Many chemical reactions go to completion. This means that the
reaction continues until all component particles (whether ions,
molecules or atoms) of one reagent have reacted, then stops.
The reaction between hydrochloric acid calcium carbonate is an
example:
2HCl(aq) + CaCO3(s) → CaCl2(aq) + H2O(l) + CO2(g)
The reaction may stop because all of the calcium carbonate, all
of the acid or all of both reagents have reacted completely.
Fuel combustion is another example: as long as a car has fuel in
its tank, the reaction will continue:
2C8H18(l) + 25O2(g) → 16CO2(g) + 18H2O(g)
As soon as the car ‘runs out of petrol’ (or diesel, LPG, etc.), it will
not move, because although the supply of oxygen is plentiful, no
fuel remains to react with the oxygen.
Some chemical reactions do not go to completion.
Instead, significant amounts of products react together
to reform the original reactants. For example, when a
solution of silver ions is added to a solution of iron(II)
ions, the following reaction occurs:
Ag+(aq) + Fe2+(aq) → Ag(s) + Fe3+(aq) (Reaction 1)
However, as soon as some products form, a few
particles react to form the original reagents:
Ag(s) + Fe3+(aq) → Ag+(aq) + Fe2+(aq) (Reaction 2)
Eventually the rate at which the products of
Reaction 1 form is the same as the rate at
which they react in Reaction 2. The rates of
the forward and reverse reactions are equal.
No further change is observed to the amount
of silver deposited. This position is called
equilibrium. The reaction is said to be in
dynamic equilibrium because the forward and
reverse reactions continue at the same rate,
so no external changes are seen.
EQUILIBRIUM
- The rates of the forward and reverse reactions are equal.
- The concentrations of products and reagents are constant.
- The concentrations of products and reagents are not necessarily
equal.
- The reaction mixture is called the equilibrium mixture.
- No further observable change occurs.
The main condition keeping a reaction at dynamic equilibrium is
that the whole set-up remains isolated, that is, the reaction
occurs in a closed system without intervention from the
atmosphere such as changes in pressure, or temperature, or any
other reagent.
Equilibrium does not mean that products and
reactants are present in equal amounts.
The forward or reverse reaction may dominate,
so products and reactants are not necessarily
in equal concentrations
‘No further change’ does not mean ‘no further
reaction’ or ‘everything stops’ when a system
reaches equilibrium. Students need to know that
particles react continuously, albeit unseen.
‘Equal rates’ does not mean ‘equal
concentrations’. Students need to understand that
this is not so.
Forward and reverse reactions are not separate
and occur alternately.
The reactions are two parts of one, reversible
reaction.
Dynamic equilibrium
One way of helping students understand dynamic equilibrium is to
show someone staying in the same position on a moving
escalator.
To stay still, the person must be walking, either taking steps up or
down, depending on the direction. We cannot see this movement,
but know it must be occurring at the same rate the escalator is
moving, or else the person’s position would change. This is
analogous to a dynamic equilibrium, in which invisible ‘movement’
is going on. Two reactions occur at the same rate, resulting in no
overall change.
Factors affecting rate of
reaction: concentration
The iodine ‘clock’ reaction can be used as a
demonstration to show that the time taken for
a product to be made can change, or in a
more advanced way as an investigation of
precise effects of changing the concentration
of a single reagent.
The clock reaction involves the detection of iodine being
produced in this reaction:
S2O82–(aq) + 2I–(aq) → 2SO42–(aq) + I2(aq) (reaction
1)
The S2O8 2– ion is called peroxodisulfate(VI) or
persulfate. Iodine is the only coloured substance in the
reaction. Iodine forms a blue–black complex with
starch, which is added as an indicator.
A clock reaction measures the time required to produce
a small, fixed amount of product, iodine in this example
To do this, a second reaction is added longside the
first one to ‘hide’ the product for a certain amount
of time. The second reaction used here is:
2S2O32–(aq) + I2(aq) → S4O62–(aq) + 2I–(aq)
(reaction 2)
S2O32– is called the thiosulfate ion and S4O6
2– is called the tetrathionate ion. Reaction 2 has
iodine as a reagent.
As soon as iodine is formed in reaction 1, it reacts with thiosulfate ions to
make colourless iodide ions again. Thiosulfate is called the limiting
reagent because this limits when the iodine is seen: when all
thiosulfate ions have reacted, no more are available to react with any
iodine produced. Instead, iodine produced in reaction 1 will make
the blue–black complex with starch and this colour will appear. The
time taken for iodine to appear depends on the concentration of
iodide ions. If this is high, iodine is produced in reaction 1 more
rapidly, so thiosulfate ions in reaction 2 react and reach their ‘limit’
quickly. If the iodide ion concentration is low, the rate at which
iodine is produced in reaction 1 is slow, so the thiosulfate ions also
react more slowly.
