Unit 6
Chemical Bonding
Chemistry I
Mr. Patel
SWHS
Topic Outline
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MUST know all assigned ions and elements!!!
Review Ions and Octet Rule (7.1)
Ionic Bonding (7.2)
Naming Ionic Compounds (9.2)
Metallic Bonding (7.3)
Covalent Bonding (8.1, 8.2)
Polarity (8.4)
Naming Covalent Molecules (9.3)
Ionic Bonding Intro
Metallic Bonding Intro
Covalent Bonding Intro
Ions
• Ion – charged species
– Must show the sign and value of charge
• Valence electrons – electrons in highest
occupied energy level
– How do we find the number of valence electrons?
• For elements 1A-8A = Group Number (except He)
– Depict using Lewis Dot Structures
EX: Consider the element Aluminum.
a) How many valence electrons does Al have?
b) Draw the Lewis Dot Structure for Al.
The Octet Rule
• Octet Rule – Atoms try to have 8 valence
electrons
– Goal: Be like a noble gas = stable
– Will lose or gain electrons
– Results in ions…What do we call these ions?
– Cations – Positive charge species (metals)
– Anions – Negative charge species (nonmetals)
EX: Consider the element Phosphorus.
a) How many valence electrons does P have?
b) Draw the Lewis Dot Structure for P.
c) Draw the Lewis Dot Structure for the ion
form of phosphorus.
d) Will it form a cation or anion? Name it.
Ionic Bonding
• Bond between Metal and Nonmetal
– Actually, it is between cations and anions
– Metal always comes first
• Ionic bonding is due to the transfer of
electrons
• Important: The compound is always neutral
– Positive = Negative
Ionic Bonding
• Consider sodium chloride, CaCl2
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–
Metal first then nonmetal
Subscript tells you number of ions
1 calcium ion for 2 chloride ions
Repeated array of ions – crystal
• Chemical Formula – shows # of ions
and smallest unit
• Formula unit – lowest wholenumber ratio of ions
CaCl2
NaCl
CaCl2
Ionic Bond Formation
Ionic Bond Formation
Ionic Bond Formation
• There are 4 steps to diagram ionic bonding
1. Draw neutral Lewis Dot Structures
(one with dots and other with x)
2. Show transfer of electrons (follow Octet Rule)
3. Show Resulting Ions
4. Write Formula
Ex: Show the Ionic Bond formation between
sodium and chlorine.
Step 1:
Draw Lewis
Dot Structure
Step 2:
Transfer the
Electron(s)
Step 3:
Resulting
Ions
Step 4:
Chemical
Formula
Na
x
Nax
1+
Use x and dots or different
colors to show differences in
valence electrons (VE).
Cl
- Metal lose e-, Nonmetal gains.
- Metal must lose all VE
- Nonmetal must have 8 VE
Cl
Na x Cl
NaCl
1-
Show resulting ions – must
have all charges. Anion must
show the transferred electron.
Only show element symbols and
subscripts – no charges, dots
Ex: Show the Ionic Bond formation between
calcium and fluorine.
Step 1:
Draw Lewis
Dot Structure
Step 2:
Transfer the
Electron(s)
Step 3:
Resulting
Ions
Step 4:
Chemical
Formula
x
x
Ca
x
F
Ca
x
F
2+
Ca 2 x F
CaF2
- Metal lose e-, Nonmetal gains.
- Calcium must lose 2 VE
- Fluorine has 7 VE, can only take
1 more = Problem
F
1-
We need to add another
fluorine atom to take other VE
from Calcium = Solution
There is one calcium and two
fluoride ions in this bond.
Ex: Show the Ionic Bond formation between
elements X (Group 3A) and Z (Group 6A).
