The best examples of shifts in equilibrium with changes in concentration

Before we jump into the best examples of shifts in equilibrium with changes in concentration, let’s anchor the logic in one sentence:

If you disturb an equilibrium by changing a concentration, the system shifts in the direction that reduces that disturbance.

Mathematically, that’s the story of the reaction quotient \(Q\) racing to catch up with the equilibrium constant \(K\). When you add a reactant, \(Q < K\), so the reaction moves forward. When you add a product, \(Q > K\), so the reaction runs backward.

Now let’s look at real examples, from classroom classics to modern industrial and biological systems.

Gas‑phase reaction examples of shifts in equilibrium with changes in concentration

Gas reactions are the most familiar examples of shifts in equilibrium with changes in concentration, because the species are easy to measure and manipulate.

Example: Nitrogen dioxide and dinitrogen tetroxide color change

Consider the classic brown–colorless equilibrium:

\[
\text{N}_2\text{O}_4(g) \rightleftharpoons 2\text{NO}_2(g)
\]

• \(\text{N}_2\text{O}_4\) is colorless.
• \(\text{NO}_2\) is brown.

If you inject extra \(\text{NO}_2\) into an equilibrium mixture, you immediately increase the product concentration. That makes \(Q > K\), so the reaction shifts **left**, forming more \(\text{N}_2\text{O}_4\) and lightening the color over time.

If instead you remove \(\text{NO}_2\) (for example, by quickly cooling a portion and condensing some out), you decrease the product concentration. Now \(Q < K\), and the system shifts **right**, generating more \(\text{NO}_2\) and deepening the brown color.

This is one of the simplest real examples of how visible properties (color) track shifts in equilibrium with concentration changes.

Example: The Haber process and ammonia production

The Haber–Bosch process is the workhorse for synthetic fertilizer:

\[
\text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g)
\]

In modern plants (still being optimized as of 2024 to cut emissions), engineers constantly tweak feed composition to keep the reaction mixture near optimal equilibrium conditions.

• Adding more \(\text{N}_2\) or \(\text{H}_2\) increases reactant concentration, making \(Q < K\) and pushing the equilibrium right, toward more ammonia.
• Removing \(\text{NH}_3\) as it forms (for example, by cooling and liquefying it) lowers product concentration, again making \(Q < K\) and driving the reaction forward.

Industrial control systems monitor gas compositions and use these examples of shifts in equilibrium with changes in concentration to decide when to recycle unreacted gases, when to purge, and how to maintain high ammonia yield while minimizing energy use.

For a broader view of industrial nitrogen use and environmental impacts, the U.S. Geological Survey and EPA provide periodic updates on fertilizer production and emissions trends (USGS.gov and EPA.gov).

Acid–base examples of shifts in equilibrium with changes in concentration

Acid–base equilibria give some of the best examples for students, because you can literally watch the pH respond as you add acid or base.

Example: Weak acid equilibrium – acetic acid and acetate

Take acetic acid in water:

\[
\text{CH}_3\text{COOH}(aq) + \text{H}_2\text{O}(l) \rightleftharpoons \text{H}_3\text{O}^+(aq) + \text{CH}_3\text{COO}^-(aq)
\]

Now imagine a buffer solution containing both acetic acid and sodium acetate.

• If you add strong acid (say \(\text{HCl}\)), you increase \(\text{H}_3\text{O}^+\). That makes \(Q > K_a\), and the equilibrium shifts left, converting \(\text{CH}_3\text{COO}^-\) into \(\text{CH}_3\text{COOH}\). The pH drops only slightly.
• If you add strong base (like \(\text{NaOH}\)), you remove \(\text{H}_3\text{O}^+\) by forming water. Now \(Q < K_a\), and the equilibrium shifts right, creating more \(\text{CH}_3\text{COO}^-\) and restoring \(\text{H}_3\text{O}^+\) levels.

This is a textbook example of shifts in equilibrium with changes in concentration that underpins how biological and industrial buffers work.

The chemistry of buffers and acid–base equilibria is summarized well in open educational resources from institutions like MIT and UC Davis (MIT OpenCourseWare, UC Davis ChemWiki at LibreTexts.org).

Example: Blood buffer system – carbonic acid and bicarbonate

Now for one of the most important real examples: your own blood. The primary buffer system is often written as:

\[
\text{CO}_2(aq) + \text{H}_2\text{O}(l) \rightleftharpoons \text{H}_2\text{CO}_3(aq) \rightleftharpoons \text{H}^+(aq) + \text{HCO}_3^-(aq)
\]

and in the lungs:

\[
\text{CO}_2(g) \rightleftharpoons \text{CO}_2(aq)
\]

Here’s how changes in concentration shift this equilibrium:

• When you hold your breath, \(\text{CO}_2\) builds up in your blood. Dissolved \(\text{CO}_2\) concentration rises, driving \(Q < K\) for the forward reaction and shifting the equilibrium *right*. More \(\text{H}^+\) is produced, and blood pH drops (respiratory acidosis).
• When you hyperventilate, you blow off \(\text{CO}_2\). Dissolved \(\text{CO}_2\) falls, making \(Q > K\) and shifting the equilibrium *left*, consuming \(\text{H}^+\) and raising blood pH (respiratory alkalosis).

