Electrons in outermost shell are called valence electrons and the valence shell is the outermost

occupied shell

Electron Configuration of IonsElectron Configuration of Ions

CATIONS CATIONS are formed when a neutral atom loses one or more electrons:

np electrons are lost first

ns electrons are lost second

nd electrons of previous shell lost third (if there’s any)

Fe [Ar] 3d6 4s2

Fe3+ [Ar] 3d5

Anions: Anions: to form monoatomic anions, add electrons until the next noble gas configuration

has been reached

N [He] 2s2 2p3 (room for 3 more electrons)

N3- [He] 2s2 2p6

O [He] 2s2 2p4 (room for 2 more electrons)

O2- [He] 2s2 2p6

When you gain an electron, you gain a positive charge

When you lose an electron, you lose a negative charge

Periodic Table and Electronic Configuration

Periodic table is divided into s,p,d & f blocks which are named for the last subshell of the element

that’s occupied

Groups 1 & 2 – s block [no. of group]

Groups 13-18 – p block [no. of group minus 10]

Transition metals – d block

Lanthanides & actinides – f block

PeriodsPeriods (horizontal rows) are numbered according to the principle quantum number of the valence

shell

Exceptions: Exceptions:

Helium 1sHelium 1s2 2 put in with the noble gases in group 18put in with the noble gases in group 18

Hydrogen 1sHydrogen 1s11 ( can act like group 1 & 17 ) ( can act like group 1 & 17 )

Periodicity of Atomic Properties

Atomic RadiusAtomic Radius of an element is defined as half the distance between the centres of neighbouring atoms.

•Values generally increase down a group and decrease from left to right across a period

Ionic RadiusIonic Radius of an element is its share of the distance between neighbouring ions in an ionic solution

•Values generally increase down a group and decrease from left to right across a period

• Cations are smaller and anions larger than their parent atoms

Ionisation energyIonisation energy of an element is the energy needed to remove an electron from an atom of an element in

the gas-phase

• 1st & 2nd ionisation energies

• generally higher for elements close to He and lower for elements close to Cs

•Very high value if electron is expelled from a closed shell

Electron Affinity, EElectron Affinity, Eeaea,, of an element is the energy released when an electron is added to a gas-phase atom

• generally positive values - generally energetically favourable

•elements with highest values are those close to O, F, Cl.

Chemical Bonds

Chemical bondChemical bond is a link between atoms

•A bond forms if the resulting arrangement of atoms has a lower energy than the sum of the energies of the separate atoms.

•Chemical bonds result from changes in location of electrons in the outermost shells of atoms (valence shells)

Ionic Bonds

An ionic bond is the electrical attraction between the opposite charges of cations and anions.

Lewis Symbols for atoms and ions - keeps track of valence electrons

Lewis Symbol : Symbol of Element + dot for each valence electron

H 1s1

He 1s2

O 1s2 2s2 2px2 2py

1 2pz1

N 1s2 2s2 2px1 2py

1 2pz1

A single dot: represents an electron alone in an orbital

A pair of dots: represents two paired electrons in the same orbital

To work out the formula of an ionic compound:

1. remove dots from Lewis symbol for metal atom.

2. transfer the dots to the Lewis symbol for the nonmetal atom and complete its valence shell

3. adjust the no.s of atoms to ensure all dots removed from metals are accommodated by nonmetals.

4. add charges of ions.

Ca [Ar] 4s2 Ca Cl [Ne] 3s2 3p5 Cl

Cl + Ca + Cl Ca2+Cl

-Cl

-

Chemical formula for calcium chloride is CaCl2

•Octet of electrons formed in each case

OCTET RULE: atoms of the reactive representative elements tend to undergo those chemical reactions that

most directly give them the electronic configuration of the nearest noble gas

Energy is needed to produce ions and when they pack together there’s a net overall lowering of energy - lattice

enthalpy is a measure of this attraction

Properties of Ionic Solids

• ions stack together in regular crystalline structures (crystalline solids)

• High m.p.s and b.p.s

• Brittle

• Form electrolyte solutions when dissolved in water

Covalent Covalent BondsBonds

• Compounds of nonmetals are not ionic - ionic compounds need ions of both positive and negative

charge.

•Atoms of nonmetals don’t become cations - too many electrons have to be lost to achieve noble gas

configurations

• Nonmetals form covalent bonds to one another by sharing pairs of electrons

• A covalent bond is a pair of electrons between two atoms.

• In covalent bonds, atoms share electrons to reach a noble gas configuration.

