Question

Write two chemical equations, one that shows \(\mathrm{H}_{2} \mathrm{O}\) reacting (with HBr) as a Bronsted base and a second that shows \(\mathrm{H}_{2} \mathrm{O}\) reacting (with \(\mathrm{NH}_{3}\) ) as a Bronsted acid.

Short answer

Water acts as a base with HBr and as an acid with NH3.

Step-by-step solution

Understanding Bronsted-Lowry Theory

The Bronsted-Lowry theory defines acids as proton (H+) donors and bases as proton (H+) acceptors. Water, therefore, can act either as an acid or a base depending on the substance it reacts with.

Identifying Water as a Base

When water \(\mathrm{H}_{2}\mathrm{O}\) acts as a base, it accepts a proton from an acid. In the case of hydrobromic acid (HBr), \(\mathrm{H}_{2}\mathrm{O}\) accepts a proton from HBr, forming hydronium ion \(\mathrm{H}_{3}\mathrm{O}^{+}\) and bromide ion \(\mathrm{Br}^{-}\): \\[ \\mathrm{H}_{2}\mathrm{O} + \mathrm{HBr} \rightarrow \mathrm{H}_{3}\mathrm{O}^{+} + \mathrm{Br}^{-} \]

Identifying Water as an Acid

When water \(\mathrm{H}_{2}\mathrm{O}\) acts as an acid, it donates a proton to a base. In the case with ammonia (NH3), \(\mathrm{H}_{2}\mathrm{O}\) donates a proton to NH3, forming hydroxide ion \(\mathrm{OH}^{-}\) and ammonium ion \(\mathrm{NH}_{4}^{+}\): \\[ \\mathrm{H}_{2}\mathrm{O} + \mathrm{NH}_{3} \rightarrow \mathrm{OH}^{-} + \mathrm{NH}_{4}^{+} \]

Review of Reactions

Based on the equations provided, water can be seen acting as both a Bronsted base and a Bronsted acid. With HBr, it accepts a proton, and with NH3, it donates a proton.

Key concepts

Water as Bronsted Base

Water, a versatile substance in chemistry, has the ability to act as either an acid or a base, thanks to the Bronsted-Lowry theory. As a Bronsted base, water accepts protons (H extsuperscript{+}) from acids. This attribute is showcased when water interacts with hydrobromic acid (HBr).

In this reaction, water (\(\mathrm{H}_{2}\mathrm{O}\)) acts as a base by accepting a proton from the acid \(\mathrm{HBr}\). This leads to the formation of a hydronium ion \(\mathrm{H}_{3}\mathrm{O}^{+}\) and a bromide ion \(\mathrm{Br}^{-}\).

The chemical equation for this reaction is:

\[ \mathrm{H}_{2}\mathrm{O} + \mathrm{HBr} \rightarrow \mathrm{H}_{3}\mathrm{O}^{+} + \mathrm{Br}^{-} \]

Key points to remember:

• The base (water) must have lone pair electrons to accept the proton.
• The acid (HBr) donates the proton to the base.
• This results in a conjugate acid-base pair: \(\mathrm{H}_{3}\mathrm{O}^{+}\) (conjugate acid of water) and \(\mathrm{Br}^{-}\) (conjugate base of HBr).

Water as Bronsted Acid

Water can also act as a Bronsted acid, which means it can donate a proton (H extsuperscript{+}) to a base. In the context of its reaction with ammonia \(\mathrm{NH}_{3}\), water exhibits its acidic property.

When water meets ammonia, it donates a proton to \(\mathrm{NH}_{3}\), transforming water into hydroxide \(\mathrm{OH}^{-}\) and ammonia into the ammonium ion \(\mathrm{NH}_{4}^{+}\).

Here's the chemical equation for clarity:

\[ \mathrm{H}_{2}\mathrm{O} + \mathrm{NH}_{3} \rightarrow \mathrm{OH}^{-} + \mathrm{NH}_{4}^{+} \]

Important highlights:

• Ammonia acts as the base by accepting the proton.
• Water, in donating a proton, becomes the conjugate base, \(\mathrm{OH}^{-}\).
• The product \(\mathrm{NH}_{4}^{+}\) is the conjugate acid of ammonia.

Chemical Equations

Chemical equations are crucial in understanding and illustrating chemical reactions. They tell us exactly what substances are involved and how they transform during the reaction. When depicting reactions such as water acting as a Bronsted base or acid, these equations become incredibly informative.

A balanced chemical equation ensures that the number of atoms for each element is the same on both sides of the equation, preserving mass and charge. Let's explore the previously discussed reactions involving water:

For water acting as a base:

\[ \mathrm{H}_{2}\mathrm{O} + \mathrm{HBr} \rightarrow \mathrm{H}_{3}\mathrm{O}^{+} + \mathrm{Br}^{-} \]

And for water acting as an acid:

\[ \mathrm{H}_{2}\mathrm{O} + \mathrm{NH}_{3} \rightarrow \mathrm{OH}^{-} + \mathrm{NH}_{4}^{+} \]

• Each equation illustrates the transfer of protons, a central theme in Bronsted-Lowry theory.
• Visualizing these equations helps us understand the role of each substance: proton donor or acceptor.
• Practicing writing and balancing chemical equations is key to mastering chemical reactions.

Understanding how to write and balance chemical equations is essential to grasp the dynamic nature of chemical reactions and correctly predict the outcome of reactions!