Chapter 8 Notes
Chapter 8
Molecular Compounds
&
Covalent Bonding
H
H
Both H atoms
have 1 unpaired
electron
H
The electrons pair up.
H
H H
Covalent Bond formed.
Does it contain ionic or covalent bonds?
Formula
Ionic or Covalent
Explanation
Why do covalent bonds form?
• If only group 5A, 6A, 7A atoms existed, ionic bonds can’t form.
NONMETALS
• Each atom needs electrons so they are not willing to lose any.
• If two Hydrogen atoms are locked in a room together, what happens?
8.1 Molecular Compounds
• Molecule: “neutral” group of atoms joined together by covalent bonds. (Sharing electrons)
• Consists of two or more nonmetals!!!
• Diatomic Molecule: a molecule consisting of two identical atoms
Properties of Molecular Compounds
• Lower melting and boiling points than ionic compounds.
CaCl2
• Most are gases or liquids at room temperature.
CO2
• Atoms are attached by more than just electrical attraction.
CaSO4
H2O2
Why, No metals?
Mg3(PO4)2
NaBr
1
Chapter 8 Notes
Molecular Formulas
•
Molecular compounds are made of molecules, not IONS!
•
Ionic compounds are expressed as formula units, not molecules.
• Chemical formula for molecular compounds.
• Shows how many atoms of each element a molecule contains.
•
A molecular formula is the chemical formula of a molecular compound.
• Subscripts are not always lowest whole number ratios. > (No simplification)
• The chemical formulas of covalent compounds are correctly described as molecular formulas
• Does not give the structure of the molecule.
Ionic vs. Covalent
8.2 The Nature of Covalent Bonding
Formula Unit
Molecule
Transfer electrons
Share electrons
Metal ­ Nonmetal
Nonmetal­Nonmetal
Solid Crystals
Solid, liquid, gas
Good electrical conductor
Poor electrical conductor
High melting point
Low melting point
Element
Electron Distribution (Show Boxes)
1s2 2s2 2p4 Oxygen
1s2 2s2 2p3 Nitrogen
1s2 2s2 2p2 Carbon
Dot
Structure
Electrons needed
Unpaired Eelctrons
• A single covalent bond is formed when a pair of electrons is shared between two atoms.
• Electron configurations are slightly different when atoms form covalent bonds. • Remember, a covalent bond is formed by the unpaired electrons in two atoms.
• For example, Carbon needs to form 4 bonds with Hydrogen. So it must have 4 half­filled orbitals instead of the neutral electron configuration. • 1s2 2s1 2p3 , not 1s2 2s2 2p2
1s2 2s1 2p3 Carbon
2
Chapter 8 Notes
Gilbert Lewis Stated...
• Sharing of electrons occurs if the atoms involved acquire the electron configurations of noble gases.
• Become stable by sharing.
Single Covalent Bond
• One shared pair of electrons.
• Each atom donates 1 electron to the bond.
• Represented by 1 dash.
H H F F
Bonding Rules
F F
Shared Pair
Bonding Rules
Nitrogen: 3 unpaired electrons
needs 3 electrons to be stable
must form 3 covalent bonds
Fluorine: 1 unpaired electron
needs 1 electron to be stable
must form 1 covalent bond
Carbon: 4 unpaired electrons
needs 4 electrons to be stable
must form 4 covalent bonds
Oxygen: 2 unpaired electrons
needs 2 electrons to be stable
must form 2 covalent bonds
Bonding Rules
Hydrogen: 1 unpaired electron
needs 1 electron to be stable
must form 1 covalent bond
Chlorine: 1 unpaired electrons
needs 1 electron to be stable
must form 1 covalent bonds
3
Chapter 8 Notes
Structural formulas show the arrangement of atoms in molecules and polyatomic ions. Chlorine bonding to Chlorine
Dot Formula
Structural Formula
Dashes are used
• 1 dash: 2 shared electrons
• 2 dashes: 4 shared electrons
• 3 dashes: 6 shared electrons
Dot Formula
Structural Formula
Cl Cl
Cl Cl
How many electrons are donated by each chlorine? _____
How many unshared pairs are in the molecule? _____
How many electrons are being shared? _____
How many shared pairs are in the molecule? _____
• The pairs of valence electrons that are not shared between atoms are called unshared pairs of electrons, or unshared pairs. • They are also called lone pairs or nonbonding pairs.
