Ionic Bond Electrostatic Attraction Between Two Oppositely Charged Ions

BONDING

A. IONIC BONDS

Ionic bond – electrostatic attraction between two oppositely charged ions

+ _ + _ + _

cation + anion or metal + nonmetal or metal + polyatomic ion

· Very strong bond

· Involves a transfer of electrons

· Results in the formation of salts (or group 1A hydroxides)

(All salts are ionic compounds but not all ionic compounds are salts.)

Ionic compound – compound composed of positive and negative ions

+ +

Ex: Na + Cl à Na+Cl- Na + Cl à Na Cl Na Cl

Characteristics of Ionic Compounds:

· Crystalline solids at room temperature

· Have high melting points (MP indicates bond strength)

· Tend to dissolve in water

· Conduct electricity in the molten state

Look at the reaction of Na + Cl:

1. Nao + ionization energy à Na+1 + 1 e- ENDO (bonds breaking requires energy)

2. Clo + 1e- à Cl- + electron affinity energy EXO (bonds forming releases energy)

3. Na+1 + Cl- à NaCl + lattice energy EXO (bonds forming releases energy)

Lattice energy – the energy given off when two oppositely charged ions bond to form

a crystal lattice or solid salt.

Octet Rule – atoms react to achieve the electron configuration of a noble gas

One endothermic step and two exothermic steps results in a NET EXOTHERMIC process.

Compounds form to achieve the lowest possible energy.

Ionic Crystal Structure of Compounds

NaCl crystal: In the 3 D structure of a NaCl

crystal, each ion is surrounded

by 6 oppositely charged ions.*

There are, therefore, many strong

ionic bonds between all adjacent

oppositely charged ions.

These large attractive forces result

in a stable structure with a very

high melting point.

A crystal’s arrangement is a regular repeating pattern called a lattice.

· the shape of a crystal depends on the internal arrangement of the ions (or atoms)

· the many possible crystal arrangements of ions depends on their sizes

Coordination number – the number of ions of opposite charge that surround each ion in a crystal

*the coordination number of NaCl is 6

B. METALLIC BONDS

Metallic bonds – attractions of free-floating valence electrons for positively charged metal ions

- may form between atoms of the same metal or atoms of different metals.

Alloys – solutions of different metals made by melting metals together and cooling them

Atoms that form metallic bonds must have:

· vacant valence orbitals that can hold additional electrons

· low ionization energies so that loosely held electrons are available for bonding

Atoms in metals are packed in compact crystal forms that are most commonly cubic in shape.

Characteristics of Metals:

· malleable and ductile

· good conductors of electricity

C. COVALENT BONDS

Covalent bond – mutual attraction of different nuclei to the same electron’s orbital

- occurs between nonmetal atoms

· Very strong bond

· Involves a sharing of electrons

· Results in the formation of molecular compounds

Molecular compound – compound composed of nonmetals held together by covalent bonds

Ex: Cl + F à ClF Cl + F à Cl F Cl F

Characteristics of Molecular Compounds:

· Tend to be liquids or gases at room temperature (though some are solids)

· Have low melting points

· Contain 2 types of forces:

1. intramolecular forces – strong covalent bonds within the molecules

2. intermolecular forces – weaker attractive forces between the molecules

Ex: Water molecules in an ice crystal: + INTERMOLECULAR:

The positive end of one H2O mc

+ _ + _ is attracted to the negative end

INTRAMOLECULAR: of an adjacent H2O molecule.

Strong covalent bonds + _

in which electrons are +

shared by the O and H

Diatomic Molecule – molecule composed of 2 atoms (this includes all HONClBrIF’s)

Both nuclei attract the same

Look at H2 : electrons

Both atoms strive

H x H to fill their 1s orbital + + + +

so both H’s attract

the pair of bonding

electrons H – H atomic orbitals overlap

1s1 1s1 Shape of H2 is linear to form new molecular orbital

s to s bonding is non-directional because s orbitals can approach each other from any direction

and overlap to form a bond

Look at F2 : 2pz is where bonding occurs py py

F 1s2 2s2 2px2y2z1 x

x F x F + pz pz + x

py new molecular py

F 1s2 2s2 2px2y2z1 orbital forms

2pz orbitals overlap

so now both atoms F – F

achieve a stable octet Shape of F2 is linear

directional bonds - bonding in which atom’s orbitals must approach at a specific direction in order for the

orbitals to overlap to form a bond (s to p, p to p, s to d, & p to d)

non-polar covalent bond – covalent bond in which there is an equal attraction for the shared electrons

