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Chemical Bonding I: Basic Concepts. Chapter 9 – (Topic 4 and 14). 4.1 – Ionic Bonding. Group. e - configuration. # of valence e -. ns 1. 1. 1A. 2A. ns 2. 2. 3A. ns 2 np 1. 3. 4A. ns 2 np 2. 4. 5A. ns 2 np 3. 5. 6A. ns 2 np 4. 6. 7A. ns 2 np 5. 7.

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Chemical Bonding I: Basic Concepts

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Chemical bonding i basic concepts

Chemical Bonding I:Basic Concepts

Chapter 9 –

(Topic 4 and 14)


4 1 ionic bonding

4.1 – Ionic Bonding


Chemical bonding i basic concepts

Group

e- configuration

# of valence e-

ns1

1

1A

2A

ns2

2

3A

ns2np1

3

4A

ns2np2

4

5A

ns2np3

5

6A

ns2np4

6

7A

ns2np5

7

Valence electrons are the outer shell electrons of an

atom. The valence electrons are the electrons that

particpate in chemical bonding.

9.1


Chemical bonding i basic concepts

Lewis Dot Symbols for the Representative Elements &

Noble Gases

9.1


Chemical bonding i basic concepts

  • Why do substances bond?

    • More stability

    • Atoms want to achieve a lower energy state


Ionic bonding

Ionic Bonding

  • Between a metal and a non-metal.

  • 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.


Chemical bonding i basic concepts

-

-

-

-

+

Li+

Li

Li

Li+ + e-

e- +

Li+

Li+ +

F

F

F

F

F

F

The Ionic Bond

[He]

[Ne]

1s22s1

1s22s22p5

1s2

1s22s22p6

9.2


Chemical bonding i basic concepts

E = k

Q+Q-

r

lattice energy

cmpd

MgF2

2957

Q= +2,-1

MgO

3938

Q= +2,-2

LiF

1036

LiCl

853

Electrostatic (Lattice) Energy

Lattice energy (E) is the energy required to completely separate one mole of a solid ionic compound into gaseous ions.

Q+ is the charge on the cation

Q- is the charge on the anion

r is the distance between the ions

Lattice energy (E) increases as Q increases and/or

as r decreases.

r F- < r Cl-

9.3


Chemical bonding i basic concepts

9.3


Ionic structures

Ionic Structures

  • In 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)


4 5 physical properties

4.5 Physical Properties


General physical properties

General physical properties

  • Depend on the forces between the particles

  • The stronger the bonding between the particles, the higher the M.P and BP

    • MP tends to depend on the existence of a regular lattice structure


Impurities and melting points

Impurities and Melting points

  • An impurity disrupts the regular lattice that its particle adopts in the solid state, so it weakens the bonding.

    • They always LOWER melting points

    • Its often used to check purity of a known molecular covalent compound because its MP will be off, proving its contamination


4 2 covalent bonding

4.2 – Covalent Bonding


Chemical bonding i basic concepts

Why should two atoms share electrons?

+

8e-

7e-

7e-

8e-

F

F

F

F

F

F

F

F

lonepairs

lonepairs

single covalent bond

single covalent bond

lonepairs

lonepairs

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

Lewis structure of F2

9.4


Chemical bonding i basic concepts

single covalent bonds

H

H

H

H

or

H

H

O

2e-

2e-

O

8e-

O

C

O

C

O

O

double bonds

8e-

8e-

8e-

double bonds

O

N

N

triple bond

N

N

triple bond

8e-

8e-

Lewis structure of water

+

+

Double bond – two atoms share two pairs of electrons

or

Triple bond – two atoms share three pairs of electrons

or

9.4


Chemical bonding i basic concepts

Lengths of Covalent Bonds

Bond Lengths

Triple bond < Double Bond < Single Bond

9.4


Chemical bonding i basic concepts

9.4


Chemical bonding i basic concepts

F

H

F

H

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

region

electron poor

region

e- poor

e- rich

d+

d-

9.5


Chemical bonding i basic concepts

F

H

Electronegativity is the ability of an atom to attract toward itself the electrons in a chemical bond.

