chapter 8 notes -...

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

CaSO4H2O2

Mg3(PO4)2NaBr

Properties of Molecular Compounds

Why, No metals?

• 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

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• 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.

• The chemical formulas of covalent compounds are correctly described as 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 of the molecule.

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)Dot

StructureElectrons needed

Unpaired Eelctrons

Oxygen 1s2 2s2 2p4

Nitrogen 1s2 2s2 2p3

Carbon 1s2 2s2 2p2

Carbon 1s2 2s1 2p3

• 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

Chapter 8 Notes

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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 FBonding Rules

Carbon: 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

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

Chapter 8 Notes

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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.

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

Chapter 8 Notes

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Oxygen bonding to OxygenDot Formula Structural Formula


O OOOHow many electrons are donated by each oxygen? _____




H2ODot Formula Structural Formula


HHO

HHO

HHO

How many electrons are donated by each hydrogen? _____

How many electrons are donated by the oxygen? _____




CH4Dot Formula Structural Formula HH C

H

HHH C

H

H

HH CH

HHow many electrons are donated by each hydrogen? _____

How many electrons are donated by the carbon? _____




Chapter 8 Notes

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HH CH

H

HHO

HHO

HH CH

H

Dot Formula Structural Dot Formula Fluidity of Shared Electrons

CO2Dot Formula Structural Formula


OO COO CHow many electrons are donated by each oxygen? _____

How many electrons are donated by the carbon? _____


How many electrons are being shared? _____How many shared pairs are in the molecule? _____

C2H4Dot Formula 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.

Chapter 8 Notes

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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

Bonding of Diatomic Molecules

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

Chapter 8 Notes

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Coordinate Covalent Bonds CO

C O

C O

C O

C 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.

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, SO42

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+

Chapter 8 Notes

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Negative Polyatomic Ions

O H

NO

O O 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

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

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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 FF

F

N HHH

B F

N HHH

F

F

P SCO NH4+

Chapter 8 Notes

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Chapter 8 Part 2Molecular Shapes &

Intermolecular Forces

8.3 VSEPR TheoryVSEPR theory states that because electron pairs repel, molecules adjust their shapes so that the valence electron pairs are as far apart as possible.

VSEPR Theory (cont.)• Valence Shell Electron Pair Repulsion.• Bond angles are created by this

repulsion of electrons

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

Chapter 8 Notes

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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 & 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

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PH3 CF4 H2S AlH3

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?

Chapter 8 Notes

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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.

• 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

Chapter 8 Notes

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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

Is a water molecule polar? ____

HHO

Is a CH4 molecule polar? _____

HH CH

H

Is a CO2 molecule polar? _____

OO C

Chapter 8 Notes

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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

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

Chapter 8 Notes

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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

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 theoppositely 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.

Chapter 8 Notes

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Hydrogen Bonds (cont.) 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

Chapter 8 Notes

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Organic CompoundsAll Carbon containing compounds

Except carbon oxides, carbides, and carbonates which are inorganic.

HydrocarbonsSimple organic compounds

Contain carbon and hydrogen

Carbon forms 4 bonds and hydrogen 1 bond

AlkanesHydrocarbons that have only single bonds between atoms.

Carbons are saturated with Hydrogen atoms

Alkenes• Unsaturated Hydrocarbons that have one or more double bonds between carbon atoms.

• Carbons are unsaturated with Hydrogen atoms

Chapter 8 Notes

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Alkynes• Unsaturated Hydrocarbons that have one or more triple bonds between carbon atoms.

• Carbons are unsaturated with Hydrogen atoms

IsomersTwo or more compounds that have the same molecular formula but different molecular structure.

Structural IsomersTwo or more compounds that have the same molecular formula but are bonded in a completely different order, therefore changing its properties.

C3H8O

End of Part 2 Intermolecular Forces

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