Lecture principles of inorganic chemistry

the bond order and the number of lone pairs and these maybe used to derive structures.. Bond OrderBond order is a measure of the number of bonding electron pairs between atoms.. Single

2P32 Winter Term 2015-16 Principles of Inorganic Chemistry

Dr M Pilkington

Lecture 1 – Recapping Important Concepts

σ and π bonds in CH2=CH2

Structure, VSEPR theory and VBT.

Assignment 1 – Drawing Lewis structures and predicting the

shapes/geometries of molecules due after class Tuesday 12thJanuary

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 with

itself and other p-block elements especially N and O.

Wide range of electronegativity, oxidation states, coordination

numbers, 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 Law’

1860-1870 (Mendeleev and Meyer): A periodic repetition of

physical and chemical properties occurs when the elements are

arranged in order of increasing atomic weight [number]’

With the development of atomic theory and spectroscopic techniques

the modern Periodic Table has evolved:

2 Bonding Models:

In covalent species, electrons are shared between atoms.

In an ionic species, one or more electrons are transferred between

atoms to form bonds.

Modern views of molecular structure are, based on applying wave

mechanics to molecules; such studies provide answers as to how and

why atoms combine Two such methods are:

1 Valence Bond (VB) approach- overlap of valence orbitals on

atoms to form bonds

2 Molecular Orbital theory (MO) of bond formation – allocates

electrons to molecular orbitals formed by the overlap (interaction)

of atomic orbitals.

Familiarity with both VB and MO concepts is necessary as it is

often the case that a given situation can be approached using one

or the other of these models.

Lewis structures – you need to be able to draw these.

arrangement of valence electrons in molecules.

the bond order and the number of lone pairs and these maybe

used to derive structures.

3 Shapes of Molecules

being able to discuss and predict chemical properties Although

here we discuss the shapes of “simple” molecules, this topic has

also important applications in the understanding of the behavior

of much larger molecules, e.g the shape of macromolecules in

biology is often important with respect to their biochemical

function

Test Question

 Draw the Lewis Structure of the Nitrato ion NO3-.

Bond Order

Bond order is a measure of the number of bonding electron pairs

between atoms Single bonds have a bond order of 1, double bonds have

a bond order of 2 and triple bonds (the maximum number) have a bond

order of 3 A fractional bond order is possible in molecules and ions that

have resonance structures In the example of ozone, the bond order

would be the average of a double bond and a single bond or 1.5 (3

divided by 2) As the bond order becomes larger, the bond length

becomes smaller

Remember atoms in the 3rdperiod or below e.g P, I do not always obey

the Octet rule!

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

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 negative

charge Molecular shape is important with respect to determining if

a molecule is polar or not

4 Valence Bond Theory

Look at Valence Bond Theory (VBT)

The actual shape of NH3is trigonal pyramidal (approximately tetrahedral

minus one atom).

Hybridization of N = sp 3

N [He] 2s 2 2p 3

2s 2p

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

H H H

N [He] 2s 2 2p 3

H 1s 1

We know that sp 3 hybrids have a 109.5 0 angle

N H H H

N H H 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 the

tetrahedron are missing because they are occupied by lone pairs, not

atoms The shape is called bent The H-O-H angle is less than NH3, due

to the greater repulsions felt with two lone pairs

 Other molecules with 2 bond plus 2 lone pairs include OF2, H2S and SF2

Bond angles vary, but all are significantly less than 109.50.

 Treat this as an exception to the octet rule.

(An atom obeys the octet rule when it gains, looses or shares

electrons to give an outer shell containing eight electrons with

the configuration ns2np6) Many molecules such as neutral

compounds of Boron simply do not contain enough valence

electrons for each atom to be associated with eight electrons.

The Shape of BF3

B F

F F

Six electrons around the Boron

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

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

F

BFF

However

In this structure with a double bond the fluorine atom is

sharing 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.

