Chemistry 121: Topic 4 - Chemical Bonding Topic 4 ... · PDF fileIB IIB 13 Al 26.98 14 Si 28.09 15 P ... Topic 4 - Chemical Bonding Topic 5: ... Chemistry 121: Topic 4 - Chemical Bonding

Chemistry 121: Topic 4 - Chemical Bonding

Topic 4: Chemical Bonding 4.0 Ionic and covalent bonds; Properties of covalent and ionic compounds

4.1 Lewis structures, the octet rule.

4.2 Molecular geometry: the VSEPR approach. Molecular polarity.

4.3 Valence bond theory, hybridization and geometry, multiple bonds.

4.4 Molecular orbital theory.


Types of Bonding:

• Covalent o Sharing of electrons o Discrete molecular species o Generally nonconductive, Low boiling point o Usually between a two or more non-metals

• Ionic o Discrete Anions and Cations held together by Electrostatic Attraction o Empirical Formula only / No discrete units o High Boiling points o Usually between one or more metal and one or more non-metal

• Metallic o Delocalized bonding / electron sharing o Low to high boiling points o Malleable


Key Concept: Coulomb’s Law: Describes the energy of interaction between charged species.

Where: V is in Joules

r is the distance between the charge centers in nm

Q1 and Q2, are the numerical unit charges

ε0 is a constant called the Vacuum Permittivity

For Example: for Sodium Chloride

Note: negative signs indicate an attraction, ion pair has less energy.

V = Q1 Q2

4πε0r = 2.31 x 10-19 J nm ( ) Q1 Q2

r

= - 8.37 x 10-19 JV = 2.31 x 10-19 J nm ( )(+1) (-1)

0.276 nm


Covalent Bonds are the result of sharing of electrons, usually result in a “completion” of unfilled electronic shells or in other words becoming isoelectronic with a noble gas.


IA THE PERIODIC TABLE OF THE ELEMENTS

VIII

1 H

1.008

IIA

VIIIA

Group Number

IIIA

IVA

VA

VIA

VIIA

2

He 4.003

3

Li 6.941

4

Be 9.012

26

Fe 55.85

Atomic Number Symbol* Atomic Weight

*Synthetic elements are hollow faced. The most stable isotope is shown.

5

B 10.81

6

C 12.01

7

N 14.01

8

O 16.00

9

F 19.00

10

Ne 20.18

11 Na 22.99

12 Mg 24.31

IIIB

IVB

VB

VIB

VIIB

|←

VIIIB →|

IB

IIB

13

Al 26.98

14

Si 28.09

15

P 30.97

16

S 32.06

17

Cl 35.45

18

Ar 39.95

19

K 39.10

20

Ca 40.08

21 Sc 44.96

22

Ti 47.88

23

V 50.94

24

Cr 52.00

25

Mn 54.94

26

Fe 55.85

27

Co 58.93

28

Ni 58.69

29

Cu 63.55

30

Zn 65.38

31 Ga 69.72

32

Ge 72.59

33

As 74.92

34

Se 78.96

35

Br 79.90

36

Kr 83.80

Lewis Dot symbols and Valance electrons: Lewis Dot Symbols are a way of

showing valence e- symbolically. Valence e- are those electrons outside of the

filled [Noble gas core] electrons.

Except for Helium, the number of valence electrons correspond to the group

number.

•H He •• F•• •• •

• •Cl•• •• •

• •O••• •

• •S••• •

• •


1s2 2s1 1s2 2s0

1s2 2s2 2p5 1s2 2s2 2p6

Ionic Compounds and the Ionic Bond: An ionic bond is the electrostatic force that holds ions together in an ionic

compound. The most common elements to form ionic compounds are Group 1A

and 2A metal atoms (cations) and group 6A and 7A (anions).

Example: Consider the following Reaction

•Li F •• •• •

• •+ Li+ + •F•• •

• •

• •

-


Covalent Bonding:

A Covalent bond occurs when two electrons are shared by two atoms. A Covalent Compound is one that contains only covalent bonds. In this case in the molecule F2 there is one covalent bond and each F has 3 pairs of lone pairs. A Lewis Structure is a representation of covalent bonding in which shared electron pairs are shown either as lines or as pairs of dots between two atoms, and lone pairs are shown as pairs of dots on individual atoms.

