General Chemistry/Properties and Theories of Acids and Bases
Acid-Base Reaction Theories[edit | edit source]
Acids and bases are everywhere. Some foods contain acid, like the citric acid in lemons and the lactic acid in dairy. Cleaning products like bleach and ammonia are bases. Chemicals that are acidic or basic are an important part of chemistry.
|You may need to refresh your memory on naming acids.|
Several different theories explain what composes an acid and a base. The first scientific definition of an acid was proposed by the French chemist Antoine Lavoisier in the eighteenth century. He proposed that acids contained oxygen, although he did not know the dual composition of acids such as hydrochloric acid (HCl). Over the years, much more accurate definitions of acids and bases have been created.
Arrhenius Theory[edit | edit source]
The Swedish chemist Svante Arrhenius published his theory of acids and bases in 1887. It can be simply explained by these two points:
Based on this definition, you can see that Arrhenius acids must be soluble in water. Arrhenius acid-base reactions can be summarized with three generic equations:
|An acid will dissociate in water producing hydrogen ions.|
|A base (usually containing a metal) will dissociate in water to produce hydroxide ions.|
|Acids and bases will neutralize each other when mixed. They produce water and an ionic salt, neither of which are acidic or basic.|
The Arrhenius theory is simple and useful. It explains many properties and reactions of acids and bases. For instance, mixing hydrochloric acid (HCl) with sodium hydroxide (NaOH) results in a neutral solution containing table salt (NaCl).
However, the Arrhenius theory is not without flaws. There are many well known bases, such as ammonia (NH3) that do not contain the hydroxide ion. Furthermore, acid-base reactions are observed in solutions that do not contain water. To resolve these problems, there is a more advanced acid-base theory.
Brønsted-Lowry Theory[edit | edit source]
The Brønsted-Lowry theory was proposed in 1923. It is more general than the Arrhenius theory—all Arrhenius acids/bases are also Brønsted-Lowry acids/bases (but not necessarily vice versa).
Acids that can donate only one proton are monoprotic, and acids that can donate more than one proton are polyprotic.
These reactions demonstrate the behavior of Brønsted-Lowry acids and bases:
|An acid (in this case, hydrochloric acid) will donate a proton to a base (in this case, water is the base). The acid loses its proton and the base gains it.|
|Water is not necessary. In this case, hydrochloric acid is still the acid, but ammonia acts as the base.|
|The same reaction is happening, but now in reverse. What was once an acid is now a base (HCl → Cl-) and what was once a base is now an acid (NH3 → NH4+). This concept is called conjugates, and it will be explained in more detail later.|
|Two examples of acids (HCl and H3O+) mixing with bases (NaOH and OH-) to form neutral substances (NaCl and H2O).|
|A base (sodium hydroxide) will accept a proton from an acid (ammonia). A neutral substance is produced (water), which is not necessarily a part of every reaction. Compare this reaction to the second one. Ammonia was a base, and now it is an acid. This concept, called amphoterism, is explained later.|
The Brønsted-Lowry theory is by far the most useful and commonly-used definition. For the remainder of General Chemistry, you can assume that any acids/bases use the Brønsted-Lowry definition, unless stated otherwise.
Lewis Theory[edit | edit source]
The Lewis definition is the most general theory, having no requirements for solubility or protons.
Lewis acids and bases react to create an adduct, a compound in which the acid and base have bonded by sharing the electron pair. Lewis acid/base reactions are different from redox reactions because there is no change in oxidation state.
Amphoterism and Water[edit | edit source]
Substances capable of acting as either an acid or a base are amphoteric. Water is the most important amphoteric substance. It can ionize into hydroxide (OH-, a base) or hydronium (H3O+, an acid). By doing so, water is
- Increasing the H+ or OH- concentration (Arrhenius),
- Donating or accepting a proton (Brønsted-Lowry), and
- Accepting or donating an electron pair (Lewis).
|A bare proton (H+ ion) cannot exist in water. It will form a hydrogen bond to the nearest water molecule, creating the hydronium ion (H3O+). Although many equations and definitions may refer to the "concentration of H+ ions", that is a misleading abbreviation. Technically, there are no H+ ions, only hydronium (H3O+) ions. Fortunately, the number of hydronium ions formed is exactly equal to the number of hydrogen ions, so the two can be used interchangeably.|
Water will dissociate very slightly (which further explains its amphoteric properties).
