Lecture outline: Chapter 8 Chemical...

Lecture outline: Chapter 8Chemical bonding

1 Lewis symbols and atoms1. Lewis symbols and atoms2. Ionic bonding3. Lattice energy4. Isoelectronic series5. Covalent bonding6 Electronegativity and bond polarity6. Electronegativity and bond polarity7. Lewis structures8. Formal charges9. Resonance, octet violations10. Bond strengths11 O idation n mber11. Oxidation number

1S. Ensign, chemical bonding

Molecular structure: the way atoms are yarranged in molecules and polyatomic ionsionsMolecular bonding: the forces that hold adjacent atoms in a molecule together


Three types of bonding:Three types of bonding:Composition Type of molecule Type of bond

metal + nonmetal

ionic compound ionic bond

nonmetal + nonmetal

molecular compound

covalent bondnonmetal compound

metal + metal solid metal metallic bond

The type of chemical bond formed between atoms, and the structure of the resulting molecule, depend on the electronic structures of the atoms that are combining 3

S. Ensign, chemical bonding

Core electrons and valence electrons

Li: 1s2 2s1 Na: 1s2 2s22p6 3s1

n = 3

n = 1

n = 2

n = 1

n = 2

n = 3

3 p+ 11p+n 1 n 1p p


Electron dot symbols (Lewis symbols)• Valence electrons only• Place electrons (dots) on four sides• Place electrons (dots) on four sides• Place single electrons on all sides

b f i ibefore pairing• Don’t discriminate “s” and “p” electrons• # valence electrons = group# for s and

p block elementsp


Electron dot symbols (Lewis symbols)• Valence electrons only• Place electrons (dots) on four sides• Place electrons (dots) on four sides• Place single electrons on all sides before

pairing• Don’t discriminate “s” and “p” electrons• # valence electrons = group# for s and p• # valence electrons = group# for s and p

block elements


The Octet RuleAtoms gain, lose, or share electrons in ways that allow them to be surroundedways that allow them to be surrounded by 8 valence electrons (obtain a Noble gas configuration)gas configuration)

Apply to s and p block elementsApply to s and p block elements

H and He (period 1) want only 2 electrons, not 8

Transition block elements are more complicated


Satisfying the octet rule through ionic bonding 2 Na( ) + Cl2( ) 2NaCl( )2 Na(s) + Cl2(g) 2NaCl(s)


Satisfying the octet rule through ionic bonding

Na: 1s2 2s22p6 3s1 Cl: 1s2 2s22p6 3s23p5

1

n = 2

n = 3

11p+ 1

n = 2

n = 3

17 p+n = 111p+ n = 117 p+


Satisfying the octet rule through ionic bonding

Na+: 1s2 2s22p6 Cl-: 1s2 2s22p6 3s23p6

+ -

1

n = 211p+ 1

n = 2

n = 3

17 p+n = 111p+ n = 117 p+


Lattice energyClassic definition: The energy required to separate one mol of an ionic compound into its constituent gaseous ions

NaCl(s) Na+(g) + Cl-(g) ΔH = +788 kJ/mol( ) (g) (g)

Alternate definition: The energy released when f i l f i i d fforming one mole of an ionic compound from the constituent gaseous ionsNa+

(g) + Cl-(g) NaCl(s) ΔH = -788 kJ/mol

When lattice energies are reported as positive numbers they refer to the classicWhen lattice energies are reported as positive numbers, they refer to the classic definition . When reported as negative numbers, they refer to the alternate definition. It always takes energy to break chemical bonds, and energy is always released when forming chemical bonds 11


QQThe potential energy of two interacting charged particles

Q Qd

dQQkE 21= Q1 Q2

d


Common ion charges that can be predicted based on proximity to Noble gasesbased on proximity to Noble gases


Lattice energy“Greater lattice energy” means more positive forGreater lattice energy means more positive for separation of the ionic compound, and more negative for formation of the ionic compoundnegative for formation of the ionic compound

NaCl(s) Na+(g) + Cl-(g) ΔH = +788 kJ/mol

Na+(g) + Cl-(g) NaCl(s) ΔH = -788 kJ/mol

LiF(s) Li+(g) + F-(g) ΔH = +1046 kJ/mol

Li+ + F- LiF(s) ΔH = -1046 kJ/molLi (g) + F (g) LiF(s) ΔH = -1046 kJ/mol

LiF has a greater (larger) lattice energy than NaClLiF has a greater (larger) lattice energy than NaCl14