The valence of an element refers to the number of covalent bonds an atom of the element forms

H has one valence electron 1s1

H H+ H H H HLine represents shared pair of

electrons

H forms one bond, therefore its valence is 1

F has 7 valence electrons [He] 2s2 2p5

F F F F

F forms one bond, therefore its valence is 1

F2 has lone pairs of electrons - pairs of valence electrons not

involved in bonding. H2 has no lone pairs

Structure of Polyatomic Species

Methane CH4

HH H HC

Carbon is

tetravalent

Single bond: single shared pair of electrons (2)

Double bond: double shared pair of electrons (4)

Triple bond: triple shared pair of electrons (6)

Bond Order: the number of electron pair bonds that link 2 atoms

Single bond has a bond order of 1

C

H

H HH

• choose atom with lowest I.E. for central atom

• H is never central

•arrange atoms symmetrically around central atom

•central atom often written first in chemical formula (CH4)

1. Count total no. of valence electrons on each atom & divide by 2 to get no. of electron pairs

Ammonia NH3 N 1 x 5 = 5

H 3 x 1 = 3

Total = 8

8 2 = 4

4 valence e- pairs

2. Write chemical symbols of atoms to show layout in molecule

H

NH H

3. Place one electron pair between each pair of bonded atoms

( one pair remains)

H

NH H

4. Complete octet of each atom (or duplet for H) by placing

electron pairs as lone pairs around atoms

H

NH H

• If don’t have enough electron pairs to form octets, form multiple bonds

• Check each atom has an octet or duplet

Water H20 H 2 x 1 = 2

O 1 x 6 = 6

Total = 8

8 2 = 4

4 valence e- pairs

O

H H

O

H H

O

H H

Polyatomic Ions

A shared electron pair can originate from one atom

NH3 + H+ NH4+ (ammonium ion)

•If both shared electrons come from one atom- called a coordinate covalent bond

H

NH H + H+

H

NH H

H+

Resonance

Benzene C6H6

Kekulé structure

2 structures have exactly same energy - blend together as a resonance hybrid - electron

density is spread evenly around the ring

Resonance stabilises a molecule by lowering its total energy

Resonance occurs between structures with the same arrangement of atoms, but

different arrangement of electron pairs

C

CC

C

CC

H

H

H

H

H

H

Lewis Acids & Bases

Lewis Acid: electron pair acceptor

Lewis Base : electron pair donor

Acid + Base Complex

O

H H

H+ + O HH

H

Lewis acid Lewis base Complex

+

ElectronegativitElectronegativityy

• most bonds lie somewhere between pure ionic and pure covalent

ElectronegativityElectronegativity is the electron withdrawing power of an atom

A measure of the electron pulling power of an atom on an electron pair in a molecule

Highest at upper right hand corner of periodic table and lowest at bottom left hand corner

The greater the difference in electronegativities of 2 elements, the greater the extent of ionic character

The greater the difference in electronegativities of 2 elements, the greater the extent of ionic character

Molecular StructureMolecular Structure

Lewis structures : 2D diagrams & generally don’t show how molecules are arranged in space

VSEPR Model: Valence Shell Electron-Pair Repulsion Model

• molecules consist of central atom & attached atoms

• attached atoms lie to corners of different shapes - describe the shapes of the molecules

• Bond Angles: angles between the bonds

• regions of high electron concentration (found in bonds and lone pairs) repel one another and take up positions as

far away from each other as possible

PredictingPredicting Shapes of Molecules• write down Lewis structure & decide how electron pairs can be arranged around each “central” atom to minimise

repulsions

CASE A: Central atom with no lone pairs

1. CaCl2 CaCl Cl Ca ClCl

Linear with bond angle of 180o

2. BF3

B

F

FF

Trigonal Planar

120o

3. CH4

Tetrahedral

109.5o

5. SF6

Trigonal bipyramidal

120o & 90o

4. PCl5

Octahedral

90o

C

H

H HH

CASE B: Molecules with multiple bonds

• VESPR theory doesn’t distinguish between single & double bonds

• Electron pairs in a double bond act as a single unit of high electron concentration

CO2Linear

CASE C: Molecules with lone pairs on central atom

VESPR formula; A = central atom

X = atom bonded to central atom

E = lone pair of electrons on central atom

NB: lone pairs on attached atoms not included

CO O

e.g. BF3 = AX3 species (no lone pairs on central atom)

NH3 = AX3E species (1 lone pair on nitrogen)

Electron arrangement:Electron arrangement: 3D arrangement of all regions of high electron concentration (bonds and lone pairs) around

central atom.

