Chemistry (IUPAC) nomenclature rules;
 7 (B) write the chemical formulas of common polyatomic ions, ionic
compounds containing main group or transition metals, covalent
compounds, acids, and bases;
 7 (C) construct electron dot formulas to illustrate ionic and covalent
bonds;
 7 (D) describe the nature of metallic bonding and apply the theory to
explain metallic properties such as thermal and electrical conductivity,
malleability, and ductility;
Objectives
 Explain why atoms form bonds
 Define chemical bond & name three types of
chemical bonds
 Compare and contrast the advantages and
disadvantages of varying molecular models
Bonding Atoms
 Why do atoms bond?
- each atom wants a full outermost energy level
- gain, lose, and share valence electrons to achieve
the duet or octet rule aka: “being happy”
- gives each atom an electron configuration similar
to that of a noble gas
ex. Group 18: He, Ne, Ar
Chemical Bonds
 Chemical Bonds
- attractive force that holds atoms or ions
together
- 3 types
ionic, covalent, metallic
- determines the structure of compound
- structure affects properties
- melting/boiling pts, conductivity etc.
Chemical Structure/Models
 Chemical Structure/Molecular Models
- arrangement of bonded atoms or ions
- bond length: the average distance between the
nuclei of two bonded atoms
- bond angles: the angle formed by two
bonds to the same atom
Molecular Models of Compounds
 Ball and stick
- atoms are represented by balls
- bonds are represented by sticks
* good for “seeing” angles
 Structural
- chemical symbols represents atoms
- lines are used to represent bonds
* good for “seeing” angles
H
H
O
Molecular Models Cont.
 Space filling
- colored circles represent atoms, and the space they
take up
- no bonds, no bond angles
 Electron Dot/Lewis Structure
- chemical symbol represent atom
- dots represent valence electrons
- 2 center dots represent a bond
- no bond angles, no bond length
Objectives
 Describe how an ionic, covalent and metallic bonds
forms
 Relate the properties of ionic compounds to the
structure of crystal lattices
 Compare polar and non polar bonds, and
demonstrate how polar bonds affect polarity of a
molecule
 Describe the structure and strength of bonds in
metals & relate their properties to their structure
Ionic Bonds / Ionic Compounds
 Definition
- bond formed by the attraction between oppositely
charged ions
cation: positive: lost e-’s
anion: negative: gained e-’s
- oppositely charged ions attract each other and form
an ionic bond
ex. Na+ + Cl- = NaCl
- electrons are transferred from one atom to another
- negative ions attract more positive ions, and soon a
Ionic Bonds Cont.
Networks / Crystal Lattices
 Networks
- repeating pattern of multiple ions
ex. NaCl
- every Na ion is next to 6 Cl ions
- strong attraction between ions creates a rigid
framework, or lattice structure: aka: crystals
ex, cubes, hexagons, tetragons
Properties of Ionic Compounds
 Structure affects properties
- strong attractions between ions: strong bonds
- high melting/boiling pt
- shatter when struck (think of it as one unit)
- conductivity
solid: ions are so close together, fixed
positions, (can’t move)
NO conductivity
liquid: ions are freely moving due to a
broken lattice structure
Good conductivity
Covalent Bonds
 Definition
- chemical bond in which two atoms share a pair of
valence electrons
- can be a single, double, or triple bond
single, 2e-’s (-); double, 4e-’s (=); triple, 6e-’s( )
- always formed between nonmetals
- mostly low melting/boiling points
 2 types of bonds
- polar
- non polar
Covalent Bond Cont.
 Non Polar
- bonded atoms that share e-’s equally
- same atoms bonded
ex. Cl – Cl: Cl2
 Polar
- bonded atoms that do not share e-’s equally
- different atoms bonded
H
ex. H – N – H: NH3
Covalent Bonds Cont.
Metallic Bonds
 Definition
- a bond formed by the attraction between positively
charged metal ion (cation) and the shared electrons
that surround it (sea of electrons)
ex. Cu
 Properties
- Conductivity: Good: electrons can move freely
- Malleable: lattice structure is flexible
Metallic Bonds Cont.
