BONDINGUnit 4

OVERVIEW Valence electrons Ions (cations/anions) Types of bonds

Ionic, covalent Properties of bonds

Bond strength, hardness/texture, melting point/boiling point, conductivity

Metallic Bonds Lattice Energy Electronegativity & Bonds Lewis Dot Structures

Octet Rule, covalent compounds, exceptions, ions, resonance, formal charge, isomers

Bond order, energy, length

Ionic Lewis Structures VSEPR Theory

Molecular Geometry Intermolecular Forces

Hydrogen Ion-dipole Dipole-dipole London dispersion

Valence Bond Theory Hybridization Sigma/pi bonds

HOW ARE COMPOUNDS FORMED?

Electrons are lost, gained, or shared.

Valence electrons: the electrons available to be lost, gained, or shared Number of valence electrons is electrons in s and

p orbitals of highest energy level (outermost shell)

Main groups on table correspond to number of valence electrons.

IONS

Atom(s) that has a negative (-) or positive (+) charge

Cation: positively charged ion

Anion: negatively charged ion

Na Na+ + e-

e- + Cl Cl-

CATIONS

Forms by losing electrons Atoms with less than 4 valence electrons

generally lose electrons

ANIONS

Forms by gaining electrons Atoms with more than 4 valence electrons

generally gain electrons

IONIC BOND Electrostatic attraction between cation and anion

(opposite charges attract) For example, an uncharged chlorine atom can pull

one electron from an uncharged sodium atom, yielding Cl−and Na+.

Cation becomes smaller and anion becomes larger Typically between a metal and nonmetal

IONIC BONDS

Form crystal lattice Solid crystal at ordinary

temperatures Organized in a characteristic

pattern of alternating positive and negative ions

All salts are ionic compounds and form crystals

Ionic compounds do not exist as single molecules

COVALENT BOND

Bond involves sharing of electrons

Forms between atoms of two nonmetallic elements

Forms molecules

Nonpolar covalent bond: electrons are shared equally resulting in an even distribution of the negative charge (no partial negative charge)

Polar covalent bond: one atom in the bond attracts electrons more than the other atom making the electron negative charge shift to that atom giving it a partial negative charge

BOND TYPES

BOND STRENGTH

Strong bonds Crystal lattice

Repeating symmetrical pattern

Weaker bonds Covalent Network Molecules

Strong forces within molecule but weak between molecules

IONIC COVALENT

MELTING & BOILING POINT

High melting/boiling points Sturdy crystal lattice

Low volatility

Low melting/boiling points Weak forces between

molecules Most are gases High volatility

IONIC COVALENT

TEXTURE & HARDNESS

Brittle Sturdy but collapses

easily if disrupted Hard structure

Soft compounds Most are gases

IONIC COVALENT

CONDUCTIVITY – IONIC

Solids – not conductors Ions can’t move

Molten state – conductors Solution – conductors (ions

separate)

Not good conductors

IONIC COVALENT

METALLIC BOND

Results from attraction between metal cations and the surrounding “sea of electrons”

Vacant p and d orbitals in metal's outer energy levels overlap, and allow outer electrons to move freely throughout the metal

Valence electrons do not belong to any one atom

METALLIC BONDS Very strong bonds

Due to “sea of electrons” Highest melting/boiling points

Electrons result in strong forces holding together Low volatility Very hard

METALLIC BONDS

Malleable/ductile When struck, one plane of metal atoms can slide past

another without breaking Great conductors (heat and electricity)

Freedom of electrons carries current Shiny and have luster

Electrons absorb light, get excited, fall, re-radiate the light

LATTICE ENERGY

The energy required to separate ions of an ionic solid

Magnitude of lattice energy depends on Charges of ions Sizes of ions Arrangement in ions

in the solid

Lattice energyCompound kJ/molLiCl 834NaCl 769KCl 701NaBr 732Na2O 2481Na2S 2192MgCl2 2326MgO 3795

Lattice energyCompound kJ/molLiCl 834NaCl 769KCl 701NaBr 732Na2O 2481Na2S 2192MgCl2 2326MgO 3795