AX3E species has 4 regions of high electron density

Strengths of repulsions are in the order:

lone pair-lone pair > lone-pair-bonding pair

> bonding pair - bonding pair

Water H2O

H O H AX2E2

• 4 electron pairs adopt a tetrahedral arrangement

• only 2 positions are occupied by atoms - classified as angular or “bent”

• lone pairs push away from each other

• bonding atoms forced closer together

O

H H

Bond angle is less than that of tetrahedron

( 104.5o compared to 109o)

Ammonia NH3Ammonia NH3

H

NH H AX3E

• 4 pairs adopt a tetrahedral arrangement

• 3 positions occupied by atoms

• H atoms have moved slightly towards each other

H

NH HTrigonal pyramidal

( 107o compared to 109.5o )

VSEPR typeNo. of e - pairs in bonds

No. of lone pairs

Structure

Example

AX2

AX2E2

AX3

AX3E

AX4

AX5

AX6

2

2

3

3

4

5

6

0

2

0

1

0

0

0

linear

nonlinear bent

planar triangular

trigonal pyramidal

tetrahedral

trigonal bipyramida

l

octahedral

BeCl2

CaCl2H2O

H2SBCl3

BF3

NH3

CH4

PCl5

SF6

Polar Polar MoleculesMolecules

Polar Covalent Bond: electron pair shared unequally between 2 atoms resulting in some partial ionic character

• arises from differences in electronegativities of atoms

O H Polar bondO more electronegative than H &

gains greater share in the bonding e- pair

O- H+- = O has a partial negative charge

+ = H has a partial positive charge

Two atoms in a polar bond with partial charges give rise to an electric dipole moment

When bonded atoms have different electronegativities, the bond is polar

Nonpolar molecule has zero dipole moment e.g. H2 & Cl2

O OC

+- - Polar bonds, but overall a non-polar molecule

O

H H

+

+

-

Polar bonds and polar molecule

AX2, AX3, AX4, AX5 & AX6 - non-polar (where X = same element)

Bond Strengths

• strength of bonds is measured by bond enthalpy, HB

• bond breaking requires energy, so all HB values are positive

• HB typically increase as bond order increases and decrease as atomic radius and no. of lone pairs increases

Bond Lengths

• Bond length: distance between the centres of 2 atoms joined by a chemical bond

• for bonds between same atoms, the shorter the bond, the stronger it is

e.g. C C is shorter & stronger than C C

• covalent radii- added to estimate bond lengths in molecules

1s 1s

Molecular Orbitals - another view of the covalent bond

H2

bond

(sigma bond)

Looking along internuclear axis,

e- distribution resembles that of an s orbital

Molecular Orbitals - another view of the covalent bond

• a shared e- pair resides in a molecular orbital formed by the partial overlapping of two atomic orbitals

• space created by the overlapping of 2 atomic orbitals is called a molecular orbital

• M.O. Can hold a maximum of 2 e- with opposite spins (like an A.O.)

• A.O. That overlap are generally those of valence shell electrons

• p-orbitals can also form sigma bonds

N2

N 1s2 2s2 2px1 2py

1 2pz1

Two 2pz orbitals form a “head-on” sigma bond

Two 2px orbitals form a pi bondLooking along internuclear

axis,

e- distribution resembles that of an p orbital

• remaining two 2py orbitals form another bond

• N2 has one sigma bond ( from both 2pz

orbitals) and two pi bonds (from both 2px and 2py orbitals)

A single bond is a bond

A double bond is a & a bond

A triple bond is a & two bonds

Promotion & HybridisationBonding in polyatomic orbitals more complex to explain e.g. CH4

• C 1s2 2s2 2px1 2py

1 (looks like C can only form 2 bonds)

• Know carbon can nearly always form 4 bonds. HOW?HOW?

• NB NB carbon has one empty p orbital (2pcarbon has one empty p orbital (2pzz))

• If we promote an eIf we promote an e-- from the 2s orbital into this empty from the 2s orbital into this empty 2p2pzz orbital, get 4 unpaired electrons. orbital, get 4 unpaired electrons.

2s

2p

1s

2s

2p

1s

Before promotion

(can only form 2 bonds)

After promotion

(can form 4 bonds)

BUT now looks like there are 2 different bond types in CH4

Type 1 H 1s C 2s

H 1s C 2px

H 1s C 2py

H 1s C 2pz

Type 2

•In fact, four sp3 hybrid (mixture) orbitals formed from combination

of s orbital and 3 p orbitals

• These orbitals have equal energies lying between the energies of s and p orbitals

sp3 orbital

Bonding in Methane (CH4)

• One unpaired electron occupies each of carbon’s sp3 hybrid orbitals

• Each of these 4 electrons can pair with an electron in a H 1s orbital

• 4 sigma () bonds are formed

• Structure of methane is tetrahedral