Using the Reaction Between Iron (III) Ions and Thiocyanate Ions to Illustrate Le Chatelier’s
Principle (Lab Report IX)
I. INTRODUCTION
Le Chatelier’s principle is a set of principles that allow us to predict which way the
position of equilibrium in a mixture will shift when the mixture is subjected to a change
in pressure, temperature and concentration. It is built on the notion that, “when a
system at equilibrium is subjected to a change, the system will respond in a way that
minimizes the effect of this change”. In other words, when an equilibrium mixture is
exposed to a certain “stress”, whether it be a change in temperature, change in
pressure, or a change in the concentration of one of the reactants or products, the
position of equilibrium will shift in the direction that opposes the effect of that stress.
A system is in equilibrium when the concentrations of the reactants and products are
both constant. Taking the change in concentration as an example to prove Le
Chatelier’s principle, if we increase the concentration of the product in a chemical
reaction in equilibrium, the system will react by moving the equilibrium position to
the left. This means that the reverse reaction will be favored now, and more product
will be converted back to its constituent reactants. This is done to make up for the
increased concentration of the product, so that the concentrations of both the
reactants and the products return to the way they were in equilibrium.
Iron (III) ions in aqueous solution react reversibly with thiocyanate ions to form the
blood red iron (III) thiocyanate ion in the following chemical reaction:
Fe3+ (aq) + SCN- ⇌ [FeSCN]2+ (aq)
The aim of this lab experiment is to see how changing the concentrations of the reactants in
the mixture will affect the position of equilibrium. Since the product formed has a very visible
red color, the intensity of this color can act as a suitable indicator of the position of
equilibrium in the mixture – the greater the concentration of the iron (III) thiocyanate ion, the
more intense the color, and the further to the right the position of equilibrium lies (towards
the products side). We will observe the changes that happen to the red color of iron (III)
thiocyanate in four individual mixtures where the “condition” that is changed is the
concentration. We will then check to see whether these changes can be explained by the
principle of Le Chatelier.
II. METHODOLOGY
Our dependent variable is the equilibrium position. This will be determined by the
change in the intensity of the red color in the initial mixture.
, Our independent variable is the concentration that is being manipulated. The
concentrations that will be changed are those of the iron (III) ions, thiocyanate ions,
and sodium hydroxide solution. The sodium hydroxide is not a reactant in the reaction,
but its function in this case would be to bond with the …
Here are our control variables:
Variable Why it was controlled How it was controlled
1. Initial mixture of 1. Before adding the 1. Placing the same
iron (III) ions and other reagents, we amount of
thiocyanate ions. must start with the potassium
2. Comparing the same initial thiocyanate
red color in all the mixture, or “stock solution and iron
mixtures to the solution”. We are (III) nitrate
same reference. only observing the solution (5 ml of a
3. Adding just one changes that 0.01 M solution),
drop of the given happen when the then dividing the
reagent in each concentrations are mixture into four
test tube mixture. changed according test tubes before
to the procedure. starting the
2. So that we are experiment.
correctly judging 2. Using one
the intensity of the reference mixture
red color in each – the initial
mixture, by mixture without
comparing it to the added
how the color reagents.
would look when 3. Using a pipette to
there is only the add precisely one
iron (III) drop of each
thiocyanate reagent.
complex in the text
tube.
3. Only one drop of
each reagent
should be added so
that the changes
can be observed at
the same scale.
Table 1.0 – Control variables in this experiment
Principle (Lab Report IX)
I. INTRODUCTION
Le Chatelier’s principle is a set of principles that allow us to predict which way the
position of equilibrium in a mixture will shift when the mixture is subjected to a change
in pressure, temperature and concentration. It is built on the notion that, “when a
system at equilibrium is subjected to a change, the system will respond in a way that
minimizes the effect of this change”. In other words, when an equilibrium mixture is
exposed to a certain “stress”, whether it be a change in temperature, change in
pressure, or a change in the concentration of one of the reactants or products, the
position of equilibrium will shift in the direction that opposes the effect of that stress.
A system is in equilibrium when the concentrations of the reactants and products are
both constant. Taking the change in concentration as an example to prove Le
Chatelier’s principle, if we increase the concentration of the product in a chemical
reaction in equilibrium, the system will react by moving the equilibrium position to
the left. This means that the reverse reaction will be favored now, and more product
will be converted back to its constituent reactants. This is done to make up for the
increased concentration of the product, so that the concentrations of both the
reactants and the products return to the way they were in equilibrium.
Iron (III) ions in aqueous solution react reversibly with thiocyanate ions to form the
blood red iron (III) thiocyanate ion in the following chemical reaction:
Fe3+ (aq) + SCN- ⇌ [FeSCN]2+ (aq)
The aim of this lab experiment is to see how changing the concentrations of the reactants in
the mixture will affect the position of equilibrium. Since the product formed has a very visible
red color, the intensity of this color can act as a suitable indicator of the position of
equilibrium in the mixture – the greater the concentration of the iron (III) thiocyanate ion, the
more intense the color, and the further to the right the position of equilibrium lies (towards
the products side). We will observe the changes that happen to the red color of iron (III)
thiocyanate in four individual mixtures where the “condition” that is changed is the
concentration. We will then check to see whether these changes can be explained by the
principle of Le Chatelier.
II. METHODOLOGY
Our dependent variable is the equilibrium position. This will be determined by the
change in the intensity of the red color in the initial mixture.
, Our independent variable is the concentration that is being manipulated. The
concentrations that will be changed are those of the iron (III) ions, thiocyanate ions,
and sodium hydroxide solution. The sodium hydroxide is not a reactant in the reaction,
but its function in this case would be to bond with the …
Here are our control variables:
Variable Why it was controlled How it was controlled
1. Initial mixture of 1. Before adding the 1. Placing the same
iron (III) ions and other reagents, we amount of
thiocyanate ions. must start with the potassium
2. Comparing the same initial thiocyanate
red color in all the mixture, or “stock solution and iron
mixtures to the solution”. We are (III) nitrate
same reference. only observing the solution (5 ml of a
3. Adding just one changes that 0.01 M solution),
drop of the given happen when the then dividing the
reagent in each concentrations are mixture into four
test tube mixture. changed according test tubes before
to the procedure. starting the
2. So that we are experiment.
correctly judging 2. Using one
the intensity of the reference mixture
red color in each – the initial
mixture, by mixture without
comparing it to the added
how the color reagents.
would look when 3. Using a pipette to
there is only the add precisely one
iron (III) drop of each
thiocyanate reagent.
complex in the text
tube.
3. Only one drop of
each reagent
should be added so
that the changes
can be observed at
the same scale.
Table 1.0 – Control variables in this experiment
