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BONDING. TOPIC 4. Terms. Covalent Bonding. Bonds Breaking them takes energy Making them gives off energy. Exothermic More energy is given off than put in Endothermic More energy is absorbed than given off Intra molecular Forces Forces within molecules (ionic, covalent and metallic)

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Bonding

BONDING

TOPIC 4


Terms

Covalent Bonding

  • Bonds

    • Breaking them takes energy

    • Making them gives off energy


  • Exothermic

    • More energy is given off than put in

  • Endothermic

    • More energy is absorbed than given off

  • Intramolecular Forces

    • Forces within molecules (ionic, covalent and metallic)

  • Intermolecular Forces (IMF)

    • Forces between particles


Ionic Bonding

+

-

Less e- = Less e- repulsion

More e- = e- more repulsion.

Metal: K

Non-Metal: Cl

  • If the electronegative difference between the atoms involved is =>1.8

    • There are always exceptions to this rule!

  • Will conduct electricity in its molten or aqueous state (This test proves ionic)


Intramolecular Forces

Drawing Ionic Bonding

Lewis Dot Diagram

-

+

X

Electrons are in pairs

Na

Cl

Special Note: The ionic bond is the electrostaticattractionbetween oppositely charged ions!

Ionic Bonding

  • Just use the valence shell

  • Be sure to include square brackets and charge after electron exchange.


Lewis Diagrams

Combine

C

Al

Fe

Mg

Be

Cl

F

Cl

O

Br

Lewis Dot diagrams us the atoms valance shell electrons


Decomposition

Intramolecular Forces

2Na+(aq) + 2Cl-(aq)  2Na(s) + Cl2(g)

+

+

+

+

+

+

-

-

CATHODE (-)

ANODE (+)

+

-

+

+

+

-

-

+

+

Conductivity is FINITE

NaCl

  • When in molten or aqueous state, ionic substances WILL conduct electricity, by the movement of (+) and (-) ions.

  • This is different from how METALS conduct electricity!


Ionic Compounds

Giant Ionic Lattices

Force

Cation

Anion

+

-

+

-

+

-

+

-

+

-

+

-

+

-

+

-

+

-

+

-

+

-

+

-

+

Like charges repel

Metal: K

Non-Metal: Cl

  • No bonds are made!!!

  • Static attractions holds them together. (opposites attract)

  • When a force is applied, ionic compounds will make a clean break.

  • Physical characteristics

    • Hard and brittle

    • Solid doesn’t conduct Electricity

    • More soluble in water than other solvents

    • High MP and BP


Cubic or Isometric

Giant Ionic Lattices

Table Salt

NaCl


Tetragonal

Giant Ionic Lattices

Cassiterite

SnO2


Orthorhombic

Giant Ionic Lattices

Agagonite

CaCO3

  • Also found in mollusk shells and coral


Hexagonal

Giant Ionic Lattices

Beryl

Be3Al2(SiO 3)6


Trigonal

Giant Ionic Lattices

Quartz

SiO2


Ionic Bonding

Giant Ionic Lattices

Beryl

Be3Al2(SiO 3)6


Triclinic

Giant Ionic Lattices

Copper(II) Sulfate

CuSO4


Intramolecular Forces

Transition Metals

Fe2+

Cu+

Fe3+

Cu2+

Iron(II) Oxide

Copper(I) Oxide

Iron (III) Oxide

Copper(II) Oxide

Multiple Ions

  • Transition metals can have multiple ions.

  • Ones you should know.


Ions

Reminder

SO4-2

PO4-3

NO3-

CO3-2

OH-

HCO3-

NH4+

Polyatomic Ions

  • Be sure to review your polyatomic ions!!!



Intramolecular Forces

Covalent Bonding

2.1

3.0

X

Differences

|3-2.1|

=0.9

H

Cl

Special Note: The covalent bond is the electrostaticattractionbetween pairs of e- and positively charged nuclei!

COVALENT BONDING

  • If the electronegative difference between the atoms involved is <1.8

  • Will NOT conduct electricity

  • Electrons are shared


Questions

Review

Na

Ca

+

CO3

Cl

+

Li

+

O

+

Na

SO3

NO3

K

+

For ionic compounds to form the valance shells of both metal and non-metal must be full!!