Double and Triple Covalent Bonds
• Double covalent bonds involve two shared pairs of electrons. • Represented by 2 dashes
• Triple covalent bonds involve three shared pairs of electrons.
• Represented by 3 dashes
4
Chapter 8 Notes
Oxygen bonding to Oxygen
Structural Formula
Dot Formula
Dot Formula
Structural Formula
O O
O O
How many electrons are donated by each oxygen? _____
How many unshared pairs are in the molecule? _____
How many electrons are being shared? _____
How many shared pairs are in the molecule? _____
H2O
Dot Formula
Dot Formula
O
Structural Formula
O
H
H
Structural Formula
O
H
H
H
H
How many electrons are donated by each hydrogen? _____
How many electrons are donated by the oxygen? _____
How many unshared pairs are in the molecule? _____
How many electrons are being shared? _____
How many shared pairs are in the molecule? _____
H
CH4
Dot Formula
Structural Formula
H
C H
H
H
H
H
C H H C H
H
H
How many electrons are donated by each hydrogen? _____
How many electrons are donated by the carbon? _____
How many unshared pairs are in the molecule? _____
How many electrons are being shared? _____
How many shared pairs are in the molecule? _____
5
Chapter 8 Notes
Dot Formula
Structural Dot Formula
O
O
H
H
H
H
H
H
Fluidity of Shared Electrons
H
C H
H
H
CH
H
CO2
Dot Formula
Structural Formula
Dot Formula
O C O
Structural Formula
O C O
How many electrons are donated by each oxygen? _____
How many electrons are donated by the carbon? _____
How many unshared pairs are in the molecule? _____
How many electrons are being shared? _____
How many shared pairs are in the molecule? _____
Dot Formula
C2 H4
Structural Formula
Two chemists go into a restaurant. The first one says "I think I'll have an H2O." The second one says "I think I'll have an H2O too" ­­ and he died. 6
Chapter 8 Notes
Why don’t metals usually form covalent bonds?
• Mg has 2 valence electrons. ­ How many covalent bonds must it form to be stable? Why don't metals form covalent bonds?
Mg
K
­ How many electrons does it have to donate?
How many more electrons does each atom need to be stable?______
• How about Aluminum?
How many covalent bonds can each atom form?______
• Bonding of diatomic molecules
• Diatomic molecules are more stable together than apart.
• F, I, N, H, Br, Cl, O
• Examples page 222
• Electron Dot Structures
C 2H 6O
C4H6O2
Bonding of Diatomic Molecules
Coordinate Covalent Bonds
Emergency Bonds
• Carbon monoxide example
(mobile)
• Electrons are “fluid”
• Once formed, they act as normal covalent bonds.
• Polyatomic ion formation.
7
Chapter 8 Notes
Fluidity of Shared Electrons
Bond 1 carbon with 1 oxygen
Carbon is unstable. Only 6 surrounding electrons.
Carbon needs 2 more electrons, but Oxygen is stable.
Carbon is sharing 2 more electrons, but didn’t have to donate any of them.
C=O
C=O
C=O
Oxygen is stable! 8 valence electrons & 2 unshared pairs.
Oxygen lets carbon use 1 of it’s unshared pairs.
Coordinate Covalent Bonds CO
C
O
C
O
C
O
O
Coordinate Covalent Bonds SO
O SO
O S
O S O
O S O
Oxygen is still stable. It donated both electrons being shared in the Coordinate Bond.
• A coordinate covalent bond is formed when one atom contributes both bonding electrons in a covalent bond.
• Arrows are used to indicate a coordinate covalent bond
• Ex.) CO, NH4+, H3O+, SO3, SO42­
C
O S
2
O
O
NH4 +
H NH
H
H+
H+
H N H
H N H
H
H
The unshared pair is now a bond, not an unshared pair.
