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 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 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 N2 + 3 H2  2NH3 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 Henri Le Chatelier (1850 - 1936) Un sistema all’equilibrio, soggetto ad una perturbazione, risponde in modo da minimizzare l’effetto della perturbazione Si puo’ razionalizzare considerando l’espressione della costante di equilibrio e di come varia cambiando P e T

Il principio di Le Chatelier-Braun 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.