All bonds in HONClBrIFs are nonpolar (purely) covalent bonds where there is an equal sharing of electrons.

electronegativity – a number that indicates the relative attraction an atom has for a shared pair of

electrons in a bond; the higher the number, the greater the attraction

electronegativity difference ( EN) is an indication of the type of bond formed:

electronegativity difference 0 – .3 = nonpolar covalent bond (equal sharing of electrons)

electronegativity difference .4 – 1.6 = polar covalent bond (unequal sharing of electrons)

electronegativity difference 1.7 = ionic bond (transfer of electrons)

polar covalent bond – a covalent bond between 2 atoms in which one atom has a greater pull (attraction)

on the shared electrons (unequal sharing) + O

Look at H2O : _

x . . + +

O 1s2 2s2 2px2y1z1 H x O x H O : H

1 H’s 1s overlaps the H2O mc’s shape + H

H 1s1 O’s 2px lobe and the is bent b/c the bonds are polar b/c O has

other H’s 1s overlaps y and z axes are a greater pull on e- than H

H 1s1 the O’s 2pz lobe perpendicular O = 3.5; H = 2.1 EN = 1.4)

MOLECULAR POLARITY

polar molecule – a molecule in which one region of the molecule has a higher electron density causing

the molecule to have positive and negative regions or ends.

The polarity of a molecule is determined by the following:

· the polarity of the bonds

· the overall distribution of the valence electrons

· the symmetry/shape of the molecule

Compare the following molecules:

Methane Ammonia Water

CH4 NH3 H2O

H -- _

H C H H N H H O

H H H

Bonds: C - H Bonds: N - H Bonds: O - H

C = 2.5 N = 3.0 O = 3.5

H = 2.1 H = 2.1 H = 2.1

EN = .4 EN = .9 EN = 1.4

bonds are slightly polar bonds are moderately polar bonds are very polar

shape is tetrahedral shape is pyramidal shape is bent

molecule is non-polar molecule is moderately polar molecule is very polar

b/c mc is symmetrical b/c mc is asymmetrical b/c mc is asymmetrical

and has 1 exposed e- pair and has 2 exposed e- pairs

gas at room temperature liquid that evaporates easily liquid at room temp

(no attractive forces b/t mc) (weak intermolecular forces) (strong intermolecular forces)

low boiling point high boiling point

MOLECULAR GEOMETRY

VSEPR (valence shell electron pair repulsion) Theory – because electron pairs repel, molecules adjust

their shape so that the valance (outer shell) electron pairs are as far apart as possible.

Shape Bond Angle Example

. . . .

Linear 180o : F : F : F2

. . . .

. .

Bent 105o H : O : H2O

. .

(The asymmetric water molecule has 2 exposed (unshared) electron pairs on one side of the

molecule. Due to their repulsion of like charges, they require more space and consequently

decrease the angle between the two hydrogen atoms.)

Trigonal Planar 120o F BF3

. .

F F

. .

Pyramidal 107o H : N : H NH3

. .

Tetrahedral 109.5o H CH4

. .

H : C : H

. .

POLYATOMIC IONS

Polyatomic ion – a group of atoms that act as a unit and carry a charge

· Made up of nonmetals

· Bonds WITHIN polyatomic ions are covalent bonds

· Cannot exist independently (can only exist as part of a compound)

Examples: Hydroxide Hydronium Ammonium

OH- H3O+1 NH4+1

O H -1 H O H +1 H +1

H H N H

[ O - H ] -1 H–O–H H +1

H H-N-H

(linear) (pyramidal)

(tetrahedral)

Sulfite Sulfate Phosphate

SO3-2 SO4-2 PO4-3

O -2 O -2 O -3

O S O O S O O P O

O O

O -2 O -2 O -3

O S O O S O O P O

O O

(pyramidal) (tetrahedral) (tetrahedral)

Coordinate covalent bond – a covalent bond in which one atom contributes both electrons in the

shared pair

MOLECULAR ORBITALS (orbitals of molecules)

Molecular orbitals are the result of the overlap of atomic orbitals when 2 atoms form a bond.