Clinicians interpret blood gas measurements using these examples of shifts in equilibrium with changes in concentration to diagnose and manage acid–base disorders. For an accessible medical discussion, see resources from the National Institutes of Health and MedlinePlus (NIH.gov, MedlinePlus.gov).

Solubility and precipitation: examples include salts and complex ions

Shifts in solubility equilibria offer some of the best examples of how concentration changes can start or stop precipitation.

Example: Silver chloride solubility and common ion effect

Consider the sparingly soluble salt:

\[
\text{AgCl}(s) \rightleftharpoons \text{Ag}^+(aq) + \text{Cl}^-(aq)
\]

The solubility product is:

\[
K_{sp} = [\text{Ag}^+][\text{Cl}^-]
\]

If you add solid NaCl to a saturated \(\text{AgCl}\) solution, you increase \(\text{Cl}^-\) concentration from a completely different source.

• Immediately, \(Q = [\text{Ag}^+][\text{Cl}^-]\) becomes greater than \(K_{sp}\).
• The system responds by precipitating more \(\text{AgCl}\), reducing both \(\text{Ag}^+\) and \(\text{Cl}^-\) until \(Q = K_{sp}\) again.

This common ion effect is a clear example of a shift in equilibrium with a concentration change that you can see as cloudiness or a solid forming.

Example: Complex ion formation – dissolving silver chloride with ammonia

Now flip the previous example of \(\text{AgCl}\) on its head. Add aqueous ammonia:

\[
\text{Ag}^+(aq) + 2\text{NH}_3(aq) \rightleftharpoons [\text{Ag(NH}_3)_2]^+(aq)
\]

As \(\text{NH}_3\) binds \(\text{Ag}^+\), the free \(\text{Ag}^+\) concentration drops.

• That makes \(Q = [\text{Ag}^+][\text{Cl}^-]\) less than \(K_{sp}\) for \(\text{AgCl}\).
• The equilibrium for \(\text{AgCl}(s) \rightleftharpoons \text{Ag}^+ + \text{Cl}^-\) shifts right, dissolving more solid \(\text{AgCl}\).

So by changing the concentration of one species (here, \(\text{NH}_3\), which indirectly lowers \(\text{Ag}^+\)), you force the dissolution equilibrium to shift. This is a favorite example of shifts in equilibrium with changes in concentration in qualitative analysis labs.

Redox and industrial examples of shifts in equilibrium with changes in concentration

Redox systems and industrial reactors provide real examples that go beyond the beaker.

Example: Dichromate–chromate color shift

The equilibrium between dichromate and chromate ions in water is pH‑dependent:

\[
2\text{CrO}_4^{2-}(aq) + 2\text{H}^+(aq) \rightleftharpoons \text{Cr}_2\text{O}_7^{2-}(aq) + \text{H}_2\text{O}(l)
\]

• \(\text{CrO}_4^{2-}\) is yellow.
• \(\text{Cr}_2\text{O}_7^{2-}\) is orange.

If you add acid, \(\text{H}^+\) concentration increases, so \(Q < K\) and the equilibrium shifts right. The solution turns more orange.

If you add base, you reduce \(\text{H}^+\) (effectively removing it), making \(Q > K\) and shifting the equilibrium left, turning the solution more yellow.

This is a visually striking example of how concentration changes drive equilibrium shifts in inorganic systems.

Example: Industrial esterification – making ethyl acetate

In many chemical plants, esterification reactions are tuned by manipulating reactant and product concentrations. Take the formation of ethyl acetate:

\[
\text{CH}_3\text{COOH}(l) + \text{C}_2\text{H}_5\text{OH}(l) \rightleftharpoons \text{CH}_3\text{COOC}_2\text{H}_5(l) + \text{H}_2\text{O}(l)
\]

To increase ester yield, operators use several tactics based on examples of shifts in equilibrium with changes in concentration:

• Use an excess of ethanol or acetic acid to keep \(Q < K\) and push the equilibrium right.
• Continuously remove water (for example, by azeotropic distillation) to keep the product side concentrations low, again maintaining \(Q < K\).

These strategies are standard in process design and are still being refined in 2024–2025 as green chemistry pushes for lower‑waste, lower‑energy processes.

Biological and environmental real examples of equilibrium shifts

Modern chemistry education increasingly leans on real examples from biology and the environment to show that these are not just classroom abstractions.

