Acids and Bases Introduction

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Introduction

Before the 1800s, acids and bases were simply categorized by their characteristics:

  • Acids are corrosive to metals, cause litmus paper to change red, taste sour, and become less acidic when mixed with bases.
  • Bases changes litmus blue, feel slippery, and become less basic when mixed with acids.

Arrhenius, Brønsted-Lowry, and Lewis Definitions of Acids/Bases

In the 1800s, Svante Arrhenius proposed that many compounds could dissociate in water, and that acids and bases were the result of either hydrogen ions or hydroxide ions, respectively, being released from these compounds. This theory explained both the common characteristics of acids and bases and why adding an acid to a bases, reduced the effectiveness of the base, and vice versa (neutralization).

Johannes Brønsted later expanded this theory to help explain why some substances, which do not have a free hydroxide, can still act as a base. The Brønsted-Lowry definition of an acid is similar to the Arrhenius version, an acid is a hydrogen ion (proton) donor. The Brønsted-Lowry definition of a base on the other hand, is that bases act as a proton acceptor (the opposite of an acid).

Gilbert Lewis finalized the definition of an acid by generalizing it to the electron chemistry of substances, not just the previous specific definition involving proton acceptor/donors. A Lewis acid is a substance that acts as an electron acceptor, while a base is one that acts as an electron donator.

One can thus consider an acid as a substance which is electron deficient and is trying to aquire more electrons (usually a pair) while a base is electron rich and is willing to release them.

Lewis acids can fit into several categories: 1. positive ions 2. having less than a full octet in the valence shell 3. polar double bonds (one end) 4. expandable valence shells

Lewis bases can fit into several categories: 1. negative ions 2. one of more unshared pairs in the valence shell 3. polar double bonds (the other end) 4. the presence of a double bond


1. Which of the following is not a Brønsted-Lowry acid/base pair?

CH3COOH / NH3
→ This is acetic acid and ammonia, which are both BL.
H20 / H2PO4-
→ Water can act as a BL acid or base, and H2PO4- is a base here.
ZnCl2 / H20
→ ZnCl2 is a Lewis acid but it does not fit the description of a BL acid.
H3AsO4 / NaOH
→ H3AsO4 is a polyprotic acid and NaOH is a strong base.

Your score is 0 / 0

Ka

When an acid, HA, dissolves in water, some molecules of the acid 'dissociate' to form hydronium ions (H+) and the conjugate base, (A-), of the acid.

H A \rightleftharpoons H^{+} + A^{-}

It is understood that H+ stands for the hydronium ion and that the equilibrium position will be based on how well the acid dissociates, as each will solvate to a different extent. If we express the Keq of this equilibrium we will be able to represent this dissociation constant. We refer to this Keq as the acid dissociation constant, and give it the name Ka just as the Keq for solubility equilibriums is given the name Ksp.

Ka is then defined as

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

The larger the Ka the more the equilibrium is shifted to the right. This shift implies that there are more hydronium ions released and thus the solution is more acidic. Thus by comparing Ka values we can easily determine which acid is stronger.

Kb

In the same sense that an acid can dissociate to form hydronium ions and create an equilibrium, so can a base. When a base, B, is placed in water it can combine with H20 to form HB+ and hydroxide (OH-).

B + H_{2}O \rightleftharpoons HB^{+} + OH^{-}
B^{+} \rightleftharpoons HB + OH^{-}

The above two expressions represent the same process, however it is often the case that the base is either neutral or positively charged. (Both are exactly the same but the later is negatively charged to begin with). If again, we write the Keq, we will be representing the equilibrium constant, but for a base this time, and so we call it Kb. From the two equilibriums above it follows that Kb is simply:


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

Remember, just like any other Keq it is simply products/reactants and we do not include the H20 because it is a pure liquid.



1. Based on their Ka values below, which acid would be the weakest at 25C?

1.40 x 10-3
4.00 x 10-10
1.40 x 10-5
2.30 x 10-11
→ The acid dissociation constant, or Ka, tells you where the equilibrium is. The larger the value, the more products exist, which also implies more hydronium ions. Thus the larger the Ka the stronger the acid. Of these, 1.40 x 10-3 is the strongest and 2.30 x 10-11 would be the weakest.

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Strong Acids and Bases

Strong acids are acids which, when dissolved in water, dissociate completely. The substance exists entirely as ions in solution and no equilibrium with its original form is achieved. HCl, HNO3, H2SO4, HBr, HI, and HClO4 are all Strong acids.

Because an equilibrium does not exist (as it is all to the product side and NEVER to the reactant side) Ka is meaningless and often omitted, undefined, or stated to be infinity (because 100 products/ 0 reactants is an infinitely big fraction!)

The same principles are true for strong bases. Some examples of strong bases are NaOH, KOH, CsOH, and Ca(OH)2 and all of these have undefined Kb values.


When a problem arises which gives a certain number of moles of a strong acid (or base) it is sufficient to assume that it completely dissociates and that an equivalent number of moles of H+ are released.



1. Given their Kb values below, which most likely represents a strong base?

2.40 x 1029
1.22 x 104
6.25 x 10-5
8.5 x 10-79
→ The base dissociation constant, or Kb, tells you where the equilibrium is. The larger the value, the more products exist, which also implies more hydroxide ions. Thus the larger the Kb the stronger the base. Strong bases have kb values that are extremely large, often so large they are undefined or excluded, thus of these, 2.40 x 1029 is the probable strong base, while 8.5 x 10-79 is so weak it would most likely not act like a base at all.

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next

Check out The pH Scale, for the next section to this topic.

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