Download Unit 4 - Dorman High School

Unit 4
Chapters
Chemical Bonding and Formulas
I. Types of Chemical Bonds
What is a bond?
How is the bond held together?
What are the two types of bonds?
1.
2.
There are two types of covalent bonds:
1.
2.
II. Electronegativity
What is electronegativity?
How does electronegativity affect the bond type?
What is the difference in electronegativity that needs to occur in order for an ionic
bond to occur?
III. Bond Polarity
If the difference in electronegativity is large enough the bond between the atoms
will have polarity. What is this called?
Any diatomic molecule will have a dipole moment. When will they not?
Some polyatomic molecules can have dipole moments. When will this occur?
IV. Stable Electron Configurations and Charges on Ions
Look at the electron configuration of sodium:
Look at the electron configuration of the sodium ion:
Look at the electron configuration of neon:
Notice anything?
In almost all stable chemical compounds of the representative elements, all of the
atoms have achieved a noble gas electron configuration.
V. Ionic Bonding and Structures of Ionic Compounds
When an ionic compound is formed the bond is extremely strong.
We write the formulas for these compounds, but they are empirical formulas
because the compound is composed of a very tightly packed and ordered
arrangement of ions.
Ionic compounds can be formed with monatomic ions as well as polyatomic ions.
Polyatomic ions are formed with atoms that are covalently bonded
together.
Ex: Sulfate
VI. Lewis Structures
It is a representation of a molecule that shows how the valence electrons are
arranged among the atoms in the molecule.
These structures use “dots” to represent the valence electrons in bonding atoms.
Ex: KBr
More examples:
1. H2
2. He
3. F2
Octet Rule:
Bonding Pairs:
Lone Pairs:
4. Ne
Now for more complicated models:
1. All valence electrons must be included for all atoms in the compound.
2. Atoms that are bonded share one or more pairs of electrons.
3. Each atom must have a full valence shell.
Know the steps
VII. Lewis Structures of Molecules with Multiple Bonds
We will need to form multiple bonds between two atoms.
Single Bond:
Double Bond:
Triple Bond:
Resonance:
Resonance structures:
Exceptions to the octet rule
Boron, Sulfur, Phosphorus
VIII. Molecular Structure
Molecular Structure (geometric structure):
When referring to the structure it is often determined by the angle of the atoms in
the molecule called the bond angle.
Examples: Bent, Linear, trigonal planar, tetrahedral
IX. Molecular Structure: The VSEPR Model
Determining the structure of the molecule is important because its structure plays
a part in the molecules chemical properties.
VSEPR – Valence Shell Electron Pair Repulsion Model
The main idea:
Example: BeCl2
Steps for Predicting Molecular Structure Using the VSEPR Model
1.
2.
3.
4.
Examples:
# of Bonding
Pairs
# of NonBonding Pairs
Electron Pair
Arrangement
Angle between
Pairs
Molecular
Structure
X. Molecular Structure: Molecules with Double Bonds
Drawing molecular structures with double bonds.
When using VSEPR model to predict the molecular geometry of a molecule, a
double bond is counted the same as a single electron pair.
Ex: CO2, [NO3]-1, [SO4]2Chemical Formulas
I. Binary Compounds –
*there are two main classes of binary compounds
A. Metal + Nonmetal
B. Nonmetal + Nonmetal
Section 4.4 – Formulas of Compounds
A. A compound is:
*the word relative here refers to ratios
The smallest piece of a compound that still has the identity and properties of the
compound is known as a _______________________or _______________________.
B. The identity of a compound can be expressed as a chemical formula.
All chemical fomulas tell you:
C. Rules for Writing Formulas
1. Each atom present is represented by its element symbol.
2- The number of each type of atom is indicated by a subscript written to the right of the
element symbol.