Example: Oxygen binding to hemoglobin

Oxygen transport in blood can be simplified as:

\[
\text{Hb}(aq) + \text{O}_2(g) \rightleftharpoons \text{HbO}_2(aq)
\]

At the lungs:

• \(\text{O}_2\) partial pressure is high, so dissolved \(\text{O}_2\) concentration is high.
• That makes \(Q < K\) and shifts the equilibrium right, forming more oxyhemoglobin \(\text{HbO}_2\).

In tissues where \(\text{O}_2\) is being consumed:

• Dissolved \(\text{O}_2\) concentration falls.
• This drives \(Q > K\), shifting the equilibrium left, and \(\text{O}_2\) is released from hemoglobin.

This is a beautifully efficient example of shifts in equilibrium with changes in concentration that underlies aerobic life. For deeper reading, see physiology resources from institutions such as Harvard and the National Library of Medicine (Harvard.edu, NCBI.NLM.NIH.gov).

Example: Ocean carbonate system and ocean acidification

The ocean’s carbonate system is a large‑scale version of acid–base equilibrium:

\[
\text{CO}_2(g) \rightleftharpoons \text{CO}_2(aq) + \text{H}_2\text{O}(l) \rightleftharpoons \text{H}_2\text{CO}_3(aq) \rightleftharpoons \text{H}^+(aq) + \text{HCO}_3^-(aq) \rightleftharpoons 2\text{H}^+(aq) + \text{CO}_3^{2-}(aq)
\]

As atmospheric \(\text{CO}_2\) levels rise (a well‑documented trend through 2024), more \(\text{CO}_2\) dissolves in seawater. This increases the concentration of dissolved \(\text{CO}_2\), pushing \(Q < K\) for the forward reactions and shifting the equilibria right, producing more \(\text{H}^+\) and lowering pH.

That shift also reduces \(\text{CO}_3^{2-}\) concentration, which is needed by marine organisms to build calcium carbonate shells. This is a real example of equilibrium shifts with concentration changes that now shows up in climate reports and marine chemistry research.

Authoritative summaries are available from agencies like NOAA and the IPCC, often hosted or linked through .gov and .org sites (NOAA.gov).

Pulling it together: how to think about any new example

By now, you’ve seen many examples of examples of shifts in equilibrium with changes in concentration:

• Gas‑phase reactions (\(\text{N}_2\text{O}_4/\text{NO}_2\), Haber process)
• Acid–base and buffer systems (acetic acid/acetate, blood \(\text{CO}_2\)/bicarbonate)
• Solubility and complex ion equilibria (\(\text{AgCl}\), \([\text{Ag(NH}_3)_2]^+\))
• Inorganic and redox systems (dichromate/chromate)
• Industrial organic synthesis (esterification)
• Biological and environmental systems (hemoglobin, ocean carbonate)

To analyze any new system, the workflow is the same:

1. Write the balanced equilibrium expression and the expression for \(Q\) and \(K\).
2. Ask: does the concentration change make \(Q\) larger or smaller than \(K\)?
3. Predict the direction of the shift so that \(Q\) moves back toward \(K\).

Once you train yourself on these real examples of shifts in equilibrium with changes in concentration, unfamiliar systems start to feel routine. You’re not guessing; you’re reading the direction of the inequality between \(Q\) and \(K\).

FAQ: examples of concentration‑driven equilibrium shifts

Q: Can you give a simple classroom example of an equilibrium shift when a product is added?
A: The \(\text{N}_2\text{O}_4/\text{NO}_2\) system is a classic example of this. Add more \(\text{NO}_2\) to an equilibrium mixture, and the reaction shifts left to form more \(\text{N}_2\text{O}_4\), lightening the brown color over time.

Q: What is a good example of concentration affecting a biological equilibrium?
A: The blood \(\text{CO}_2\)–bicarbonate buffer is one of the best examples. Changes in dissolved \(\text{CO}_2\) concentration during breathing directly shift the equilibrium that generates or consumes \(\text{H}^+\), which in turn changes blood pH.

Q: Do changes in concentration always change the value of K?
A: No. The equilibrium constant \(K\) depends only on temperature (and, for gases, sometimes pressure conventions). When you change a concentration, you change \(Q\), not \(K\). The system shifts so that \(Q\) returns to the same \(K\) at that temperature.

Q: Are there real industrial examples of using concentration to control equilibrium?
A: Yes. The Haber process and industrial esterification are prominent examples of shifts in equilibrium with changes in concentration. Plants feed reactants in excess or remove products as they form specifically to keep \(Q < K\) and drive the desired reaction direction.

Q: How can I quickly decide which way an equilibrium will shift in a new problem?
A: After any concentration change, mentally recompute \(Q\). If \(Q < K\), the system shifts right (toward products). If \(Q > K\), it shifts left (toward reactants). Practicing with the examples of examples of shifts in equilibrium with changes in concentration in this article makes that process much faster.