1 chapter 8 covalent bonding ball-and-stick model of epinephrine (adrenaline) key: c o h n

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Chapter 8

Covalent Bonding

Ball-and-stick model of epinephrine (adrenaline)

Key: C O H N

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8.1 Molecular Compounds

OBJECTIVES:– Distinguish between the melting

points and boiling points of molecular compounds and ionic compounds

– Describe the information provided by a molecular formula

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Bonds are…Bonds are… Forces that hold groups of atoms Forces that hold groups of atoms

together and make them function as together and make them function as a unit. Two types:a unit. Two types:

1)1) Ionic bondsIonic bonds – – transfertransfer of electrons results in of electrons results in oppositely charged particles that attract one another. oppositely charged particles that attract one another. a)a) Ionic solids are crystals with regular Ionic solids are crystals with regular

arrangements of +/- ionsarrangements of +/- ionsb)b) In solutions, cations and anions exist separately In solutions, cations and anions exist separately

– there are– there are NO molecules NO molecules!!2)2) Covalent bondsCovalent bonds – – sharingsharing of electrons. The of electrons. The

resulting particle is called a resulting particle is called a moleculemolecule = = individual individual units of a covalent compoundunits of a covalent compound

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Sodium Chloride Crystal LatticeSodium Chloride Crystal Lattice

•Ionic compounds Ionic compounds organize in a organize in a characteristic characteristic crystal latticecrystal lattice of of alternating alternating positive and positive and negative ions, negative ions, repeated over and repeated over and over.over.

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MoleculesMolecules Many elements found in nature

are in the form of molecules: a neutral group of atoms joined

together by covalent bonds. For example, air contains oxygen For example, air contains oxygen

molecules, consisting of two molecules, consisting of two oxygen atoms joined covalentlyoxygen atoms joined covalently

Called a “Called a “diatomicdiatomic molecule molecule” (O” (O22))

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Formation of a Covalent Bond(Video: Formation of H2)

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Covalent Bonds

Nonmetals hold on to their valence electrons.

They acquire a stable, noble gas configuration by sharing valence electrons with each other = covalent bonding

By sharing, both atoms “count” the shared electrons toward a noble gas configuration

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Covalent Bonding

Fluorine has seven valence electrons – (but would more stable with 8)

F

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Covalent bonding Fluorine has seven valence

electrons A second F atom with seven…

F F

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electrons A second atom also has seven If they share a pair of electrons…

F F

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electrons A second atom also has seven By sharing electrons…

F F

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F F

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F F

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F F

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Covalent bonding Fluorine has seven valence electrons A second atom also has seven By sharing electrons… …both end up with full valence

shells

F F

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Covalent bonding Fluorine has seven valence electrons A second atom also has seven By sharing electrons… …both end up with full valence shells

F F8 Valence electrons

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Covalent bonding Fluorine has seven valence electrons A second atom also has seven By sharing electrons… …both end up with full valence shells

F F8 Valence electrons

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Molecular Compounds Compounds that are bonded covalently (like

H2O, CO2) are called molecular compounds Molecular compounds tend to have relatively

lower melting points than ionic compounds Ions are held together in a crystal solid by

relatively strong ionic bonds Covalent bonds are very strong but this is an

intramolecular force It’s the forces between molecules that holds a

solid together (intermolecular forces) and these are often rather weak

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Molecular Compounds

Thus, molecular compounds are often gases or liquids at room temperature

– Ionic compounds are solids A molecular compound has a molecular

formula:

– Shows how many atoms of each element a molecule contains

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Molecular Compounds The formula for water is H2O

–The subscript “2” after hydrogen means there are 2 atoms of hydrogen; if there is only one atom, the subscript (1) is omitted

Molecular formulas do not tell any information about the structure (the arrangement of the various atoms).

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These are some of the different ways to represent ammonia:

1. The molecular formula shows how many atoms of each element are present

2. The structural formula ALSO shows the arrangement of these atoms!

3. The ball and stick model is the BEST, because it shows a 3-dimensional arrangement.

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8.2 Nature of Covalent Bonding

OBJECTIVES– Describe how electrons are shared to

form covalent bonds, and identify exceptions to the octet rule

– Demonstrate how electron dot structures represent shared electrons

– Describe how atoms form double or triple covalent bonds

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8.2 Nature of Covalent Bonding

OBJECTIVES:

–Distinguish between a covalent bond and a coordinate covalent bond, and describe how the strength of a covalent bond is related to its bond dissociation energy

–Describe how oxygen atoms are bonded in ozone

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A Single Covalent Bond is...

