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2P32 Winter Term 2007-8 Principles of Inorganic ChemistryDr. M. Pilkington

Lecture 1 Recapping Important Concepts

Inorganic Chemistry and the Periodic Table

Bonding Models

Shapes of Molecules - Lewis Structures

Valence bond theory: cases of NH3. H2O and BF3

Lewis Acids and Bases

and bonds in CH2=CH2

The Shapes of Molecules Relationship between Lewis

Structure, VSEPR theory and VBT.Assignment 1 Drawing Lewis structures and predicting theshapes/geometries of molecules due in 4.30pm Monday 14th January

1. Inorganic Chemistry and the Periodic Table

Carbon is only one element and has limited bonding modes,oxidation states and coordination numbers.

But it does CATENATE well and forms MULTIPLE BONDS withitself and otherp-block elements especially N and O.

For the rest of the elements:

Wide range of electronegativity, oxidation states, coordinationnumbers, ability to form multiple bonds and catenate etc

How can we make sense of such wide ranging behaviors?

We have a system called the Periodic Table. The Periodic Law1860-1870 (Mendeleev and Meyer): A periodic repetition ofphysical and chemical properties occurs when the elements arearranged in order of increasing atomic weight [number]

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With the development of atomic theory and spectroscopic techniquesthe modern Periodic Table has evolved:

2P32 Course Outline:

Lectures 1-16Coordination Chemistryof transition Metal ions

Lectures 17 34Descriptive InorganicChemistry Main GroupElements.

2. Bonding Models:

In covalent species, electrons are shared between atoms.In an ionic species, one or more electrons are transferred betweenatoms to form bonds.

Modern views of molecular structure are, based on applying wavemechanics to molecules; such studies provide answers as to how andwhy atoms combine. Two such methods are:1. Valence Bond (VB) approach- overlap of valence orbitals onatoms to form bonds.2. Molecular Orbital theory (MO) of bond formation allocateselectrons to molecular orbitals formed by the overlap (interaction)of atomic orbitals.

Familiarity with both VB and MO concepts is necessary as it isoften the case that a given situation can be approached using oneor the other of these models.

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Lewis structures you need to be able to draw these.

Lewis presented a simple but useful method of describing thearrangement of valence electrons in molecules.

Lewis structures give the connectivity of an atom in a molecule,

the bond order and the number of lone pairs and these maybeused to derive structures.

Revise your first year notes.

3. Shapes of Molecules

Understanding the shapes of molecules is an important step inbeing able to discuss and predict chemical properties. Althoughhere we discuss the shapes of simple molecules, this topic hasalso important applications in the understanding of the behaviorof much larger molecules, e.g the shape of macromolecules inbiology is often important with respect to their biochemicalfunction

Test Question

Draw the Lewis Structure of the Nitrato ion NO3-.

How many bonds, how many bonds?

What is the nitrogen-oxygen bond order?

Are there possible resonance structures, can you draw them?

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Bond Order

Single bond - first order

Double bond = second order

Triple bond third order

Bond order is a measure of the number of bonding electron pairsbetween atoms. Single bonds have a bond order of 1, double bonds havea bond order of 2 and triple bonds (the maximum number) have a bondorder of 3. A fractional bond order is possible in molecules and ions thathave resonance structures. In the example of ozone, the bond orderwould be the average of a double bond and a single bond or 1.5 (3divided by 2). As the bond order becomes larger, the bond lengthbecomes smaller.

Remember atoms in the 3rd period or below e.g. P, I do not always obeythe Octet rule!

The Shape of Ammonia (NH3) VSEPR is important here.

N HH

H

Lone Pair

Lewis Structure

We have to consider repulsions between the lone pair and valence electronsactual structure:

N

HH

H

H-N-H angle is just slightly smaller than 109.50

The Nitrogen atom is Pyramidal

But why isnt the NHN angle 900?

Ammonia is a polar molecule with N carrying a partial negativecharge. Molecular shape is important with respect to determining ifa molecule is polar or not.

4. Valence Bond Theory

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Look at Valence Bond Theory (VBT)

The actual shape of NH3 is trigonal pyramidal (approximately tetrahedralminus one atom).