Only valence electrons are shown in a Lewis structure

H HH + H → H H →

+ F → FF F


Covalent Bonding and Lewis Structures: Lewis’ Octet rule: An atom other than hydrogen tends to form bonds until it is surrounded by eight valence electrons. In other words, a covalent bond forms when there are not enough electrons for each individual atom to have a complete octet. In a single bond, two atoms are held together by one electron pair. Many compounds are held together by multiple bonds, that is, bonds formed when two atoms share two or more pairs of electrons. A double bond occurs when two atoms share two pairs of electrons. A triple bond when two atoms share three pairs of electrons. For Example: CO2 and C2H4


For Example: N2 and C2H2

Multiple bonds are shorter than single covalent bonds. Bond length is defined as the distance between the nuclei of two covalently bonded atoms in a molecule

For a given pair of atoms triple bonds are shorter than double bonds which are shorter than single bonds

The shorter multiple bonds are also more stable.


Comparison of Ionic and Covalent Properties:


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

Elements with high electronegativity have a greater tendency to attract electrons than do elements with low electronegativity. Electronegativity is related to electron affinity and ionization energy. Fluorine has a high electron affinity and a high ionization energy and thus has a high electronegativity. Sodium has a low electronegativity Electronegativity is a relative concept. An element's electronegativity can be measured only in relation to the electronegativity of other elements.

∆En > 2 → Ionic compound ∆En < 2 → Covalent



Electronegativity Differences result in polar molecules:


Lewis Structure Corrections; PCl5 (40 VE), SF4 (34 VE)

P ClCl

Cl• •

• •

• •

• •

• •

••

••

••

Cl• •

••

••

••

• •Cl •

•••

•• S FF

F

F

• •

• •

• •

• •

• •

••

••

••

••

••

••

• •


VSEPR: Valance Shell Electron Pair Repulsion: The Structure around a given atom is determined principally by minimizing electron pair repulsion. Generic Structure type: Use letter A for the “central” atom, X for any atoms associated with it and E for any electron pairs. Examples: Be (AX2); F (AXE3) C (AX2); O (AXE2) C (AXE); N (AXE)

Be F F • •

• •

• •

• •••

••

C N••

••

-C O•

•••O

• • • •


Classifications: Type Structure Orbital Hybrid Angle(s) AX2 ; AXE Linear sp 180° Type Structure Orbital Hybrid Angle(s) AX3 Trigonal Planar sp2 120° AX2E Bent sp2 <120° AXE2 Linear sp2 180° Type Structure Orbital Hybrid Angle(s) AX4 Tetrahedral sp3 109.5° AX3E Trigonal Pyramidal sp3 <109.5° AX2E2 Bent sp3 <<109.5° AXE3 Linear sp3 180° Type Structure Orbital Hybrid Angle(s) AX5 Trigonal Bipyramidal sp3d 90° & 120° AX4E See-Saw sp3d 90° & 120° AX3E2 T- shaped sp3d 90° AX2E3 Linear sp3d 180° AXE4 Linear sp3d 180°


Classifications (continued): Type Structure Orbital Hybrid Angle(s) AX6 Octahedral sp3d2 90° AX5E Square Pyramidal sp3d2 90° AX4E2 Square Planar sp3d2 90° AX3E3 T-Shaped sp3d2 90° AX2E4 Bent sp3d2 <90° AXE5 Linear sp3d2 180°


AX2 ; AXE AX3 AX2E AXE2 AX4 AX3E AX2E2 AXE3


AX5 AX4E AX3E2 AX2E3 AXE4 AX6 AX5E AX4E2 AX3E3 AX2E4 AXE5





Assignment: Draw all the Example Lewis Structures listed with Formal Charge.