8
Chapter 8 Notes
Negative Polyatomic Ions
H3O +
H+
O
H
H H
H+
O
H
O
O
H
H
O N
O
O
H
O
H
Resonance
• Resonance structures occur when two or more valid electron dot formulas can be written for a molecule.
• Ex. O3, CO32­
• Same formula, different Structures
9
Chapter 8 Notes
Exceptions to the Octet Rule
• Sometimes it is impossible to write electron dot structures that fulfill the octet rule. Occurs whenever the total number of valence electrons in the species is an odd number or less than eight.
• Only certain metals: Be, Al, B
Exceptions to the Octet Rule
• Some metals do form covalent bonds, but result in a shortage of valence electrons.
• Why is BF3 attracted to NH3?
Exceptions to the octet
F
F B
F
F
F
B F
H N H
H N H
H
H
P
S
10
Chapter 8 Notes
CO
NH4+
Chapter 8 ­ Part 2
Molecular Shapes &
Intermolecular Forces
8.3 VSEPR Theory
VSEPR theory states that because electron pairs repel, molecules adjust their shapes so that the valence electron pairs are as far apart as possible.
More about shapes…
• Molecules are 3 dimensional.
• Molecular shape is effected by unshared pairs of electrons.
• Each shape has a specific bond angle.
VSEPR Theory (cont.)
• Valence Shell Electron Pair Repulsion.
• Bond angles are created by this repulsion of electrons
Bond Angles
•
•
•
•
•
Tetrahedral = 109.5°
Linear = 180°
Bent = 105°
Pyrimad = 107°
Trigonal Planar = 120°
11
Chapter 8 Notes
Common Molecular Shapes
Molecular Shapes
Pyramidal
Bent
H
O
H
Tetrahedral
H N H
H
H
H
C H
H
Linear Triatomic: HCN, CO2
All binary compounds
2 Bonds & 0 unshared pairs
No unshared pairs to bend molecule
Trigonal Planar: BH3, COH2
3 bonds & 0 unshared pair
Bent triatomic: H2O
2 bonds & 2 unshared pair
Unshared pairs bend the molecule
2 unshared pair is bent most
Trigonal Pyramidal: NH3
3 bonds & 1 unshared pair
Tetrahedral: CH4
105
4 bonds & 0 unshared pair
12
Chapter 8 Notes
PH3
CF4
H 2S
AlH3
Ionic vs. Molecular Compounds
Ionic Compounds form solid crystals. Why?
Molecular Compounds form gases, liquids, & solids. Why?
13
Chapter 8 Notes
How many electrons are shared?
Which atom has a greater electronegativity?
Which atom has become more negative?
H
F
H
F
H
F
8.4 Polar Bonds and Molecules
• When the atoms in a bond are the same, the bonding electrons are shared equally and the bond is a nonpolar covalent bond
• Ex. diatomics
• The atom with stronger electronegativity in a polar bond acquires a slightly negative charge. The less electronegative atom acquires a slightly positive charge.
• Ex. HCl, H2O
Polar or Non­Polar
H
Cl
Cl
Cl
• When two different atoms are joined by a covalent bond and the bonding electrons are shared unequally, the bond is a polar covalent bond, or simply a polar bond.
Electronegativity
• Ability of atoms to attract electrons.
• Determines the reactivity and strength of polar covalent bonds.
• HCl: Moderately polar covalent
• HF: Very polar covalent (Reactive)
• See page 177.
14
Chapter 8 Notes
Electronegativity of Atoms
F = 4.0
Br = 2.8
O = 3.5
I = 2.5
N = 3.0
C = 2.5
Cl = 3.0
S = 2.5
Which bond is the most polar?
H
Cl
H
F
H
I
Hydrogen = 2.1
• A molecule that has two poles is called a dipolar molecule, or dipole.
• Not every molecule with polar bonds is itself polar.
Is a water molecule polar? ____
• In a polar molecule one end of the molecule is slightly negative and the other is slightly positive.
• Dipolar molecules
• Ex.) HCl,H2O, HF
Is a CH4 molecule polar? _____
H
O
H
Polar Molecules
H
H C H
H
15
Chapter 8 Notes
Non­Polar Molecules
Is a CO2 molecule polar? _____
• When a molecule has no difference in charge between opposite ends or sides of the molecule.