 Evidence for a resonance structure comes from the B-F

distances measured in the solid state They are shorter by ~15

pm’s compared to the B-F distances in BF4- Generally as we

move from a single bond towards a double bond our bond

lengths shorten by approximately 15 ppm’s.

FB

empty 'p' on B filled 'p' on F

F B F

F

empty 'p' on B filled 'p' on F

The MO diagram is complex but the result for BF3 is one

π-bond spread over 3 B-F links.

To Summarize: BF3

The B atom has three bond pairs in its outer shell Minimizing the

repulsion causes this molecule to have a trigonal planar shape, with the

F atoms forming an equilateral triangle about the B atom The F-B-F

bond angles are all 120°, and all the atoms are in the same plane.

of electrons which can be used to form a bond with the boron:

5 Lewis Acids and Bases

p orbital not used in hybridization

possible - which is at 120° to each other in a plane The remaining p

orbital 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 like

this before they joined

together:

The various atomic orbitals which are pointing towards each other

now merge to give molecular orbitals, each containing a bonding pair

of electrons

σ orbital – no nodal planes

Π orbital one nodal plane containing the nuclei.

 Notice that the p orbitals are so close that they are overlapping

sideways.

different kind In this one the electrons aren't held on the line

between the two nuclei, but above and below the plane of the

molecule A bond formed in this way is called a pi bond.

Π-orbital above and below nodal plane The σ -bond is protected but the Π -bond is sticking up and is

not protected by the rest of the molecule, hence these

electrons are exposed to reacting species and it is why alkenes

and alkynes are reactive.

7 Relationship between Lewis Structure, VBT,VSEPR

 Valence Shell Electron Pair Repulsion Theory (VSEPR) enables us to

predict the shape of the central atoms electron pairs and in turn the

hybridization of the central atom.

Lewis Structure

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

hybridization molecular geometry - only looks at shape of

atoms; not lone pairs

bond overlap(VBT)

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

Number of Bonded Atoms

and Lone Pairs on Central

2 XeF4(36 electrons)

Xe

F

six pairs of electrons around Xe

lone pair geometry - octahedral

A typical midterm/exam question would be:

2 Give (i) the molecular shape, (ii) the electron pair geometry at

the central atom and (iii) the hybridization of the central atom.

Practice Exercise

atom and its molecular geometry.

Lecture 2 - Introduction to Metal Complexes

isomers?

2P32 – Principles of Inorganic Chemistry Dr M Pilkington

Assignment 1 due in next Monday at 4.30pm, please write your name and

student ID on your work and staple all loose sheets together.

Transition Metals – located in the d-block of the periodic

table

transition metals d-block

(f-block) inner transition metals

Rows are across are called Periods Columns down are called Groups

Transition Elements begin in the 4th period

K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga

[Ar]4s1 [Ar]4s2

d-block

Ground State Electronic Configuration for the first

row Transition Metals

Hg – Mercury (hydrargyram) “liquid silver”

ƒ Mn exists in 11 oxidation states -3 upto +7

colored.

“Metal Complexes” (coordination compounds)

metal had “valencies” (oxidation numbers) that could be

satisfied by combination with elements having opposite

valencies.

„ Cr3+valence of +3

„ O2-valence of -2

„ Cl-valence of -1

Examples of metal complexes CrCl3, Cr2O3

BUT CrCl3reacts with ammonia (NH3) to form a new compound.

„ Studied in Switzerland at the University of Zurich

„ He lectured in both organic and inorganic chemistry

„ He developed the theory of coordination chemistry

„ He prepared and studied coordination compounds and discovered optically active forms of 6-coordinate octahedral complexes

„ His coordination chemistry extended through a whole range of systematic inorganic chemistry and into organic chemistry and he was awarded the Nobel Prize in Chemistryin 1913

http://nobelprize.org/chemistry/laureates/1913/werner-lecture.html

Nobel Lecture

Werner studied the following metal complexes:

1 In this series of compounds, cobalt has a constant

coordination number of 6 (coordination number is the

number of groups that can bond directly to the metal).