For given arrangement of ions, lattice energy

increases as the charges on the ions increase and as their radii decrease

ELECTRONEGATIVITY

Atom’s ability to attract electrons Difference in electronegativity values

between two elements can determine type of bond Ionic Bonding = Over 1.7 Polar Covalent = 1.7 - 0.4 Non-polar Covalent = 0-0.3

Examples KCl = 2.2 (ionic) CuS = 0.6 (polar covalent) O2 = 0.0 (nonpolar covalent)

LEWIS DOT NOTATIONS

Valence electrons are represented by dots drawn around the symbol of an element

1 valence e-

X2 valence e-

X3 valence e-

X4 valence e-

X5 valence e-

X6 valence e-

X7 valence e-

X8 valence e-

X

OCTET RULE

Compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has 8 electrons in its outer shell.

Hydrogen is exception and only needs 2 electrons for a complete shell

8Electrons!

LEWIS DOT STRUCTURES

Show electron distribution in compounds Lone pair: nonbonding pair of valence electrons Bond pair: valence electrons shared between

two elements Shared electrons pairs represented by two

dots (:) or by a single line ( - )

HCllone pair

shared orbond pair

•••• ••••

STEPS FOR BUILDING A DOT STRUCTURESTEPS FOR BUILDING A DOT STRUCTURE

Ammonia, NH3

1. Add up the total number of valence electrons that can be used.

H = 1 and N = 5

Total = 1 + 1 + 1 + 5 = 8 electrons

2. Decide on the central atom;

If carbon it is always central atom

If no carbon, choose the least electronegative atom

Never use hydrogen!

Therefore, N is central on this one

3. Draw the basic skeletal structure of the molecule.

H H

H

N

Building a Dot Structure

H H

H

N4. Draw in the electrons around the central atom. Then place the remaining electrons to form complete octets and lone pairs as necessary.

N has 3 BOND PAIRS and 1 LONE PAIR.

..

.. .. ..

5. Check to make sure there are 8 electrons around each atom except H.

Building a Dot Structure

6. Count electrons! Make sure that the number of electrons in your structure equals the total number from the beginning.

H HHN.... .. ..

H HHN.... .. ..Total Electrons = 8

MULTIPLE BONDS (DOUBLE)

C C

H

H

H

H

Two pairs of shared electrons

Example: ethene

C C

H

H

H

H

OR

MULTIPLE BONDS (TRIPLE)

C C HH

Three pairs of shared electrons

Example: ethyne

C C HH

OR

EXCEPTIONS TO THE OCTET RULE

Element is 3rd period or higher, is the central atom, AND is bonded to electronegative atoms (such as O, F, Cl, Br) may have more than 8 electrons

Electrons use empty valence d orbitals

Be is stable with 4 electrons

B is stable with 6 electrons

LEWIS DOT STRUCTURES FOR IONS

Add or remove electrons based on charge of ion If the ion has a negative (-) charge add electrons

to the Lewis structure. If the ion has a positive (+) charge, then

subtract electrons from the Lewis structure.

Try CO3-2 O

O

C O

-2

SPECIES WITH AN ODD TOTAL NUMBER OF ELECTRONS A very few species exist where the total

number of valence electrons is an odd number

This must mean that there is an unpaired electron which is usually very reactive.

Radical Species that has one or more unpaired electrons They are believed to play significant roles in

aging and cancer. Example – NO

It has a total of 11 valence electrons: 6 from oxygen and 5 from nitrogen.

:N::O:

:.

O O O O O O

O

O

O

Draw O3

O

O

O

OO

O

Which are the same and which are different?

RESONANCE STRUCTURES Occurs when more than one valid Lewis

structure can be written for a particular molecule.

The resonance structures are the same except for the placement of the electrons (meaning the bonds).

RESONANCE BONDS

Resonance bonds are shorter and stronger than single bonds.

Resonance bonds are longer and weaker than double bonds.

H

H

H

H

H

H

H

H

H

H

H

H

FORMAL CHARGES A bookkeeping system for electrons.