Hybridization of N = sp3N [He] 2s2 2p3

2s 2p

Hybridizationmix the orbitals -" like mixing together a red and white plant"

H HH

N [He] 2s2 2p3

H 1s1

We know that sp3 hybrids have a 109.50 angle

N

H

H

H

N

HH

H

Molecular Structure of NH3- cannot see the lone pair on N but

there is a flattened lone pair

Compared to H20

The O in H2O has 2 bond pairs and 2 lone pairs. Two corners of thetetrahedron are missing because they are occupied by lone pairs, notatoms. The shape is called bent. The H-O-H angle is less than NH3, dueto the greater repulsions felt with two lone pairs

Other molecules with 2 bond plus 2 lone pairs include OF 2, H2S and SF2.Bond angles vary, but all are significantly less than 109.50.

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Treat this as an exception to the octet rule.(An atom obeys the octet rule when it gains, looses or shareselectrons to give an outer shell containing eight electrons withthe configuration ns2np6). Many molecules such as neutralcompounds of Boron simply do not contain enough valenceelectrons for each atom to be associated with eight electrons.

The Shape of BF3

B

F

F

F

Six electrons around the Boron

2s 2p

B 2s2 2pF 2s2 2p5

sp2 this leaves an empty 2p orbital

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This leaves an unused "p orbital" perpendicular to the plane of BF3

F B

F

F

But if we want B to have an octet how can we achieve this?

A hybrid of 4 resonance structures is the best Lewis representation for the real

strucure of BF3.

F

B

F F

F

B

F F F

B

F

F F

B

F

F

However...In this structure with a double bond the fluorine atom issharing extra electrons with the boron.The fluorine would have a '+' partial charge, and the boron a'-' partial charge, this is inconsistent with the

electronegativities of fluorine and boron.Conclusion - the Octet Rule breaks down here.

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6. versus -bonding

Ethene, C2H4, sp2

HH

H H

p orbital not used in hybridization

The three sp2 hybrid orbitals arrange themselves as far apart as

possible - which is at 120 to each other in a plane. The remaining porbital is at right angles to them.

C-H overlap to give sigma bonds.

Nodal Plane

fn = 0 (wave function)

i.e. no electron density

Two lobes one with apositive sign the otherwith a negative sign gothough a node.

The two carbon atoms and four

hydrogen atoms would look likethis before they joinedtogether:

The various atomic orbitals which are pointing towards each othernow merge to give molecular orbitals, each containing a bonding pairof electrons.

orbital no nodal planes

orbital one nodal plane containing the nuclei.

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Notice that the p orbitals are so close that they are overlappingsideways.

This sideways overlap also creates a molecular orbital, but of adifferent kind. In this one the electrons aren't held on the linebetween the two nuclei, but above and below the plane of themolecule. A bond formed in this way is called api bond.

-orbital above andbelow nodal plane

The -bond is protected but the -bond is sticking up and is

not protected by the rest of the molecule, hence theseelectrons are exposed to reacting species and it is why alkenesand alkynes are reactive.

7. Relationship between Lewis Structure, VBT,VSEPR

Valence Shell Electron Pair Repulsion Theory (VSEPR) enables us topredict the shape of the central atoms electron pairs and in turn thehybridization of the central atom.

Lewis Structure

Electron Pair Geometry (VSEPR) - non bonding electrons and bonded atoms

hybridization molecular geometry - only looks at shape ofatoms; not lone pairs

bond overlap

(VBT)

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Methane, Ammonia, Water

Electron pair = nonbonding electrons + bonded atoms

Molecular only looks at shape of atoms; not lone pairs

Electron pair geometry: Tetrahedral Tetrahedral Tetrahedral

Molecular geometry : Tetrahedral Triangular pyramidal Bent/Angular

109.5 107.5 104.5

sp3d2Octahedral6

sp3dTrigonal Bipyramidal5

sp3Tetrahedral4

sp2Trigonal Planar/ triangular3

spLinear2

HybridizationShape (e-pair Geometry)Number of Bonded Atomsand Lone Pairs on CentralAtom

Examples:

1. H2O

- electron pair geometry = tetrahedral (2 lpand 2 bp)

- molecular shape = bent

- O hybrization = sp3

H O H

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2. XeF4 (36 electrons)

Xe

F

F F

F

six pairs of electrons around Xe

lone pair geometry - octahedral

Xe = sp3d2 hybridized

F = sp3 hybridized

the lone pairs are far appart therefore the compound as a SQUARE PLANARmolecular geometry.

Xe

F F

F F

A typical midterm/exam question would be:

1. Draw the Lewis Structure of XeF4

2. Give (i) the molecular shape, (ii) the electron pair geometry atthe central atom and (iii) the hybridization of the central atom.