Chapter 8 Concepts of Chemical Bonding

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Chapter 8Concepts of Chemical

Bonding

Chemistry, The Central Science, 10th editionTheodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten

John D. BookstaverSt. Charles Community College

St. Peters, MO 2006, Prentice Hall, Inc.

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November 15Chapter 8 – Chemical Bonding

Section 1 and 2Bonding typesLewis StructuresIonic bondingHW: 1,3,4,7 to 19 odd ,24,25 There will be a quiz any day on hw

questions

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Chemical Bonds

• Three basic types of bonds:Ionic

• Electrostatic attraction between ions

Covalent• Sharing of electrons

Metallic• Metal atoms bonded to

several other atoms

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Ionic BondingElectron Transfer

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Energetics of Ionic Bonding

As we saw in the last chapter, it takes 495 kJ/mol to remove electrons from sodium.

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We get 349 kJ/mol back by giving electrons to chlorine.

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• But these numbers don’t explain why the reaction of sodium metal and chlorine gas to form sodium chloride is so exothermic!

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• The reaction is violently exothermic.• We infer that the NaCl is more stable than its constituent

elements. Why?• Na has lost an electron to become Na+ and chlorine has

gained the electron to become Cl. Note: Na+ has an Ne electron configuration and Cl has an Ar configuration.

• That is, both Na+ and Cl have an octet of electrons surrounding the central ion.

C:\Documents and Settings\Administrator\My Documents\IB chemistry\IB Powerpoints\IB 2003\Media_Portfolio\FormationofSodiumChloride\FormationofSodiumChloride.MOV

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• There must be a third piece to the puzzle.

• What is as yet unaccounted for is the electrostatic attraction between the newly formed sodium cation and chloride anion.

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Lattice Energy

• This third piece of the puzzle is the lattice energy:The energy required to completely separate a mole of

a solid ionic compound into its gaseous ions.• The energy associated with electrostatic

interactions is governed by Coulomb’s law:

Eel = Q1Q2

d

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• Lattice energy depends on the charges of the ions and the sizes of the ions

is a constant (8.99 x 10 9 J·m/C2), Q1 and Q2 are the charges on the ions, and d is the distance between ions.

Lattice energy increases as* The charges on the ions increase* The distance between the ions decreases.

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• NaCl forms a very regular structure in which each Na+ ion is surrounded by 6 Cl ions.

• Similarly, each Cl ion is surrounded by six Na+ ions.• There is a regular arrangement of Na+ and Cl in 3D.• Note that the ions are packed as closely as possible.• Note that it is not easy to find a molecular formula to

describe the ionic lattice.

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Energetics of Ionic Bond Formation• The formation of Na+(g) and Cl(g) from Na(g) and Cl(g)

is endothermic.• Why is the formation of NaCl(s) exothermic?• The reaction NaCl(s) Na+(g) + Cl(g) is endothermic

(H = +788 kJ/mol).• The formation of a crystal lattice from the ions in the gas

phase is exothermic:Na+(g) + Cl(g) NaCl(s) H = 788 kJ/mol

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By accounting for all three energies (ionization energy, electron affinity, and lattice energy), we can get a good idea of the energetics involved in such a process.

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1. Vaporize the metal (enthalpy of vaporization)

• Na (s) Na (g)

2. Break diatomic nonmetal molecules (if applicable) (bond enthalpy)

• ½ Cl2 (g) Cl-

3. Remove electron(s) from metal (ionization energy)

• Na (g) Na+ (g) + e-

Born-Haber CycleApplication of Hess’s Law

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• 4 Add electron(s) to nonmetal (electron affinity)

Cl (g) + e- Cl- (g)

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• 5 Put ions together to form compound

(lattice energy)• Na+ + Cl- NaCl (s)

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• Overall Reaction:Na (s) + ½ Cl2 (g) NaCl (s)

• This is useful because all quantities are directly measurable except lattice energy. The Born-Haber cycle can be used to calculate lattice energy from the other values.