2 As the NH3molecules are removed they are replaced by

Cl-which acts as if it is covalently bonded to cobalt.

3 Chloride and Ammonia are now called ligands.

4 Ligands are a Lewis base/electron pair donors that can

bind to a metal ion.

5 A metal complex – metal ion combined with ligands.

6 Coordination complexes are neutral and counter ions are not

bonded to the central metal ion but balance the charge.

For example:

[Co(NH3)6]Cl3

within the square bracket.

they balance the charge (Co3+) they are “free” to react with

[Co(NH3)5Cl]Cl2 rewrite as [Co(NH3)5Cl]2++ 2Cl

-Lavender (now only two reactive Cl-’s).

[Co(NH3)4Cl2]Cl rewrite as [Co(NH3)4Cl2]++ Cl

React with 3 moles of AgNO 3

React with 2 moles of AgNO 3

React with 1 mole of AgNO 3

Isomers - have the same formula but different structures, i.e

different spatial arrangements

What is the geometry of [Co(NH3)4Cl2]Cl ?

ortho 1,3 isomermeta 1,4 isomerpara

There are three possibilities so this does not fit with

„ Consider Octahedral

http://www.iumsc.indiana.edu/morphology/symmetry/octahedral.html

6 vertices, 8 sides

M

Metal ion in the center, ligands are on the vertices, all six

vertices are identical.

Geometry of [Co(NH3)4Cl2]Cl is Octahedral

Cl

Cl

Co

ClCl

Co or

1800

900

trans isomer the two Cl

ligands are far apart (1800)

cis isomer - the two Cl ligands are close to each other (900)

Many other examples of complexes of this coordination geometry

known.

This geometry reduces the steric crowding that is a problem in

other geometries and makes them unfavourable.

We accept Werner’s conclusions, today further evidence to

confirm his conclusions is provided by X-ray crystallography.

Why is the coordination number 6 so common?

ligands.

„ Plot the ionic radii of transition metal ions, most of them are in

the range 75-90 pm which can accommodate 6 ligands and hence

Radius Ratio Rules

The structures of many crystals can be rationalised to a first approximation

by considering the relative sizes and numbers of ions present The radius

ratio r+/r- can be used to make a first guess at the likely coordination

number and geometry around the cation using a set of simple rules:

Cubic 8

> 0.73

Octahedral 6

0.41-0.73

Tetrahedral 4

0.22-0.41

Trigonal Planar 3

0.15-0.22

Linear 2

< 0.15

Predicted Coordination Geometry of Cation

Predicted Coordination Number of Cation Value of r+/r-

„ For [M(H20)6]n+in order for the metal ion to accommodate six ligands it

must have at least a 52 pm radius.

„ Six waters fit exactly around a metal ion with 52 pm radius (not

crowded).

„ At 92 pm’s the six waters have moved far enough apart that more

waters can fit.

„ If the metal ion is smaller than 52pm you will fit fewer H2O’s, if it is

larger then you have more space and can go to a larger coordination

number.

„ Hence transition metal ions display a number of coordination numbers

but 6 is very common.

„ N and O donors are the most abundant in biology (amino acids), S is also

present but it is allot larger than O and as a consequence Fe fits 6 O’s

but only 4 S’s This is apparent in thiocluster compounds of Fe.

To Summarize:

Lecture 3 - Classification and Nomenclature

1 Ligand Classification:

Rodgers Chapter 2.

2P32 – Principles of Inorganic Chemistry Dr M Pilkington

Theory of Coordination Chemistry

Alfred Werner (1866-1919)

 1893, age 26: coordination theory

 Nobel prize for Chemistry, 1913

 Addition of 6 mol NH3to CoCl3(aq)

Ligand – Lewis base (electron pair donor) that is bonded

to a metal ion Ligands are anionic or neutral.

