Unit 2 - Bonding

9.1

Valence electrons are the outer shell electrons of an atom. The valence electrons are the electrons thatparticpate in chemical bonding.

I 1ns1

II 2ns2

III 3ns2np1

IV 4ns2np2

V 5ns2np3

VI 6ns2np4

VII 7ns2np5

Group # of valence e-e- configuration

9.1

Lewis Dot Symbols for the Representative Elements &Noble Gases

Trends in the Periodic TableThe elements have a regular (periodic)

recurrence of physical and chemical properties

The horizontal rows are called periodsThe vertical columns are called groups

Atomic SizeCovalent Atomic Radius – half of the distance

between the nuclei of two atoms in a homonuclear diatomic molecule

Group trend – atomic size increases as we move down a groupThe farther away the e- from the nucleus, the less

strongly they are held Periodic trend – atomic size decreases as we

move from left to rightNuclear charge increases as we move left to

right, pulling the e- closer to the nucleus

Ionic SizeCations are always smaller than neutral

atoms because they have less e-

Loss of outer shell e- results in increased attraction by the nucleus for the remaining e-

Anions are always larger than neutral atoms because they have more e-

Addition of outer shell e- results in less attraction to the protons of the nucleus

Ionization EnergyThe energy required to remove an electron from

a gaseous atomNa (g) + Energy Na+

(g) + e-

Group trend – Ionization energy decreases as we move down a group because the size of the atom is increasing, with more orbitals of e-

, allowing the outermost e- to be easily removed

Periodic trend – Ionization energy increases from left to right. Nuclear charge is increasing so more energy is required to remove e-

Electron AffinityThe energy change (release) that accompanies the

addition of an electron to a gaseous atomF (g) + e- F-

(g) + EnergyGroup trend – Electron affinity generally decreases

with increasing atomic size (releases less energy as you go down)

Periodic trend – Generally increases from left to right because atoms become smaller due to increased nuclear charge (releases more energy as you go left to right)

ElectronegativityThe tendency for an atom to attract electrons

to itself. It is affected by the distance from the atom’s valence electrons to the nucleus.

Group trend – Electronegativity usually decreases as you move down a group

Periodic trend – Electronegativity usually increases as you move across a period

Do the Noble Gases have electronegativity?

Why do substances bond? More stability Atoms want to achieve a lower energy state

Chemical bonds: an attempt to fill electron shells

1. Ionic bonds – 2. Covalent bonds –3. Metallic bonds

Ionic BondingBetween a metal and a non-metal with very different

electronegativity.

Metals lose electrons becoming a cations, while non-metals gain electrons becoming anions.

An ionic bond is an electrostatic attraction between the oppositely charged ions.

Ionic Bonds: One Big Greedy Thief Dog!

. Ionic bond – electron from Na is transferred to Cl, this causes a charge imbalance in each atom. The Na becomes (Na+) and the Cl becomes (Cl-), charged particles or ions.

9.2

Li + F Li+ F -

The Ionic Bond

1s22s11s22s22p5 1s21s22s22p6[He][Ne]

Li Li+ + e-

e- + F F -

F -Li+ + Li+ F -

9.2

Li + F Li+ F -

The Ionic Bond

1s22s11s22s22p5 1s21s22s22p6[He][Ne]

Li Li+ + e-

e- + F F -

F -Li+ + Li+ F -

Ionic StructuresIn an ionic compound (solid), the ions are packed

together into a repeating array called a crystal lattice.

The simplest arrangement is one in which the spheres in the base are packed side by side. Opposite charges are attracted to each other.

Its called simple cubic packing (NaCl is an example)Ionic formulas are always Empirical Formulas

(simplest)

Properties of Ionic Compounds•All ionic compounds form crystals.

• Ionic compounds tend to have high melting and boiling points. To break the positive and negative charges apart, it takes a huge amount of energy.

• Ionic compounds are very hard and very brittle. Again, this is because of the way that they're held together – strong attraction of oppositely charged ions. These ions simply don't move around - so they don't bend at all.  This also explains the brittleness of ionic compounds. If we give a big crystal a strong enough whack with a hammer, we usually end up using so much energy to break the crystal that the crystal doesn't break in just one spot, but in a whole bunch of places. Instead of a clean break, it shatters. 

• Ionic Compounds are poor conductors of electricity in the solid state.Ions are held tightly together and cannot move. Therefore ions cannot conduct electricity

• Ionic compounds conduct electricity when molten (melted).Ions are mobile and can therefore conduct electricity.

