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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

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slide1

UNIT 7:

CHEMICAL BONDING

slide2

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 brokenabsorbs (requires) energy
  • [endo, so DH = (+)]
  • being madereleases energy
  • [exo, so DH = (-)]
slide3

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

slide4

BONDING TYPES

  • IONICbond is the
  • electrostatic attraction between
  • (+) cations and (-) anions
  • following the transfer of electrons
  • from metal atoms to non-metalatoms
  • called “ion-ion interactions
  • COVALENT bond is the
  • sharing of electron pairs between non-metal atoms
  • NON-POLARPOLAR
  • equalsharing of unequal sharing of
  • bonding electrons bonding electrons

distorted cloud

symmetrical cloud

slide5

BOND TYPEdepends on the

  • ELECTRONEGATIVITY OF THE ATOMSin 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 leftright across a period,
  • decreases down a group
electronegativity
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.
slide8

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

Actually more a continuum than clear-cut boundaries:

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

polar covalent bonds
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).
polar covalent bonds1
Polar Covalent Bonds
  • \

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

slide11

d+

d-

d-

Dipole moment is 1.90 Debye

Polar molecule

Dipole moment is zero

Non-polar molecule

energetics of ionic bonding
Energetics of Ionic Bonding

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

energetics of ionic bonding1
Energetics of Ionic Bonding

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

energetics of ionic bonding2
Energetics of Ionic Bonding

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

energetics of ionic bonding3
Energetics of Ionic 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.
lattice energy

Q1Q2

d

Eel = 

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:
slide19

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

slide21

Attractive forces between ions described by

COULOMB’S LAW:

charges on the ions

E = k

Q1+ Q2-

+

-

Lattice Energy:

attractive force

between ions

r

distance between ion centers

in the lattice

constant

lattice energy1
Lattice Energy
  • Lattice energy, then, increases with the charge on the ions.
  • It also increases with decreasing size of ions.
lattice energy2
Lattice Energy
  • Lattice energy, then, increases with the charge on the ions.
  • It also increases with decreasing size of ions.
slide26

Which compound has

a) more ionic bond character? ______

CaO

the one with the greater E.N. difference

CaO

b) a stronger ionic bond? ______

the one with the greater lattice energy

1) CaO or NaCl

MPt: 801oC

MPt: 2613oC

= -1

= -4

Ca2+ O2-

Na+ Cl-

larger chargesgreater lattice energy so stronger bond

(the more heat released, the lower the PE)

E.N Diff:

3.5-1.0

= 2.5

E.N Diff:

3.0-0.9

= 2.1

slide27

2) CaO or MgO

MPt: 2825oC

2+

EN Diff:

3.5-1.2

= 2.3

2-

2-

2+

E = -4

r

E = -4

r

So MgO has stronger ionic bond

But CaO has more ionic bond character

  • CaCl2 or K2S

both ions smaller

E = -2

r

E = -2

r

EN Diff:

2.5-0.8

= 1.7

EN Diff:

3.0-1.0

= 2.0

So CaCl2 has stronger ionic bond

and more ionic bond character

slide28

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

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

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

to form the cation

3. Dissociation EnergyN2(g) -----> 2N(g)

for diatomic (also called

Bond Energy)

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

to form the anion

5. Lattice EnergyM+(g)+ N-(g)---> MN(s)

when gaseous ions come together

energetics of ionic bonding4
Energetics of Ionic 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.

slide31

Use Table of Bond Energies!!

Chemical reactions involve bond-breakingand bond-making.

Each chemical bond has an ___________________ in kJ/mol.

BOND ENERGY is __________________________________

Always ______ (_____)

BOND LENGTH is __________________________________

Like leggos!!

See Handout with Table

“average” bond energy

the energy required to break a bond

(+) endo

the distance between 2 nuclei connected by

a bond

BOND ENERGY is __________________________________

slide32

-

Bond Order 1 2 3

atoms

pulled closer

# e-prs

in a bond

slide33

-

shortens

1) As # of shared pairs increases, the bond length ___________

2) The _______the bond energy, the ________ the bond.

greater

stronger

“more stable”

slide34

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

  • initial minus final

“D” means bond energy

[Note: The only time we use “initial minus final”,

rather than “final minus initial”!]

average bond enthalpies
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.
enthalpies of reaction
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

slide37

making

bonds

breaking

bonds

436 kJ/mol

239 kJ/mol

427 kJ/mol

Moles cancel!!