Properties of Ionic Compounds
• Arranged into a crystal lattice
– Large attractive forces = stable, strong structure
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Solid at room temperature
High melting points
Poor conductor as a solid
Good conductor when
molten or in solution
• Overall exothermic
Colligative Properties
• Ionic Compounds are known as salts
• Salts will ionize when in solution
– Split into ions
• Colligative Properties
– Properties that depend on the nature of the solute
but not the quantity
– Boiling Point Elevation – salts increase Boiling Point
– Freezing Point Depression – salts decrease Freezing
Point
• Number of ions determine the effect: CaCl2 vs
NaCl
Covalent Bonding
• Bond between Nonmetal and Nonmetal
– Can also include semimetals
– NO IONS (cations/anions)
• Covalent bonding is due to the sharing of
electrons
• Molecule – group of neutral atoms held
together by covalent bonds
Covalent Bonding
• Covalent molecules are defined structures
– No crystal lattice
– Has a specific 3-D structure
• Molecular Formula – shows how many atoms
of each element are in a molecule
– We do not reduce formulas like ionic compounds
– Ex: H2O, CO, CH4, C6H12O6
Depictions of Covalent Molecules
Covalent Molecule Shapes
• Sharing of electrons are caused by overlapping
and hybridizing orbitals (electron location)
• VSEPR Theory – Valence Shell Electron Pair
Repulsion Theory
• VSEPR helps explain and predict the shape of
molecules
– Theory states that shape of molecules based on
minimizing the repulsion of valence electron pairs
– Keep electrons as far apart as possible
Methane = CH4 - Tetrahedral
Properties of Covalent Molecules
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Distinct groupings of atoms = molecule
Solid, liquid or gas at room temperature
Low melting points
Poor conductor
Polar or Nonpolar
Diatomic Molecules
• There are 7 elements that can not be found as
individual atoms – found in pairs
• Diatomic molecule – two atoms
• H2 , N2 , O2 , F2 , Cl2 , Br2 , I2 (Group 7 + HON)
Types of bonds
• Covalent molecules share bonds to complete
octets – octet rule still applies!
• Three types of bonds: single, double, triple
Comparing Bonds
Bond
Electrons
Involved
Bond
Length
Bond Strength
Diatomics
Single
Bond
2 electrons
Longest
Weakest
H2 , F2 , Cl2 , Br2 , I2
Double
Bond
4 electrons
Moderate
Intermediate
O2
Triple
Bond
6 electrons
Shortest
Strongest
N2
Valence Electrons not participating in bonding are called
non-bonding electrons or lone pairs.
Polarity
• Covalent Bonding is sharing of electrons
• Electrons can be shared equally or unequally
depending on the strengths of the atoms
• If electrons have different electronegativities,
the molecule will be polar
• Like dissolves Like
Polarity
• Polar – electrons shared unequally
– Align themselves with an electric field
– Ex: Water
• Nonpolar – electrons shared equally
– All diatomics are nonpolar
Metallic Bonding
• These are the forces that hold metals together
• Valence electrons are a sea of electrons
around nuclei
– Excellent conductors
• Metals atoms arranged in compact and
orderly patterns.
Comparing Ionic and Covalent Bonding
Characteristic
Ionic Bonding
Covalent Bonding
Elements
Metal and Nonmetal
Nonmetal and Nonmetal
Bond Formation
Transfer electron
Share electron
Product of bond
Formula Unit (salt)
Molecule
Physical State
Solid
Solid, Liquid, Gas
Melting Point
High
Low
Conductivity
Good Conductor
Poor to Non-conductor
Nomenclature
No Prefixes
Can Have Roman Numerals
Always Prefixes
No Roman Numerals
Follow Octet Rule
YES
YES
Force
Intramolecular
Intramolecular
NOMENCLATURE
RULES
Nomenclature: Type 1
Ionic Compounds with Fixed Charges
• Groups 1A-7A have fixed charges…memorize
these charges (Skip 4A and 8A)
1A: 1+
5A: 3-
2A: 2+
6A: 2-
3A: 3+
7A: 1-
• Must be able to go from formula to name
AND name to formula
Nomenclature: Type 1
Ionic Compounds with Fixed Charges
• Rules for Formula  Name:
– Write down full name of the cation
– Write down name of the anion (-ide)
– Ex: K2O = potassium oxide
