Who accepts H+ ions and who donates H+ ions?

Let's look at the whole picture now. There is a scale for acids and bases just like everything else. Here are a couple of definitions you should know:

Acid: A solution that has an excess of H+ ions. It comes from the Latin word acidus, which means "sharp" or "sour".

Base: A solution that has an excess of OH- ions. Another word for base is alkali.
Aqueous: A solution that is mainly water. Think about the word aquarium. AQUA means water.
Strong Acid: An acid that has a very low pH (0-4).
Strong Base: A base that has a very high pH (10-14).
Weak Acid: An acid that only partially ionizes in an aqueous solution. This means that not every molecule breaks apart. Weak acids usually have a pH close to 7 (3-6).
Weak Base: A base that only partially ionizes in an aqueous solution. This means that not every molecule breaks apart. Weak bases usually have a pH close to 7 (8-10).
Neutral: A solution that has a pH of 7. It is neither acidic nor basic.
We told you about that guy Arrhenius and his ideas about concentrations of hydrogen and hydroxide ions. You're also going to learn about Brønsted-Lowry ideas. These two chemists from Denmark and England looked at acids as donors and bases as acceptors. What were they donating and accepting? Hydrogen ions. It's a lot like the first definition we gave, where an acid breaks up and releases/donates a hydrogen ion. This newer definition is a little bit more detailed. Scientists used the new definition to describe more bases, such as ammonia (NH3). Since bases are proton acceptors, when ammonia was seen accepting an H+ and creating an ammonium ion (NH4+), it could be labeled as a base. You didn't have to worry about hydroxide ions anymore. If it got the H+ from a water molecule, then the water (H2O) was the proton donor. Does that mean the water was the acid in this situation? Yes.

A chemist named Lewis offered a third way to look at acids and bases. Instead of looking at hydrogen ions, he looked at pairs of electrons (remember our pictures with dot structures in Atoms and Elements?). In Lewis' view, acids accept pairs of electrons and bases donate pairs of electrons. We know that both of these descriptions of acids and bases use completely opposite terms, but the idea is the same. Hydrogen ions still want to accept two electrons to form a bond. Bases want to give them up. Overall, Lewis' definition was able to classify even more compounds as acids or bases.
What really happens in those solutions? It gets a little tricky here. Let's look at the breakup of molecules in aqueous (water-based) solutions one more time for good measure. Acids are compounds that dissociate (break) into hydrogen (H+) ions and another compound when placed in an aqueous solution. Remember that acetic acid example? Bases are compounds that break up into hydroxide (OH-) ions and another compound when placed in an aqueous solution. We'll talk about baking soda in a few paragraphs.

Let's change the wording a bit. If you have an ionic/electrovalent compound and you put it in water, it will break apart into two ions. If one of those ions is H+, the solution is acidic. The strong acid hydrogen chloride (HCl) is one example. If one of the ions is OH-, the solution is basic. An example of a strong base is sodium hydroxide (NaOH). There are other ions that make acidic and basic solutions, but we won't be talking about them here.

That pH scale we talked about is actually a measure of the number of H+ ions in a solution. If there are a lot of H+ ions, the pH is very low. If there are a lot of OH- ions compared to the number of H+ ions, the pH is high.

Think about this idea for a second: Why would a liquid with high levels of NaOH be very basic, yet dangerous at the same time? The Na-OH bond breaks in solution and you have sodium ions (positive) and hydroxide ions (negative). The sodium ions don't really pose a danger in solution, but there are a huge number of hydroxide ions in solution compared to the hydrogen ions that might be floating around as H3O+ (a hydronium ion).All of those excess OH- ions make the pH super high, and the solution will readily react with many compounds. The same thing happens on a less dangerous scale when you add baking soda to water. During the dissociation, OH- ions and carbonic acid are released in the solution. The number of OH- ions is greater than the number of H3O+ ions (H+ and H2O), and the pH increases. It's just not as strong a difference as in sodium hydroxide. That's basically it. (Ha ha! Get it?)

More information in part one.

In 1923, chemists Johannes Nicolaus Brønsted and Thomas Martin Lowry independently developed definitions of acids and bases based on the compounds' abilities to either donate or accept protons (\(H^+\) ions). In this theory, acids are defined as proton donors; whereas bases are defined as proton acceptors. A compound that acts as both a Brønsted-Lowry acid and base together is called amphoteric.