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Step Energetics1

2

3

4

5

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Step Energetics1 endothermic

2 endothermic

3 endothermic

4 Exothermic

5 Exothermic (highly)

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Lattice Energy

• Lattice energy, then, increases with the charge on the ions.

• It also increases with decreasing size of ions.

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• Examples – Which in each pair would have the greatest lattice energy?1. LiF or LiCl2. NaCl or MgCl23. KBr or KI4. MgCl2 or MgO

5. CaO or NaF

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• These phenomena also helps explain the “octet rule.”

• Metals, for instance, tend to stop losing electrons once they attain a noble gas configuration because energy would be expended that cannot be overcome by lattice energies.

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November 17

• Properties of ionic compounds/ Naming ionic compounds

• COVALENT BONDS• LEWIS STRUCTURES• RESONANCE STRUCTURES• EXCEPTIONS TO OCTECT RULE• STRENGTHS OF COVALENT BONDS

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Properties of Ionic Substances• Crystalline structure, result from electrostatic

forces that keep ions in a rigid well defined three dimensional arrangement.

• hard and brittle.• Represented by empirical formulas• High melting points• Do not conduct electricity in the solid state

but they conduct electricity when are dissolved in water or in the liquid state

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Naming Ionic Compounds

• Binary compounds – ending ide• When metal has several oxidation

states indicate the o.n. with the stock system. Examples

• Lead (IV) oxide • Manganese (III) Nitrate• Iron (II) Phosphate• Calcium Sulfate

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COVALENT BONDING

Sharing of electrons.Typical bond between non metals.Can be polar or non polar.Can be single, double or triple

bonds.

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General properties of substances containing covalent bonds

• They are molecular substances.• Generally they have low melting points.• They exists in the 3 states.• In the solid state many are pliable (like

plastics).• They are not conductors of heat or

electricity.

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Naming Covalent Compounds

• Use IUPAC nomenclature with roman numerals for the oxidation numbers

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Bond Types and Nomenclature• In general, the least electronegative element is named

first.• The name of the more electronegative element ends in

–ide.• Compounds are named according to their ions (or

perceived ions), including the charge on the cation if it is variable.

• Compounds of metals with high O.N. have properties similar to covalent compounds and sometimes are named using the convention for molecular compounds

• Mn2O7 can be called Dimanganese heptoxide or Manganese(VII)oxide

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Bond Types and Nomenclature

Ionic MolecularMgH2 H2SFeF2 OF2

Mn2O3 Cl2O3

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Bond Types and Nomenclature

Ionic MolecularMgH2 Magnesium

hydrideH2S Hydrogen

sulfideFeF2 Iron(II) fluoride OF2 Oxygen (II)

fluorideMn2O3 Manganese(III)

oxideCl2O3 Chlorine (III)

oxide

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Covalent Bonding• In these bonds atoms share

electrons. Pairs of electrons become the glue that bind to atoms together.

• There are several electrostatic interactions in these bonds:Attractions between electrons

and nucleiRepulsions between electronsRepulsions between nuclei

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Polar Covalent Bonds

• Although atoms often form compounds by sharing electrons, the electrons are not always shared equally.

• Fluorine pulls harder on the electrons it shares with hydrogen than hydrogen does.

• Therefore, the fluorine end of the molecule has more electron density than the hydrogen end.

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Electronegativity:

• The ability of atoms in a molecule to attract electrons to itself.

• On the periodic chart, electronegativity increases as you go……from left to right

across a row.…from the bottom to

the top of a column.

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Electronegativity and Bond Polarity

• Difference in electronegativity is a gauge of bond polarity:• electronegativity differences between 0 and 0.5 result in non-

polar covalent bonds (equal or almost equal sharing of electrons);

• electronegativity differences around between 0.5 and 2 result in polar covalent bonds (unequal sharing of electrons);

• electronegativity differences greater that 2 result in ionic bonds (transfer of electrons).