• Ionic compounds conduct electricity when aqueous (dissolved in water). If we take a salt and dissolve it in water, the water molecules pull the positive and negative ions apart from each other. Instead of the ions being right next to each other, they are able to move around in the water and conduct electricity. View here

Covalent Bonding

COVALENT BONDbond formed by the sharing of electrons

Covalent BondBetween nonmetallic elements of similar

electronegativity.Formed by sharing electron pairsStable non-ionizing particles, they are not

conductors at any stateExamples; O2, CO2, C2H6, H2O, SiC

Covalent bonds- Two atoms share one or more pairs of outer-shell electrons. Fluorine Atom Fluorine Atom

Fluorine Molecule (F2)

A covalent bond is a chemical bond in which two or more electrons are shared by two atoms.

Why should two atoms share electrons?

F F+

7e- 7e-

F F

8e- 8e-

F F

F F

Lewis structure of F2

lone pairslone pairs

lone pairslone pairs

single covalent bond

single covalent bond

9.4Dot diagram of F2

8e-

H HO+ + OH H O HHor

2e- 2e-

Lewis structure of water

Double bond – two atoms share two pairs of electrons

single covalent bonds

O C O or O C O

8e- 8e-8e-double bonds double bonds

Triple bond – two atoms share three pairs of electrons

N N8e-8e-

N N

triple bondtriple bond

or

9.4

Properties of Covalent Compounds• Covalent compounds tend to have low melting and boiling points.Many simple molecular substances are gases or liquids at room temperature.

• Most molecular compounds are poor conductors of electricity in all states.There are no free electrons available to move and conduct an electric current.

• Covalent compounds are usually much softer than ionic material.

• Covalent compounds tend to be flammable than ionic compounds.

• Most molecular compounds are insoluble in water, but will dissolve in nonpolar organic solvents.

9.4

Drawing Lewis Structures

1.Write the dot diagram for each atom present in the compound.

2. Take the TWO atoms with the MOST unpaired dots and join (bond) them together.

3.Add the remaining atoms where they are needed the most. (starting with remaining atoms with most unpaired dots)

N2H4HOF HCNH4COCH2OTry the following:

Bonds in all the polyatomic

ions and diatomics are all covalent

bonds

In covalent bonding, one or more pair of electrons are shared. However, all ‘sharing’ is NOT the same. Therefore, we can have Nonpolar (Pure) Covalent Bonds and Polar Covalent Bonds.

when electrons are shared equally

NONPOLAR COVALENT

BONDS

H2 or Cl2

Nonpolar Covalent bonds- Two atoms equally share one or more pairs of outer-shell electrons.

Fluorine Atom Fluorine Atom

Fluorine Molecule (F2)

when electrons are shared but shared unequally.

POLAR COVALENT

BONDS

H2O

Polar Covalent Bonds: Unevenly matched, but willing

to share.

- water is a polar molecule because oxygen is more electronegative than hydrogen, and therefore electrons are pulled closer to oxygen.

H F

FH

Polar covalent bond or polar bond is a covalent bond with greater electron density around one of the two atoms (electrons are shared unequally)electron rich

regionelectron poor

region e- riche- poor

d+ d-

To determine the type of Bonding present in a compound, we compare the electonegativities of the bonding atoms. The difference in electronegativity between these atoms predicts the type of bond present

The Electronegativities of Common Elements

Electronegativity is the ability of an atom to attract toward itself the electrons in a chemical bond.Electronegativity - relative, F is highest

H F

electron poor

region

electron richregion

Nonpolar (Pure)Covalent

share e- equally

Polar Covalent

partial transfer of e-

Unequal sharing

Ionic

transfer e-

Increasing difference in electronegativity

Classification of bonds by difference in electronegativity

Difference Bond Type

0 0.4 Nonpolar (Pure) Covalent> 1.7 Ionic

0.4 < and ≤1.7 Polar Covalent

Creates a “Bonding Continuum”

Classify the following bonds as ionic, polar covalent, or covalent: The bond in CsCl; the bond in H2S; andthe NN bond in H2NNH2.Cs – 0.7 Cl – 3.0 3.0 – 0.7 = 2.3 Ionic

H – 2.1 S – 2.5 2.5 – 2.1 = 0.4Nonpolar Covalent

N – 3.0 N – 3.0 3.0 – 3.0 = 0NonPolar Covalent

9.5

Predicting Molecular Geometry

1. Draw Lewis structure for molecule.2. Count number of lone pairs on the

central atom and number of atoms bonded to the central atom.

3. Use VSEPR to predict the geometry of the molecule.

Valence shell electron pair repulsion (VSEPR) model:

Predict the geometry of the molecule from the electrostatic repulsions between the electron (bonding and nonbonding) pairs.