Electronegativity - relative, F is highest

electron rich

region

electron poor

region

9.5


Chemical bonding i basic concepts

The Electronegativities of Common Elements

9.5


Chemical bonding i basic concepts

Increasing difference in electronegativity

Nonpolar Covalent

Polar Covalent

Ionic

partial transfer of e-

Unequal sharing

share e- equally

transfer e-

Classification of bonds by difference in electronegativity

Difference

Bond Type

0

Nonpolar Covalent

 2

Ionic

0 < and <2

Polar Covalent

9.5


Chemical bonding i basic concepts

Classify the following bonds as ionic, polar covalent,

or covalent: The bond in CsCl; the bond in H2S; and

the 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 = 0.4

Polar Covalent

N – 3.0

N – 3.0

3.0 – 3.0 = 0

NonPolar Covalent

9.5


Chemical bonding i basic concepts

DH0 = 436.4 kJ

H2 (g)

H (g)

+

H (g)

DH0 = 242.7 kJ

Cl2 (g)

Cl (g)

+

Cl (g)

DH0 = 431.9 kJ

HCl (g)

H (g)

+

Cl (g)

DH0 = 498.7 kJ

O2 (g)

O (g)

+

O (g)

O

DH0 = 941.4 kJ

N2 (g)

N (g)

+

N (g)

O

Bond Energies

Single bond < Double bond < Triple bond

N

N

The enthalpy change required to break a particular bond in one mole of gaseous molecules is the bond energy.

Bond Energy

9.10


Coordinate covalent or dative bond

Coordinate Covalent or Dative Bond

  • A covalent bond in which one of the atoms donates both electrons.

  • Properties do not differ from those of a normal covalent bond.


Coordinate covalent bonds dative

Coordinate covalent bonds (dative)

  • A covalent bond that occurs between two atoms in which both electrons shared in the bond come from the same atom.

  • Both electrons from the nitrogen are shared with the upper hydrogen

  • Ammonium has 3 polar covalent bonds and 1 coordinate (dative) covalent bond.


Examples

Hydronium (H3O+)

Carbon monoxide (CO)

Examples


Chemical bonding i basic concepts

Writing Lewis Structures

  • Draw skeletal structure of compound showing what atoms are bonded to each other. Put least electronegative element in the center.

  • Count total number of valence e-. Add 1 for each negative charge. Subtract 1 for each positive charge.

  • Complete an octet for all atoms except hydrogen

  • If structure contains too many electrons, form double and triple bonds on central atom as needed.

9.6


Chemical bonding i basic concepts

Write the Lewis structure of nitrogen trifluoride (NF3).

F

N

F

F

Step 1 – N is less electronegative than F, put N in center

Step 2 – Count valence electrons N - 5 (2s22p3) and F - 7 (2s22p5)

5 + (3 x 7) = 26 valence electrons

Step 3 – Draw single bonds between N and F atoms and complete

octets on N and F atoms.

Step 4 - Check, are # of e- in structure equal to number of valence e- ?

3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons

9.6


Chemical bonding i basic concepts

Write the Lewis structure of the carbonate ion (CO32-).

O

C

O

2 single bonds (2x2) = 4

1 double bond = 4

8 lone pairs (8x2) = 16

O

Total = 24

Step 1 – C is less electronegative than O, put C in center

  • Step 2 – Count valence electrons C - 4 (2s22p2) and O - 6 (2s22p4)

    • -2 charge – 2e-

4 + (3 x 6) + 2 = 24 valence electrons

Step 3 – Draw single bonds between C and O atoms and complete

octet on C and O atoms.

Step 4 - Check, are # of e- in structure equal to number of valence e- ?

3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons

Step 5 - Too many electrons, form double bond and re-check # of e-

9.6


Chemical bonding i basic concepts

What are the resonance structures of the

carbonate (CO32-) ion?

-

-

+

+

O

O

O

O

O

O

O

O

O

C

C

C

O

O

O

-

-

-

-

O

O

O

-

-

A resonance structure is one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure.

9.8


Chemical bonding i basic concepts

  • Resonance is possible whenever a Lewis structure has a multiple bond and an adjacent atom with at least one lone pair.

  • The following is the general form for resonance in a structure of this type.