Topic 5: Chemical Bonding 5.0 Ionic and covalent bonds; Properties of covalent and ionic compounds

5.1 Lewis structures, the octet rule.

5.2 Molecular geometry: the VSEPR approach. Molecular polarity. 5.3 Valence bond theory, hybridization and geometry, multiple bonds.

5.4 Molecular orbital theory.


Dipole Moments and Polarity: In molecules containing atoms with different electronegativities, there is an unequal sharing of electron density. In other words, a polar covalent bond is formed. This can be represented by: Dipole Moment (µ): Is a quantitative measure of polarity. Magnitude of charge Q multiplied by the distance between charges.

µ = Q x r µ is normally expressed in Debye units (D)

1 D = 3.336 x 10-30 C m

F H FHδ+ δ-


Dipole Moments and Polarity: Dipoles are vector quantities that are additive to produce an overall molecular dipole moment.

For Diatomic molecules from atoms of different electronegativity there is always a dipole moment.

For polyatomic molecules the resultant dipole moment is a vector sum of the individual bond dipoles. As a result not all molecules with polar covalent bonds have a net molecular dipole.

CO2 is a non-polar molecule EN: O (3.5); C (2.5) N (3.0); H (2.1); F (4.0)

OC O


Dipole Moments and Polarity:

Other Examples: C2Cl2H2 (connectivity ClHC2HCl)


Molecular Orbital Theory: The previous MO model, Valence Bond Theory, does not account for all of the observed properties of molecules, for example the paramagnetic characteristics of Oxygen. Rather than considering an interaction of Atomic Orbitals centered on atoms, MO theory develops a new set of molecular orbitals. For H2 formation: In a sigma molecular orbital, σ, (bonding or antibonding) the electron density is concentrated symmetrically around a line between the two nuclei of the bonding atoms. A single covalent bond (such as H-H or F-F) is almost always a σ bond.






Bond Order: Fundamentally, a molecule forms because it has lower energy

than the separated atoms. In the simple MO model this is reflected by the

number of bonding electrons (those that achieve lower energy in going from the

free atoms to the molecule) versus the number of antibonding electrons (those

that are higher in energy in the molecule than in the free atoms). If the number of

bonding electrons is greater than the number of antibonding electrons in a given

molecule, the molecule is predicted to be stable.

The quantitative indicator of molecular stability (bond strength) for a diatomic

molecule is the bond order: the difference between the number of bonding

electrons and the number of antibonding electrons, divided by 2.

Bond Order = (# Bonding Electrons) – (# Antibonding Electrons)

2


Combinations of p-orbitals: Consider the shape of p orbitals and how they might be able to combine.


In a pi molecular orbital (bonding or antibonding), the electron density is concentrated above and below an imaginary line joining the two nuclei of the bonding atoms. Two electrons in a p molecular orbital form a pi bond




Summary of Diatomic Molecular orbital Formation:


Measuring Paramagnetic Properties:




Rules Governing Molecular Electron Configuration and Stability: The number of molecular orbitals formed is always equal to the number of atomic orbitals combined

The more stable the bonding molecular orbital, the less stable the corresponding antibonding molecular orbital

The filling of molecular orbitals proceeds from low to high energies. In a stable molecule, the number of electrons in bonding molecular orbitals is always greater than that in antibonding molecular orbitals

Each molecular orbital can accommodate up to two electrons with opposite spins in accordance with the Pauli exclusion principle

Electrons enter molecular orbitals with parallel spins (Hund's rule)

The number of electrons in the molecular orbitals is equal to the sum of all the electrons on the bonding atoms

Bond Order = ½ ((# e- in Bonding MO’s)-( # e- in Antibonding MO’s))


Examples: Comment on Stability of H2+, H2, He2

+, He2 Recall:


Heterogenous Diatomic Molecular Orbitals:

NO (11 VE) NO+ or CN- (10 VE)

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Chemistry 121: Topic 4 - Chemical Bonding Topic 4 ... · PDF fileIB IIB 13 Al 26.98 14 Si 28.09 15 P ... Topic 4 - Chemical Bonding Topic 5: ... Chemistry 121: Topic 4 - Chemical Bonding