Does not give actual charge of atoms Helps decide which Lewis structure is more preferred

Used to show the approximate distribution of electron density in a molecule or polyatomic ion

Assign an atom electrons: Each atom gets half of the bonding electrons it has

(single, double, and triple bonds) Each atom gets all unshared (nonbonding) electrons

that are found on it

FORMAL CHARGES

To solve for formal charge:

(Total valence electrons) – (electrons assigned to atom)

To decide which structure is preferred:

Choose the Lewis structure in which the atoms bear formal charges closest to zero

Choose the Lewis structure in which any negative charges reside on the more electronegative atoms

FORMAL CHARGES

Example: CO2

For each oxygen4 electrons from unshared electrons2 from the bonds (1/2 of 4)6 total

Formal charge = 6 - 6 = 0

For carbon 0 unshared electrons

4 from the bonds (1/2 of 8)4 total

Formal charge = 4 - 4 = 0

O=C=O

ISOMERSIsomers – compounds whose molecules have

the same overall molecules but different structures

C C O

H

H

H

H

H

H C CO

H

H

H

H

H

H

C CO

H

H

H

H

H

HC C O

H

H

H

H

H

H

BOND ORDER

Refers to the average number of bonds that an atom makes in all of its bonds to other atoms Draw Lewis structures and determine bonds Single bond counts as 1 bond, double counts as

2, and triple counts as 3

Bond Order =

Examples: CH3Cl = 4/4 = 1 CO2 = 4/2 = 2

CO3-2 = 4/3 = 4/3 ClO4

-1 = 4/4 = 1

number of bondsnumber of atoms bonded

BOND ENERGIES Bond energies and lengths differ between single,

double, and triple bonds Bonds between elements become shorter and

stronger as multiplicity increases. The greater the bond energy, the shorter the

bond length

Bond Bond Length Bond energy type order pm kJ/mol

C C 1 154 347C C 2 134 615C C 3 120 812

BOND LENGTH AND ENERGYBond Length

(pm)Energy

(kJ/mol)

C - C 154 346

C=C 134 612

CC 120 835

C - N 147 305

C=N 132 615

CN 116 887

C - O 143 358

C=O 120 799

CO 113 1072

N - N 145 180

N=N 125 418

NN 110 942

IONIC COMPOUNDS & LEWIS STRUCTURES

Na + Cl Na + Cl

The electron from Na is given to the Cl. Now both satisfy the octet rule.

Below, two electrons are given to S, one from each K.

+

VSEPR model

Most important factor in determining geometry is relative repulsion between electron pairs.

Electrons arranged so that pairs are as far apart from each other as possible.

Occurs in 3-dimensional space

Molecule adopts the shape that minimizes the electron

pair repulsions.

Valence Shell Electron Pair Repulsion theory

MOLECULAR GEOMETRY

Molecules have specific shapes.

Determined by the number of electron pairs around the central species

• Bonded and unshared pairs count (unshared pairs take up slightly more space)

• Multiple bonds are treated as a single bond for geometry.

• Geometry affects factors like polarity and solubility.

1 atom bonded to another atom

Electron Domains

Basic Geometr

y

0 lone pair

1 lone pair

2 lone pairs

3 lone pairs

4 lone pairs

1 Linear180°

       

1 ELECTRON DOMAIN

2 atoms, or lone electron pairs, or a combination of the two, bonded to a central atom.

Electron Domains

Basic Geometry

0 lone pair 1 lone pair 2 lone pairs 3 lone pairs

2 Linear180°

Linear180°

   

2 ELECTRON DOMAINS

3 atoms, or lone electron pairs, or a combination of the two, bonded to a central atom.

Electron Domains

Basic Geometry

0 lone pair 1 lone pair 2 lone pairs 3 lone pairs

3trigonal planar120°

bent / angular<120°

 

3 ELECTRON DOMAINS

4 ELECTRON DOMAINS 4 atoms, or lone electron pairs, or a

combination of the two, bonded to a central atom.

Electron Domains

Basic Geometry

0 lone pair 1 lone pair 2 lone pairs

4Tetrahedral

109.5°Pyramidal<109.5°

bent / angular<109.5°

5 ELECTRON DOMAINS 5 atoms, or lone electron pairs, or a

combination of the two, bonded to a central atom.