Transition metal ions• Not as easily predictable as s

3d4s

4pNot as easily predictable as sand p block elements (e.g., Fe2+ and Fe3+, Cu+ and Cu2+)

3s

3p

ncre

asin

g E• The first electrons lost are from

the subshell with the largest value of n (contrast this with

2s

2p

Invalue of n (contrast this with order of electron addition to orbitals)

1s


1s 2s 2p 3s 3p 4s 4p3d

Transition metal ions• The first electrons lost are from the subshell with the largest value of n g

(contrast this with order of electron addition to orbitals)

1s 2s 2p 3s 3p 4s 4p3d1s 2s 2p 3s 3p 4s 4p3d

F 2+

Fe:

Fe2+:

C 3

Cr:

Cr3+:


d0,d5, and d10 are good electron configurationsconfigurations

d0

10

dd5

d10

S. Ensign, chemical bonding17

Transition metal ions with d0, d5, and d10

electron configurationselectron configurations

d0

10

dd5

d10


Common ion charges for some important elementselements

Noblee gases


Transition metals can have a wide range of oxidation states in compounds, e.g MnO (+2), Mn2O3 (+3), MnO2 (+4), MnO4

-(+7)

Polyatomic ions



1 Lewis symbols and atoms1. Lewis symbols and atoms2. Ionic bonding3. Lattice energy4. Isoelectronic series5. Covalent bonding6 Electronegativity and bond polarity6. Electronegativity and bond polarity7. Lewis structures8. Formal charges9. Resonance, octet violations10. Bond strengths11 O idation n mber11. Oxidation number


Effects of removing an electron from or adding an electron to an atomadding an electron to an atom

• Z ff felt by the valence electrons?Zeff felt by the valence electrons?• Electron-electron repulsion of valence

electrons?electrons?• Principal quantum number (n) of valence

l t ?electrons?


Na: 1s2 2s22p6 3s1 Na+: 1s2 2s22p6

1

n = 2

n = 3

11p+ n = 1

n = 211p+

n = 111p+ n 1


Cl: 1s2 2s22p6 3s23p5 Cl-: 1s2 2s22p6 3s23p6Cl: 1s 2s 2p 3s 3p Cl : 1s 2s 2p 3s 3p

n = 3 n = 3

n = 1

n = 217 p+ n = 1

n = 217 p+


Cl: 1s2 2s22p6 3s23p5 Cl-: 1s2 2s22p6 3s23p6Cl: 1s 2s 2p 3s 3p Cl : 1s 2s 2p 3s 3p

n = 3n = 3

n = 1

n = 217 p+ n = 1

n = 217 p+


Author Popnose, http://commons.wikimedia.org/wiki/File:Atomic_%26_ionic_radii.svgThis file is licensed under the Creative Commons Attribution-Share Alike 3.0 Unported license.

Satisfying the octet rule through ionic bonding 2 Na( ) + Cl2( ) 2NaCl( )2 Na(s) + Cl2(g) 2NaCl(s)


Isoelectronic species: same number of TOTAL electronsnumber of TOTAL electrons

• What monoatomic ions are isoelectronicWhat monoatomic ions are isoelectronic with Ne?

• What is Z felt by the valence electrons of• What is Zeff felt by the valence electrons of each species?A th i f l t t ll t• Arrange the ions from largest to smallest based on Zeff considerations


Isoelectronic species: same number of electrons• What monoatomic ions are

isoelectronic with Ne?• What is Zeff felt by the

valence electrons of each species?

• Arrange the ions from largest to smallest based on Zeff considerations


This week in chemistryThis week in chemistry• No on line quiz this week ☺

Recitations meet re ie e am 2 and re ie• Recitations meet, review exam 2, and review beginning of chapter 8, recitation quiz

• Read materials for chapter 8• work problems in chapter 8 self test, watch

chapter 8 self test tutorials to help in developing problem-solving skills.

S. Ensign, Chem. 121030

S. Ensign, Chem. 121030S. Ensign, Chem. 1210 3030


Isoelectronic species: same number of electrons• What monoatomic ions are

isoelectronic with Ne?• What is Zeff felt by the

valence electrons of each species?

