-Brønsted-Lowry acids are proton donors and bases are proton acceptors.
-In water, an acid can donate a proton to form aqueous H+ and the conjugate base; a base can accept a proton from water to form OH– and the conjugate acid.
-A buffer solution makes able to add a strong acid or base to a solution without causing a large change in the pH
-A buffer solution contains both a weak acid (HA) and its conjugate base (A-).
A conjugate acid is the acid compound, usually indicated with HX, of a pair of compounds that differ from each other by gain or loss of a proton. A conjugate acid can release or donate a proton.
To understand it better, we can say it like this:
Acid + Base → Conjugate Base + Conjugate Acid
Example: The ammonium cation is the conjugate acid that is formed when the base ammonia reacts with water:
In water, an acid can donate a proton to form aqueous H+ and the conjugate base while a base can accept a proton from water to form OH– and the conjugate acid.
[H+] is higher the [OH–] in acidic solutions and the pH is less than 7, while the opposite is true in basic solutions ([H+] < [OH– ], pH > 7). Consequently, aqueous solutions of substances such as HCl are acidic, while aqueous solutions of substances such as NaOH or NH3 are basic.
The dissolution of some salts into water can affect pH. Anions that are the conjugate bases of weak acids react with water to form OH–. Cations that are the conjugate acids of weak bases can release H+. In the equations, HA represents a weak acid and A– its conjugate base. A weak acid is an acid what is not completely dissociated in water, but forms an equilibrium instead. The equilibrium constant of this equilibrium is called the acid dissociation constant, Ka, and is defined as follow:
HA(aq) A–(aq) + H+(aq)
Ka = [H+][A− ]/[HA]
We can write exaclty the same for a weak base in solution and the equilibrium constant is called the base dissociation constant, Kb.
A–(aq) + H2O(l) HA(aq) + OH–(aq)
Kb = [OH− ][HA]/[A− ]
The stronger an acid, the larger its Ka and the weaker its conjugate base (smaller Kb).
The weaker an acid, the smaller its Ka and the stronger its conjugate base (larger Kb).
If we now add together the above equations, we will find that the net ionic equation is:
H2O(l) H+(aq) + OH–(l)
KaKb = [H+][A− ]/[HA] x [HA][OH− ]/[A− ] = [H+][OH− ] = Kw.
Kw, the water's auto-dissociation constant, has a value of 1.0 × 10-14 at 25°C.
Anions derived from strong acids,which dissociate completely in water, such as Cl–, Br–, I–, HSO4–, ClO4–, and NO3–, do not react with water and do not affect pH. The parent acids are so strong in water that the conjugate bases (the derived anions) are too weak to react with the water. In addition, cations of the group 1A metals (Li+, Na+, K+, Rb+, Cs+ ) and the group 2A metals (Ca2+, Sr2+, Ba2+) do not react with water and are non acids and consequently, they do not affect the pH of the solution in any way. Anions with ionizable protons such as HCO3–, H2PO4–, and HPO42– are amphoteric: they may be acidic or basic, depending on the values of Ka and Kb for the ion, also depending on the chemical species present in the water solution.
A buffer is an aqueous solution that has a highly stable pH. If you add an acid or a base to a buffered solution, its pH will not change significantly. Adding water to a buffer or allowing water to evaporate will not change the pH of a buffer.
We can say that a buffer solution resists large changes in pH upon the addition of small amounts of strong acid or strong base. A buffer has two components: one that will react with added H+ and one that will react with added OH–. When hydrogen ions are added to a buffer, they will be neutralized by the base present in the buffer. If hydroxide ions are added to a buffer solution, they will be neutralized by the acid. These neutralization reactions will not have much effect on the overall pH of the buffer solution, because, as said, they neutralize the excess or acid or base.
Usually these two parts forming a buffer are a weak acid and its conjugate base (or vice versa) and are often prepared by mixing a weak acid (or weak base) with a salt of that acid (or base).
When preparing a buffer solution and selecting an acid for it, you need to choose an acid that has a pKa close to your desired pH. This will give your buffer nearly equivalent amounts of acid and conjugate base so it will be able to neutralize as much H+ and OH- as possible.
added OH– reacts with the weak acid: HA + OH–H2O + A– K1 = 1/Kb, A–
added H+ reacts with the conjugate base A– + H+ HA K2 = 1/Ka, HA.
Example of buffers in real life:
Vinegar is a solution of a weak acid called acetic acid, CH3COOH; its conjugate base is the acetate ion, CH3COO- .
Since sodium acetate (CH3COO- Na+) dissociates in water to yield acetate ions and sodium ions, adding sodium acetate to an acetic acid solution is one way to prepare an acetic acid buffer. The pH of this solution is equal to the pKa of acetic acid, which is 4.76, so acetic acid buffer solutions are best if the desired pH is around 4.76.
Citric acid is found lemons and other citrus fruits. Like acetic acid, it's a weak acid; citric acid is polyprotic, meaning that each molecule can donate more than one hydrogen ion in solution. Citric acid buffers are best if the desired pH is in the 3 to 6.2.