Category: 4. Acid-Base Reactions

A gas evolution reaction is a chemical process that produces a gas, such as oxygen or carbon dioxide.

A gas evolution reaction is a chemical process that produces a gas, such as oxygen or carbon dioxide. In the following examples, an acid reacts with a carbonate, producing salt, carbon dioxide, and water, respectively.

Nitric acid reacts with sodium carbonate to form sodium nitrate, carbon dioxide, and water:

2HNO3(aq)+Na2CO3(aq)→2NaNO3(aq)+CO2(g)+H2O(l)

Sulfuric acid reacts with calcium carbonate to form calcium sulfate, carbon dioxide, and water:

H2SO4(aq)+CaCO3(aq)→CaSO4(aq)+CO2(g)+H2O(l)

Hydrochloric acid reacts with calcium carbonate to form calcium chloride, carbon dioxide, and water:

2HCl(aq)+CaCO3(aq)→CaCl2(aq)+CO2(g)+H2O(l)

The following setup demonstrates this type of reaction:

Reaction of acids with carbonates: In this reaction setup, lime water is poured into one of the test tubes and sealed with a stopper. A small amount of hydrochloric acid is carefully poured into the remaining test tube. A small amount of sodium carbonate is added to the acid, and the tube is sealed with a rubber stopper. The two tubes are connected. As a result of the acid-carbonate reaction, carbon dioxide is produced and the lime water turns milky.

The test tube on the right contains limewater (a solution of calcium hydroxide, Ca(OH)₂). On the left, a solution of hydrochloric acid has been added to a solution of sodium carbonate to generate CO2(g)CO2(g). The test tubes are sealed with rubber stoppers and connected with a delivery tube. As the reaction proceeds, the limewater on the right turns from clear to milky; this is due to the CO2(g)CO2(g) reacting with the aqueous calcium hydroxide to form calcium carbonate, which is only slightly soluble in water.

1. Rinse the burette with the standard solution, the pipette with the unknown solution, and the conical flask with distilled water.
2. Place an accurately measured volume of the analyte into the Erlenmeyer flask using the pipette, along with a few drops of indicator. Place the standardized solution into the burette, and indicate its initial volume in a lab notebook. At this stage, we want a rough estimate of the amount of known solution necessary to neutralize the unknown solution. Let the solution out of the burette until the indicator changes color, and record the value on the burette. This is the first titration and it is not very precise; it should be excluded from any calculations.
3. Perform at least three more titrations, this time more accurately, taking into account where the end point will roughly occur. Record the initial and final readings on the burette, prior to starting the titration and at the end point, respectively. (Subtracting the initial volume from the final volume will yield the amount of titrant used to reach the endpoint.)
4. The end point is reached when the indicator permanently changes color.

The resulting solution at the equivalence point will have a pH dependent on the acid and base’s relative strengths. You can estimate the equivalence point’s pH using the following rules:

• A strong acid will react with a weak base to form an acidic (pH < 7) solution.
• A strong acid will react with a strong base to form a neutral (pH = 7) solution.
• A weak acid will react with a strong base to form a basic (pH > 7) solution.

When a weak acid reacts with a weak base, the equivalence point solution will be basic if the base is stronger and acidic if the acid is stronger; if both are of equal strength, then the equivalence pH will be neutral. Weak acids are not often titrated against weak bases, however, because the color change is brief and therefore very difficult to observe.

You can determine the pH of a weak acid solution being titrated with a strong base solution at various points; these fall into four different categories: (1) initial pH; (2) pH before the equivalence point; (3) pH at the equivalence point; and (4) pH after the equivalence point.

Titration of a weak acid by a strong base: The pH of a weak acid solution being titrated with a strong base solution can be found at each indicated point.

Acid-base titration setup: The pink color is caused by the phenolphthalein indicator.

Before you begin the titration, you must choose a suitable pH indicator, preferably one that will experience a color change (known as the “end point”) close to the reaction’s equivalence point; this is the point at which equivalent amounts of the reactants and products have reacted. Below are some common equivalence point indicators:

• strong acid – strong base titration: phenolphthalein indicator
• weak acid – weak base titration: bromthymol blue indicator
• strong acid-weak base titration: methyl orange indicator the base is off the scale (e.g., pH > 13.5) and the acid has pH > 5.5: alizarine yellow indicator
• the base is off the scale (e.g., pH > 13.5) and the acid has pH > 5.5: alizarine yellow indicator
• the base is off the scale (e.g., pH > 13.5) and the acid has pH > 5.5: alizarine yellow indicator
• the acid is off the scale (e.g., pH < 0.5) and the base has pH < 8.5: thymol blue indicator

Acid-base titration can determine the concentrations of unknown acid or base solutions.

Setting up an Acid-Base Titration

An acid-base titration is an experimental procedure used to determined the unknown concentration of an acid or base by precisely neutralizing it with an acid or base of known concentration. This lets us quantitatively analyze the concentration of the unknown solution. Acid-base titrations can also be used to quantify the purity of chemicals.

Acid-base titration: The solution in the flask contains an unknown number of equivalents of base (or acid). The burette is calibrated to show volume to the nearest 0.001 cm³. It is filled with a solution of strong acid (or base) of known concentration. Small increments are added from the burette until, at the end point, one drop changes the indicator color permanently. (An indication of the approaching equivalence point is that the indicator changes color but changes back after stirring.) At the equivalence point, the total amount of acid (or base) is recorded from the burette readings. The number of equivalents of acid and base must be equal at the equivalence point.

