Upload
others
View
1
Download
0
Embed Size (px)
Citation preview
Chapter 8 Notes
1
Chapter 8Molecular Compounds
&Covalent Bonding
Why do covalent bonds form?• If only group 5A, 6A, 7A atoms existed, ionic bonds can’t form.• Each atom needs electrons so they are not willing to lose any.• If two Hydrogen atoms are locked in a room together, what happens?
NONMETALS
H HBoth H atoms have 1 unpaired
electron
H H H HThe electrons pair up. Covalent Bond formed.
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
Does it contain ionic or covalent bonds?Formula Ionic or Covalent ExplanationCaCl2CO2
CaSO4
H2O2
Mg3(PO4)2NaBr
Properties of Molecular Compounds• Lower melting and boiling points than ionic compounds.• Most are gases or liquids at room temperature.• Atoms are attached by more than just electrical attraction.
Chapter 8 Notes
2
• Molecular compounds are made of molecules, not IONS!
• Ionic compounds are expressed as formula units, not molecules.
• A molecular formula is the chemical formula of a molecular compound.
Molecular Formulas• Chemical formula for molecular compounds.• Shows how many atoms of each element a molecule contains.• Subscripts are not always lowest whole number ratios. (No simplification)• Does not give the structure.
• The chemical formulas of covalent compounds are correctly described as molecular formulas
Ionic vs. CovalentFormula 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
8.2 The Nature of Covalent Bonding
• A single covalent bond is formed when a pair of electrons is shared between two atoms.
Element Electron Distribution (Show Boxes)
DotStructure
Electrons needed
Unpaired Eelctrons
Oxygen 1s2 2s2 2p4
Nitrogen 1s2 2s2 2p3
Carbon 1s2 2s2 2p2
Carbon 1s2 2s1 2p3
Chapter 8 Notes
3
• 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.
• 1s22s12p3, not 1s22s2 2p2
• 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 Shared Pair
F F
Bonding RulesCarbon: 4 unpaired electronsneeds 4 electrons to be stablemust form 4 covalent bonds
Oxygen: 2 unpaired electronsneeds 2 electrons to be stablemust form 2 covalent bonds
Bonding RulesNitrogen: 3 unpaired electrons needs 3 electrons to be stable must form 3 covalent bonds
Fluorine: 1 unpaired electronneeds 1 electron to be stablemust form 1 covalent bond
Chapter 8 Notes
4
Bonding RulesHydrogen: 1 unpaired electronneeds 1 electron to be stablemust form 1 covalent bond
Chlorine: 1 unpaired electronsneeds 1 electron to be stablemust form 1 covalent bonds
Structural formulas show the arrangement of atoms in molecules and polyatomic ions.
Dashes are used• 1 dash: 2 shared electrons• 2 dashes: 4 shared electrons• 3 dashes: 6 shared electrons
Chlorine bonding to ChlorineDot Formula Structural Formula
Dot Formula Structural Formula
Cl ClClClHow 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.
Chapter 8 Notes
5
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
Oxygen bonding to OxygenDot Formula Structural Formula
Dot Formula Structural Formula
O OOOHow 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? _____
H2ODot Formula Structural Formula
Dot Formula Structural Formula
HHO
HHO
HHO
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? _____
CH4Dot Formula Structural Formula
Chapter 8 Notes
6
HH CH
HHH C
H
H
HH CH
HHow 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? _____
HH CH
H
HHO
HHO
HH CH
H
Dot Formula Structural Dot Formula
Fluidity of Shared Electrons CO2Dot Formula Structural Formula
Dot Formula Structural Formula
OO COO CHow 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? _____
C2H4Dot Formula Structural Formula
Chapter 8 Notes
7
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.
Why don’t metals usually form covalent bonds?
• Mg has 2 valence electrons. How many covalent bonds must it form to be stable?
How many electrons does it have to donate?• How about Aluminum?
Why don't metals form covalent bonds?
Mg KHow many more electrons does each atom need to be stable?______
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
Diatomic Molecules C2H6O C4H6O2
Chapter 8 Notes
8
Bonding Trick• Total electrons all atoms need to be stable? • (CH4 = 16)• Total valence electrons that all atoms have? • (CH4 = 8)• Subtract the number they have from the number needed.• 16 – 8 = 8• Divide that number by 2. (8 ÷2 = 4)• CH4 needs to form 4 covalent bonds to be stable.
Coordinate Covalent BondsEmergency Bonds
• Carbon monoxide example• Electrons are “fluid”• Once formed, they act as normal
covalent bonds.• Polyatomic ion formation.
(mobile)
Fluidity of Shared Electrons C O
C OC O
Bond 1 carbon with 1 oxygen
C=OC=OC=O
Carbon is unstable. Only 6 surrounding electrons.
Oxygen is stable! 8 valence electrons & 2 unshared pairs.
Carbon needs 2 more electrons, but Oxygen is stable.
Oxygen lets carbon use 1 of it’s unshared pairs.
Oxygen is still stable. It donated both electrons being shared in the Coordinate Bond.
Carbon is sharing 2 more electrons, but didn’t have to donate any of them.
SO O
SO O
SO O
Chapter 8 Notes
9
Coordinate Covalent Bonds SO 2
SO O SO O
SO O SO OSO O
• 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, SO4
2
NH4 +
N HHH N HH
H
H+
N HHH
H+
The unshared pair is now a bond, not an unshared pair.