• Arrange the ions from largest to smallest based on Zeff considerations


Which of the following sets contains species th t ll i l t i ?that are all isoelectronic?

a O F Nea. O, F, Neb. C, Si, Ge

P3 S2 Clc. P3-, S2-, Cl-

d. Na, Mg, Al


Covalent bondingSharing of valence electrons between atoms• Sharing of valence electrons between atoms

• Usually nonmetal + nonmetal• Share electrons such that each nonmetal has

an octet of electrons (ex., H)


H2 bond formation

H HH H

H H 34S. Ensign, chemical bonding

Potential energy diagram for H-H bond formation

400

mol

) 200

ergy

(kJ/

m

0

Ene -200

0 0 0 5 1 0 1 5 2 0 2 5 3 0

-400

radius (angstroms)

0.0 0.5 1.0 1.5 2.0 2.5 3.0


Only the valence electrons are important in covalent bond formation

Cl: 1s2 2s22p6 3s23p5F: 1s2 2s22p5

n = 3n = 2

n = 1

n = 217 p+n = 1

n = 217 p+

FClCl


A simple rule (that works a lot of the time) for predicting the number of covalent bonds necessary

f f i l to form an octet for an atom in a covalent compound: 8 - group#


A simple rule (that works a lot of the time) for predicting the number of covalent bonds necessary

f f i l to form an octet for an atom in a covalent compound: 8 - group#


Electron dot structures (Lewis structures)• Electron dot notation for covalent cmpds.p• Valence electrons only• Place electrons (dots) on all atoms such that the “octet

rule” is satisfied for all atoms (ex. H, which needs only two electrons)

• An electron pair between two atoms represents a covalent bond : =covalent bond :

• An electron pair on an atom that is not bonded to another atom is called a “lone pair” or a “nonbonded pair”

• Some compounds will need “multiple bonds” in order for all atoms to have an octet 39


Some compounds will need “multiple bonds” in order for all atoms to have an octetin order for all atoms to have an octet

O2O2

NN2


Multiple bonds are shorter (and stronger) than single bondsthan single bonds


Sharing of bonded electron pairs in covalent compoundscompounds

X X

X Y


ElectronegativityThe ability of an atom in a molecule to attract the shared electrons in a bondattract the shared electrons in a bond

Electronegativity increases fromincreases from left to right and from bottom to ftop in the Periodic Table


Ionization energy: a measure of how gystrongly an atom holds on to its electronselectronsElectron affinity: a measure of how strongly an atom attracts an external electronElectronegativity: a measure of how strongly an atom in a molecule attractsstrongly an atom in a molecule attracts shared electrons in a covalent bond


Some electronegativity values

Increases from left to right in a period45


Differences in electronegativity of atoms dictate h t ki d f b d f b t th twhat kind of bond forms between the atoms

ΔEN = 0 a nonpolar covalent bondΔEN 0, a nonpolar covalent bond

ΔEN = 0 2, a polar covalent bond

ΔEN > 2, an ionic bond


ΔEN = 0, a nonpolar covalent bond

ΔEN = 0 2, a polar covalent bond

ΔEN > 2, an ionic bond

What type of bond is present in each of the following molecules?

N ClNaClCl2HClHClH2O


“Electrostatic potential maps” show the charge densities on molecules with atoms c a ge de s t es o o ecu es t ato s

having different electronegativities

H F

less EN more EN

C OO

S. Ensign, chemical bonding48

Drawing Lewis structures of molecules(1) Sum the valence electrons from all atoms(1) Sum the valence electrons from all atoms(2) Decide on atom connections connect with a single

bondbond(“central atom” is usually written first in formula)

(3) Determine the number of valence electrons(3) Determine the number of valence electrons remaining:total valence electrons minus #used in bondingtotal valence electrons minus #used in bonding

(4) Complete octets of terminal atoms (ex. H)(5) Place leftover electrons on the central atom(5) Place leftover electrons on the central atom(6) if no unassigned electrons remain, and the central

atom doesn’t have an octet use multiple bondsatom doesn t have an octet, use multiple bonds


Drawing Lewis structures of molecules: PH31) Sum the valence electrons )

from all atoms2) Decide on atom

connections connect with a (“ ”single bond (“central atom”

is usually written first in formula)

