ap chem 2 nd semester 7-bonding 8-liq/soln 9-kinetics 10-equilibrium 11-acid-base 12-electrochem...

UNIT 7:CHEMICAL BONDING

BONDS

are forces which hold atoms and ions together in compounds

form because this gives the “lowest possible energy for the system”

(lower P.E. means more stable bond thermodynamically)

being broken absorbs (requires) energy [endo, so DH = (+)]

being made releases energy [exo, so DH = (-)]

OCTET RULE

atoms will lose or gain valence electrons, or share electrons so as to have a total of eight valence (outermost energy level) electrons (s2p6)

in highest “n” value sublevels Ex: 3s23p6, 5s25p6

BONDING TYPES

IONIC bond is the electrostatic attraction between (+) cations and (-) anions

following the transfer of electrons from metal atoms to non-metal atoms

called “ion-ion interactions

COVALENT bond is thesharing of electron pairs between non-metal atoms

NON-POLAR POLARequal sharing of unequal sharing of

bonding electrons bonding electrons symmetrical

clouddistorted cloud

BOND TYPE depends on the ELECTRONEGATIVITY OF THE ATOMS in the bond:

ELECTRONEGATIVITY

is the attraction an atom has for another atom’s electrons in a bond

(or an atom’s ability to attract bonding electrons to itself, “greediness”)

is a relative value and has no units

values range from 0.7 (Cs most active metal) to 4.0 (F most active non-metal)

generally increases leftright across a period, decreases down a group

Basic Concepts of Chemical

Bonding

© 2012 Pearson Education, Inc.

Electronegativity• Electronegativity is the ability of atoms in a molecule

to attract electrons to themselves.• 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.


Bonding


ELECTRONEGATIVITY “difference” indicates BOND TYPE EN DIFF

0------0.4-------------------------------------2.0-------------3.3 TYPE: non-polar -----> polar covalent -------> ionic covalent

0.5 1.9

The greater the EN DIFF, the more the “ionic character” of the bond

The lower the EN DIFF, the more the “non-polar character” of the bond

Actually more a continuum than clear-cut boundaries:


Bonding


Polar Covalent Bonds

• 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).


Bonding


Polar Covalent Bonds

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

• \

Dipole moment is zero Non-polar molecule

Dipole moment is 1.90 Debye Polar molecule

Dipole moment can be defined as the product of magnitude of the partial charge and the distance separating them.

d-

d+

d-

http://en.wikipedia.org/wiki/File:Cis-1,2-dichloroethene.png

Let’s go to handout


Bonding


Energetics of Ionic Bonding

As we saw in the last chapter, it takes 496 kJ/mol to remove electrons from sodium.


Bonding



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


Bonding



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


Bonding



• 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.


Bonding


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

IONIC COMPOUNDS• formed from metal cations & non-metal anions• the attractive force between these ions is expressed

as “LATTICE ENERGY”

Def: energy released when gaseous ions form an ionic compound

M+(g) + N-

(g) ---> MN(s) + heat

Attractive forces between ions described by

COULOMB’S LAW:

E = k Q1+ Q2

-

rLattice Energy:

attractive force

between ions

constant

charges on the ions

distance between ion centers in the lattice

+ -


Bonding


Lattice Energy

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

• It also increases with decreasing size of ions.


Bonding



Bonding


Lattice Energy

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

• It also increases with decreasing size of ions.

the one with the greater lattice energy

Which compound has a) more ionic bond character? ______

the one with the greater E.N. differenceb) a stronger ionic bond? ______

1) CaO or NaClCa2+ O2- Na+ Cl-= -4 = -1

E.N Diff: 3.5-1.0 = 2.5

E.N Diff: 3.0-0.9 = 2.1

larger charges greater lattice energy so stronger bond (the more heat released, the lower the PE)