H3O +
HHO
HHO
H+
HHO
H+
Negative Polyatomic Ions
O H
NO
O O O H
Molecular Formula
Dot Formula
Structural Formula
UnsharedPairs
3 from Homework and notes
1 Coordinate Covalent(CO, SO2, BF3)
1 Polyatomic Ion(H3O+, NH4
+, OH, CN)
Homework Quiz 8.2
Chapter 8 Notes
10
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
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
Chapter 8 Notes
11
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
B F
F
F
N HHH
B F
N HHH
F
F
• Diamagnetic: when all of the electrons are paired (Magnetic effects cancel)
• Paramagnetic: contain one or more unpaired electrons (Strong attraction to a magnetic field)
• Ferromagnetism: attraction of iron, cobalt, and nickel for magnetic field
Very similar to paramagnetic, but stronger
Diamagnetism
• Each orbital has an electron pair. • The electrons are always pulling in the opposite direction.
• These elements are not magnetic!!
Paramagnetism
• All electrons are pointing in the same direction.• They are free to move in one direction.
• Strong magnetism
Chapter 8 Notes
12
Expanding the Octet• when atoms are surrounded by more than 8
electrons• Some atoms have empty "d" orbitals that can be
used to hold extra electrons.
• Ex.) PCl5, SF6
P S
CO NH4+
8.3 VSEPR Theory• VSEPR theory states that because electron pairs repel, molecules adjust their shapes so that the valenceelectron pairs are as far apart as possible.
8.3 VSEPR Theory (cont.)• Valence Shell Electron Pair Repulsion.• Bond angles are created by this
repulsion of electrons
Chapter 8 Notes
13
More about shapes…• Molecules are 3 dimensional.• Molecular shape is effected by
unshared pairs of electrons.• Each shape has a specific bond angle.
Bond Angles• Tetrahedral = 109.5°• Linear = 180°• Bent = 105°• Pyrimad = 107°• Trigonal Planar = 120°
Molecular ShapesBent Pyramidal Tetrahedral
HHO N HH
HHH C
H
H
Common Molecular ShapesLinear Triatomic: HCN, CO2All binary compounds2 Bonds & 0 unshared pairsNo unshared pairs to bend molecule
Trigonal Planar: BH3, COH23 bonds & 0 unshared pair
Bent triatomic: H2O2 bonds & 1 or 2 unshared pairUnshared pairs bend the molecule2 unshared pair is bent most
Trigonal Pyramidal: NH33 bonds & 1 unshared pair
Tetrahedral: CH44 bonds & 0 unshared pair
105
Chapter 8 Notes
14
PH3 CF4 H2S AlH3
Chapter 8 Notes
15
Ionic Compounds form solid crystals. Why?
Molecular Compounds form gases, liquids, & solids. Why?
Ionic vs. Molecular Compounds
H F
H F
H F
How many electrons are shared?
Which atom has a greater electronegativity?
Which atom has become more negative?
Polar or NonPolar
ClH ClCl
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
• 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.
Chapter 8 Notes
16
• 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
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.
Electronegativity of Atoms
F = 4.0 Br = 2.8O = 3.5 I = 2.5N = 3.0 C = 2.5Cl = 3.0 S = 2.5Hydrogen = 2.1
Which bond is the most polar?
ClH FH IH
• A molecule that has two poles is called a dipolar molecule, or dipole.• Not every molecule with polar bonds is itself polar.
Polar Molecules• In a polar molecule one end of the molecule is slightly negative and the other is slightly positive.• Dipolar molecules• Ex.) HCl,H2O, HF
Chapter 8 Notes
17
Is a water molecule polar? ____
HHO
Is a CH4 molecule polar? _____
HH CH
H
Is a CO2 molecule polar? _____
OO C
NonPolar Molecules• When a molecule has no difference in charge between opposite ends or sides of the molecule.• Not very reactive!• H2, F2, CO2, Cl2, CCl4
• Water is only polar due to it’s shape
Molecular Formula Structural Shape Bond Angle Polar or Nonpolar
BF3, AlH3, CH4, NH3, H2O, HCN, H2S, N2
Molecular Formula Dot Structural Unshared Pair
CO, CO2, C2H4, H3O+
C2H2, NH4+, C4H6, C2H6O
Homework Quiz 8.3Which bond is most polar?
Ionic vs. CovalentFormula 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
Chapter 8 Notes
18
Attractions Between Molecules• In addition to covalent bonds in molecules, there
are attractions between molecules, or intermolecular attractions
• Covalently bonded atoms attracted to each other.
Gases
No attraction
Nonpolar Molecules
Liquids
Dipole Attraction
Polar Molecules
Solids
Ionic Attraction
Ions Form Crystals
Nonpolar molecules are usually gases!
O O
H H O O
O O
OO
O OOO
C
OO
C
O
OC
OO C
O
OC H
HHH
H H
HH
Intermolecular Attractions• 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
(Between)
Van der Waals forces• The weakest attractions between molecules. Not Bonds!!!!!!• Three types are Dispersion forces, Dipole interations, and Hydrogen bonds• Hydrogen > Dipole > Dispersion
Attractions between polarized molecules
Dispersion Forces• 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
Chapter 8 Notes
19
Due to movement, the electrons move to one side and create a separation of charge.
Dispersion forces Dipole Interactions• 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.
A covalent bond with a dipole.
A cation attracted to a dipole.
A dipole attracted to a dipole
Most dipoles involve hydrogen.
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.
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 Bonds (cont.)
Chapter 8 Notes
20
Hydrogen Bonding in Water
Attraction Hydrogen Bonding is the attraction between polar molecules with hydrogen.
Why is there so much water?Water molecules are polar.
• 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.
Properties of Molecular Substances
• The physical properties of a compound depend on the type of bonding it displays.
• Ionic or Covalent
• 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
Coordinate Covalent Bonds CO
C O
C O
C O
C O