Predicting Bond Type
Objectives
 Recognize monoatomic ions, metals with multiple
ions and polyatomic ions
 Name and determine chemical formulas for
monoatomic ions, metals with multiple ion and
polyatomic ions
Naming Ions
 Monoatomic Ions
- cation
-name of element with ion
ex. (Na) Sodium (Na+) Sodium ion
- anion
- name of element with the suffix –ide
ex. (Br) Bromine (Br-) Bromide
 Ions with multiple cations
- transition metals
- most form 2+, 3+ and 4+
ex. Cu+, Cu2+
Naming Metals with Multiple Ions
 Transition Metals
- form multiple ions
- in order to name the ion use a roman numeral to
indicate the charge
ex. Cu2+: Copper (II), Titanium (III): Ti3+
Practice Problems:
Fe3+: Iron (III)
Mercury (III): Hg3+
Pb4+: Lead (IV)
Chromium (II): Cr2+
Polyatomic Ions
 Definition
- an ion made of one or more atoms that are
covalently bonded and that act as a unit
(atoms that have lost or gained electrons)
ex. CO3 2- , NH4+
- behave the same as other ions
- polyatomic ions can combined like
any other ion (as a unit)
ex. NH4NO3
1:1 ratio
(NH4)2SO4 2:1 ratio
Polyatomic Ions
 Naming polyatomic ions
- not logical
- rules for some compounds
 -ite & -ate endings
- indicates the presence of oxygen
- called oxyanions
- if (-) does not specify how many oxygen atoms
are present
ex. Sulfate:4, Nitrate:3, Acetate:2
Polyatomic Ions Cont.
- often several oxyanions differ only in
the number of oxygen atoms present
ex. Sulfur
- ion with more oxygen takes the –ate
ending
ex. SO4
- ion with less takes the –ite ending
ex. SO3
 Common Oxyanions
* Make sure you know these: memorize
Polyatomic Ions Cont.
 Common Polyatomic Ions
Objectives
 Name ionic compounds from formulas
 Determine the chemical formulas for ionic
compounds from compound name
Naming Ionic Compounds
 Naming ionic compounds (binary)
Formula to Name
- name of cation followed by the name of the
anion
ex. NaCl: Sodium Chloride
ZnO: Zinc (II) Oxide
CuCl2: Copper (II) Chloride
- formulas must indicate the relative number of
cations and ions if transitional
Naming Ionic Compounds
 Practice Problems
MgBr2
Magnesium Bromide
KI
Potassium Iodide
CuCl2
Copper (II) Chloride
Fe2S3
Iron (III) Sulfide
Formulas of Ionic Compounds
 Writing formulas for ionic compounds
Name to Formula
- balance the cation charge and anion charge,
leaving NO net charge
- use subscripts to denote the number of atoms in the
formula
ex. NaCl: Na+ Cl- : NaCl
CaCl: Ca2+ Cl- : CaCl2
**1 to 1 ratios do not designate charge**
**Criss-Cross charges into subscripts**
Practice Problems
 Write the formula for the following atoms
a. lithium oxide
Li2O
b. beryllium chloride
BeCl2
c. titanium (III) nitride
TiN
d. cobalt (III) hydroxide
Co(OH)3
Objectives
 Name Covalent compounds from formulas
 Determine the chemical formulas for covalent
compounds from compound name
Naming Covalent Compounds
 Prefix System
# of atoms
1
2
3
4
5
6
7
8
9
10
prefix
mono
di
tri
tetra
penta
hexa
hepta
octa
nona
deca
Naming Covalent Compounds Cont.
 Rules for the prefix system
1. less electronegative element is given first. It is
given a prefix only if it contributes more than one
atom to a molecule of the compound
2. The second element is named by combining (a) a
prefix indicating the number of atoms contributed by
the atom (b) the root of the name of the second
element, and (c) the ending –ide
3. The o or a at the end of a prefix is usually
dropped when the word following the prefix begins
with another vowel
ex. Monoxide or pentoxide
Naming Covalent Compounds Cont.
Naming covalent compounds from formula
1. SiO2
Silicon dioxide
2. PBr3
Phosphorus tribromide
3. CI4
Carbon tetraiodide
4. N2O3
Dinitrogen trioxide
Writing Formulas for Covalent
Compunds
 Writing formulas from names
1. Carbon Dioxide
CO2
2. Dinitrogen Pentoxide
N 2O 5
3. Triphosphorus monosulfide
P3S
4. Sulfur Monobromide
SBr