Chemical bonding i basic concepts

Be – 2e-

2H – 2x1e-

H

Be

H

4e-

B – 3e-

3 single bonds (3x2) = 6

3F – 3x7e-

F

B

F

24e-

Total = 24

9 lone pairs (9x2) = 18

F

Exceptions to the Octet Rule

The Incomplete Octet

BeH2

BF3

9.9


Chemical bonding i basic concepts

N – 5e-

S – 6e-

N

O

6F – 42e-

O – 6e-

48e-

11e-

F

6 single bonds (6x2) = 12

F

F

Total = 48

S

18 lone pairs (18x2) = 36

F

F

F

Exceptions to the Octet Rule

Odd-Electron Molecules (radicals -very reactive)

NO

The Expanded Octet (central atom with principal quantum number n > 2)

SF6

9.9


Valence shell electron pair repulsion vsepr model

# of atoms

bonded tocentral atom

# lone

pairs on central atom

Arrangement ofelectron pairs

Molecular

Geometry

Class

linear

linear

B

B

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

10.1


Chemical bonding i basic concepts

Cl

0 lone pairs on central atom

Be

Cl

2 atoms bonded to central atom

10.1


Chemical bonding i basic concepts

# of atoms

bonded tocentral atom

# lone

pairs on central atom

trigonal planar

trigonal planar

Arrangement ofelectron pairs

Molecular

Geometry

Class

VSEPR

AB2

2

0

linear

linear

AB3

3

0

10.1


Chemical bonding i basic concepts

10.1


Chemical bonding i basic concepts

# of atoms

bonded tocentral atom

# lone

pairs on central atom

trigonal planar

trigonal planar

AB3

3

0

Arrangement ofelectron pairs

Molecular

Geometry

Class

tetrahedral

tetrahedral

VSEPR

AB2

2

0

linear

linear

AB4

4

0

10.1


Chemical bonding i basic concepts

10.1


14 1 molecules with more than 4 electron pairs

14.1 - Molecules with more than 4 electron pairs

  • Molecules with more than 8 valence electrons [expanded valence shell]

  • Form when an atom can ‘promote’ one of more electron from a doubly filled s- or p-orbital into an unfilled low energy d-orbital

  • Only in period 3 or higher because that is where unused d-orbitals begin


Why does this promotion occur

Why does this ‘promotion’ occur?

  • When atoms absorb energy (heat, electricity, etc…)their electrons become excited and move from a lower energy level orbital to a slightly higher one.

  • How many new bonding sites formed depends on how many valence electrons are excited.