A sharing of two valence electrons. Only nonmetals (including H). Different from ionic bonds because

discrete molecules are formed Two atoms are joined

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Sodium Chloride Crystal LatticeSodium Chloride Crystal Lattice

•Ionic compounds Ionic compounds organize in a organize in a characteristic characteristic crystal latticecrystal lattice of of alternating alternating positive and positive and negative ions, negative ions, repeated over and repeated over and over.over.

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How to show the formation… It’s like a jigsaw puzzle. You put the pieces together to end up

with the right formula. Carbon is a special example – it can

share 4 electrons: 1s22s22p2? Let’s show how water is formed with

covalent bonds, by using an electron dot diagram…

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Water

H

O

Each hydrogen has 1 valence electron- Each hydrogen would be

stable if it had 1 more The oxygen has 6 valence

electrons- Each would be stable with

2 more electrons They share to give each other

complete octets

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Water Put the pieces together The first hydrogen is stable The oxygen still needs one more

H O

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Water So, a second hydrogen attaches Every atom has full energy levels

H OH

Note the “unshared” pairs of electrons (also called lone pairs)

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Examples:

1. Conceptual Problem 8.1 on page 220

2. Do PCl3

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Multiple Bonds Sometimes atoms share more than

one pair of valence electrons. A double bond is when atoms share

two pairs of electrons (4 total) A triple bond is when atoms share

three pairs of electrons (6 total) Table 8.1, p.222 - Know these 7

elements as diatomic:

Br2 I2 N2 Cl2 H2 O2 F2 What’s the deal with the oxygen dot diagram?

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Dot diagram for Carbon dioxide CO2 - Carbon is central

atom ( more metallic ) Carbon has 4 valence

electrons

– 4 more to be stable Oxygen has 6 valence

electrons

– 2 more to be stable

O

C

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Carbon dioxide Attaching 1 oxygen leaves the

oxygen 1 short, and the carbon 3 short

OC

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Carbon dioxide Attaching the second oxygen

leaves both of the oxygens 1 short, and the carbon 2 short

OCO

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Carbon dioxide The only solution is to share more

OCO

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OCO

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OCO

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OCO

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Carbon dioxide The only solution is to share more Requires two double bonds all the electrons in each double

bond “count” for each atom

OCO

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Carbon dioxide The only solution is to share more Requires two double bonds Each atom can count all the electrons in

the bond

OCO8 valence electrons

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the bond

OCO8 valence electrons

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the bond

OCO

8 valence electrons

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Covalent BondingNote how the formation of covalent bonds fills the ½ full s and p orbitals of the

bonding atoms.

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How to draw Lewis Structures?1) Add up all the valence electrons.2) Count up the total number of

electrons needed to give all the atoms a noble gas configuration.

3) Subtract #2 - 1; then Divide by 24) Tells you how many bonds to draw5) Fill in the rest of the valence

electrons to “fill up” the atoms’ valence shells.

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Example NH3, which is ammonia N – central atom; has 5

valence electrons, needs 8 H - has 1 (x3) valence

electrons, needs 2 (x3) NH3 has 5+3 = 8

NH3 needs 8+6 = 14 (14-8)/2= 3 bonds 4 atoms with 3 bonds

N

H

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N HHH

Examples Draw in the bonds; start with singles All 8 electrons are accounted for Everything is full – done with this one.

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Example: HCN HCN: C is central atom N - has 5 valence electrons, needs 8 C - has 4 valence electrons, needs 8 H - has 1 valence electron, needs 2 HCN has 5+4+1 = 10

HCN needs 8+8+2 = 18

(18-10)/2= 4 bonds 3 atoms with 4 bonds – this will require

multiple bonds - not to H though

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HCN Put single bond between each atom Need to add 2 more bonds Must go between C and N (Hydrogen is full)

NH C

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HCN Put in single bonds Must have 2 more bonds Must go between C and N, not the H Uses 8 electrons – need 2 more to

equal the 10 it has

NH C

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HCN Put in single bonds Need 2 more bonds Must go between C and N Uses 8 electrons - 2 more to add Must go on the N to fill its octet

NH C

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Another way of indicating bonds

Often use a line to indicate a bond Called a structural formula Each line is 2 valence electrons

H HO = H HO

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Other Structural Examples

H C N

C OH

H

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A Coordinate Covalent Bond... When one atom donates both

electrons in a covalent bond. Carbon monoxide (CO) is a good

example:

OCBoth the carbon and oxygen give another single electron to share

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Coordinate Covalent Bond When one atom donates both

electrons in a covalent bond. Carbon monoxide (CO) is a good

example:

OC

Oxygen gives both of these electrons, since it has no more singles to share.