Brønsted-Lowry theory of acid and bases took the Arrhenius definition one step further, as a substance no longer needed to be composed of hydrogen (H+) or hydroxide (OH-) ions in order to be classified as an acid or base. For example , consider the following chemical equation:

\[ HCl \; (aq) + NH_3 \; (aq) \rightarrow NH_4^+ \; (aq) + Cl^- \; (aq) \]

Here, hydrochloric acid (HCl) "donates" a proton (H+) to ammonia (NH3) which "accepts" it , forming a positively charged ammonium ion (NH4+) and a negatively charged chloride ion (Cl-). Therefore, HCl is a Brønsted-Lowry acid (donates a proton) while the ammonia is a Brønsted-Lowry base (accepts a proton). Also, Cl- is called the conjugate base of the acid HCl and NH4+ is called the conjugate acid of the base NH3.

A Brønsted-Lowry acid is a proton (hydrogen ion) donor.
A Brønsted-Lowry base is a proton (hydrogen ion) acceptor.

In this theory, an acid is a substance that can release a proton (like in the Arrhenius theory) and a base is a substance that can accept a proton. A basic salt, such as Na+F-, generates OH- ions in water by taking protons from water itself (to make HF):

\[F^-_{(aq)} + H_2O_{(l)} \rightleftharpoons HF_{(aq)} + OH^-\]

When a Brønsted acid dissociates, it increases the concentration of hydrogen ions in the solution, \([H^+]\); conversely, Brønsted bases dissociate by taking a proton from the solvent (water) to generate \([OH^-]\).

\[HA_{(aq)} \rightleftharpoons A^-_{(aq)} + H^+_{(aq)}\]

Acid Ionization Constant:

\[K_a=\dfrac{[A^-][H^+]}{[HA]}\]

\[B_{(aq)} + H_2O_{(l)} \rightleftharpoons HB^+_{(aq)} + OH^-_{(aq)}\]

\[K_b = \dfrac{[HB^+][OH^-]}{[B]}\]

The determination of a substance as a Brønsted-Lowery acid or base can only be done by observing the reaction. In the case of the HOH it is a base in the first case and an acid in the second case.

To determine whether a substance is an acid or a base, count the hydrogens on each substance before and after the reaction. If the number of hydrogens has decreased that substance is the acid (donates hydrogen ions). If the number of hydrogens has increased that substance is the base (accepts hydrogen ions). These definitions are normally applied to the reactants on the left. If the reaction is viewed in reverse a new acid and base can be identified. The substances on the right side of the equation are called conjugate acid and conjugate base compared to those on the left. Also note that the original acid turns in the conjugate base after the reaction is over.

Note

Acids are Proton Donors and Bases are Proton Acceptors

For a reaction to be in equilibrium a transfer of electrons needs to occur. The acid will give an electron away and the base will receive the electron. Acids and Bases that work together in this fashion are called a conjugate pair made up of conjugate acids and conjugate bases.

\[ HA + Z \rightleftharpoons A^- + HZ^+ \]

A stands for an Acidic compound and Z stands for a Basic compound

A Donates H to form HZ+.
Z Accepts H from A which forms HZ+
A- becomes conjugate base of HA and in the reverse reaction it accepts a H from HZ to recreate HA in order to remain in equilibrium
HZ+ becomes a conjugate acid of Z and in the reverse reaction it donates a H to A- recreating Z in order to remain in equilibrium

Why is \(HA\) an Acid?
Why is \(Z^-\) a Base?
How can A- be a base when HA was and Acid?
How can HZ+ be an acid when Z used to be a Base?
Now that we understand the concept, let's look at an an example with actual compounds! \[ HCl + H_2O \rightleftharpoons H_3O^+ + Cl^¯ \]

HCL is the acid because it is donating a proton to H2O
H2O is the base because H2O is accepting a proton from HCL
H3O+ is the conjugate acid because it is donating an acid to CL turn into it's conjugate acid H2O
Cl¯ is the conjugate base because it accepts an H from H3O to return to it's conjugate acid HCl

How can H2O be a base? I thought it was neutral?