3) Determine the number of3) Determine the number of valence electrons remaining: total valence electrons minus #used in bonding

4) Complete octets of terminal atoms (ex. H)

5) Pl l ft l t5) Place leftover electrons on the central atom

6) if no unassigned electrons remain and the centralremain, and the central atom doesn’t have an octet, use multiple bonds


Drawing Lewis structures of molecules: NF31) Sum the valence electrons )









Drawing Lewis structures of molecules: CO21) Sum the valence electrons )









Drawing Lewis structures of molecules: CO21) Sum the valence electrons )









Drawing Lewis structures of molecules: HCN1) Sum the valence electrons )









Drawing Lewis structures of molecules: NH4+ ion

1) Sum the valence electrons )from all atoms

2) Decide on atom connections connect with a

(“ ”single bond (“central atom” is usually written first in formula)






Drawing Lewis structures of molecules: PO43- ion

1) Sum the valence electrons )from all atoms

2) Decide on atom connections connect with a

(“ ”single bond (“central atom” is usually written first in formula)







1 Lewis symbols and atoms1. Lewis symbols and atoms2. Ionic bonding3. Lattice energy4. Isoelectronic series5. Covalent bonding6 Electronegativity and bond polarity6. Electronegativity and bond polarity7. Lewis structures8. Formal charges9. Resonance, octet violations10. Bond strengths11 O idation n mber


11. Oxidation number


As we are about to see, there are many molecules for which multiple Lewis

structures can be drawn that satisfy thestructures can be drawn that satisfy the octet rule for each bonded atom


Formal charge• “bookkeeping” for valence electrons• bookkeeping for valence electrons

– Draw a reasonable Lewis structureAssign electrons to atoms:– Assign electrons to atoms:

• All unshared (nonbonding; lone pair) e- are assigned to the atom on which they are foundthe atom on which they are found

• ½ of bonding e- are assigned to each atom in the bond

• Formal charge = # of valence e- in isolatedFormal charge # of valence e in isolated atom minus # of assigned e-

• The most stable Lewis structure is that in• The most stable Lewis structure is that in which atoms bear the smallest formal charges and where (-) formal charges residecharges, and where (-) formal charges reside on the more electronegative atoms 59


Formal chargeAssign electrons to atoms:Assign electrons to atoms:

• All unshared (nonbonding; lone pair) e- are assigned to the atom on which they are found

• ½ of bonding e- are assigned to each atom in the bond• ½ of bonding e are assigned to each atom in the bond

)e nonbonding (# - )e bonding(#21 - )e valence(#F.C. --

freeatomin −=

Apply concept of formal charge to carbon monoxide (CO):


The most stable Lewis structure is that in which atoms bear the smallest formal charges, and where (-) formal charges reside on the more electronegative atomscharges reside on the more electronegative atoms

Apply this concept to the cyanate anion (OCN-):


Resonance structures: ozone (O3)


Resonance structures and bond order• Ozone can be drawn as two equivalent resonance structures:Ozone can be drawn as two equivalent resonance structures:

• Ozone is actually a “hybrid” of structures (i) and (ii), where the double bond is “delocalized” over the entire molecule. It is not easy to represent delocalization using Lewis structures aseasy to represent delocalization using Lewis structures as illustrated below:– Too may electrons on O3:

– Too many electrons on O1::

– Too many e- on central O and not enough on terminal oxygens:

– Unpaired electrons on terminal O (not case):63


Resonance structures and bond order• It is best to draw ozone as the two equivalent• It is best to draw ozone as the two equivalent

resonance structures:

• The actual bond order is given by the formula:

pairsbondingnumbertotallocations bond ofnumber

pairs bondingnumber total order bond =

51pairs bonded 3Oorderbond 3 == 5.1locations bond O-O 2

O order, bond 3 ==64


Resonance structures: nitrate ion


Exceptions to the octet rule• Molecules with an odd number of electronsMolecules with an odd number of electrons

• Molecules with an atom that doesn’t have an octet (n=1, some B B d )B, Be compounds)

• Molecules with an atom that has more than an octetn = 3 and below: availability of d orbitals allow octet rule to be exceeded, if

necessary66


Exceptions to the octet ruleMolecules with an atom that has more than an octetMolecules with an atom that has more than an octet

n = 3 and below: availability of d orbitals allow octet rule to be exceeded, if necessary


For the sulfate ion (SO42-)

1) Draw the most favorable structure with regard to formal1) Draw the most favorable structure with regard to formal charge (you can violate the octet rule for S)

2) How many resonance structures are possible for the l l d i (1)?molecule you drew in (1)?