AB2 2 0

Class

# of atomsbonded to

central atom

# lonepairs on central atom

Arrangement of electron pairs

MolecularGeometry

linear linear

B B

Cl ClBe

2 atoms bonded to central atom

0 lone pairs on central atom

10.1

AB2 2 0 linear linearClass

# of atomsbonded to

central atom

# lonepairs on central atom

Arrangement of electron pairs

MolecularGeometry

VSEPR

AB3 3 0 trigonal planar

trigonal planar

AB2 2 0 linear linearClass

# of atomsbonded to

central atom

# lonepairs on central atom

Arrangement of electron pairs

MolecularGeometry

VSEPR

AB3 3 0 trigonal planar

trigonal planar

AB4 4 0 tetrahedral

tetrahedral

Class

# of atomsbonded to

central atom

# lonepairs on central atom

Arrangement of electron pairs

MolecularGeometry

VSEPR

AB3 3 0 trigonal planar

trigonal planar

AB2E 2 1 trigonal planar bent

Class

# of atomsbonded to

central atom

# lonepairs on central atom

Arrangement of electron pairs

MolecularGeometry

VSEPR

AB3E 3 1

AB4 4 0 tetrahedral

tetrahedral

tetrahedral

trigonal pyramid

al

Class

# of atomsbonded to

central atom

# lonepairs on central atom

Arrangement of electron pairs

MolecularGeometry

VSEPR

AB4 4 0 tetrahedral

tetrahedral

AB3E 3 1 tetrahedral

trigonalpyramidal

AB2E2 2 2 tetrahedral

bent

HO

H

bonding-pair vs. bondingpair repulsion

lone-pair vs. lone pairrepulsion

lone-pair vs. bondingpair repulsion< <

Polarity and shapeThe shape of the molecule directly influences

the overall polarity of the molecule.If there is symmetry the charges cancel each

other out, making the molecule non-polarIf there is no symmetry, then its polar

Polar bonds do not guarantee a polar molecule

Ex: CCl4 and CO2 both have polar bonds, but both are NON-POLAR molecules. They have a dipole moment of zero (no net dipole or ‘resultant vector’.

The greater the dipole moment, the more polar the molecule

Why is molecular polarity important?Polar molecules have higher melting

and boiling points (for example the BP of HF is 19.5° C, and the BP of F2 is –188° C).

Polar solvents dissolve ionic and polar molecules more efficiently than non-polar solvents

Dipole Moments and Polar Molecules

10.2

H F

electron richregionelectron poor

region

d+ d-

‘polar molecule’

‘polar molecule’

Which of the following molecules have a dipole moment?H2O, CO2, SO2, and CH4

O HHdipole momentpolar molecule

SO O

CO O

no dipole momentnonpolar molecule

dipole momentpolar molecule

C

H

H

HH

no dipole momentnonpolar molecule

Does BF3 have a dipole moment?

Symmetrical Molecule – No Resultant Dipole NONPOLAR MOLECULE

The bent shape creates an overall positive end and negative endof the molecule = POLAR

The symetry of the molecule Cancels out the “charges” Making this NON-POLARNo overall DIPOLE

Examples to TryDetermine whether the following

molecules will be polar or non-polar molecules.SI2

CH3FAsI3

H2O2

Summary of Polarity of Molecules

Linear:When two atoms attached to central atom

are the same, the molecule will be Non-Polar (CO2)

When the two atoms are different the dipoles will not cancel, and the molecule will be Polar (HCN)

Bent:The dipoles created from this molecule will

not cancel creating a net dipole moment and the molecule will be Polar (H2O)

Summary of Polarity of Molecules

Pyramidal:The dipoles created from this molecule will

not cancel creating a net dipole and the molecule will be Polar (NH3)

Trigonal Planar:When the three atoms attached to central

atom are the same, the molecule will be Non-Polar (BF3)

When the three atoms are different the dipoles will not cancel, resulting in a net dipole, and the molecule will be Polar (CH2O)

TetrahedralWhen the four atoms

attached to the central atom are the same the molecule will be Non-Polar

When three atoms are the same, and one is different, the dipoles will not cancel, and the molecule will be Polar

METALLIC BONDbond found in metals; holds metal atoms together very strongly

Metallic BondFormed between atoms of metallic elementsElectron cloud around atoms Good conductors at all states, lustrous, very

high melting pointsExamples; Na, Fe, Al, Au, Co

Metallic Bonds: Mellow dogs with plenty of bones to go

around.

Ionic Bond, A Sea of Electrons

This ‘sea of mobile electrons’ allows metals to conduct an electricity current – electrons are free to travel.

‘Sea of Electrons’ is also the reason why most metals are malleable and ductile. As the electrons are free to move, it allows the metal to bend without breaking.