These two electrons make a lone pair.

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Coordinate Covalent Bond When one atom donates both

electrons in a covalent bond. Carbon monoxide (CO)

OCC O

The coordinate covalent bond is shown with an arrow as:

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Coordinate covalent bondMost polyatomic cations and

anions contain covalent and coordinate covalent bonds

Table 8.2, p.224Sample Problem 8.2, p.225 The ammonium ion (NH4

1+) can be shown as another example

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Bond Dissociation Energies...

The total energy required to break the bond between 2 covalently bonded atoms

High dissociation energy usually means the chemical is relatively

unreactive, because it takes a lot of energy to break it down.

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Resonance is... When more than one valid dot

diagram is possible. Consider the two ways to draw ozone

(O3) Which one is it? Does it go back and

forth? It is a hybrid of both, like a mule; and

shown by a double-headed arrow found in double-bond structures!

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Resonance in OzoneResonance in Ozone

Neither structure is correct, it is actually a hybrid of the two. To show it, draw all varieties possible, and join them with a double-headed arrow.

Note the different location of the double bond

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ResonanceResonanceOccurs when more than one valid Lewis structure can be written for a particular molecule (due to position of double bond)

•These are resonance structures of benzene.•The actual structure is an average (or hybrid) of these structures.

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Resonance in a carbonate ion (CO3

2-):

Resonance in an acetate ion (C2H3O2

1-):

Polyatomic ions – note the different positions of the double bond.

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The 3 Exceptions to Octet rule For some molecules, it is

impossible to satisfy the octet rule#1. usually when there is an odd

number of valence electrons–NO2 has 17 valence electrons,

because the N has 5, and each O contributes 6. Note “N” page 228

It is impossible to satisfy octet rule, yet the stable molecule does exist

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Exceptions to Octet rule• Another exception: Boron

• Page 228 shows boron trifluoride, and note that one of the fluorides might be able to make a coordinate covalent bond to fulfill the boron

• #2 -But fluorine has a high electronegativity (it is greedy), so this coordinate bond does not form

• #3 -Top page 229 examples exist because they are in period 3 or beyond

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8.4 Polar Bonds & Molecules OBJECTIVES:

– Describe how electronegativity values determine the distribution of charge in a polar molecule

– Describe what happens to polar molecules when they are placed between oppositely charged metal plates

– Evaluate the strength of intermolecular attractions compared with the strength of ionic and covalent bonds

– Identify the reason why network solids have high melting points

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Bond Polarity Covalent bonding means shared

electrons–but, do they share equally?

Electrons are pulled, as in a tug-of-war, between the atoms’ nuclei–In equal sharing (such as

diatomic molecules), the bond that results is called a nonpolar covalent bond

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Bond Polarity When two different atoms bond

covalently, there is almost always an unequal sharing of the bonding e-

– the more electronegative atom has a stronger attraction for the bonding e-

– Thus the bonding pair of e- spend most time near the more electronegative atom acquire a partial (−) charge

– called a polar covalent bond, or simply polar bond.

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Electronegativity?Electronegativity?

The ability of an The ability of an atom in a molecule atom in a molecule to attract shared to attract shared electrons to itself.electrons to itself.

Linus Pauling1901 - 1994

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Table of Electronegativities

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Bond Polarity Consider HCl

H: electroneg = 2.1

Cl: electroneg = 3.0

– the bond is polar

– the chlorine acquires a slight negative charge, and the hydrogen a slight positive charge

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Bond Polarity Only partial charges, less than a

true 1+ or 1− as in ionic bond Written as:

HCl the positive and minus signs (with

the lower case delta: ) denote partial charges.

and

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Bond Polarity Can also be shown:

*arrow points to more electronegative atom

H Cl

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Polar molecules Sample Problem 8.3, p.239 A polar bond may make the entire

molecule “polar”

– Depends on shape of molecule HCl has polar bonds, and one side is

partially +, the other partially –, so the whole molecule is polar

– A molecule like HCl that has two poles is said to have a dipole

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Practice Problems

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Polar molecules

The effect of polar bonds on the polarity of the entire molecule depends on the molecule shape

–carbon dioxide has two polar bonds, and is linear = nonpolar molecule!