What is the chemical formula?

What is the names for each?


Intramolecular Forces

Covalent Bonding

H

H

X

X

C

Cl

H

H

H

X

H

X

H

H

COVALENT BONDING

  • Structural formula

  • Lewis structure


Lewis Structures

Intramolecular Forces

H2O

1

1

6

H

H

Hydrogen can only hold 2e- remaining must be paired on Oxygen

O

-

4 = 4

8

COVALENT BONDING

  • 1) Sum all valence e-

  • 2) Subtract 2e- for every bond

  • 3) Place e- around periphery atoms to form octets. The remaining around central atom

  • 4) All atoms MUST be paired!!!!!!


Intramolecular Forces

Lewis Structures

Covalent Bonding

HL: PCl5, PCl4+, PCl6-

and XeF4

  • Draw the following Lewis structures

  • H2 Cl2

  • O2 N2

  • HCN C2H6

  • C2H4 C2H2


Special Lewis Structures

Intramolecular Forces

+

+

Lone pair of e-

H

N

H

H

Electrophile

H

Covalent Bonding

  • Coordinate or dative covalent bonds

  • When both e- are shared from the same atom. (Not one from each as before)

  • Occurs when a non bonding e- pair donates an e- to an e- deficient atom.


Intramolecular Forces

Special Lewis Structures

Covalent Bonding

  • Draw the following Lewis structures

  • CO

  • H3O+


Length, Strength & Hybrid Resonance

Intramolecular Forces

O

2-

2-

2-

O

O

C

O

O

C

Don’t forget to show the e- pairs!!

C

O

O

O

O

CO32-

  • More bonds = more strength & shorter bonds

  • Resonance structures

    • Bond length is longer than a double bond but shorter than a single bond


Length & Strength

Intramolecular Forces

Ethene

Carboxylic Acid

Ethyne

O

H

H

C

R = Functional Group

C

C

H

C

C

H

R

OH

H

H

CO32-

  • Compare the two molecules

  • Ethyne has stronger and shorter bonds

  • C=O bond is stronger and shorter due to Oxygen being more electronegative


Bond Polarity

Intramolecular Forces

δ+

δ-

Dipole Moment

H

X

Cl

Covalent Bonding

  • Non-Metals are fighting for e-

  • Atom with larger electronegativity will hold the e- closer to itself.

  • Atoms become slightly charged.


Exceptions to the Octet Rule

Intramolecular Forces

F

B

F

F

Covalent Bonding

  • BF3

  • Actual structure: Boron is e- deficient

    • This is known because of its reactivity towards electron rich molecules such as NH3

  • CNOF all obey the octet rule.


Intramolecular Forces

Formal Charge

Covalent Bonding

  • SO42-

    • Single bonds (8 e- around S)

    • Double bonds (12 e- around S)

  • Formal Charge = (# valence e- on free atom) – (# valence e- assigned to the atom in the molecule)

  • (Valence e-)assigned = (# lone pair e-) + ½ (# of shared e-)

  • 1) Molecules attempt to achieve Formal Charge as close to 0 as possible.

  • 2) Any negative Formal charge will reside on most electronegative atom.


VSEPR (shape)

Intramolecular Forces

3 Pairs of e-

120o

2-

O

F

C

B

2 Pairs of e-

180o

C

O

O

O

O

F

F

Covalent Bonding

  • VSEPR (Valence Shell Electron Pair Repulsion)

  • Paired e- attempt to get as far away from each other as possible.

  • Multiple bonds still count only as 1 pair!!


VSEPR

Intramolecular Forces

4 Pairs of e-

109.5o

H

Lone pair

107o

Lone pair

104.5o

O

N

C

H

H

H

H

H

H

H

H

Covalent Bonding

  • Tetrahedral

  • Lone pair e- have increased charge density and require more room

  • More repulsion from lone pair will decrease bond angle.


Intramolecular Forces

Home Work

Covalent Bonding

  • Predict the shape AND bond angles

  • H2S PbCl4 H2CO SO2

  • NO3- PH3 NO2-

  • NH2- POCl3 CO2




Expanded valance shell 14 1
Expanded Valance Shell (14.1)

  • Molecules with more than 8 electrons

  • Electron promotion:


Dipole Moment

Molecule Polarity (4.2.6)

2δ-

Cl

Cl

Cl

δ-

H

δ-

Non Polar

O

δ+

C

C

H

H

H

H

δ+

δ+

δ+

H

H

H

Covalent Bonding

  • Polarity effects state change (physical change)

  • Unequal sharing causes a dipole moment to form

  • Q: Why is BF3 non-polar whereas PF3 is polar?