Chemical bonding i basic concepts

# of atoms

bonded tocentral atom

# lone

pairs on central atom

trigonal planar

trigonal planar

AB3

3

0

Arrangement ofelectron pairs

Molecular

Geometry

Class

trigonal

bipyramidal

trigonal

bipyramidal

VSEPR

AB2

2

0

linear

linear

tetrahedral

tetrahedral

AB4

4

0

AB5

5

0

10.1


Chemical bonding i basic concepts

10.1


Chemical bonding i basic concepts

# of atoms

bonded tocentral atom

# lone

pairs on central atom

trigonal planar

trigonal planar

AB3

3

0

Arrangement ofelectron pairs

Molecular

Geometry

Class

trigonal

bipyramidal

trigonal

bipyramidal

AB5

5

0

octahedral

octahedral

VSEPR

AB2

2

0

linear

linear

tetrahedral

tetrahedral

AB4

4

0

AB6

6

0

10.1


Chemical bonding i basic concepts

10.1


Chemical bonding i basic concepts

10.1


Chemical bonding i basic concepts

lone-pair vs. lone pair

repulsion

lone-pair vs. bonding

pair repulsion

bonding-pair vs. bonding

pair repulsion

>

>


Chemical bonding i basic concepts

# of atoms

bonded tocentral atom

# lone

pairs on central atom

Arrangement ofelectron pairs

Molecular

Geometry

Class

trigonal planar

bent

VSEPR

trigonal planar

trigonal planar

AB3

3

0

AB2E

2

1

10.1


Chemical bonding i basic concepts

# of atoms

bonded tocentral atom

# lone

pairs on central atom

Arrangement ofelectron pairs

Molecular

Geometry

Class

trigonal pyramidal

tetrahedral

VSEPR

tetrahedral

tetrahedral

AB4

4

0

AB3E

3

1

10.1


Chemical bonding i basic concepts

# of atoms

bonded tocentral atom

# lone

pairs on central atom

trigonal

pyramidal

Arrangement ofelectron pairs

Molecular

Geometry

AB3E

3

1

tetrahedral

Class

bent

tetrahedral

O

H

H

VSEPR

tetrahedral

tetrahedral

AB4

4

0

AB2E2

2

2

10.1


Chemical bonding i basic concepts

# of atoms

bonded tocentral atom

# lone

pairs on central atom

trigonal

bipyramidal

distorted tetrahedron

Arrangement ofelectron pairs

Molecular

Geometry

Class

VSEPR

trigonal

bipyramidal

trigonal

bipyramidal

AB5

5

0

AB4E

4

1

10.1


Chemical bonding i basic concepts

# of atoms

bonded tocentral atom

# lone

pairs on central atom

trigonal

bipyramidal

distorted tetrahedron

Arrangement ofelectron pairs

Molecular

Geometry

AB4E

4

1

Class

trigonal

bipyramidal

T-shaped

F

F

Cl

F

VSEPR

trigonal

bipyramidal

trigonal

bipyramidal

AB5

5

0

AB3E2

3

2

10.1


Chemical bonding i basic concepts

# of atoms

bonded tocentral atom

# lone

pairs on central atom

trigonal

bipyramidal

distorted tetrahedron

Arrangement ofelectron pairs

Molecular

Geometry

AB4E

4

1

Class

trigonal

bipyramidal

T-shaped

AB3E2

3

2

trigonal

bipyramidal

linear

I

I

I

VSEPR

trigonal

bipyramidal

trigonal

bipyramidal

AB5

5

0

AB2E3

2

3

10.1


Chemical bonding i basic concepts

octahedral

octahedral

AB6

6

0

# of atoms

bonded tocentral atom

# lone

pairs on central atom

square pyramidal

octahedral

Arrangement ofelectron pairs

Molecular

Geometry

Class

F

F

F

Br

F

F

VSEPR

AB5E

5

1

10.1


Chemical bonding i basic concepts

octahedral

octahedral

AB6

6

0

# of atoms

bonded tocentral atom

# lone

pairs on central atom

square pyramidal

octahedral

AB5E

5

1

Arrangement ofelectron pairs

Molecular

Geometry

Class

square planar

octahedral

F

F

Xe

F

F

VSEPR

AB4E2

4

2

10.1


Chemical bonding i basic concepts

10.1


Predicting molecular geometry

What are the molecular geometries of SO2 and SF4?

F

S

F

F

O

O

S

F

Predicting Molecular Geometry

  • Draw Lewis structure for molecule.

  • Count number of lone pairs on the central atom and number of atoms bonded to the central atom.

  • Use VSEPR to predict the geometry of the molecule.

AB4E

AB2E

distorted

tetrahedron

bent

10.1


Polarity and shape

Polarity and shape

  • The 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-polar

  • If there is no symmetry, then its polar


Chemical bonding i basic concepts

  • 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

  • The greater the dipole moment, the more polar the molecule


Why is molecular polarity important

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


Chemical bonding i basic concepts

F

H

d-

d+

Dipole Moments and Polar Molecules

electron rich

region

electron poor

region

10.2


Chemical bonding i basic concepts

10.2


Chemical bonding i basic concepts

10.2


Chemical bonding i basic concepts

Which of the following molecules have a dipole moment?

H2O, CO2, SO2, and CH4

O

O

S

H

H

H

O

O

O

C

H

H

C

H

dipole moment

polar molecule

dipole moment

polar molecule

no dipole moment

nonpolar molecule

no dipole moment

nonpolar molecule

10.2


Chemical bonding i basic concepts

Does BF3 have a dipole moment?

10.2


Chemical bonding i basic concepts

The bent shape creates an

overall positive end and negative end

of the molecule = POLAR

The symetry of the molecule

Cancels out the “charges”

Making this NON-POLAR

No overall DIPOLE


Examples to try

Examples to Try

  • Determine whether the following molecules will be polar or non-polar

    • SI2

    • CH3F

    • AsI3

    • H2O2


Summary of polarity of molecules

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 molecules1

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)


Tetrahedral

Tetrahedral

When 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


Allotropes of carbon

Allotropes of Carbon


Allotropes

Allotropes

  • Carbon can bond with itself in at least three different ways giving us 3 different materials

    • Diamond

    • Graphite

    • Buckyballs and nanotubes


Diamond

Diamond

  • Carbons are bonded via sp3 hybridization to 4 other carbon atoms forming a giant network covalent compound.


