Introduction to acid-base theory

In this lesson, we will learn:

• To understand how acids, bases and salts were originally and currently defined in chemistry.
• To understand how acid-base theory is related to pH and the amphoteric nature of water.
• To understand what H⁺ ions do in solution and how acid-base reactions depend on them.

• Even though acids, bases and salts are common and well-studied in chemistry, there are different sets of definitions that depend on what you focus on in an acid-base reaction. The chemicals that we call ‘acids’ (e.g. hydrochloric acid) and ‘bases’ (such as sodium hydroxide) all have something in common which leads to the acidic/basic properties that we know.


• The original definitions of acid, base and salt come from Arrhenius theory (named after Svante Arrhenius), which contains a few basic ideas about how acids and bases are related:
• An acid is any substance which, in water, produces hydrogen ions (H⁺).
• A base is any substance which, in water, produces hydroxide ions (OH^-).
• A salt is any substance which is the product of an acid-base reaction.
  Practically, this means any ionic compound that isn't an acid or base is a salt.


• These definitions were the first of their kind in describing the properties of acids and bases (Arrhenius did a lot for chemistry!) and are the reason why many chemistry teachers give the following (not perfect) simple rule to 'spot' acids and bases:
• Any ionic compound beginning with H is an acid.
• Any ionic compound with an OH group is a base.


• Today there are two theories or 'points of view’ to acid-base reactions. Brønsted-Lowry theory which is closely related to the original Arrhenius theory is much more common:
• Brønsted-Lowry acid-base theory is about protons.
• A Brønsted-Lowry acid is a proton (H⁺) donor such as HCl.
• A Brønsted-Lowry base is a proton acceptor.


• There is also Lewis theory is about electrons.
• A Lewis acid is an electron pair acceptor such as BH₃.
• A Lewis base is an electron pair donor.


• Looking at the Arrhenius and Brønsted-Lowry definitions of acids and bases, you will see they are about the hydrogen ion (H⁺). This is actually not what causes the properties acids are known for. The properties that ‘acids’ have are from the hydronium ion, H₃O⁺. This gets produced rapidly when protons mix with water.
• Hydrogen atoms have only one proton and one electron. Therefore when 'hydrogen ions (H⁺) are produced' by a substance (like hydrochloric acid, HCl) dissolving in water, what is actually released is just a single proton without the electron it used to have – quite literally just one tiny proton with nothing surrounding it!
• The charge density of a lone proton with nothing surrounding it is INCREDIBLY high – an individual proton of +1 charge is extremely small compared to a hydrogen atom with the same positive nucleus with the electron 'cloud' orbiting it. This makes H⁺ extremely reactive and it will interact with anything remotely negative nearby. Remember, this is all happening dissolved in water.
• This tiny ultra-concentrated proton immediately interacts with a negative lone pair on the oxygen atom of a nearby water molecule. In doing this, a hydronium ion, H₃O⁺, is formed. This process can be described with the equation:
H⁺_(aq) + H₂O _(l) → H₃O⁺
• The process can be explained in two ways:
• Acids release H⁺ ions in solution which 'protonate' water molecules.
• Acids release H⁺ in solution which are ‘hydrated’ (added to by water).


• The reverse process occurs when bases are added to solution. When bases (that produce OH^-) are dissolved, their strong negative charge means they react to neutralize H⁺ and H₃O⁺ species, which decreases their concentration in the solution. This can also deprotonate neutral water molecules. See the equation below: H₃O⁺ _(aq) + OH^- _(aq) → 2H₂O (l)
  • This process of decreasing hydronium ion concentration is what makes the pH rise. Eventually a lack of hydronium ions to neutralize the hydroxide ions means there will be free hydroxide (OH^-) ions in solution, which causes the basic properties chemists observe.
  
  
• We have seen above that water can act as a base in acidic conditions; it accepts H⁺ ions to form the hydronium ion. However, water also can act as an acid in basic conditions; it donates H⁺ to bases in solution, forming a hydroxide ion as a result.
  • Water’s ability to act as both acid and base makes it an amphoteric molecule.
  
  
• Acid-base theory is important for understanding how pH is measured. Remember, the definition of acidic is pH < 7, while basic is pH > 7.
  • The equation to find pH of a solution is: pH = -log[H₃O⁺]
    
  • This is an inverse logarithmic expression which means the following:
    • Inverse: as the concentration of hydronium ions increases, pH decreases.
    • Logarithmic: To change the pH by 1, you have to change [H₃O⁺] by a factor or ten. For example, a solution of pH 4 is ten times more concentrated with H₃O⁺ than a solution of pH 5.
  • This equation explains why acidic solutions have a low pH value, while basic solutions (where [H₃O⁺] is low) have a high pH value.