Hybridization 14 2 2
Hybridization (14.2.2)

  • Sigma bond: σ (single bond)

    • Axial overlap of orbital’s

1s1

2px2

py2

pz2

H

Cl


Hybridization 14 2
Hybridization (14.2)

  • Sigma bond: σ (single bond)

    • Axial overlap of orbital’s

Cl

Cl


Hybridization 14 21
Hybridization (14.2)

  • Pi bond: π(Double bond, one σ bond)

    • Parallel overlap of orbital’s

O

O

N

N


Hybridization 14 2 3
Hybridization (14.2.3)

  • Hybridization electron promotion

    • New Orbital sp3

2s2

2px2

py2

pz2

Excited State

Ground State

C

4 Equal orbital`s capable of holding a maximum of 2 electrons each


Hybridization 14 22
Hybridization (14.2)

  • How to determine Hybridized orbital`s

    • Look at the shape


Carbon

Allotropes

C

C

C

C

C

Giant Covalent

  • 1) Diamond (Tetrahedron, localized e-)

    • Very hard and does not conduct electricity

  • 2) Fullerenes (C60) Hexagonal and pentagonal rings

    • Nanotubes


Allotropes

Carbon

C

C

C

C

Weak Pi Bonds

C

C

HL: sp hybrid

Delocalized electrons able to move

Giant Covalent

  • 3) Graphite (Planar, delocalized e-)

    • Weak pi bonding between sheets cause it to conduct electricity and be slippery.

    • Bonds are shorter than a tetrahedral due to the pi bonding


Benzene (14.3)

Pi bonds overlap allowing for electrons to be delocalized over the entire molecule.

C

C

C

C

C

C

C6H6

  • Planar, delocalized e-

    • Regular bonding would predict an alternating double bond (Resonance structure)

  • Hybrid theory shows sp2 configuration


Intramolecular Forces

Silicon

Si

Si

Si

Si

Si

Si

Si

Si

Si

Si

Silicon

Tetrahedron Configuration

Similar to diamond


Intramolecular Forces

Silicon & Silicon dioxide

SiO2 but based on a network of SiO4

O

Si

O

O

O

Quartz

  • Single bonds formed between Oxygen to satisfy the octet.

  • HL: Less overlap in the P-sub orbital due to atomic size difference therefore Pi bonds do not form.



Metallic Bonding

Intramolecular Forces

+

+

+

+

-

+

+

+

+

-

-

-

Sea of electrons

+

+

+

+

+

-

-

+

+

+

+

-

-

Conductivity is INFINITE

Metallic Bonding

  • In solid state

  • Outer e- are delocalized and free to move about

  • Bond is a result of electrostatic attraction between Fixed positive metal ions and delocalized e-


Physical Properties

Malleability

+

+

+

+

-

+

+

+

+

-

-

-

+

+

+

+

+

-

-

-

+

+

+

+

-

Metallic Bonding

  • The ability for a material to be pounded into thin sheets.

  • Aluminum Foil

  • Swords and Folding


Physical Properties

Ductility

+

+

+

+

+

+

+

+

+

+

+

+

Electrons have been excluded

Metallic Bonding

  • The ability for a material to be pulled into wire

  • Or in this case extruded into a wire


Physical Properties

Metallic Bonding

  • Because e- can move easily it can conduct energy. (Heat or electricity)

  • MP related to attractive force (between atoms)

    • 1) Size of Cation(+)

    • 2) # of valence e-

    • 3) Atom packing

  • Size increases MP decreases:

  • Giant Covalent substances have very high mp


Allotrops

Metallic Bonding

  • Same element but different structure

  • Carbon

    • Diamond

    • Graphite

    • Fulluron


Inter molecular forces

INTERMOLECULAR FORCES

Topic 4


Intermolecular Forces (4.3.1)

van der Waals’ Forces

Intermolecular Forces

Charge Induction

Charge Induction

d+

d-

d+

d-

d+

d-

IMF

  • Van der Waals Forces


Dipole-Dipole (4.3.1)

Intermolecular Forces

Cl

Cl

Cl

Cl

d-

d-

C

C

H

H

H

H

d+

d+

IMF

  • Polar molecules (polar covalent) have slightly charged ends

    • Opposites attract.