Properties of diamond

Properties of Diamond

  • High melting point due to strong directional covalent bonds (3550 C)

  • Extremely hard because it is difficult to break atoms apart or move them in relation to one another

  • No electrical conductivity because electrons are localized in specific bonds

  • Insoluble in polar and non-polar solvents because molecular bonds are stronger than any intermolecular forces


Graphite

Graphite

  • Carbon atoms are bonded via sp2 hybridization.

  • Carbon atoms form sheets of six sided rings with p-orbitals perpendicular from plane of ring.


Graphite structure

Graphite Structure

  • Carbon has 4 valence e- to bond with. 3 are used for closest atoms in rings. 1 is delocalized in p-orbitals

  • The presence of p-orbitals allows for strong van der waals forces that hold the sheets together


Properties of graphite

Properties of Graphite

  • Different from Diamond

    • Conducts electricity because of delocalized electrons

    • Slippery can be used as lubricant, sheets can easily slip past each other (think of a deck of cards)

  • Same as Diamond

    • High melting point (higher actually because of delocalized e-, 3653C)

    • Insoluble (same reason)


Fullerenes

Fullerenes

  • Buckyballs: spherical

  • Nanotubes: tube shaped

  • Both have very interesting properties

    • Super strong

    • Conduct electricity and heat with low resistance

    • Free radical scavenger


Buckyballs

Buckyballs

  • Carbon atoms bond in units of 60 atoms (C-60) forming a structure similar to a soccerball with interlocking six sided and five sided rings.

  • sp2 hybridization

  • Extra p-orbitals form pi bonds resulting in

    • Electrical conductivity

    • Stronger covalent bonds, therefore stronger materials


Chemical bonding i basic concepts

Generally, intermolecular forces are much weaker than intramolecular forces.

4.3 - Intermolecular Forces

Intermolecular forces are attractive forces between molecules.

Intramolecular forces hold atoms together in a molecule.

  • Intermolecular vs Intramolecular

  • 41 kJ to vaporize 1 mole of water (inter)

  • 930 kJ to break all O-H bonds in 1 mole of water (intra)

“Measure” of intermolecular force

boiling point

melting point

DHvap

DHfus

DHsub

11.2


Chemical bonding i basic concepts

Orientation of Polar Molecules in a Solid

Intermolecular Forces

Dipole-Dipole Forces

Attractive forces between polar molecules

11.2


Chemical bonding i basic concepts

Ion-Dipole Interaction

Intermolecular Forces

Ion-Dipole Forces

Attractive forces between an ion and a polar molecule

11.2


Chemical bonding i basic concepts

11.2


Chemical bonding i basic concepts

Intermolecular Forces

Dispersion Forces

Attractive forces that arise as a result of temporary dipoles induced in atoms or molecules

ion-induced dipole interaction

dipole-induced dipole interaction

11.2


Chemical bonding i basic concepts

Induced Dipoles Interacting With Each Other

11.2


Chemical bonding i basic concepts

Dispersion forces usually increase with molar mass.

Intermolecular Forces

Dispersion Forces Continued

Polarizability is the ease with which the electron distribution in the atom or molecule can be distorted.

  • Polarizability increases with:

  • greater number of electrons

  • more diffuse electron cloud

11.2


Chemical bonding i basic concepts

O

O

S

What type(s) of intermolecular forces exist between each of the following molecules?

HBr

HBr is a polar molecule: dipole-dipole forces. There are also dispersion forces between HBr molecules.

CH4

CH4 is nonpolar: dispersion forces.

SO2

SO2 is a polar molecule: dipole-dipole forces. There are also dispersion forces between SO2 molecules.

11.2


Chemical bonding i basic concepts

or

H

H

B

A

A

A

Intermolecular Forces

Hydrogen Bond

The hydrogen bond is a special dipole-dipole interaction between the hydrogen atom in a polar N-H, O-H, or F-H bond and an electronegative O, N, or F atom.

A & B are N, O, or F

11.2


Chemical bonding i basic concepts

Decreasing molar mass

Decreasing boiling point

Why is the hydrogen bond considered a “special” dipole-dipole interaction?