|The presence of hydrogen ions indicates an acid, whereas the presence of hydroxide ions indicates a base. Being neutral, water dissociates into both equally.|
|This equation is more accurate—hydrogen ions do not exist in water because they bond to form hydronium.|
Ammonia[edit | edit source]
Another common example of an amphoteric substance is ammonia. Ammonia is normally a base, but in some reactions it can act like an acid.
|Ammonia acts as a base. It accepts a proton to form ammonium.|
|Ammonia also acts as an acid. Here, it donates a proton to form amide.|
Ammonia's amphoteric properties are not often seen because ammonia typically acts like a base. Water, on the other hand, is completely neutral, so its acid and base behaviors are both observed commonly.
Conjugate Acids and Bases[edit | edit source]
In all the theories, the products of an acid-base reaction are related to the initial reactants of the reaction. For example, in the Brønsted-Lowry theory, this relationship is the difference of a proton between a reactant and product. Two substances which exhibit this relationship form a conjugate acid-base pair.
|Hydroiodic acid reacts with water (which serves as a base). The conjugate base is the iodide ion and the conjugate acid is the hydronium ion. The acids are written in red, and the bases are written in blue. One conjugate pair is written bold and the other conjugate pair is in cursive.|
|Ammonia (basic) reacts with water (the acid). The conjugate acid is ammonium and the conjugate base is hydroxide. Again, acids are written in red, and the bases are written in blue. The conjugate pairs are distinguished with matching fonts.|
Strong and Weak Acids/Bases[edit | edit source]
A strong acid is an acid which dissociates completely in water. That is, all the acid molecules break up into ions and solvate (attach) to water molecules. Therefore, the concentration of hydronium ions in a strong acid solution is equal to the concentration of the acid.
The majority of acids exist as weak acids, an acid which dissociates only partially. On average, only about 1% of a weak acid solution dissociates in water in a 0.1 mol/L solution. Therefore, the concentration of hydronium ions in a weak acid solution is always less than the concentration of the dissolved acid.
Strong bases and weak bases do not require additional explanation; the concept is the same.
|The conjugate of a strong acid/base is very weak. The conjugate of a weak acid/base is not necessarily strong.|
This explains why, in all of the above example reactions, the reverse chemical reaction does not occur. The stronger acid/base will prevail, and the weaker one will not contribute to the overall acidity/basicity. For example, hydrochloric acid is strong, and upon dissociation chloride ions are formed. Chloride ions are a weak base, but the solution is not basic because the acidity of HCl is overwhelmingly stronger than basicity of Cl-.
|Although the other halogens make strong acids, hydrofluoric acid (HF) is a weak acid. Despite being weak, it is incredibly corrosive—hydrofluoric acid dissolves glass and metal!|
Most acids and bases are weak. You should be familiar with the most common strong acids and assume that any other acids are weak.
|HCl, HBr, HI||Hydrohalic acids|
Within a series of oxyacids, the ions with the greatest number of oxygen molecules are the strongest. For example, nitric acid (HNO3) is strong, but nitrous acid (HNO2) is weak. Perchloric acid (HClO4) is stronger than chloric acid (HClO3), which is stronger than the weak chlorous acid (HClO2). Hypochlorous acid (HClO) is the weakest of the four.
Common strong bases are the hydroxides of Group 1 and most Group 2 metals. For example, potassium hydroxide and calcium hydroxide are some of the strongest bases. You can assume that any other bases (including ammonia and ammonium hydroxide) are weak.
|Acids and bases that are strong are not necessarily concentrated, and weak acids/bases are not necessarily dilute. Concentration has nothing to do with the ability of a substance to dissociate. Furthermore, polyprotic acids are not necessarily stronger than monoprotic acids.|
Properties of Acids and Bases[edit | edit source]
Now that you are aware of the acid-base theories, you can learn about the physical and chemical properties of acids and bases. Acids and bases have very different properties, allowing them to be distinguished by observation.