HLigand - Lewis Base Has one pair of electrons

1 Ligand Classification

1 Monodentate Ligands

 For example :NH3 is a monodentate ligand.

“one toothed” – bind to a metal ion through a single

Fe can exist in number of oxidation states.

A biologically important metal ion.

[(NC)5Fe(III)CNFe(III)(CN)5]

5-A Biological 5-Application of Fe

Iron-sulfur proteins are proteins characterized by the presence

of iron-sulfur clusters containing sulfide -linked iron centers in

variable oxidation states

Structural motifs

the thiolato sulfur centers, from cysteinyl residues, are terminal

ligands The sulfide groups are either two- or

three-coordinated A common motif features a four iron ions and four

sulfide ions placed at the vertices of a cubane -type structure.

4Fe-4S clusters

Aconitase - aconitate hydratase; is an enzyme that catalyses the

stereo-specific isomerization of citrate to isocitrate via cis- aconitate in the

tricarboxylic acid cycle , a non- redox -active process.

Illustration of pig aconitase in complex with the [Fe4S4] cluster The

protein is colored by secondary structure, and iron atoms are blue and the

sulfur red

3 Ambidentate Ligands

:C N: Ambidentate

NO2- O N O lone pairs on O and N

can bind through either lone pair

It is not possible for N to bind to the same Fe2+

Fe2+ CN

:N C S Thiocyanide ligand can bind through S or N but not both at the

same time to one metal ion

Two kinds of binding sites – the ligand can bind one metal ion

through one or the other but not both simultaneously

You will need to be able to draw the Lewis acid structures correctly so you

can figure out how a ligand will bind.

4 Multi-/Polydentate Chelating Ligands

Multidentate – “multitooth”; chelating - “crab” claw

 Example 1 Ethylenediamine H2NCH2CH2NH2 (en)

Forms a 5-membered chelate ring; you can think of it havings “2

claws” coming in to grab the metal

Ligands that are bound to a metal through several donor sites.

Chelate – “Chelos” (greek) meaning “crab” (crabs grab their food

with two claws, in the same way a metal can be attracted by two

lone pairs from different groups on the same ligand).

H2C CH2

M

HH

HH

The Chelate Effect

complex involving bidentate or polydentate ligands is greater

than that of a complex containing a corresponding number of

comparable monodentate ligands

 This is called the chelate effect.

strain.

Example 2 Ethylenediaminetetraacetate (ETDA)

NCH2CH2N

C

OO

COO

two Nitrogen's and four Oxygens bond to

a single metal ion - Hexadentate Ligand.

N

OMO

O

[M(EDTA)]

n-When we draw the ligand chelated to the metal ion we

do not draw all of the carbons.

N

O N

Cr

-Cr O N

[Cr(edta)]–

Examples of Multidentate Chelating Ligands

OH R'

equimolar with another form

Enol form

the alcohol has an OH ending

enol

C O

O R'

R'R

H

C CO

R'R

H

Mwhere R = CH3

acetylacetone

 -Diketones – bidentate ligands

Make sure you can draw the structures and metal complexes

of all of the ligands on your ligand sheet.

Problem – Test 2007 worth 8 marks

1 Draw Lewis structures for the following ligands:

(i) Pyridine C5H5N (2 marks)

(ii) Nitrato NO3- (2 marks)

(iii) Nitro NO2- (2 marks)

Which ligand(s) above is/are ambidentate and why?

(2 marks)

2 Naming Metal Complexes – refer to handout and Rodgers

1. For complex ions write the cation first and anion last e.g K2[PtCl4] –

(note you do not use the prefix mono, di or tri etc here to

indicate the number of cations)

2. Name the ligands first in alphabetical order, the metals last.

3. Prefixes to indicate numbers (di, tri, tetra, penta, hexa etc…)

for all monoatomic ligands, polyatomic ligands with short names and

neutral ligands with special names (see Table 2.4 Rodgers).

4. Prefixes bis-, tris-, tetrakis-, pentakis-, hexakis- for ligands whose

names contain a prefix of the first type, neutral ligands without

special names, ionic ligands with particularly long names.