Factors affecting rate of
reaction: temperature
Every reaction has an energy barrier that must be overcome if reactant
molecules are to combine and become product molecules.
Most reactions we show in school chemistry do not need much encouragement,
as room temperature often provides sufficient energy for reactant molecules to
cross the relevant energy barrier. The level of energy required is called the
activation energy . The activation energy for any reaction is fixed – nothing can
be done to raise or lower it.
However, a chemist can change the number of
reactant particles that have the necessary energy
level. Changing temperature is one way of
achieving this. Reaction rate doubles for every 10
°C rise in temperature. The main factor influencing
the rise in reaction rate is that more successful
collisions occur between reactant particles. This is
because increasing the energy level of reagents
means more collisions occur. As the individual
energy level of the particles is raised, more
collisions involve reactant particles with the
activation energy (or enthalpy) to react.
Factors affecting rate of reaction: the
presence of
a catalyst
A catalyst provides an alternative route to the products. The
alternative route has a lower energy barrier than the original
reaction. Students need to understand that the catalyst does
not lower the activation energy. The activation energy for the
original reaction remains unchanged. The alternative routes
vary depending on the catalyst type.
Why do reactions ‘go’
and what happens when reactions
occur?
Measuring energy changes associated with chemical
reactions enables us to understand why they happen
and what exactly occurs between reactant particles
during a reaction. The main reason chemical reactions
occur is that the amount of ‘disorder’ involved always
increases as a result. Disorder can increase in two
ways: either energy is more spread out after the
reaction than before, or the number of ways the
particles in the reaction are arranged is greater than
before. Some chemical reactions involve both of these,
others just one.
2HCl(aq) + CaCO3(s) → CaCl2(aq) + H2O(l) + CO2(g)
The numbers of moles on both sides of the arrow are
the same, three moles of reactants and three moles of
product are shown.
But the state symbols show the reaction produces one
mole of gas. The particles of carbon dioxide can be
arranged in more ways than those of calcium
carbonate, which are fixed in a solid lattice
structure. So, the amount of disorder is increased in the
products.
The second way of thinking about disorder relates to energy distribution. Many reactions are
exothermic. This means heat is lost to the surroundings as products form. This increases the
distribution of the energy. A good example is the combustion of fuel:
2C8H18(l) + 25O2(g) → 16CO2(g) + 18H2O(g)
In this reaction, not only is there a large increase in the number of moles (27 reactant
compared to 34 product) but energy is given out. This is distributed to the car engine to make
movement. Energy is released when chemical bonds form between carbon and oxygen
atoms, making carbon dioxide, and hydrogen and oxygen atoms, making water. More energy
is released when these bonds are made than was required to break bonds between carbon
and hydrogen atoms and oxygen atoms in the reactants.
Chemists use the word entropy to describe
disorder.
Increasing entropy is the most likely reason
that chemical reactions occur.
Students’ misconceptions about
chemical
reactions
1) Students think endothermic reactions cannot be
spontaneous
There is a strong association between chemical
reactions occurring and energy being released.
Students become used to reactions that ‘get
hotter’ when they occur, such as that between
hydrochloric acid and magnesium, or acid and
carbonate reactions. Finding appropriate
spontaneous reactions that get colder is much
harder.
Hence, students develop the notion that only
exothermic reactions occur without extra ‘help’.
2) Students think that bond breaking releases energy
A natural consequence of thinking of fuels as energy
stores is to think that fuel molecules release energy
when they burn. The image is rather like cracking an
egg – when the shell is broken, the contents are
released. Even students who know ‘bond breaking
requires energy’ may think this, on the grounds that a
small amount of energy may be needed to start off the
break, but the actual breaking releases far more. This
misconception makes it very difficult to understand
energetics properly.
Equilibrio chimico
• L’equilibrio chimico è uno stato di
equilibrio dinamico in cui la velocità di
formazione dei prodotti è uguale alla
loro velocità di decomposizione nei
reagenti
N2 + 3 H2  2NH3
La legge di azione di massa
• La composizione della miscela di reazione
all’equilibrio è descritta dalla sua costante di
equilibrio Keq.
