Animated Solution for Chemistry - Ionic Equilibrium: An aqueous solution of 1M NaCl and 1M HCl is

When we mix different chemicals in a beaker, predicting the behavior of the resulting solution requires a solid understanding of ionic equilibrium. In this problem, we are asked to evaluate a mixture of $1\text{M NaCl}$ and $1\text{M HCl}$. The core question is twofold: Is it a buffer, and what is its pH?

To answer the first part, we must recall the strict definition of a buffer solution. A buffer is a special type of solution that fiercely resists changes in its pH when small amounts of strong acid or strong base are added.

The golden rule of buffers: They must contain a weak component. Specifically, an acidic buffer requires a weak acid and its conjugate base (usually provided by a salt of that weak acid). A basic buffer requires a weak base and its conjugate acid.

Let's look at what we poured into our beaker:

1. Hydrochloric Acid ($\text{HCl}$): This is a classic strong acid. In water, it dissociates completely into ${\text{H}}^{+}$ and ${\text{Cl}}^{-}$ ions. There are virtually no intact $\text{HCl}$ molecules left. 2. Sodium Chloride ($\text{NaCl}$): This is a neutral salt. It is the product of a neutralization reaction between a strong acid ($\text{HCl}$) and a strong base ($\text{NaOH}$). It also dissociates completely into ${\text{Na}}^{+}$ and ${\text{Cl}}^{-}$ ions.

Because both components are strong electrolytes that dissociate completely, we lack the crucial 'weak' component. There is no reservoir of undissociated molecules to absorb any extra ${\text{H}}^{+}$ or ${\text{OH}}^{-}$ ions that might be added later. Therefore, this solution cannot act as a buffer.

Now, what about the acidity of the solution?

The presence of $1\text{M HCl}$ means we have a massive concentration of hydrogen ions ($[{\text{H}}^{+}]=1\text{M}$). The $\text{NaCl}$ does not undergo hydrolysis, meaning neither ${\text{Na}}^{+}$ nor ${\text{Cl}}^{-}$ reacts with water to alter the ${\text{H}}^{+}$ or ${\text{OH}}^{-}$ balance.

Since the solution is dominated by the strong acid, it is highly acidic. By definition, an acidic solution at standard temperature has a pH strictly less than 7. In fact, for a $1\text{M}$ solution of a strong monoprotic acid, the pH would theoretically be $-\log (1)=0$.

Final Conclusion: The mixture is not a buffer, and its pH is less than 7.