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Non Polar Covalent bondsEqual electron sharing

• Happens between atoms with same electronegativity.

• All diatomic molecules have non polar covalent bonds!

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• When two atoms share electrons unequally, a bond dipole results.

• The dipole moment, , produced by two equal but opposite charges separated by a distance, r, is calculated:

= Qr• It is measured in debyes (D).

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The greater the difference in electronegativity, the more polar is the bond.

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Lewis Structures(Lewis electron-dot structures)

Lewis structures are representations of molecules showing all electrons, bonding and nonbonding.

A line between 2 atoms represents a pair of electrons (2 electrons = 1 line )

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Multiple Bonds• It is possible for more than one pair of electrons to be

shared between two atoms (multiple bonds):• One shared pair of electrons = single bond (e.g. H2);

• Two shared pairs of electrons = double bond (e.g. O2);

• Three shared pairs of electrons = triple bond (e.g. N2).

• Generally, bond distances decrease as we move from single through double to triple bonds.

H H O O N N

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Drawing Lewis StructuresDrawing Lewis Structures

1. Add the valence electrons.2. Write symbols for the atoms and show which atoms are

connected to which. 3. Complete the octet for the central atom, then complete

the octets of the other atoms.4. Place leftover electrons on the central atom.5. If there are not enough electrons to give the central atom

an octet, try multiple bonds.

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Writing Lewis Structures

1. Find the sum of valence electrons of all atoms in the polyatomic ion or molecule. If it is an anion, add one

electron for each negative charge.

If it is a cation, subtract one electron for each positive charge.

PCl3

5 + 3(7) = 26

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2. The central atom is the least electronegative element that isn’t hydrogen. Connect the outer atoms to it by single bonds.

Keep track of the electrons:

26 6 = 20

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3. Fill the octets of the outer atoms.


26 6 = 20 18 = 2

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4. Fill the octet of the central atom.


26 6 = 20 18 = 2 2 = 0

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5. If you run out of electrons before the central atom has an octet…

…form multiple bonds until it does.

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• Examples – Draw Lewis Structures for each:

1. H2O

2. CO2

3. NCl34. SO2

5. SO3

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• PRACTICE LEWIS STRUCTURES• FORMAL CHARGE• EXCEPTIONS TO OCTECT RULE• a) Atoms with fewer than 8 valence e-

• b) Atoms with more than 8 valence e-

• c) Molecules with odd number of e-

• RESONANCE

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Formal Charge• It is possible to draw more than one Lewis structure with

the octet rule obeyed for all the atoms.• To determine which structure is most reasonable, we use

formal charge.• Formal charge is the charge on an atom that it would

have if all the atoms had the same electronegativity.

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• To calculate formal charge: • All nonbonding electrons are assigned to the atom on

which they are found.• Half the bonding electrons are assigned to each atom

in a bond.• Formal charge is:

valence electrons - number of bonds - lone pair electrons

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• To assign formal charges.For each atom, count the electrons in lone pairs and

half the electrons it shares with other atoms.Subtract that from the number of valence electrons for

that atom: The difference is its formal charge.

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• The best Lewis structure……is the one with the fewest charges.…puts a negative charge on the most

electronegative atom.

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Formal Charge• Consider:

• For C: • There are 4 valence electrons (from periodic table).• In the Lewis structure there are 2 nonbonding electrons and 3

from the triple bond. There are 5 electrons from the Lewis structure.

• Formal charge: 4 - 5 = -1.

C N

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Formal Charge• Consider:

• For N:• There are 5 valence electrons.• In the Lewis structure there are 2 nonbonding electrons and 3

from the triple bond. There are 5 electrons from the Lewis structure.