Types of molecular orbitals:

1. sigma bond – bond formed when 2 atomic orbitals combine to form a molecular orbital

that is symmetrical along the axis connecting 2 nuclei

· END-TO-END orbital overlap is extensive therefore bonds are very strong

Two “s” atomic orbitals can overlap to form a sigma-bond molecular orbital:

s atomic orbital + s atomic orbital sigma-bonding molecular orbital

Two “p” atomic orbitals can overlap head to head to form a sigma-bonding molecular orbital:

p atomic orbital + p atomic orbital sigma-bonding molecular orbital

2. pi bond - bond formed when 2 atomic orbitals combine to form a molecular orbital in

which the bonding electrons are located in sausage shaped regions above and

below the bond axis

· SIDE-TO-SIDE orbital overlap is not as extensive (as in a sigma bond) therefore

the bonds tend to be weaker

Two “p” atomic orbitals can overlap side-to-side to form a pi-bonding molecular orbital:

p atomic orbital + p atomic orbital pi-bonding molecular orbital

ORBITAL HYBRIDIZATION

Hybridization – the combining of 2 or more orbitals of nearly the same energy into orbitals of equal energy

· occurs when atoms promote electrons into nearby orbitals to increase their bonding capacity

(occurs in atoms of Gr. 2, 3, 4 or 6A)

· occurs when several atomic orbitals mix to form the same number of hybrid orbitals

· hybrid orbitals are equivalent because they have the same size, shape and energy

· all sigma bonds and lone pair electrons (exposed or nonbonding e-) require hybrid orbitals

Look at BeH2: Beryllium promotes one of its2s electrons to its empty 2p orbital

Be’s one 2s orbital and one 2p orbitals then mix to form two sp hybrid orbitals

2s 2p hybrid sp orbitals unhybridized 2p orbitals

Be mix to form and

Now the 1s orbitals of the Hydrogen atoms can overlap with the sp hybrid orbitals of the Be:

The resulting molecule is linear: H – Be – H

because the sp hybrid orbitals form 180o angles

H H and one hydrogen overlaps on each end.

s s

H Be H Bonding Beryllium Hydride

atomic hybrid atomic molecule

orbital orbitals orbital

The H – Be bonds are sigma bonds because they are the result of the end-to-end orbital overlap.

Hybridization of Carbon - Carbon promotes one of its 2s electrons to its empty 2p orbital in order to

s p p p increase its bonding capacity.

Look at Methane: When methane forms – Carbon’s one 2s orbital and three 2p orbitals mix to

( CH4 ) form four sp3 hybrid orbitals to provide four bonding sites for the four H atoms.

s p p p s s s s sp3

C: H’s:

H H H H

+ +

Hydrogen Carbon Hydrogen

atomic orbitals hybrid orbitals atomic orbitals Methane molecule

All four bonds formed are sigma bonds because they are the result of end-to-end orbital overlap.

The shape of the methane molecule is tetrahedral because the sp3 bond angles are 109.5o

Look at Ammonia (NH3) :

s p p p When H bonds with N, four sp3 molecular orbitals must form:

N Three of the orbitals provide

bonding sites for the H’s

H’s One of the orbitals provides a

H H H site for the lone pair of e-

. .

· Not all of the hybrid orbitals of the central atom must be used for bonding, N

H H

That is, lone pair electrons can also be accommodated in hybrid orbitals. H

· The 3 bonds formed are sigma bonds because they are the result of end-to-end orbital overlap.

· The sp3 bond angles here are 107o, because the bonding electrons are repeled by the unshared

(lone) pair of electrons and require more room, therefore the molecular shape is pyramidal.

MULTIPLE BONDS

Look at Ethene: ( C2H4 ) : Each carbon will form two C – H single bonds and one C = C double bond

H H single bond - 1 shared pair of electrons