11.2


4 4 metallic bonding

4.4 Metallic bonding


Chemical bonding i basic concepts

nucleus &

inner shell e-

mobile “sea”

of e-

Types of Crystals

  • Metallic Bonds- electron sea model - metal cations in a sea of valence electrons

  • Lattice points in crystal are occupied by metal atoms

  • the valence electrons do not “belong” to a single cation, but are delocalized and may move about

  • Good conductors of heat and electricity

Cross Section of a Metallic Crystal

11.6


Metallic bond

Metallic bond

  • Occurs between atoms with low electronegativities

  • Metal atoms pack close together in 3-D, like oranges in a box.

  • Close-packed lattice formation


Chemical bonding i basic concepts

  • Many metals have an unfilled outer orbital

  • In an effort to be energy stable, their outer electrons become delocalised amongst all atoms

  • No electron belongs to one atom

  • They move around throughout the piece of metal.

  • Metallic bonds are not ions, but nuclei with moving electrons


Physical properties

Physical Properties

Conductivity

  • Delocalised electrons are free to move so when a potential difference is applied they can carry the current along

  • Mobile electrons also mean they can transfer heat well

  • Their interaction with light makes them shiny (lustre)


Malleability

Malleability

  • The electrons are attracted the nuclei and are moving around constantly.

  • The layers of the metal atoms can easily slide past each other without the need to break the bonds in the metal

  • Gold is extremely malleable that 1 gram can be hammered into a sheet that is only 230 atoms thick (70 nm)


Melting points

Melting points

  • Related to the energy required to deform (MP) or break (BP) the metallic bond

  • BP requires the cations and its electrons to break away from the others so BP are very high.

  • The greater the amount of valence electrons, the stronger the metallic bond.

  • Gallium can melt in your hand at 29.8 oC, but it boils at 2400 oC!


4 5 physical properties1

4.5 Physical Properties


General physical properties1

General physical properties

  • Depend on the forces between the particles

  • The stronger the bonding between the particles, the higher the M.P and BP

    • MP tends to depend on the existence of a regular lattice structure


Impurities and melting points1

Impurities and Melting points

  • An impurity disrupts the regular lattice that its particle adopts in the solid state, so it weakens the bonding.

    • They always LOWER melting points

    • Its often used to check purity of a known molecular covalent compound because its MP will be off, proving its contamination


How would this ideal heat curve look different if the substance was contaminated

How would this ideal heat curve look different if the substance was contaminated?


Volatility

Volatility

  • A qualitative measure of how readily a liquid or solid is vaporised upon heating or evaporation

    • It is a measure of the tendency of molecules and atoms to escape from a liquid or a solid.

    • Relationship between vapour pressure and temperature (B.P)

  • Mostly dealing with liquids to gas, however can occur from solid directly to gas (dry ice).

  • The weaker the intermolecular bonds, the more volatile


Conductivity

Conductivity

  • Generally molecules have poor solubility in polar solvents like water, but if they do dissolve they do not for ions

  • There are no charged particles to carry the electrical charge across the solution.

  • Example: sugar dissolves in water

  • C12H22O11(s) C12H22O11(aq)


Dissolving sugar covalent compound

Dissolving sugar (covalent compound)

  • It takes energy to break the bonds between the C12H22O11 molecules in sucrose crystal structure.

  • It also takes energy to break the hydrogen bonds in water so that one of these sucrose molecules can fit into solution.

  • In order for sugar to dissolve, there must be a greater release of energy when the dissolution occurs than when the breaking of bonds occur.


Ionic compounds

Ionic compounds

  • The energy needed to break the ionic bond must be less than the energy that is released when ions interact with water.

  • The intermolecular ion-dipole force is stronger than the electrostatic ionic bond

  • Breaks up the compound into its ions in solution.


Chemical bonding i basic concepts

  • Soluble salt in water breaks up as

    NaCl (s) Na+ (aq) + Cl-(aq)

  • http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/molvie1.swf


Ionic compounds1

Ionic compounds

  • Held together by strong 3-d electrostatic forces.

  • They are solid at room temperature and pressure

  • If one layer moves a fraction, the ions charges are off and now repulsion occurs. This is the reason they are strong, yet brittle.