Indicators[edit | edit source]
Made with special chemical compounds that react slightly with an acid or base, indicators will change color in the presence of an acid or base. A common indicator is litmus paper. Litmus paper turns red in acidic conditions and blue in basic conditions. Phenolphthalein purple is colorless in acidic and neutral solutions, but it turns purple once the solution becomes basic. It is useful when attempting to neutralize an acidic solution; once the indicator turns purple, enough base has been added.
Conductivity[edit | edit source]
A less informative method is to test for conductivity. Acids and bases in aqueous solutions will conduct electricity because they contain dissolved ions. Therefore, acids and bases are electrolytes. Strong acids and bases will be strong electrolytes. Weak acids and bases will be weak electrolytes. This affects the amount of conductivity.
However, acids will react with metal, so testing conductivity may not be plausible.
Physical properties[edit | edit source]
|The following is for informative purposes only. Do not sniff, touch, or taste any acids or bases as they may result in injury or death.|
The physical properties of acids and bases are opposites.
These properties are very general; they may not be true for every single acid or base.
Another warning: if an acid or base is spilled, it must be cleaned up immediately and properly (according to the procedures of the lab you are working in). If, for example, sodium hydroxide is spilled, the water will begin to evaporate. Sodium hydroxide does not evaporate, so the concentration of the base steadily increases until it becomes damaging to its surrounding surfaces.
Chemical Reactions[edit | edit source]
Neutralization[edit | edit source]
Acids will react with bases to form a salt and water. This is a neutralization reaction. The products of a neutralization reaction are much less acidic or basic than the reactants were. For example, sodium hydroxide (a base) is added to hydrochloric acid.
This is a double replacement reaction.
Acids[edit | edit source]
|Acids react with metal to produce a metal salt and hydrogen gas bubbles.|
|Acids react with metal carbonates to produce water, CO2 gas bubbles, and a salt.|
|Acids react with metal oxides to produce water and a salt.|
Bases[edit | edit source]
Bases are typically less reactive and violent than acids. They do still undergo many chemical reactions, especially with organic compounds. A common reaction is saponification: the reaction of a base with fat or oil to create soap.
Practice Questions[edit | edit source]
1. Name the following compounds that will form, and identify as an acid or base:
- a) Br + H
- b) 2H + SO3
- c) K + H
- d) 2H + SO6
- e) 3H + P2
- f) H + BrO2
- g) Na + Cl
2. What are the conjugate acids and bases of the following:
- a) water
- b) ammonia
- c) bisulfate ion
- d) zinc hydroxide
- e) hydrobromic acid
- f) nitrite ion
- g) dihydrogen phosphate ion
3. In a conductivity test, 5 different solutions were set up with light bulbs. The following observations were recorded:
- Solution A glowed brightly.
- Solution B glowed dimly.
- Solution C glowed dimly.
- Solution D did not glow.
- Solution E glowed brightly.
- a) Which solution(s) could contain strong bases?
- b) Which solution(s) could contain weak acids?
- c) Which solution(s) could contain ions?
- d) Which solution(s) could contain pure water?
- e) Based solely on these observations, would it be possible to distinguish between acidic and basic solutions?
4. Identity the conjugate base and conjugate acid in these following equations:
- a) HCl + H2O → H3O+ + Cl-
- b) HClO + H2O → ClO- + H3O+
- c) CH3CH2NH2 + H2O → CH3CH2NH3+ + OH-
5. Identify these bases as Arrhenius, Brønsted-Lowry, or both.
- a) strontium hydroxide
- b) butyllithium (C4H9Li)
- c) ammonia
- d) potassium hydroxide
- e) potassium iodide
6. Based on the Brønsted-Lowry Theory of Acids and Bases, would you expect pure water to have no dissolved ions whatsoever? Explain, using a balanced chemical equation.
1. 2. 3. 4. 5. 6.
Notes[edit | edit source]
- ^ Brown, Theodore E.; Lemay, H. Eugene; Bursten, Bruce E.; Murphy, Catherine; Woodward, Patrick (2009), Chemistry: The Central Science (11th ed.), New York: Prentice-Hall, ISBN 0136006175.