)e nonbonding (# - )e bonding(#21 - )e valence(#F.C. --

freeatomin −=

2


For the sulfate ion (SO42-)

1) Draw the most favorable structure with regard to formal charge (you can violate ) g g (ythe octet rule for S)

2) How many resonance structures are possible for the molecule you drew in (1)?

)enonbonding(#-)ebonding(#1-)evalence(#F.C. --freeatomin

−= )enonbonding (# )e bonding(#2

)e valence(#F.C. freeatomin


Concept checkp• Ionic bonding• Covalent bonding• Covalent bonding

– Electronegativity• Polar covalent bond• Nonpolar covalent bond

– Lewis structures• Octet rule• Octet violations

– Formal chargeFormal charge– Resonance


Concept check





Bond strengthsI i b d L tti• Ionic bonds: Lattice energy– The energy required to separate one mol of

an ionic compound into it’s constituentan ionic compound into it s constituent gaseous ions

NaCl(s) Na+ + Cl- ΔH +788 kJ/ l

• Covalent bonds: Bond Dissociation energy

NaCl(s) Na+(g) + Cl-(g) ΔH = +788 kJ/mol

Covalent bonds: Bond Dissociation energy– The energy required to break the bonds in

one mol of a gaseous molecular (covalent) g ( )compound

H H H + H ΔH = +436 kJ/molD(H H) 436 kJ/ lD(H-H) = +436 kJ/mol


Bond Dissociation energyThe energy required to break the bonds in one mol of a

l l ( l t) dgaseous molecular (covalent) compound


Table of average bond energies (D, kJ/mol)

Bond D Bond D Bond D H—H 436 C H 413 C—H 413 C—C 348 C═C 614 C≡ C 839 C—N 293 C═N 615 C≡N 891 C O 358 C O 799 C≡O 1072C—O 358 C═O 799 C≡O 1072 C—S 259 N—H 391

163 418 941N—N 163 N═N 418 N≡N 941 N—O 201 N═O 607 O—H 463

O OO—O 146 O2 495 S—H 339 S—S 266


Bond dissociation energiesBond dissociation energies

• Values are always positive It always takesValues are always positive. It always takes energy to break a chemical bond


Review concepts: Chapter 5Enthalpy of formation (ΔHf°)

The change in enthalpy (heat input or heatThe change in enthalpy (heat input or heat output) associated with the formation of one mol of a compound from its constituent elements inof a compound from its constituent elements in their “standard state” forms

Standard state ≡ 25° C and one atmosphere of pressure

The “standard state form” of an element is the form most stable at 25° C and one atmosphere of pressure


Using a table of enthalpies of formation toReview concepts: Chapter 5

Using a table of enthalpies of formation to calculate the enthalpy change (ΔH°) for a

chemical reaction of interestchemical reaction of interest

ΔH°rxn = Σ nΔHf°(products) - Σ mΔHf°(reactants) rxn f ( ) f ( )

Appropriate stoichiometric coefficients in the balanced chemical equationthe balanced chemical equation

Standard enthalpies of formation for selected compounds at 298.15 K (25°)

Substance ∆H°f (kJ/mol) Substance ∆H°f (kJ/mol)

CO (g) -393 5 NH (g) -46 19CO2(g) -393.5 NH3(g) -46.19

CH4(g) -74.8 NaCl(s) -410.9

C6H12O6(s) -1273 C(s) diamond 1.88

CH OH(l) 238 6 H O(g) 241 8


CH3OH(l) -238.6 H2O(g) -241.8

C2H2(g) 226.7 H2O(l) -285.85

Estimating ΔH°for a reaction from a table of average bond enthalpiestable of average bond enthalpies

• Determine which bonds are broken. Sum these bond enthalpies (E is required tothese bond enthalpies (E is required to break a bond: + sign)

• Determine which bonds are formed. Sum these bond enthalpies (E is released when a bond forms: - sign)

∑∑=° formed)D(bonds-broken)D(bondsΔH

• Pay attention to the number and types of

∑∑= formed) D(bonds - broken) D(bonds ΔHrxn

• Pay attention to the number and types of bonds in the reactants and products!!