Chemical bonding i basic concepts

  • Molten or dissolved ionic compounds conduct electricity

  • Insoluble in most solvents, yet H2O is polar and attracts both the + and – ions from salts


Covalent bonding properties

Giant covalent

Ex: diamond, silicon dioxide

Very hard

Very high MP (>1000oC)

Does not conduct

Insoluble in all solvents

Molecular covalent

Ex: CO2, alcohols, I2

Usually soft, malleable

Low MP (<200oC)

Does not conduct

More soluble in non-aqueous solvents, unless they can h-bond

Covalent bonding properties


Solubility of methanol in water

Solubility of methanol in water

  • http://www.mhhe.com/physsci/chemistry/animations/chang_7e_esp/clm2s3_4.swf

  • Alcohols generally become less soluble, the longer the carbon chain due to the decreasing tendency for hydrogen bonding to occur intermolecularly.


States of matter

States of matter

  • Physical state depends on intermolecular forces

  • The weaker the attraction, the more likely it’s a gas, while stronger attractions indicate solid.


14 2 hybridization

14.2 - Hybridization

  • the concept of mixing atomic orbitals to form new hybrid orbitals

  • Used to help explain some atomic bonding properties and the shape of molecular orbitals for molecules.

  • The valence orbitals (outermost s and p orbitals) are hybridised (mathematically mixed) before bonding, converting some of the dissimilar s and p orbitals into identical hybrid spn orbitals

  • We must know sp, sp2, and sp3 hydrid orbitals


Hybridization mixing of two or more atomic orbitals to form a new set of hybrid orbitals

Hybridization – mixing of two or more atomic orbitals to form a new set of hybrid orbitals.

  • Mix at least 2 nonequivalent atomic orbitals (e.g. s and p). Hybrid orbitals have very different shape from original atomic orbitals.

  • Number of hybrid orbitals is equal to number of pure atomic orbitals used in the hybridization process.

  • Covalent bonds are formed by:

    • Overlap of hybrid orbitals with atomic orbitals

    • Overlap of hybrid orbitals with other hybrid orbitals

10.4


Hybrid orbitals

Hybrid orbitals

  • Carbon has 4 valence electrons.

  • 2 electrons paired up in the s-orbital, and 2 electrons unpaired in the p-orbital.

  • So why does it commonly make 4 bonding sites?


Chemical bonding i basic concepts

  • One of carbon’s paired s-orbital electrons is ‘promoted’ to the empty p-orbital

  • This produces a carbon in an excited state which has 4 unpaired electrons (4 equivalent bonding sites)


Chemical bonding i basic concepts

Formation of sp3 Hybrid Orbitals

10.4


Chemical bonding i basic concepts

10.4


Chemical bonding i basic concepts

Formation of sp Hybrid Orbitals

10.4


Chemical bonding i basic concepts

Formation of sp2 Hybrid Orbitals

10.4


Chemical bonding i basic concepts

How do I predict the hybridization of the central atom?

  • Draw the Lewis structure of the molecule.

  • Count the number of lone pairs AND the number of atoms bonded to the central atom

# of Lone Pairs

+

# of Bonded Atoms

Hybridization

Examples

2

sp

BeCl2

3

sp2

BF3

4

sp3

CH4, NH3, H2O

5

sp3d

PCl5

6

sp3d2

SF6

10.4


Chemical bonding i basic concepts

10.4


Chemical bonding i basic concepts

10.5


Chemical bonding i basic concepts

10.5


Chemical bonding i basic concepts

Pi bond (p) – electron density above and below plane of nuclei of the bonding atoms

Sigma bond (s) – electron density between the 2 atoms

10.5


Chemical bonding i basic concepts

10.5


Chemical bonding i basic concepts

10.5


Chemical bonding i basic concepts

10.5


Chemical bonding i basic concepts

Describe the bonding in CH2O.

H

O

C

H

C – 3 bonded atoms, 0 lone pairs

C – sp2

10.5


Sigma s and pi bonds p

How many s and p bonds are in the acetic acid

(vinegar) molecule CH3COOH?

H

H

C

H

C

O

O

H

Sigma (s) and Pi Bonds (p)

1 sigma bond

Single bond

1 sigma bond and 1 pi bond

Double bond

Triple bond

1 sigma bond and 2 pi bonds

s bonds = 6

+ 1 = 7

p bonds = 1

10.5


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