For example:

[Cr(H2O)4Cl2]+ - tetraaquodichlorochromium(III) ion.

[Cr(NH2CH2CH2NH2)3]3+ - tris(ethylenediammine)chromium(III) ion.

Note the parenthesis around the organic ligand name.

5. If the anion is complex, add the suffix –ate to the name of the metal

If the symbol comes from latin/greek, then we go back to the

latin/greek for the name of the anion.

6. Put the oxidation state in Roman numerals in parantheses after the

name of the central metal ion.

7. Practice to get the hang of this – Examples from Rodgers and Practice

Handout.

For example:

[CoCl4]2- -tetrachlorocobalt ate (II) ion

[Fe(CN)6]4- - hexacyano ferrate (II) ion

 Anionic Ligands (end in –ide or –e) – add –o

Bridging Ligands – use  to indicate a bridge.

Tetraamminecobalt(III)- -amido--peroxo -tetraamminecobalt(III) ion

 If there is more than one of a given bridging ligand, the prefix

indicating the number of ligands is placed after the .

given in alphabetical order.

For example:

Naming Metal Complexes

Examples:

 [Co(NH3)6]3+ hexaamminecobalt(III) ion.

 [PtCl2(NH3)2] diamminedichloroplatinum(II)

 [Fe(CN)6]4- hexacyanoferrate(II) ion.

 [Fe(H2O)6]2+hexaaquoiron(II) ion.

We distinguish which atom is bound to the metal in the naming

e.g SCN-is named as thiocyanato, but NCS-is named as

isothiocyanato

CN-- cyano, but NC-- isocyano

NO2-- nitro, but ONO-is nitrito

Rodgers Chapter 2, page 22

2P32 – Principles of Inorganic Chemistry Dr M Pilkington

Lecture 4 - Transition Metal Complexes

and Terminology.

Structural Isomers and Stereoisomers.

1 Transition Metal Complexes: Definitions and Terminology.

General Convention

 The word ligand is derived from the Latin verb ‘ligare’ meaning to

bind.

 In a complex we have a Lewis Acid Base interaction: p

neutral ligand to an acceptor.

 A line is used to denote the interaction between an anionic ligand and

the acceptor.

 Often however, this convention is ignored and a line to denote both

types of interaction is used

Review your Acid/Base Interactions

3+

3

 Each N atom donates a pair of electrons to the Co3+metal ion, i.e each

NH3molecule is a Lewis base while the metal ion is the Lewis acid

 We think of the metal to ligand interaction as being essentially covalent,

but in reality this is not entirely true as the character of metal-ligand

interactions varies with the nature of the metal ion and the ligand

+ + +

the Co3+centre

 It implies transfer of charge from ligand to metal and figure (a)

shows the resulting charge distribution This is unrealistic since the

Co3+centre becomes more negatively charged than would be

unfavorable given its electropositive nature

 At the other extreme, consider bonding in terms of an ionic model

(b) the 3+charge remains localized on the cobalt and the six NH

(b), the 3+charge remains localized on the cobalt and the six NH3

ligands remain neutral However this model also does not agree with

experimental studies on this complex So this model is flawed.

1/2 1/2 1/2

+ + + +

 We have to apply Pauling’s Electroneutrality Principle which states

that the distribution of charge on a molecule or ion is such that the

charge on a single atom is within the range +1 to -1 (ideally close to

zero)

 In this case the net charge on the Co3+ metal centre should be close

to zero

 In order to satisfy this the Co In order to satisfy this the Co ion can accept a total of only 3 3+ion can accept a total of only 3

electrons from the six ligands, thus giving the charge distribution

above

 This model is actually 50% ionic and 50% covalent

 This representation shows that a bridging chloride ion donates two pairs

of electrons to two Co3+metal ions which are the Lewis acids, accepting

the lone pairs.