• Not very reactive!
• H2, F2, CO2, Cl2, CCl4
O C O
• Water is only polar due to it’s shape
Attractions Between Molecules
• In addition to covalent bonds in molecules, there are attractions between molecules, or intermolecular attractions
Gases
Liquids
Solids
No attraction
Dipole Attraction
Ionic Attraction
Nonpolar Molecules
Polar Molecules
Ions Form Crystals • Covalently bonded atoms attracted to each other.
Nonpolar molecules are usually gases!
O
(Between)
Intermolecular Attractions
C
O
H H
O
O
H
O
H
O
O O
O
H
O
H
O
O
C
• Hold molecules together.
• Weaker than either an ionic or covalent bond.
• They are responsible for whether a molecular compound is a gas, liquid, or solid.
• Intermolecular attractions
O
H
O
O C
O
H
C
O
O
C
O
O
H H
16
Chapter 8 Notes
Dispersion Forces
Van der Waals forces
• The weakest attractions between molecules. Not Bonds!!!!!!
• Three types are Dispersion forces, Dipole interations, and Hydrogen bonds
• Hydrogen > Dipole > Dispersion
• The weakest of all intermolecular interactions.
• Thought to be caused by the motion of electrons.
• Strength of dispersion forces increases as the number of electrons in a molecule increases
• Electrons are not lost or gained
Attractions between polarized molecules
Dispersion forces
Due to movement, the electrons move to one side and create a separation of charge.
Dipole Interactions
•
•
•
A covalent bond with a dipole.
A cation attracted to a dipole.
A dipole attracted to a dipole
Most dipoles involve hydrogen.
Occur when polar molecules are attracted to one another
Electrostatic attractions occur between the oppositely charged regions of dipolar molecules.
Similar to ionic bonding, but much weaker attraction.
Hydrogen Bonds
• Strongest of all intermolecular attractions. (Must involve hydrogen!)
• Dipole interactions with hydrogen.
• An atom or molecule is attracted to a Hydrogen atom that is already bonded to an atom with high electronegativity.
17
Chapter 8 Notes
Hydrogen Bonds (cont.)
Hydrogen Bonds (cont.)
• The covalently bonded hydrogen becomes slightly positive.
• Unshared electron pairs and atoms with high electronegativity become attracted to the slightly(+) Hydrogen.
Hydrogen Bonding in Water
Why is there so much water?
Water molecules are polar.
Attraction
Hydrogen Bonding is the attraction between polar molecules with hydrogen.
Properties of Molecular Substances
• The physical properties of a compound depend on the type of bonding it displays.
• Ionic or Covalent
• The oxygen atom becomes slightly negative and each hydrogen becomes slightly positive.
• This causes an intermolecular attraction between water molecules.
• The attraction water molecules have for one another is called Hydrogen bonding.
• Network Solid: All of the atoms are covalently bonded to each other. (Crystals)
• No intermolecular attractions.
• Most stable type of molecule.
• Very high melting point.
• Ex.) Diamonds
18
Chapter 8 Notes
Organic Compounds
All Carbon containing compounds
Except carbon oxides, carbides, and carbonates which are inorganic.
Hydrocarbons
Simple organic compounds
Alkanes
Hydrocarbons that have only single bonds between atoms.
Carbons are saturated with Hydrogen atoms
Contain carbon and hydrogen
Carbon forms 4 bonds and hydrogen 1 bond
19
Chapter 8 Notes
Alkenes
Alkynes
• Unsaturated Hydrocarbons that have one or more double bonds between carbon atoms.
• Unsaturated Hydrocarbons that have one or more triple bonds between carbon atoms.
• Carbons are unsaturated with Hydrogen atoms
• Carbons are unsaturated with Hydrogen atoms
Isomers
Two or more compounds that have the same molecular formula but different molecular structure.
Structural Isomers
Two or more compounds that have the same molecular formula but are bonded in a completely different order, therefore changing its properties.
C 3H 8O
End of Part 2 ­ Intermolecular Forces
20