CaO

CaO

MPt: 2613oC MPt: 801oC

E = -4

rE = -4 r

2+ 2+2-

2) CaO or MgO

2-

So MgO has stronger ionic bondBut CaO has more ionic bond character

3) CaCl2 or K2S

E = -2 r

E = -2

rSo CaCl2 has stronger ionic bond and more ionic bond character

EN Diff:3.5-1.2= 2.3

EN Diff:3.0-1.0= 2.0

EN Diff:2.5-0.8= 1.7

MPt: 2825oC

both ions smaller

The OVERALL energy change in “formation of an ionic solid” must be calculated in steps: (then add energies together)

1. Enthalpy of sublimation M(s) -----> M(g)

2. Ionization Energy M(g) -----> M+(g)

to form the cation

3. Dissociation Energy N2(g) -----> 2N(g)

for diatomic (also called

Bond Energy)

4. Electron Affinity N(g) + e- ----> N-

(g)

to form the anion

5. Lattice Energy M+(g) + N-

(g)---> MN(s)

when gaseous ions come together


Bonding



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.


Bonding


Use Table of Bond Energies!! Chemical reactions involve bond-breaking and bond-

making.

Each chemical bond has an ___________________ in kJ/mol.

BOND ENERGY is __________________________________

Always ______ (_____) BOND LENGTH is

__________________________________

“average” bond energy

the energy required to break a bond

(+) endo

the distance between 2 nuclei connected by a bond

Like leggos!!

See Handout with Table

BOND ENERGY is __________________________________

Ex: carbon-carbon bonds

TYPE of bond SINGLE C-C DOUBLE C=C TRIPLE C=C electron pairs 1 shared pair 2 shared pairs 3 shared pairs

Bond Energy 347 kJ/mol 614 kJ/mol 839 kJ/mol Bond Length 154 pm 134 pm 120 pm

Bond Order 1 2 3

atoms pulled closer

-

# e- prs in a bond

Ex: carbon-nitrogen bonds TYPE of bond SINGLE C-N DOUBLE C=N TRIPLE C=N electron pairs 1 shared pair 2 shared pairs 3 shared pairs

Bond Energy 305 kJ/mol 615 kJ/mol 891 kJ/mol Bond Length 143 pm 138 pm 116 pm Bond Order 1 2 3

1) As # of shared pairs increases, the bond length ___________ 2) The _______the bond energy, the ________ the bond.

shortens

greater stronger

“more stable”

-

To calculate the DH of the reaction (enthalpy change) using bond energy values:

DHrxn = SD (bonds broken) - SD (bonds made) energy absorbed energy released

energy/mol Thus if net energy change is (+) meaning “more energy absorbed than released” then DH = (+) indicating an overall endothermic process

[Note: The only time we use “initial minus final”, rather than “final minus initial”!]

“D” means bond energy

initial minus final


Bonding


Average Bond Enthalpies• Table 8.4 lists the

average bond enthalpies for many different types of bonds.

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


Bonding


Enthalpies of Reaction

So,

H = [D(C—H) + D(Cl—Cl)] − [D(C—Cl) +

D(H—Cl)]

= [(413 kJ) + (242 kJ)] − [(328 kJ) + (431 kJ)]

= (655 kJ) − (759 kJ)

= −104 kJ

breaking bonds

making bonds

436 kJ/mol 239 kJ/mol 427 kJ/mol

436 kJ 242 kJ 2 431 kJ

-184kJ/ 2mol HCl produced

[ 671 kJ ] - [854kJ ]

We can get to

f -92kJ/mol

Moles cancel!!

941 kJ

436 kJ/mol 391 kJ/mol

123

1 2 3

4 5 6

3 436 kJ 6 391 kJ

941 kJ/mol

[ 941 + 1308 ] - [2346 ]

-97kJ/ 2mol NH3 produced

f -49kJ/mol -46

Group A # gives # val e-s

NH3 PCl35 1 5 7

5 + 3 = 8 total 5 + 21 = 26 total

lost e-

gained e-

trigonal planar(polyatomic ion)

now spread bonds equidistant from each other

FORMAL CHARGE is a hypothetical charge on an atom in a molecule or

ion helps to determine the best possible Lewis structure for a molecule or ion is the difference between the total number of

valence electrons of a particular atom and the number of electrons involved in bonds and/or lone pairs

The sum of all the formal charges for a

molecule/ion is equal to the charge on that molecule/ion.