3. When only one atom of a given type is present, the subscript “1” is not written.
II. Naming Binary Ionic (metal + nonmetal) Compounds
A. There are two types of binary ionic compounds
1. Type I – metal has fixed charge (representative elements plus Ag+,
Zn2+ and Cd2+)
Examples:
2. Type II – metal has a charge that may vary (transition, inner transition)
Examples:
B. Rules for Naming Type I Ionic Compounds
1. The cation (metal) is always named first and the anion is named second.
2. The name of the cation is just the name of that element.
3. The name of the anion takes the root name of the element with an “ide”
ending.
Examples -
C. Rules for Naming Type II Ionic Compounds – all of the same rules from
above apply…with one addition: a roman numeral (I, II, III,…etc) is used
to denote the charge on the metal.
*charge of metal can be figured out from reverse criss-cross method or
just from the fact that net charge = 0
Examples –
III. Naming Binary Covalent (nonmetal + nonmetal) Compounds
Rules
1. First element in formula is named first using the full element name.
2. The second element is named by using root element name plus “ide” ending.
3. Prefixes are used to denote the number of atoms present of each type.
4. The prefix “mono” is never used on the first element in the formula.
*examples Prefix
mono
di
tri
tetra
penta
hexa
hepta
octa
Number indicated
1
2
3
4
5
6
7
8
IV. Naming Compounds That Contain Polyatomic Ions
A. Polyatomic ions are –
1. Special type: Oxyanions
B. Rules for naming compounds with polyatomic anions.
1. Cation comes first, anion second.
2. Polyatomic ion names are not altered with “ide” endings or prefixes.
3. All other rules for naming type I and type II ionic compounds apply.
Examples –
Naming Chemical Compounds Flow Chart
Is there a Metal?
Yes
No
Is there a Polyatomic Ion?
Yes
TYPE IV
TYPE III
No
Where is the Metal located?
Groups 1 & 2
TYPE I
Transition Metal
TYPE II
VI. Naming Acids
An acid is –
two main types of acids
A. Binary Acids (H plus one other element)
Rules
1. Binary acid names always start with the prefix “hydro”
2. After prefix use anion root name with “ic” ending.
`
Examples –
B. Acids Containing Polyatomic Ions
Rules
1. NEVER use prefix “hydro”
2. anion may end in “ic” or “ous”
Examples –
VII. Writing formulas from names.
A. First, determine whether or not it is ionic, covalent, or an acid.
1. Ionic compounds and acids:
Use criss-cross method (works for polyatomic ions too)
*Examples -
2. Covalent Compounds: use prefixes as subscripts in formula.
*Examples –
VIII. Molar Mass
The mass (in grams) of 1 mole of the substance.
Examples:
Formula WeightIX. Percent Composition of Compounds
Mass Fraction = _____________________________________________
Mass Percent of something is the mass of something 1 mole of the
compound.
Examples:
X. Formulas of Compounds
A. Empirical FormulaExamples:
B. Molecular FormulaXI. Calculation of Empirical Formulas
Steps:
1. Obtain the mass of each element present (in grams)
2. Determine the number of moles of each element.
3. Divide the number of moles of each element by the smallest number of
moles of the elements present.
4. If needed, multiply the numbers from Step 3 by the smallest integer that
will convert them to whole numbers.
Best explained through examples
XII. Calculation of Molecular Formulas
Can be calculated from % Compositions or from Empirical formulas
You need 3 pieces of information:
Empirical formula
Empirical formula mass
Molar Mass
Empirical Formula
(EF)
Empirical Formula
Mass
(EFM)
Molar Mass ÷
Empirical Formula
Mass
(MM – EFM)
Molecular Formula
(MF)
Atoms
Formula Units
Molecules
Divide
by 6.02 x 1023
{atoms / mole}
{formula units / mole}
{molecules / mole}
Multiply
by 6.02 x 1023
{atoms / mole}
{formula units / mole}
{molecules / mole}
moles
Multiply by
22.4
{liters/mole}
Multiply by
atomic mass
formula mass
or
molecular mass
{grams/mole}
gas at STP
only
Divide by
atomic mass
formula mass
or
molecular mass
{grams/mole}
grams
Divide by
22.4
{liters/mole}
gas at STP
only
Liters
(gas at STP only)