• Formal charge = 5 - 5 = 0.• We write:

C N

C N

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Examples – Determine the formal charge on each atom for each:

1.H2O

2.CO2

3.NCl3

4.SO2

5.SO3

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Formal Charge• The most stable structure has:

• the lowest formal charge on each atom,• the most negative formal charge on the most electronegative

atoms.

Resonance Structures• Some molecules are not well described by Lewis

Structures.• Typically, structures with multiple bonds can have

similar structures with the multiple bonds between different pairs of atoms

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Resonance

This is the Lewis structure we would draw for ozone, O3. -

+

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Resonance

• But this is at odds with the true, observed structure of ozone, in which……both O—O bonds

are the same length.…both outer

oxygens have a charge of 1/2.

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Resonance

• One Lewis structure cannot accurately depict a molecule such as ozone.

• We use multiple structures, resonance structures, to describe the molecule.

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Resonance

Just as green is a synthesis of blue and yellow…

…ozone is a synthesis of these two resonance structures.

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Resonance

• In truth, the electrons that form the second C—O bond in the double bonds below do not always sit between that C and that O, but rather can move among the two oxygens and the carbon.

• They are not localized, but rather are delocalized.

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Resonance

• The organic compound benzene, C6H6, has two resonance structures.

• It is commonly depicted as a hexagon with a circle inside to signify the delocalized electrons in the ring.

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• EXCEPTION TO THE OCTECT RULE• BOND STRENGTH

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Exceptions to the Octet Rule

• There are three types of ions or molecules that do not follow the octet rule:Ions or molecules with an odd number of

electrons.Ions or molecules with less than an octet.Ions or molecules with more than eight

valence electrons (an expanded octet).

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Odd Number of Electrons

Though relatively rare and usually quite unstable and reactive, there are ions and molecules with an odd number of electrons.

ClO2 NO NO2 O2-

superoxide ion

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Molecules with odd # of e-

• It is impossible to draw a structure obeying the octet rule.

• They are called free radicals

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Fewer Than Eight Electrons

• Consider BF3:Giving boron a filled octet places a negative charge

on the boron and a positive charge on fluorine.This would not be an accurate picture of the

distribution of electrons in BF3.

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Therefore, structures that put a double bond between boron and fluorine are much less important than the one that leaves boron with only 6 valence electrons.

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The lesson is: If filling the octet of the central atom results in a negative charge on the central atom and a positive charge on the more electronegative outer atom, don’t fill the octet of the central atom.

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• BF3 reacts very energetically with molecules having an unshared pair of electrons that can be used to form a bond with B

• Example formation of NH3BF3

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More Than Eight Electrons

• The only way PCl5 can exist is if phosphorus has 10 electrons around it.

• It is allowed to expand the octet of atoms on the 3rd row or below.Presumably d orbitals in

these atoms participate in bonding.

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Even though we can draw a Lewis structure for the phosphate ion that has only 8 electrons around the central phosphorus, the better structure puts a double bond between the phosphorus and one of the oxygens.

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• This eliminates the charge on the phosphorus and the charge on one of the oxygens.

• The lesson is: When the central atom is on the 3rd row or below and expanding its octet eliminates some formal charges, do so.

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Ions with expanded shell

• ICl4-

• XeF2

• ClF4-

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Strengths of Covalent Strengths of Covalent BondsBonds

• The energy required to break a covalent bond is called the bond dissociation enthalpy, D. That is, for the Cl2 molecule, D(Cl-Cl) is given by H for the reaction:

Cl2(g) 2Cl(g).• When more than one bond is broken:

CH4(g) C(g) + 4H(g) H = 1660 kJ• the bond enthalpy is a fraction of H for the

atomization reaction:D(C-H) = ¼H = ¼(1660 kJ) = 415 kJ.

• Bond enthalpies can either be positive or negative.

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Bond Enthalpies and the Enthalpies of Reactions

• We can use bond enthalpies to calculate the enthalpy for a chemical reaction.