Bond D Bond D Bond D H—H 436 C—H 413 C C 348 C C 614 C≡ C 839

Determine ΔH° for the combustion of hydrogen

C—C 348 C═C 614 C≡ C 839 C—N 293 C═N 615 C≡N 891 C—O 358 C═O 799 C≡O 1072 C—S 259 N H 391

gas using information in the table of average bond energies N—H 391

N—N 163 N═N 418 N≡N 941 N—O 201 N═O 607 O—H 463 O O 146 O 495

energies2H2 + O2 2H2O

O—O 146 O2 495 S—H 339 S—S 266

∑∑=° formed) D(bonds - broken) D(bonds ΔHrxn


Determine ΔH° for the combustion of propane using information in the table of

Bond D Bond D Bond D H—H 436 C—H 413 C C 348 C C 614 C≡ C 839

∑∑=° formed) D(bonds - broken) D(bonds ΔHrxn

information in the table of average bond energies

C—C 348 C═C 614 C≡ C 839 C—N 293 C═N 615 C≡N 891 C—O 358 C═O 799 C≡O 1072 C—S 259 N H 391∑∑

C CC

H H

H

H

H

N—H 391 N—N 163 N═N 418 N≡N 941 N—O 201 N═O 607 O—H 463 O O 146 O 495

H HH

O—O 146 O2 495 S—H 339 S—S 266


Bond length

Bond D (kJ/mol)

Length (Å)

Bond D (kJ/mol

Length (Å)

Bond D (kJ/mol)

Length (Å)(kJ/mol) (Å) (kJ/mol

)(Å) (kJ/mol) (Å)

H—H 436 0.74C—H 413 1.10C—C 348 1.54 C═C 614 1.34 C≡ C 839 1.21C—N 293 1.47 C═N 615 1.27 C≡N 891 1.15C—O 358 1.43 C═O 799 1.22 C≡O 1072 1.13C—S 259 1.81N—H 391 0.98

163 1 40 418 1 20 941 1 10N—N 163 1.40 N═N 418 1.20 N≡N 941 1.10N—O 201 1.36 N═O 607 1.15O—H 463 0.94O—O 146 1.32 O2 495 112S—H 339 1 32S—H 339 1.32S—S 266 2.08


Concept check





Oxidation number• The “charge” that results when electrons in a covalentThe charge that results when electrons in a covalent

bond are assigned to the more electronegative atom

Increase EN


Rules for assigning oxidation numbers1. For an atom in it’s elemental form, oxidation # = 01. For an atom in it s elemental form, oxidation # 0

2. For a monatomic ion, oxidation # = ion charge

3. For binary compounds of nonmetals: Assign the element with greater electronegativity the “ion charge” it would have if it was in an ionic compoundcharge it would have if it was in an ionic compound

4 The sum of oxidation #s is 0 in a neutral compound4. The sum of oxidation #s is 0 in a neutral compound, while the sum is the overall charge for polyatomic ions 84


Common oxidation states

Element Ox. State in cmpds

F Always -1

O Usually -2 (ex. w/ F), -1 in O Usua y (e / ),peroxides

Cl, Br, I Usually -1 (except w/ F and O), , y ( p )

H +1 w/ nonmetal; -1 w/metal


A fail proof way to assign oxidation numbers in binary compoundsy p

1. Assign the element with greater electronegativity the “ion charge” it would have if it was in an ionic compound. Write this number below the element symbol in theWrite this number below the element symbol in the chemical formula

2. Write the sum of the “ion charges” for all the atoms of gthat element in the chemical formula above the element symbol

3 Determine the “total ion charge” necessary for all atoms3. Determine the total ion charge necessary for all atoms of the other element to give the overall charge present on the molecule or polyatomic ion. Write this number

b th d l tabove the second element4. Divide the “total ion charge” written above the second

element by the number of atoms of that element in theelement by the number of atoms of that element in the formula to get its oxidation number. Write this below the element symbol 86


Oxidation numbers in molecules of nonmetals containing three or more atomsnonmetals containing three or more atoms