Day 5

Rules:

To count electrons in formal charges:

1) lone pairs = 2e-

(“unshared” or “non-bonding pairs”)2) single bonds = 1e-

(“shared” or “bonding pairs”)3) double bonds = 2e-

4) triple bonds = 3e-

Now: Go to Overhead

For 2 “non-equivalent” Lewis Structures,

choose the one with:

1) formal charges closest to zero,

2) and the (-) formal charge is on the

most electronegative atom

VSEPR Theory means Valence Shell Electron-Pair Repulsion Theory1) all electrons “paired” 2) all atoms have stable octet, H has duet3) pairs are spread equidistant from each other around atom used to determine 3-dimensional geometry of molecules and ions

Molecular Geometry:

geom. of ATOMS in the molecule

Electronic Geometry:

geom. of ELECTRON DOMAINS around an atom (used to find hybridization around a particular atom)

Day 6

KEY IDEA: bonds (shared pairs) and lone pairs (unshared pairs) arrange themselves so that repulsion is “minimized” (they are as far apart as they can get!) 2 ways electrons are positioned around an atom in a molecule or ion:

1) in bonds2) in lone pairs called “electron domains”

Molecular Shapes• The shape of a molecule plays an important role

in its reactivity.• By noting the number of bonding and nonbonding

electron pairs, we can easily predict the shape of the molecule.

MolecularGeometries

and Bonding© 2012 Pearson Education, Inc.

What Determines the Shape of a Molecule?

• Simply put, electron pairs, whether they be bonding or nonbonding, repel each other.

• By assuming the electron pairs are placed as far as possible from each other, we can predict the shape of the molecule.

Examples: Consider the central atom in each example!!! NH3 has 4 “electron domains”

1 lone pair 3 bonds

Note: Multiple bonds count as a single “electron domain”.

CO2 has 2 “electron domains”

0 lone pairs 2 bonds

EG - tetrahedral

MG - pyramidal

EG - linearMG - linear when NO LONE prs,

EG & MG the same

Lone pairs exert greater repulsion than bonding prs.

107o

180o

MolecularGeometries


Electron Domains

• We can refer to the electron pairs as electron domains.

• In a double or triple bond, all electrons shared between those two atoms are on the same side of the central atom; therefore, they count as one electron domain.

• The central atom in this molecule, A, has four electron domains.

MolecularGeometries


Valence-Shell Electron-Pair Repulsion Theory (VSEPR)

“The best arrangement of a given number of electron domains is the one that minimizes the repulsions among them.”

MolecularGeometries


Electron-Domain Geometries

Table 9.1 contains the electron-domain geometries for two through six electron domains around a central atom.

MolecularGeometries


Electron-Domain Geometries

• All one must do is count the number of electron domains in the Lewis structure.

• The geometry will be that which corresponds to the number of electron domains.

MolecularGeometries


Molecular Geometries

• The electron-domain geometry is often not the shape of the molecule, however.

• The molecular geometry is that defined by the positions of only the atoms in the molecules, not the nonbonding pairs.

MolecularGeometries


Linear Electron Domain

• In the linear domain, there is only one molecular geometry: linear.

• NOTE: If there are only two atoms in the molecule, the molecule will be linear no matter what the electron domain is.

MolecularGeometries


Trigonal Planar Electron Domain

• There are two molecular geometries:– Trigonal planar, if all the electron domains are

bonding,– Bent, if one of the domains is a nonbonding pair.

