Fe acac 2

Foundations of Chemistry Laboratory Manual EQUILIBRIUM and LE CHÂTELIER’S PRINCIPLE

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EXPERIMENT 4F

Equilibrium and

Le Châtelier’s Principle (This experiment is done in pairs. Note: you may wish to divide part 1 and 2 between partners.)

Useful background reading (this is not compulsory but may be helpful):

Tro, 4th and 5th Edition: Sections 15.3, 15.7, 15.8, 14.9 (Intro only) – Questions 1 and 2 Sections 12.1 and 12.6 – Question 3

What is the relevance of this prac…?

The prac brings together several concepts that underpin many areas of chemistry study. You will undertake your first laboratory synthesis in which you make a compound (much like cooking but you don’t get to lick the bowl!).

You will then analyse, using Le Châtelier’s Principle, how the reaction conditions may be optimised in order to maximise the amount of product you obtain. Le Châtelier’s Principle can be used to predict outcomes on a small scale such as your reaction vessel, on a miniscule scale such as in cells and on a planetary scale such as in Earth’s atmosphere.

Finally, you will examine how the charge of a species determines what solvents it can be

dissolved in. The type of possible intermolecular forces present between the solute and solvent

will dictate solubility and this is investigated during this practical. Intermolecular forces are

incredibly important and we take them for granted all the time. They are responsible for oxygen

being a gas at room temperature so we can breathe it in and water being a liquid at room

temperature so we can drink it.

Learning objectives (remember these are different to the scientific objectives):

On completion of this practical, you should have:

 Become familiar with the class of chemical compounds called “co-ordination complexes”

 Understand that a co-ordination complex consists of a metal cation at the centre

surrounded by ligands

 Recall the concept of equilibrium from lectures and consider how it relates to this

practical

A BIG Question

What is life?

Life is dependent on many things working

together in concert to give a cohesive whole.

One of the many things on which human life

is dependent is the process of equilibrium.

Equilibrium processes are involved in

controlling the acidity of our blood and the

transport of oxygen in our bodies, among

many other things.

Foundations of Chemistry Laboratory Manual EQUILIBRIUM and LE CHÂTELIER’S PRINCIPLE

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Fe3+ and 6 x O

H H

 Become familiar with Le Châtelier’s Principle and use it to predict and explain changes in

the equilibrium position based on a change of reaction conditions

Note: this practical has three parts and can be quite long. However, many of the questions

(including all of those in Part Three) do not rely on experimental results and can be answered

prior to the practical if you have an understanding of equilibrium, Le Châtelier’s Principle and

intermolecular forces. It is possible to thoroughly prepare for this practical before you step into

the lab and students who do this should not find any problems with its length.

Introduction (extra background)

CO-ORDINATION COMPLEXES

This experiment involves the synthesis and investigation of a co-ordination complex:

tris(acetylacetonato) iron (III). Co-ordination complexes are a class of compounds that most

commonly involve a central metal ion which is surrounded by a certain number of molecules or

ions called ligands. These ligands are said to “co-ordinate” to the metal centre and therefore

the bond formed between them is referred to as a co-ordinate bond.

The complex you begin with in solution at the start of your reaction in Part One is made up of

the following species:

Fe3+ is the central metal ion and it is surrounded by six water ligands.

The formula for this complex is displayed like this: [Fe(H2O)6] 3+. This

species is also referred to as the “aquated Fe3+ ion” because it is a

metal cation surrounded by only water ligands. Everything inside the

square brackets is part of the complex – the metal centre and ligands.

The charge outside is that of the overall complex. The complex is also

shown diagrammatically in two ways below:

In diagram (a) we can see the Fe3+ cation at the centre with six positions around it for six water

ligands. Diagram (b) is conventionally how complexes are displayed and corresponds to the

formula: [Fe(H2O)6] 3+. The charge on the overall complex is a result of adding all the charges on

the metal centre and ligands together. The metal in this case has a 3+ charge and all six water

Fe3+

H2O

H2O OH2

OH2

OH2

OH2

Fe

H2O

H2O OH2

OH2

OH2

OH2

3+

a b

Foundations of Chemistry Laboratory Manual EQUILIBRIUM and LE CHÂTELIER’S PRINCIPLE

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O

C H

O

H3C CH3

ligands are neutral and contribute no charge (water molecules are neutral). Therefore the

overall charge on the complex is simply 3+, as shown on the outside of the square brackets.

For a species to have the potential to act as a ligand in a complex, it must have a lone pair of

electrons in its structure that it can donate to the metal centre. The atom that has the lone

electron pair on it is known as the donor atom for that ligand. This is the oxygen atom in water

for the complex above. Notice that in the complex diagrams this is why the ligand water

formula is always written to show the oxygen bonded to the metal (ie on the right hand side it

is written OH2 rather than the usual H2O to show O as the donor atom).

All ligands will therefore contain at least one donor atom with a lone electron pair on it. Each

donor atom can form one bond to the metal centre and therefore occupy one position around

it. Some ligands contain more than one donor atom and these are called multidentate ligands.

The ligand present on your final product complex is one such example shown below:

acetylacetonate

acac¯ ion

This is the acetylacetonate ion (acac¯ in shorthand). The two oxygen atoms share the negative

charge and therefore both contain lone electron pairs and both act as donor atoms. The acac¯ is

known as a bidentate ligand because it has two donor atoms which will co-ordinate at two sites.

The chemistry of coordination complexes has been a fertile area of research for many decades.

Coordination complexes are involved in many biological and industrial processes and are the

crucial component in many biologically active compounds. Hæmoglobin, Vitamin B12 and

chlorophyll are coordination complexes of iron, cobalt and magnesium respectively.

GENERAL COMPLEX FORMATION

It is useful to assume that the formation of complexes takes place in a series of steps eg:

M?+(aq) + L?–(aq) [M(L)]?(aq)

[M(L)]?(aq) + L?–(aq) [M(L)2] ?(aq) etc...

We have already seen that the aquated metal ion M?+(aq) may itself be represented as a

coordination complex with water molecules as ligands; [M(H2O)n] ?+(aq). (Eg [Fe(H2O)6]

3+.)

The formation of [M(L)]?(aq) will occur when water molecules are replaced by the ligand, L. If L

is a mono-dentate ligand then it will displace one water molecule. If L is multidentate then it

will displace several water molecules depending on how many donor atoms it has.

Foundations of Chemistry Laboratory Manual EQUILIBRIUM and LE CHÂTELIER’S PRINCIPLE

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O

C H

O

H3C CH3

O

C

O

H3C CH3

HH

acidic hydrogens

base

acetylacteone acacH

acetylacetonate ion acac

FORMATION OF THE COMPLEX TRIS(ACETYLACETONATO) IRON(III)

Acetylacetonate, acac¯, has already been provided as an example of a multidentate ligand. It is

formed by treating the organic compound acetylacetone, acacH, with a base:

In Part One of this experiment you

will use the acacH molecule to synthesise the complex tris(acetylacetonato) iron(III) – known in

shorthand as “tris-acac iron”.

The first step in the formation of the tris-acac iron complex is the hydr