 In lit h n thinkin b t th b ndin th f m l h n th

Cl Co3+

Co3+

 In reality, when thinking about the bonding, the formal charge on the

chloride ion is not actually -1, which means also that the formal charge

on the two Co3+metal ions are not strictly +3 either.

 What is important is that we have a complex above which has an overall

charge of +5 In order to easily determine the overall charge of the

complex, the above representation is easy to use (3+3-1 = 5)

 With respect to thinking about the coordinate bond it does not however

accurately represent the formal charges on the metal ions and the

ligands.

 This is analogous to a C-Cl bond in organic chemistry, we write C-Cl but

in reality this does not accurately describe the bonding interaction

since the electrons are not evenly shared and the truth is Cδ+–Clδ-.

Determination of Formal Oxidation States of Metals

in Coordination Complexes

central metal atom in a complex is very important Proceed as

follows:

 e.g [FeCl4]2-has 4 Cl-ligands and overall 2-charge, so it must

contain Fe+2 or Fe(II).

Isomers – Compounds with the same formula but different

properties that result from different structures There are two

Review of Isomerism - Structural Isomers and Stereoisomers.

1 Structural isomers have the same molecular formula but

different molecular structures (different connectivities or

different numbers and kinds of chemical bonds.

CH3OCH3(dimethylether) and CH3CH2OH (ethanol)

i Geometric isomers have different spatial arrangement

2 Stereoisomers not only have the same formulas but also the

same connectivities of their atoms The spatial arrangements

of the atoms are different There are two examples:

results in different geometries (different bond angles or

different distances between nonbonded atoms, for example)

2 Optical isomers have the same geometrical parameters but

are related as nonsuperimposable mirror images (In other

words, the molecule or ion is chiral.) Optical isomers get

their names because they are able to rotate a

plane-polarized light beam to the left or to the right

Organic example: CHFClI A carbon atom with four

different groups attached to it has a nonsuperimposable

mirror image.

C

H

CH

F

Cl I

CI

mirror images non superimposable

2 Isomerism in Transition Metal Complexes

There are many types of structural isomers in transition metal

complexes We will explore three of them.

1 Ionization isomers - Ligands inside the coordination sphere

Structural Isomers

exchange places with ligands outside the coordination sphere

Ionization isomers are so-named because they give different ions

when dissolved in water

Example: There are three compounds with the formula

CrCl3.6H2O One is violet, one is grey-green, and the third is deep

green

The violet isomer produces 3 moles of silver chloride upon

reaction with silver nitrate, and does not lose water in a

desiccator.

[Cr(H2O)6]Cl3 (violet)

The grey-green isomer gives 2 moles of silver chloride upon

reaction with silver nitrate, and loses one mole of water when

stored in a desiccator

[Cr(H2O)5Cl]Cl2.H2O (grey-green)

The deep green isomer gives 1 mole of silver chloride upon

reaction with silver nitrate, and loses two moles of water when

Note that the chloride ions that react with silver nitrate are

the ones not bonded to the chromium(III) ion, and the water

molecules that are lost in a desiccator are the uncoordinated

ones

2 Linkage isomers - Linkage isomers can exist when one or more

A compound with the formula CoCl2(NO2) 5NH3has two isomers

NO2- An 18 electron system

ON

O an ambidentate ligand

A compound with the formula CoCl2(NO2).5NH3has two isomers,

one yellow and one red

Each precipitates two moles of silver chloride, therefore both

chloride ions are outside the cobalt(III) coordination sphere

Neither has an aqueous solution that is basic to pH paper,

therefore all the ammonias are bonded to cobalt (III)

The obvious possibility is that the ambidentate nitrite group is

differently bonded in these two complexes: [Co(NH ) N O ]Cl and

differently bonded in these two complexes: [Co(NH3)5N O2]Cl2and

[Co(NH3)5O NO]Cl2.

Today, we would assign the structures on the basis of infrared

spectra: N- and O-bonded nitrite have different N-O stretching