2 431 kJ

242 kJ

436 kJ

[ 671 kJ ] - [854kJ ]

-184kJ/ 2mol HCl produced

We can get to

-92kJ/mol

f

slide38

1 2

3

1

2

3

4 5

6

  • 391 kJ/mol
  • 941 kJ/mol
  • 436 kJ/mol
  • 6 391 kJ
  • 3 436 kJ
  • 941 kJ
  • [ 941 + 1308 ] - [2346 ]
  • -97kJ/ 2mol NH3produced
  • -49kJ/mol
  • -46
  • f
slide39

Group A # gives # val e-s

5 7

5 1

NH3

PCl3

5 + 21 = 26 total

5 + 3 = 8 total

slide40

lost e-

gained e-

slide43

now spread bonds

equidistant from

each other

trigonal planar

(polyatomic ion)

slide44

Day 5

  • 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.
slide45

Rules:

  • To count electrons in formal charges:
  • lonepairs = 2e-
  • (“unshared” or “non-bonding pairs”)
  • 2) singlebonds = 1e-
  • (“shared” or “bonding pairs”)
  • 3) doublebonds = 2e-
  • 4) triple bonds = 3e-
  • Now: Go to Overhead
slide46

For 2 “non-equivalent”Lewis Structures,

choose the one with:

formal charges closest to zero,

and the (-) formal charge is on the

most electronegative atom

slide47

Day 6

Electronic Geometry:

geom. of ELECTRON DOMAINS

around an atom

(used to find hybridization

around a particular atom)

Molecular Geometry:

geom. of ATOMS

in the molecule

VSEPR Theory

meansValence Shell Electron-Pair Repulsion Theory

1) all electrons “paired”

2) all atoms have stable octet, H has duet

3) pairs are spread equidistant from each other around atom

used to determine 3-dimensional geometry of molecules

and ions

slide48

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 bonds
  • 2) in lone pairs
  • called “electron domains”
molecular shapes
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.
what determines the shape of a molecule
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.
slide51

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”.
  • CO2has 2“electron domains”
  • 0 lone pairs
  • 2 bonds

107o

Lone pairs exert

greater repulsion

than bonding prs.

EG - tetrahedral

MG - pyramidal

180o

EG - linear

  • when NOLONE prs,
  • EG & MG the same

MG - linear

electron domains
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.
valence shell electron pair repulsion theory vsepr
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.”

valence shell electron pair repulsion theory vsepr1
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.”

electron domain geometries
Electron-Domain Geometries

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

electron domain geometries1
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.
molecular geometries
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.
linear electron domain
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.
trigonal planar electron domain
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.
slide60

not 120o

  • O3has 3“electron domains”
  • 1 lone pair
  • 2 bonds
  • NO3has 3 “electron domains”
  • 0 lone pairs
  • 3 bonds

116.8o

  • EG– trigonal planar
  • MG – bent

120o

  • EG– trigonal planar
  • EG– trigonal planar
slide61

CH4 has 4 “electron domains”

  • 0 lone pairs
  • 4 bonds
  • 109.5o
  • EG – tetrahedral
  • MG– tetrahedral
slide62

HYBRIDIZATION

  • While VSEPR Theory helps us predict the “spatial arrangement” of atoms in a molecule or ion, Valence Bond Theorydescribes “how bonding occurs” in terms of overlapping atomic orbitals.
  • These orbitals are often mixed or “hybridized” to form new “hybrid orbitals” which have
  • lower energy (stability) than the separate atomic orbitals
  • degenerate with each other (equal in that lower energy)
  • equidistant from each other (evenly spaced) about the atom
hybrid orbitals
Hybrid Orbitals
  • Consider beryllium:
    • In its ground electronic state, beryllium would not be able to form bonds, because it has no singly occupied orbitals.
hybrid orbitals1
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.

hybrid orbitals2
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.
hybrid orbitals3
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.
hybrid orbitals4
Hybrid Orbitals
  • These two degenerate orbitals would align themselves 180 from each other.
  • This is consistent with the observed geometry of beryllium compounds: linear.
hybrid orbitals5
Hybrid Orbitals

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

hybrid orbitals6
Hybrid Orbitals

With carbon, we get four degenerate sp3 orbitals.

valence bond theory
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.
valence bond theory1
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.
sigma bonds
Sigma () Bonds
  • Sigma bonds are characterized by
    • Head-to-head overlap.
    • Cylindrical symmetry of electron density about the internuclear axis.
sigma bonds1
Sigma () Bonds
  • Sigma bonds are characterized by
    • Head-to-head overlap.
    • Cylindrical symmetry of electron density about the internuclear axis.
pi bonds
Pi () Bonds
  • Pi bonds are characterized by
    • Side-to-side overlap.
    • Electron density above and below the internuclear axis.
single bonds
Single Bonds

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

multiple bonds
Multiple Bonds

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

slide77

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

slide78

(4 e- domains)

sp3

tetrahedral

sp3

(4 e- domains)

tetrahedral

sp3

(4 e- domains)

tetrahedral

sp2

(3 e- domains)

trigonal planar

slide79

sp

  • (2 e- domains)
  • linear
  • sp3d
  • (5 e- domains)
  • trigonalbipyramid
  • sp3d
  • (5 e- domains)
  • trigonalbipyramid
slide80

sp3d

  • (5 e- domains)
  • trigonalbipyramid
  • sp3d2
  • (6 e- domains)
  • octahedral
  • sp3d2
  • (6 e- domains)
  • octahedral
slide81

sp3d2

  • (6 e- domains)
  • octahedral
  • sp3d
  • (5 e- domains)
  • trigonalbipyramid
slide82

sigma bonds “ s ”

pi bonds “ p ”

overlap and bonding
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.
slide84

…C=C…

1 sigma bond

&

1 pi bond

slide85

ethene

C2H4

  • H2C=CH2

p bond

SIGMA BOND