• Practice:
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H2S
LiF
Al2O3
– hydrogen sulfide, lithium fluoride, aluminum oxide
Nomenclature: Type 1
Ionic Compounds with Fixed Charges
• Rules for Name  Formula:
– Write symbol and charge of cation and anion
– Use subscripts to make all positive = negative
(cross charges and reduce)
– EX: Lithium Phosphide = Li1+ P3-  Li3P
• Practice:
– Magnesium bromide, barium sulfide, Calcium nitride
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MgBr2
BaS
Ca3N2
Nomenclature: Type 2
Ionic Compounds with Variable Charges
• Groups 1A-7A have fixed charges- charge always
the same (Skip 4A and 8A)
• Other metals (transition metals) do not have fixed
charges – multiple possibilities for charge
– We must indicate the specific charge
• Example: Mg – always Mg2+
• Example: Mn – can be Mn1+, Mn5+, Mn6+, Mn7+
Nomenclature: Type 2
Ionic Compounds with Variable Charges
• Rules for Formula  Name:
– Write down full name of the cation and anion (-ide)
– Find total negative charge = total positive charge
– Find charge on each cation
– Write charge as Roman Numeral between cation and
anion in name
– Ex: FeCl3 = iron(III) chloride
• Each Cl is 1- charge = b/c there are 3 Cl there is total of 3• This means there is a total of 3+ so Fe must be 3+
• Write charge of Fe as roman numeral in name
Nomenclature: Type 2
Ionic Compounds with Variable Charges
Practice Formula  Name:
1. SnS2
1. Tin(IV) sulfide
2. Cu2O
2. Copper(I) oxide
3. Fe3P2
3. Iron(II) phosphide
Nomenclature: Type 2
Ionic Compounds with Variable Charges
• Rules for Name  Formula:
– Write symbol and charge of cation and anion
– Charge of cation comes from Roman Numeral
– Use subscripts to make all positive = negative
(cross charges and reduce)
– EX: Cobalt(II) nitride = Co2+ N3-  Li3N2
• Charge of cobalt came from roman numeral
• Charge of anion came from periodic table
• Cross charges (positive = negative)
Nomenclature: Type 2
Ionic Compounds with Variable Charges
Practice Name  Formula:
1. Manganese(II) chloride
1. MnCl2
2. Iron(III) oxide
2. Fe2O3
3. Copper(II) sulfide
3. CuS
Nomenclature: Type 3
Ionic Compounds with Polyatomic Ions
• The compounds have more than two elements
– Must know polyatomic ions (page 257)
• Treat the polyatomic ion as a single unit that WILL
NOT CHANGE
– Nitrate = NO31-
2 nitrates = (NO31-)2
• Must be able to go from formula to name AND
name to formula
Nomenclature: Type 3
Ionic Compounds with Polyatomic Ions
• Rules for Formula  Name:
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Write down full name of the cation
Use Roman Numerals is cation is transition metal
Write down name of anion (-ide or polyatomic ion)
Ex: Ba(OH)2 = barium hydroxide
Ex: Pb3(PO4)2 = lead(II) phosphate
• Practice:
– Fe(CN)3
Li2SO4
NH4C2H3O2
– Iron(III) cyanide, lithium sulfate, ammonium acetate
Nomenclature: Type 3
Ionic Compounds with Polyatomic Ions
• Rules for Name  Formula:
– Write symbol and charge of cation and anion
– Use subscripts to make all positive = negative
(cross charges and reduce)
– EX: Tin(IV) sulfite= Sn4+ (SO32-)  Sn(SO3)2
• Practice:
– Calcium hydroxide, copper(I) nitrite, ammonium phosphate
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Ca(OH)2
CuNO2
(NH4)3PO4
Nomenclature: Type 4
Covalent Molecules
• The molecules do not contain metals.
• Need to know Greek prefixes
Nomenclature: Type 4
Covalent Molecules
• Rules for Formula  Name:
– Write down full name of the first element
– Write down modified name of second element (-ide)
– Place Greek prefixes before each element name to
denote the number of atoms
– No mono prefix on first element
– Ex: CO2 = carbon dioxide
• Practice:
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N2O5
NO3
XeF6
– Dinitrogen pentoxide, nitrogen trioxide, xenon hexafluoride
Nomenclature: Type 4
Covalent Molecules
• Rules for Name  Formula:
– Write symbol of both elements
– Use prefixes as subscripts
– EX: phosphorus pentafluoride = PF5
• Practice:
– Dihydrogen monoxide, sulfur heptachloride
–
H2O
SCl7