Electron Domains

Basic Geometry

0 lone pair 1 lone pair 2 lone pairs 3 lone pairs

5trigonal

bipyramidal180°, 120°,

90°

seesaw/ sawhorse t-shape

linear

6 ELECTRON DOMAINS 6 atoms, or lone electron pairs, or a

combination of the two, bonded to a central atom.

Electron Domains

Basic Geometry

0 lone pair 1 lone pair 2 lone pairs 3 lone pairs

6Octahedral90°, 180°

square pyramid

square planar

INTERMOLECULAR FORCES

Forces of attraction between molecules Forces within molecules are intramolecular

Why would boiling point be a good indicator of intermolecular force strength?

HYDROGEN BONDING H bonded to N, O, F

Example: Water Strongest intermolecular force

Large electronegativity difference Size of H atom allows it to get close to unshared

pair of electrons in adjacent molecule

POLARITY

A molecule, such as HF, that has a center of positive charge and a center of negative charge is said to be polar, or to have a dipole moment.

H F+ -

Slight positive sideSmaller electronegativity

Slight negativeLarger electronegativity

DIPOLES

Dipole – created by equal but opposite charges that are separated by a short distance Molecules with dipoles are polar because of

uneven charge distribution Direction is from its positive to negative pole Molecules can have multiple dipoles Dipoles can cancel each other out if in equal and

opposite directions

DRAW THE DIRECTION OF THE DIPOLES FOR THE FOLLOWING MOLECULES:

ION-DIPOLE FORCE

Exists between ion and the partial charge on the end of a polar molecule Cations attracted to negative end of

dipole Anions attracted to positive end of

dipole Magnitude of attraction increases

as charge of ion increases as magnitude of dipole increases

Stronger than dipole-dipole forces

DIPOLE-DIPOLE FORCES

Not as strong as ion-dipole forces Strength of attraction increases with

increasing polarity

Example: HCl

LONDON DISPERSION FORCES

The constant motion of electrons sometimes creates temporary dipoles for an instant Momentary uneven charge creates a dipole Can induce a dipole in another molecule

LONDON DISPERSION FORCES

Weakest intermolecular force Sometimes called dipole induced dipole

All molecules have LDF Dependent on motion of electrons so more

electrons means more chances for LDF LDF strength increases

with increasing molar mass

Example: O2

VALENCE-BOND THEORY Covalent bonding involves sharing of

electrons Electrons exist in orbitals

Orbitals overlap so that electrons can form pairs to make a bond

As orbitals overlap, they mix and form new hybrid orbitals Mixing of the orbitals is called hybridization

HYBRIDIZATION = the blending of orbitals

Poodle

+

+ Labrador

=

=

=

=

+

+s orbital p orbital

Labradoodle

sp orbital

EXAMPLE: CH4

Orbital notation for carbon:

Carbon only has 2 electrons available to bond, but it has to make 4 bonds

Carbon promotes one of its electrons to the 2p orbital so that each electron can pair up with the 1s electron from each of the four carbons

1s orbitals of four hydrogen atoms

Promoted electrons in carbon allow 4 bonds

EXAMPLE: CH4

The overlap of carbon’s 1 electron in the s orbital and 3 electrons in the p orbital creates four hybrid orbitals This hybridization is called sp3

The new hybrid orbitals have more energy than an s orbital but less than a p orbital

HYBRID ORBITALS

To determine hybridization Draw Lewis structure Determine electron domains for target atom Hybrid orbitals correlate to number of domains

Orbital Name Orbitals Combined Electron domains

sp 1 s / 1 p orbital 2

sp2 1 s / 2 p orbitals 3

sp3 1 s / 3 p orbitals 4

sp3d 1 s / 3 p / 1 d orbital

5

sp3d2 1 s / 3 p / 2 d orbitals

6

SIGMA AND PI BONDS

Sigma () bonds exist in the region directly between two bonded atoms.

Exist on the line (internuclear axis) between two atoms

Single bonds

Sigma bonds

s s

s p

p p

SIGMA AND PI BONDS Pi () bonds exist in the region above and below a line

drawn between two bonded atoms. Exist perpendicular to the line (internuclear axis) between

two atoms (double and triple bonds)

Pi bonds

Single bond 1 sigma bond

Double Bond 1 sigma, 1 pi bond

Triple Bond 1 sigma, 2 pi bonds