Introduction to acid-base theory - Acid-Base Theory

Introduction to acid-base theory

Lessons

Notes:

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.

Notes:

  • 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.
  • Intro Lesson
    What are acids and bases?
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Introduction to acid-base theory

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