O3 has 3 “electron domains”

1 lone pair 2 bonds

NO3 has 3 “electron domains”


EG – trigonal planar

MG – bent

116.8 o

not 120o

120o



CH4 has 4 “electron domains”


109.5o

EG – tetrahedral

MG – tetrahedral

HYBRIDIZATION

While VSEPR Theory helps us predict the “spatial arrangement” of atoms in a molecule or ion, Valence Bond Theory describes “how bonding occurs” in terms of overlapping atomic orbitals.

These orbitals are often mixed or “hybridized” to form new “hybrid orbitals” which have 1)lower energy (stability) than the separate atomic

orbitals2)degenerate with each other (equal in that lower

energy) 3)equidistant from each other (evenly spaced) about

the atom

MolecularGeometries


Hybrid Orbitals

• Consider beryllium:– In its ground electronic

state, beryllium would not be able to form bonds, because it has no singly occupied orbitals.

MolecularGeometries


Hybrid Orbitals

But if it absorbs the small amount of energy needed to promote an electron from the 2s to the 2p orbital, it can form two bonds.

MolecularGeometries


Hybrid Orbitals

• With hybrid orbitals, the orbital diagram for beryllium would look like this (Fig. 9.15).

• The sp orbitals are higher in energy than the 1s orbital, but lower than the 2p.

MolecularGeometries


Hybrid Orbitals

• Mixing the s and p orbitals yields two degenerate orbitals that are hybrids of the two orbitals.– These sp hybrid orbitals have two lobes like a p orbital.– One of the lobes is larger and more rounded, as is the

s orbital.

MolecularGeometries


Hybrid Orbitals• These two degenerate orbitals would align

themselves 180 from each other.• This is consistent with the observed geometry of

beryllium compounds: linear.

MolecularGeometries


Hybrid Orbitals

Using a similar model for boron leads to three degenerate sp2 orbitals.

MolecularGeometries


Hybrid Orbitals

With carbon, we get four degenerate sp3 orbitals.

MolecularGeometries


Valence Bond Theory

• Hybridization is a major player in this approach to bonding.

• There are two ways orbitals can overlap to form bonds between atoms.

MolecularGeometries


Sigma () Bonds

• Sigma bonds are characterized by– Head-to-head overlap.– Cylindrical symmetry of electron density about the

internuclear axis.

MolecularGeometries


Pi () Bonds

• Pi bonds are characterized by– Side-to-side overlap.– Electron density above and below the internuclear

axis.

MolecularGeometries


Single BondsSingle bonds are always bonds, because overlap is greater, resulting in a stronger bond and more energy lowering.

MolecularGeometries


Multiple Bonds

In a multiple bond, one of the bonds is a bond and the rest are bonds.

In an atom: the one “s orbital” is spherical

the three “p orbitals” are pair-of-pear at 90o from each other

(Bond angles, experimentally determined, form the basis of hybridization theory

These angles did not correspond to the angles of atomic orbitals.)

Go to overhead

(4 e- domains)

tetrahedralsp3

(4 e- domains)

tetrahedralsp3

(4 e- domains)

tetrahedralsp3

(3 e- domains) sp2

trigonal planar

(2 e- domains)

linearsp

sp3d(5 e- domains)

(5 e- domains)

trigonal bipyramid

trigonal bipyramid

sp3d

(5 e- domains)

trigonal bipyramid

sp3d

(6 e- domains)

octahedralsp3d2

(6 e- domains)

octahedralsp3d2

(6 e- domains)

octahedralsp3d2

(5 e- domains)

trigonal bipyramidsp3d

sigma bonds “ s ”

pi bonds “ p ”

MolecularGeometries


Overlap and Bonding

• We think of covalent bonds forming through the sharing of electrons by adjacent atoms.

• In such an approach this can only occur when orbitals on the two atoms overlap.

…C=C…

1 sigma bond & 1 pi bond

SIGMA BOND

etheneC2H4

H2C=CH2

p bond

Solids lesson go to AP Solids pwpt

ap chem 2 nd semester 7-bonding 8-liq/soln 9-kinetics 10-equilibrium 11-acid-base 12-electrochem...

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