When an acid and a base react, we call it a neutralisation reaction. The products are water and a salt. Until now, we have looked at strong acids (like HCl) reacting with strong bases (like NaOH). The salt is neutral (pH = 7), so the final (salt) solution is neutral.
However, this is not the case when we use a weak acid or a weak base. The salts themselves react with water, so affect the pH. For example, when we react hydrochloric acid (strong acid) with ammonia (weak base), the product is acidic ammonium chloride.
The strength of an acid can be determined in a few ways:
pH
rate of reaction
conductivity
These are all linked to the same thing - the extent of dissociation.
Strong Acids
Strong acids fully dissociate. This means they completely break apart into their ions (the conjugate base and hydronium). Therefore, they have a high concentration of hydronium ions. With a higher concentration of hydronium ions, they:
have a lower pH
have more hydronium ions per unit volume to react
have a higher concentration of ions to carry a charge
Weak Acids
Weak acids only partially dissociate. They still produce hydronium ions, but there are still "unreacted" acid molecules/particles so the concentration of hydronium ions is lower than for a strong acid. This means they:
have a higher pH (but still <7)
have fewer hydronium ions per unit volume (than strong acids) to react
have a lower concentration of ions to carry a charge
Last year, we learned about the nature and reactions of acids and bases. This year, we refine and deepen this understanding, including a better understanding of pH and acid/base strength.
Brønsted-Lowry Acids and Bases
We often simplify our understanding of acids as substance with H+ ions, and bases as substance with OH- ions. Unfortunately, this understanding is inadequate and incorrect.
This year, we use the Brønsted-Lowry definitions. We can use their reactions with water to show the difference between acids and bases:
ACIDS
An acid is a proton donor. This means it loses an H+ ion.
HA + H2O⇌ H3O+ + A-
BASES
A base is a proton acceptor. This means it gains an H+ ion.
When a system is in dynamic equilibrium, the concentration of reactants and products stays the same because the rates of reaction of the forward and reverse reactions are the same.
What if we add an external influence to the system, such as heat, pressure, or more reactants...?
This is where le Chatelier's Principle comes in. It states that the equilibrium system will "shift" to minimise the impact of any external factors (such as those mentioned). What this means is that the rate of one of the reactions will increase (but not the other one!), re-establishing the equilibrium "position" - the concentration of the reactants or products will change until the KC is correct (for that specific temperature - remember that the KC for an equilibrium is different at different temperatures).
This is very well explained here (as well as being an overview of everything to do with Equilibrium so far):
We need to be able to explain the effect on an equilibrium of:
Increasing the amount (therefore concentration) of a reactant or product
Removing some/all (therefore reducing the concentration) of a reactant or product
Increasing/decreasing the temperature
Increasing/decreasing the pressure
Introducing a catalyst
We then need to link these to what we would observe (temperature changes, colour changes etc.)
Changing Concentration of Reactants or Products
When we increase the concentration of a reactant (by adding more of it into the system), the forward reaction is favoured. This means the forward reaction's rate or reaction is increased. This is to "use up" some of the reactants, so re-establishing a dynamic equilibrium system. If the forward reaction is exothermic, the system will also feel warmer. If the forward reaction is endothermic, the system will absorb heat energy from the surroundings, so feel colder.
When we reduced the concentration of a reactant (maybe by reacting some of it with another chemical that we introduce), the reverse reaction is favoured. The rate of reaction of the reverse reaction increases in order to make more reactants, so re-establishing a dynamic equilibrium system. If the forward reaction is endothermic, the system will also feel warmer as the reverse reaction is exothermic. If the forward reaction is exothermic, the system will absorb heat energy from the surroundings, so feel colder, as the reverse reaction must be endothermic.
You can apply both of these to a change in concentration of the products.
Changing the Temperature
If you add heat energy, the system will try to absorb this heat energy, so favour the endothermic reaction until a new dynamic equilibrium position is established.
If you cool the system, the exothermic reaction will be favoured.
Changing Pressure
This only affects gaseous particles. If you increase the pressure, you limit the space for gaseous particles. Therefore, an equilibrium system will favour the reaction which produces fewer gaseous particles until a new dynamic equilibrium is established.
Introducing a Catalyst
This has no effect upon the position of the equilibrium. However, it does mean that dynamic equilibrium is established quicker, as the catalyst increases the rate of both the forward and the reverse reactions.
The position of a system in dynamic equilibrium can be described mathematically, using the equilibrium constant, KC.
We need to know how to:
Write a KC expression for an equilibrium
Calculate KC (when given the concentrations of reactants and products at equilibrium).
Calculate the concentration of a reactant/product (when given the KC value and concentrations of all other reactants and products).
Explain whether the equilibrium "favours" the forward or reverse reaction, based upon the KC value.
We can then use this knowledge to apply le Chatelier's Principle to changes in the equilibrium position due to outside influences (our next key concept).
Some reactions are easily reversed. In class, we saw that blue hydrated copper sulfate crystals can be turned into anhydrous white copper sulfate crystals, by heating (which removes the crystallised water molecules). By adding a drop or two of water, this is reversed, and releases a lot of heat energy!
Equilibrium
When a reversible reaction is contained within a closed system (such as a boiling tube with a rubber/cork bung, or a bottle with a lid on it), the forward and reverse reactions "compete", until the rate of the forward reaction is equal to the rate of reaction of the reverse reaction. The concentrations of the reactants and products do not change, even though both reactions (forward and reverse) are happening. We call this dynamic equilibrium. We will explore this in more detail over the next week.
This video summarises these two concepts well (and is the one we looked at in class on Monday):
Chemical reactions occur when particles collide with enough energy and in the correct orientation. We can speed up (or slow down) reactions by changing some of the conditions that will cause reactions to occur:
more often
with more energy
in the correct orientation
This part of the topic is focused on how we affect the rate of a reaction, using our knowledge of the Collision Theory.