    • Large electronegative difference = stronger attraction.


Hydrogen Bonding (4.3.1)

van der Waals’ Forces

Intermolecular Forces

d-

d-

O

O

H

H

d+

d+

d+

d+

H

H

IMF

  • Hydrogen Bonding (F, O or N bonded to H)

    • Due to small size and high electronegativity of non metals

    • Creates a large charge difference

    • Basically a super strong dipole-dipole bond


Intermolecular Forces

Boiling Point Trends (4.3.2)

Get a picture of group 4,5,6,7 boiling points for hydrides

Key question is why does water have an abnormally high BP?

H bonding with O, F and N

IMF

  • Phase change when IMF are overcome

  • Be sure to explain using the words IMF and how they affect the bonds BETWEEN particles.

  • Van der Waals’ Forces are ALWAYS present!!!


Physical Properties

Increasing Melting Point

  • Van der Waal’s: Lowest MP, Non polar

    • Butane (C4H10)

  • Dipole-dipole: Slightly miscible

    • Propanone C3H6O

  • Hydrogen Bonding: Miscible with polar substances

    • H2O

  • Ionic Bonding: Only conducts electricity when liquid or aqueous. (Decomposition when it does)

    • NaCl

  • Metallic Bonding: Conducts electricity, not water soluble, MP regulated by, valance, size and packing.

    • Fe

  • Giant Covalent: Highest MP, Insoluble in both non-polar and polar solvents. Does not conduct electricity except for graphite.

    • Diamond and Graphite (Allotropes)


Bonding questions
Bonding Questions

  • Compare the following for B.P

  • HF and HCl

  • H2O and H2S

  • NH3 and PH3

  • CH3OCH3 and CH3CH2OH

  • CH3CH2CH3, CH3CHO and CH3CH2OH



Hybridization 14 23
Hybridization (14.2)

  • Sigma bond: σ (single bond)

    • Axial overlap of orbital’s


Hybridization 14 24
Hybridization (14.2)

  • Sigma bond: σ (single bond)

    • Axial overlap of orbital’s


Hybridization 14 25
Hybridization (14.2)

  • Sigma bond: σ (single bond)

    • Axial overlap of orbital’s


Hybridization 14 26
Hybridization (14.2)

  • Sigma bond: σ (single bond)

    • Axial overlap of orbital’s


Lattice formation
Lattice Formation

  • Where the heat comes from

  • Route 1: A + B + C + E

  • Route 2: F

  • Hess’s law: A + B +C + E = F

    +107 + 122 + 496 + (-349) + E = -411

    E = -787 kJ mol-1


Lattice Enthalpy

Intramolecular Forces

Na(s) + ½ Cl2(g) Na+Cl- or NaCl

1) Na(s) Na(g) ½ Cl2(g)  Cl(g)

2) Na(g) Na+(g) + e- Cl(g) + e-  Cl-(g)

3) Na+(g) + Cl-(g)  NaCl(s)

NaCl

  • 1) Production of Gaseous atoms

  • 2) Formation of Gaseous ions

  • 3) Production of solid ionic lattice


Born haber cycle
Born-Haber Cycle

E

Na+(g) + Cl-(g)

NaCl(s)

Lattice Enthalpy

ΔHθI.E.

1stIonization of Na

+496 kJ mol-1

ΔHθE.A.

1st electron affinityof Cl

-349 kJ mol-1

C

D

Endothermic

Exothermic

Na(g)

Cl(g)

F

ΔHθf

Formation of NaCl

-411 kJ mol-1

ΔHθat

Atomization of Cl

+122kJ mol-1

ΔHθat

Atomization of Na

-107 kJ mol-1

A

B

Na(s) + Cl2(g)


Spare Parts

+

H

N

Cl

-

O

C

C

C

C

C

O

O

H

H

C

O

H

C

C

N

Cl

H

H

H

δ+

C

C

H

H

H

H

δ-

X

H

H

H


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