• Secondo la legge di azione di massa, per una
generica reazione:
aA + bB  cC + dD
le concentrazioni all’equilibrio delle varie specie
soddisfano:
Kc = [C]c[D]d/[A]a[B]b
Qui la Keq è espressa in funzione delle
concentrazioni
La costante di equilibrio
della reazione
Kc = [C]c[D]d/[A]a[B]b
Il rapporto fra le concentrazioni molari dei
prodotti di reazione ed il prodotto delle
concentrazioni molari dei reagenti
all’equilibrio, ciascuna concentrazione
essendo elevata ad una potenza pari al
coefficiente stechiometrico con cui la specie
compare nella reazione, è costante a T
costante.
Significato di Keq
CO2  CO + ½ O2
A 100°C Kc = 10-36
Kc = [CO][O2]1/2/[CO2] =10-36
Dalla Keq risulta che all’equilibrio le
concentrazioni di CO e O2 sono
trascurabili
E’ bene notare che le
concentrazioni molari
nell’espressione della Keq
sono quelle all’equilibrio, e
non quelle iniziali.
Prevedere la direzione di
una reazione
• La conoscenza di Keq ci consente di
dire se una miscela di reazione di
composizione arbitraria evolverà verso
i prodotti o verso i reagenti
Reazione diretta e reazione
inversa
N2 + 3 H2  2NH3
Kc = [NH3]2/[N2][H2]3
2NH3  N2 + 3 H2
K’c = [N2][H2]3/ [NH3]2
Kc = 1/ K’c
Costante di equilibrio e le
pressioni parziali
• Negli equilibri in fase gassosa può essere
comodo esprimere Keq in funzione delle
pressioni parziali.
Kc = [NH3]2/[N2][H2]3
PV =nRT
CM = n/V = P/RT
Kc = P2NH3/RT2 · RT/PN2 · RT3/P3H2 =
= P2NH3/PN2P3H2 · RT2
Si può definire una nuova costante
Kp = P2NH3/PN2P3H2
•In generale:
Kp = Kc (RT) Dn
dove Dn = differenza fra le moli di prodotti e
quelle di reagenti
Perturbando l’Equilibrio
• Supponiamo di avere un sistema
all’equilibrio
• Disturbiamo ora l’equilibrio
– Aggiungendo o sottraendo reagenti e/o prodotti
– Variando le dimensioni del contenitore
– Variando la pressione
– Variando la Temperatura
• Come reagisce il sistema?
Principio di Le Chatelier
Un sistema all’equilibrio,
soggetto ad una
perturbazione, risponde
in modo da minimizzare
l’effetto della
perturbazione
• Si puo’ razionalizzare
considerando l’espressione
Henri Le Chatelier (1850 - 1936)
della costante di equilibrio e
di come varia cambiando P
eT
Il principio di Le ChatelierBraun
• Sia data una miscela di reazione
all’equilibrio.
• I parametri che determinano la condizione di
equilibrio sono T, V, P e le concentrazioni
delle varie specie.
• Quando si cambia uno di questi parametri, il
sistema evolverà per raggiungere un nuovo
stato di equilibrio che si oppone alla
modifica apportata.
Principio di Le Chatelier-Braun
e posizione dell’equilibrio
•Una variazione in P o nelle concentrazioni
provocherà una variazione nelle
concentrazioni all’equilibrio.
•L’effetto della variazione di T sulla posizione
dell’equilibrio si comprende sapendo se una
reazione è esotermica o endotermica.
Effetto dell’aggiunta di un
reagente
Kc = [C]c[D]d/[A]a[B]b
Se si aumenta la concentrazione di un
reagente la reazione procederà quindi
verso destra fino a ristabilire
concentrazioni tali da soddisfare la Kc.
Effetto opposto se si introduce un
prodotto nella miscela di reazione.
Effetto della pressione
PCl5(g)  PCl3(g) + Cl2(g)
Se si aumenta P, la miscela di equilibrio
cambia composizione nel senso di
diminuire il numero totale di molecole allo
stato gassoso presenti nel recipiente.
Per questa reazione quindi l’equilibrio si
sposterebbe a sinistra.
Non c’è effetto della P se non c’è variazione
nel numero di moli durante la reazione.
Effetto della temperatura
L’aumento di T sposta l’equilibrio
nella direzione che corrisponde alla
reazione endotermica.
Es. N2 + 3 H2  2NH3 DH° = -92 kJ
La reazione è esotermica.
Un aumento di T favorisce la
decomposizione di NH3 nei suoi
prodotti.