• In nonmetals, H always is given an oxidation b f 1 t t b i i 1 t hnumber of +1, so start by assigning +1 to each

hydrogen• Assign the other nonmetals oxidation numbers

with priorities based on highest electronegativity (e.g. O first, then N, the S, then C) using the strategy described for binary compounds


A fool proof way to assign oxidation numbers in binary compoundsy p

1. Assign the element with greater electronegativity the “ion charge” it would have if it was in an ionic compound. Write this number below the element symbol in the chemical formula

2 Write the sum of the “ion charges” for all the atoms of that element in the2. Write the sum of the “ion charges” for all the atoms of that element in the chemical formula above the element symbol

3. Determine the “total ion charge” necessary for all atoms of the other element to give the overall charge present on the molecule or polyatomic ion. Write this number above the second element

4. Divide the “total ion charge” written above the second element by the number of atoms of that element in the formula to get its oxidation number. Write this below the element symboly

Overall chargeSum of charges:

H2OOverall charge on molecule (zero if nothing written)2

Oxidation numbers:88


Oxidation numbers in molecules of nonmetals containing three or more atomsnonmetals containing three or more atoms

• In nonmetals, H always is given an oxidation number of +1 so start by assigning +1 to each hydrogenof +1, so start by assigning +1 to each hydrogen

• Assign the other nonmetals oxidation numbers with priorities based on highest electronegativity (e g Opriorities based on highest electronegativity (e.g. O first, then N, the S, then C) using the strategy described for binary compounds

Sum of charges:

y p

HCO3- Overall charge on

polyatomic ion3

Oxidation numbers:89


Determine oxidation numbers for the atoms in the following compounds:Li2O

H3PO4

g p1. Assign the element with greater electronegativity

the “ion charge” it would have if it was in an ionic compound

H3PO4

MnO4-

C O 2

2. The sum of oxidation #s is 0 in a neutral compound, while the sum is the overall charge for polyatomic ions

Cr2O72-

ClO4-

C7H8

NH3

NH4+

H2O2

NaH EN increases from left to right in a period90


Li O

Determine oxidation numbers for the atoms in:

I ENLi2O

H3PO4

Increases EN

MnO4-

1. Assign the element with greater electronegativity the “ion charge” it would have if it was in1. Assign the element with greater electronegativity the ion charge it would have if it was in an ionic compound



Cr2O72-


I EN2 7

ClO4-

C H

Increases EN

C7H8




NH3


I ENNH3

NH4+

H O

Increases EN

H2O

NaH




F l h l t b kk i i L iOxidation # ≠ formal charge

• Formal charge: electron bookkeeping in Lewis structures, keeps track of which atom(s) donate

t h d l t i l t b dor accept shared electrons in a covalent bond)e nonbonding (# - )e bonding(#

21 - )e valence(#F.C. --

freeatomin −=

• Oxidation number: Based on electronegativity differences, keeps track of which atoms more , ptightly hold onto shared electron pairs in a covalent bond

• Compare oxidation number and formal charge for some simple molecules and polyatomic ionsfor some simple molecules and polyatomic ions


•Compare oxidation number and formal charge for some simple molecules and polyatomic ions


Importance of oxidation numbers?• Keeping track of electron gain/loss inKeeping track of electron gain/loss in

oxidation/reduction reactions• An atom in a compound is “oxidized” (undergoesAn atom in a compound is oxidized (undergoes

oxidation)if it’s oxidation number increases during the course of a chemical reaction

• An atom in a compound is “reduced” (undergoes reduction) if it’s oxidation number decreases during the course of the chemical reaction


Importance of oxidation numbers?• Keeping track of electron gain/loss in oxidation/reduction reactionsg g• An atom in a compound is “oxidized” if it’s oxidation number increases during the

course of a chemical reaction• An atom in a compound is “reduced” if it’s oxidation number decreases during the

course of the chemical reaction


Concept check• Ionic bonding• Ionic bonding• Covalent bonding

Electronegativity– Electronegativity• Polar covalent bond• Nonpolar covalent bondp

– Lewis structures• Octet rule• Octet violations

– ResonanceF l h– Formal charge

– Bond strengthO id ti b• Oxidation number


Lecture outline: Chapter 8 Chemical...

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