Alkalimetry, or alkimetry, is the specialized analytic use of acid-base titration to determine the concentration of a basic (alkaline) substance; acidimetry, or acidometry, is the same concept applied to an acidic substance.

Materials for a Titration Procedure

• burette
• white tile (used to see a color change in the solution)
• pipette
• pH indicator (the type depends on the reactants )
• Erlenmeyer or conical flask
• titrant (a standard solution of known concentration; a common example is aqueous sodium carbonate)
• analyte, or titrand (the solution of unknown concentration)

Water is amphoteric, which means it can act as either an acid or a base. In the reaction between acetic acid, CH₃CO₂H, and water, H₂O, water acts as a base. The acetate ion CH₃CO₂^– is the conjugate base of acetic acid, and the hydronium ion H₃O⁺ is the conjugate acid of the base, water:

CH₃COOH + H₂O ⇌ CH₃COO^– + H₃O⁺

Water can also act as an acid, as when it reacts with ammonia. The equation given for this reaction is:

H₂O + NH₃ ⇌ OH^– + NH₄⁺

Here, H₂O donates a proton to NH₃. The hydroxide ion is the conjugate base of water, which acts as an acid, and the ammonium ion is the conjugate acid of the base, ammonia.

The Brønsted-Lowry acid-base theory has several advantages over the Arrhenius theory. Recall that the Arrhenius theory defines an acid as any species that increases the concentration of H⁺/H₃O⁺ in solution. Consider the following reactions of acetic acid (CH₃COOH), the organic acid that gives vinegar its characteristic taste:

1. CH₃COOH + H₂O ⇌ CH₃COO^– + H₃O⁺
2. CH₃COOH + NH₃ ⇌ CH3COO^– + NH₄⁺

Both theories easily describe the first reaction: CH₃COOH acts as an Arrhenius acid because it acts as a source of H₃O⁺ when dissolved in water, and it acts as a Brønsted acid by donating a proton to water. In the second example CH₃COOH undergoes the same transformation, in this case donating a proton to ammonia (NH₃); this cannot be described using the Arrhenius definition of an acid, however, because the reaction does not produce H₃O⁺.

A Brønsted acid is any species capable of donating a proton; a Brønsted base is any capable of accepting a proton.

In chemistry, the Brønsted-Lowry theory is an acid-base theory, independently proposed by Johannes Nicolaus Brønsted and Thomas Martin Lowry in 1923. In this system, acids and bases are defined as follows: an acid is any species that is able to donate a hydrogen cation (H⁺, a proton); a base is any species with the ability to accept a hydrogen cation (H⁺). To that end, if a compound is to behave as an acid by donating a proton, there must be a base to accept that proton; the Brønsted-Lowry concept is therefore defined by the reaction:

acid + base ⇌ conjugate base + conjugate acid

The conjugate base is the ion or molecule that remains after the acid has donated its proton, and the conjugate acid is the species created after the base accepts the proton. The reaction can proceed either forward backward; in each case, the acid donates a proton to the base.

How can organisms whose bodies require a near-neutral pH ingest acidic and basic substances (a human drinking orange juice, for example) and survive? Buffers are the key. Buffers usually consist of a weak acid and its conjugate base; this enables them to readily absorb excess H⁺ or OH^–, keeping the system’s pH within a narrow range.

Maintaining a constant blood pH is critical to a person’s well-being. The buffer that maintains the pH of human blood involves carbonic acid (H₂CO₃), bicarbonate ion (HCO₃^–), and carbon dioxide (CO₂). When bicarbonate ions combine with free hydrogen ions and become carbonic acid, hydrogen ions are removed, moderating pH changes. Similarly, excess carbonic acid can be converted into carbon dioxide gas and exhaled through the lungs; this prevents too many free hydrogen ions from building up in the blood and dangerously reducing its pH; likewise, if too much OH^– is introduced into the system, carbonic acid will combine with it to create bicarbonate, lowering the pH. Without this buffer system, the body’s pH would fluctuate enough to jeopardize survival.

Buffers in the body: This diagram shows the body’s buffering of blood pH levels: the blue arrows show the process of raising pH as more CO2 is made; the purple arrows indicate the reverse process, lowering pH as more bicarbonate is created.

Antacids, which combat excess stomach acid, are another example of buffers. Many over-the-counter medications work similarly to blood buffers, often with at least one ion (usually carbonate) capable of absorbing hydrogen and moderating pH, bringing relief to those that suffer “heartburn” from stomach acid after eating.

The stronger the acid, the more readily it donates H⁺. For example, hydrochloric acid (HCl) is highly acidic and completely dissociates into hydrogen and chloride ions, whereas the acids in tomato juice or vinegar do not completely dissociate and are considered weak acids; conversely, strong bases readily donate OH^– and/or react with hydrogen ions. Sodium hydroxide (NaOH) and many household cleaners are highly basic and give up OH^– rapidly when placed in water; the OH^– ions react with H⁺ in solution, creating new water molecules and lowering the amount of free H⁺ in the system, thereby raising the overall pH. An example of a weak basic solution is seawater, which has a pH near 8.0, close enough to neutral that well-adapted marine organisms thrive in this alkaline environment.