• We recognize that in any chemical reaction bonds need to be broken and then new bonds get formed.

• The enthalpy of the reaction is given by the sum of bond enthalpies for bonds broken less the sum of bond enthalpies for bonds formed.

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Strengths of Covalent Strengths of Covalent BondsBonds

The energy required to break a covalent bond is called the bond dissociation enthalpy, D. That is, for the Cl2 molecule, D(Cl-Cl) is given by H for the reaction:

Cl2(g) 2Cl(g).• When more than one bond is broken:

CH4(g) C(g) + 4H(g) H = 1660 kJ• the bond enthalpy is a fraction of H for the atomization

reaction (since 4 bonds are broken each is ¼ )

D(C-H) = ¼H = ¼(1660 kJ) = 415 kJ.• Bond enthalpies can either be positive or negative.

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Average Bond Enthalpies• This table lists the

average bond enthalpies for many different types of bonds.

• Average bond enthalpies are positive, because bond breaking is an endothermic process.

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Average Bond EnthalpiesNOTE: These are

average bond enthalpies, not absolute bond enthalpies; the C—H bonds in methane, CH4, will be a bit different than theC—H bond in chloroform, CHCl3.

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Bond Enthalpies and the Enthalpies of Reactions

• We can use bond enthalpies to calculate the enthalpy for a chemical reaction.

• We recognize that in any chemical reaction bonds need to be broken and then new bonds get formed.

• The enthalpy of the reaction is given by the sum of bond enthalpies for bonds broken less the sum of bond enthalpies for bonds formed.

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• Mathematically, if Hrxn is the enthalpy for a reaction, then

• We illustrate the concept with the reaction between methane, CH4, and chlorine:

CH4(g) + Cl2(g) CH3Cl(g) + HCl(g) Hrxn = ?

formed bondsbroken bonds DDHrxn

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• In this reaction one C-H bond and one Cl-Cl bond gets broken while one C-Cl bond and one H-Cl bond gets formed.

• The overall reaction is exothermic which means than the bonds formed are stronger than the bonds broken.

• The above result is consistent with Hess’s law.

kJ 104

Cl-HCl-CCl-ClH-C

DDDDHrxn

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• Example – Use bond energies from table 8.4 to calculate ΔH for the following reaction:

C C

C

H

H

H

H H

H + C C C

H

H

H

H

H

H

H Cl

H

Cl

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Bond Enthalpy and Bond Length• We know that multiple bonds are shorter than single

bonds.• We can show that multiple bonds are stronger than

single bonds.• As the number of bonds between atoms increases, the

atoms are held closer and more tightly together.• NOTE: A double bond between two atoms is not twice

as strong as a single bond between the same two atoms.

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• Mathematically, if Hrxn is the enthalpy for a reaction, then

• We illustrate the concept with the reaction between methane, CH4, and chlorine:

CH4(g) + Cl2(g) CH3Cl(g) + HCl(g) Hrxn = ?

formed bondsbroken bonds DDHrxn

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• In this reaction one C-H bond and one Cl-Cl bond gets broken while one C-Cl bond and one H-Cl bond gets formed.

• The overall reaction is exothermic which means than the bonds formed are stronger than the bonds broken.

• The above result is consistent with Hess’s law.

kJ 104

Cl-HCl-CCl-ClH-C

DDDDHrxn

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• Example – Use bond energies from table 8.4 to calculate ΔH for the following reaction:

C C

C

H

H

H

H H

H + C C C

H

H

H

H

H

H

H Cl

H

Cl

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Bond Enthalpy and Bond Length• We know that multiple bonds are shorter than single

bonds.• We can show that multiple bonds are stronger than

single bonds.• As the number of bonds between atoms increases, the

atoms are held closer and more tightly together.• NOTE: A double bond between two atoms is not twice

as strong as a single bond between the same two atoms.

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Chapter 8 Concepts of Chemical Bonding

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