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Polar molecules The effect of polar bonds on the

polarity of the entire molecule depends on the molecule shape– water has two polar bonds and a bent

shape; the highly electronegative oxygen pulls the e- away from H’s = very polar!

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Polar molecules When polar molecules are placed between oppositely charged

plates, they tend to become oriented with respect to the positive and negative plates.

With current off With current on

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Intermolecular Forces = attractions between molecules They are what determine if a

compound will exist as a solid liquid or gas at room temperature

The weakest are called van der Waal’s forces - there are two kinds:

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Van der Waal’s Forces

#1. Dispersion forces• Weakest, caused by random motions of e- • Strength increases as # e- increases (size

of atom/molecule)• Exist in all compounds

halogens F and Cl are gases; bromine is liquid; iodine is solid – (Group 7A)

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Dispersion Forces Animation

Link to Dispersion Forces Animation

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Van der Waal’s Forces #2 Dipole Interactions Occur when polar molecules attract each other Usually, slightly stronger than dispersion forces

++ −−

Dipole-dipole attraction (weak)

Opposites attract, but these are only partial charges (), so attractions NOT as strong as ionic bonds

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#3. Hydrogen bonding …is the attractive force between a strongly + H

atom (polar bonded to a N, O, F), and a lone pair of electrons on a N, O, or F atom of another molecule.

N, O, F, are very electronegative (so strong – charge)Strongest type of dipole attraction

The H atom “shares” a lone pair of electrons with an O, N, or F atom in another molecule

On average, the strongest of the 3 IM forces.

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#3. Hydrogen bonding defined: When a hydrogen atom is: a) covalently bonded to a highly electronegative

atom (O, N, or F), AND b) is also weakly bonded to an unshared electron

pair of a nearby highly electronegative atom– The hydrogen is very electron deficient (its

lone electron is being shared in a covalent bond with a highly electronegative atom)

• so it is strongly attracted to electrons on electronegative atoms on other molecules

– Hydrogen atom has no shielding for its nucleus, so that proton is strongly attracted to any nearby electrons

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Hydrogen Bonding

HH

O+ -

+

H HO+-

+

This H is bonded covalently to the O in the same molecule, and H-bonded to a nearby unshared pair on the O atom of another molecule.

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H-bonding causes H2O to

be a liquid at room

conditions

H-bonds form between H2O and NH3 molecules

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Order of Strength: IM ForcesOn average:

Dispersion forces are weakest Dipole interactions a little stronger Hydrogen bonding strongest

All 3 are weaker than ionic or covalent bonds All are variable: not all dispersion forces alike, not

all H-bonds the same strength For example: dispersion forces are weakest, but

in large molecules, they can be significantly stronger than the other IM forces of smaller molecules

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Attractions and properties Why are some chemicals gases, some

liquids, some solids at room temperature?

– Depends on the type of bonding!

Characteristic Ionic Compounds Covalent Compounds

Representative Unit Formula unit Molecule

Bond Formation Electrical attraction between oppositely charged ions

Attraction between atoms for shared electrons

Type of Elements Metals and Non-metals Nonmetals

Usual Physical State Solid Solid, liquid or gas

Forces Holding Particles Together in a Solid (IM Forces)

Ionic BondsDispersion forces, dipole-dipole, H-bond, covalent bond

Melting Point High (usually > 300ºC) Low (usually < 300ºC)

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Attractions and properties Network covalent solids – solids

in which all the atoms are covalently bonded to each other

The force holding the particles together in a solid (IM force) is the strongest type of chemical bond!

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Attractions and properties Network covalent solids melt at

very high temperatures, or not at all (decompose instead)

–Diamond does not really melt, but vaporizes to a gas at 3500oC

–SiC, used in grinding, has a melting point of about 2700oC

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Covalent Network CompoundsCovalent Network CompoundsSome covalently bonded substances DO NOT form discrete molecules.

Diamond, a network of covalently bonded carbon atoms

Graphite, a network of covalently bonded sheets of carbon atoms

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8.4 Polar Bonds & Molecules OBJECTIVES (Accomplished??)

– Describe how electronegativity values determine the distribution of charge in a polar molecule

– Describe what happens to polar molecules when they are placed between oppositely charged metal plates

– Evaluate the strength of intermolecular attractions compared with the strength of ionic and covalent bonds

– Identify the reason why network solids have high melting points

1 chapter 8 covalent bonding ball-and-stick model of epinephrine (adrenaline) key: c o h n

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