Volume 3
3.1. THE ALKALI METALS

Natural Occurrence and Discovery • Physical Properties • Chemical Properties • Alkali Metal Hydroxides and Their Strength as Bases • Preparation and Reactions of NaOH and KOH • Selected Examples of Alkali Metal Salts • Exercises
3.1.1. Natural Occurrence and Discovery. Of the six alkali metals (Li, Na, K, Rb, Cs, and Fr), sodium and potassium are among the top 10 most mass-abundant elements in the Earth's crust (Na 2.5%; K 2.0%). Sodium and potassium compounds have been known since ancient times. Lithium and especially rubidium and cesium are rare elements that were discovered in the 19th century. Francium is a radioactive and extremely rare element that was discovered in 1939. The total amount of Fr in the entire Earth's crust is estimated at just 20-30 g. All alkali metals occur in nature exclusively in the form of positively charged ions as constituents of a broad variety of minerals and salts. Sodium and potassium metals were made for the first time by electrolysis of NaOH and KOH (Humphry Davy, 1807). Lithium metal was first prepared in 1821, and metallic cesium and rubidium were isolated shortly after their discovery in the early 1860s.

3.1.2. Physical Properties. Table 1 lists the atomic masses, density values, atomic and ionic radii, and melting points for all of the alkali metals, except Fr. The largest quantity of Fr ever produced in the laboratory was on the order of 300,000 atoms. Such an amount is way too small for an experimental determination of physical properties of a substance, like density and melting and boiling points. The atomic size of the alkali metals increases in the order Li < Na < K < Rb < Cs (Table 1), in full accord with an increase in the number of electron shells down the group.

Table 1. Selected physical properties of the alkali metals.
Alkali metals are soft and can be easily cut with a knife. Among all metals, lithium, sodium, and potassium are the only ones that are less dense than water (Table 1). With its density of only 0.53 g/cm3 lithium is the lightest metal known, being not only almost nearly twice as light as water, but also lighter than most types of wood.

If asked to name a yellow metal, nearly everyone would immediately say, "gold". There is, however, one more metal that is yellow in color, cesium. As can be seen in this truly remarkable video, cesium is yellow, although not as bright yellow as gold. When watching the video, note that the cesium metal inside the sealed glass tube is partially molten because its melting point (28.4 oC) is below the temperature of the hand holding the tube. Except cesium, all other alkali metals are silver in color. [Note that when talking colors of metals we mean individual simple substances, not alloys. An alloy is a mix of two or more individual metals, conventionally made by co-melting them. There are many yellow copper and gold alloys. The two most widely known yellow copper alloys are bronze (Cu + Sn) and brass (Cu + Zn).]

Alkali metals are flame colorants. As shown in Figure 3-1 and in this video, each alkali metal imparts a particular color to a flame. Sodium compounds are indispensable for making yellow fireworks and signal flares. The ability of the alkali metals and some other elements to give a certain color to a flame has long been used in the so-called flame test. For the test, a chemical sample is placed in the flame of a laboratory burner. The color observed serves an indication of the presence or absence of the element the sample is analyzed for. Care should be exercised, however, when drawing conclusions from results of a flame test. For example, the rather weak purple color from potassium (Figure 3-1) may well not be seen in the presence of even small amounts of sodium due to the vastly brighter and more intense yellow color that Na emits.
Figure 3-1. Flame colors imparted by alkali metals. Left to right: Li (red; source), Na (yellow; source), K (purple; source), Rb (purplish-blue; source), and Cs (purplish-blue; source).
Digression. Why are atoms/ions of many elements flame colorants? A sufficiently high temperature provides the electrons of an atom/ion with enough energy to move from their lowest energy state (ground state) to higher energy orbitals (excited state). As the electrons return from the excited state to the ground state, the energy difference is released in the form of light. Being distinct for each element, this difference controls the color of the emitted light. As a result, compounds of different elements produce different flame colors.
3.1.3. Chemical Properties. The alkali metals are located in Group 1 on the far left of the periodic table. Their atoms have only one electron in the outermost shell, as shown below.

Li 1s2 2s1
Na 1s2 2s2 2p6 3s1
K 1s2 2s2 2p6 3s2 3p6 4s1
Rb 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s1
Cs 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s1

By discarding its single valence electron, an alkali metal atom becomes a positively charged ion with the most stable electron configuration of a noble gas. It is this stability that serves as the driving force behind reactions of the alkali metals with substances that can accept electrons. All reactions of alkali metals (but not their ions!) are redox reactions where an alkali metal is oxidized to the corresponding cation by an electron-acceptor.

Being the least electronegative elements, the alkali metals readily lose their valence electrons to even weak oxidizers, such as water. When reacting with water, alkali metals reduce protons (H+) to hydrogen atoms, which then dimerize to give H2 (Figure 3-2). The powerful electron-donating ability of the alkali metals makes this reaction fast, even though the concentration of H+ in water is low.
Figure 3-2. Reaction of sodium, an alkali metal, with water.


Watch this unique demonstration to see how all five different alkali metals react with water. A summary of these reactions is presented in Figure 3-3. The reaction of Li with water is smooth while being rather fast. Sodium reacts with water more vigorously. The reaction of potassium is even more energetic, producing enough heat for the H2 formed to catch fire. Rubidium and cesium explode on contact with water, the reaction of Cs being violent enough to shatter the glass dish used in the experiment. The order of reactivity observed, Li < Na < K < Rb < Cs, parallels the propensity of different alkali metal atoms to give away their valence electron (Figure 3-3). The larger in size (and heavier) the alkali metal atom, the farther away from the nucleus its outermost shell electron is and the weaker the electrostatic attraction between them. The weaker the attraction, the more easily the electron is discarded, which accounts for the observed order of reactivity, Li < Na < K < Rb < Cs.
Figure 3-3. Reactivity of different alkali metals toward water.


Naturally, all of the alkali metals react with stronger electron-acceptors than water, including halogens, phosphorus, and sulfur (Figure 3-4). These reactions are highly exothermic and often violent, as can be seen from the video demonstrations of reactions of Na metal with sulfur, phosphorus, and chlorine.
Figure 3-4. Reactions of sodium with sulfur, phosphorus, and chlorine.


Did you notice in the videos that in order to initiate some of the reactions, a small quantity of water had to be added to the reagents? The main role of water in promoting these reactions is the clearing of the surface of the sodium metal. All alkali metals are quickly oxidized by oxygen and moisture on exposure to air. That is why the shiny surface of just cut sodium quickly turns dull (watch this video). The products of these reactions, mainly Na2O and NaOH, form a layer on the Na surface, which blocks access of reagents such as S or Cl2 to the underlying metal. Water dissolves the protecting layer, thereby exposing the Na metal to the reagent.

One might expect reactions of an alkali metal with O2 to be simple and straightforward, giving rise to the corresponding metal oxide, such as Na2O and K2O. While quite a few sources like this one state so, this is not the case. Among all alkali metals, only lithium reacts with O2 to give the expected "normal" oxide, Li2O.

4 Li + O2 = 2 Li2O

The reaction of Na with O2 produces mixtures of Na2O and sodium peroxide, Na2O2. Burning potassium in air yields a mixture of K2O, potassium peroxide (K2O2), and potassium superoxide (KO2). Rubidium and cesium also form mixtures of binary oxides when combined with oxygen.

3.1.4. Alkali Metal Hydroxides and Their Strength as Bases. All alkali metal hydroxides are water-soluble ionic compounds and powerful electrolytes strongly dissociated in solution. To paraphrase George Orwell's famous quote, "All animals are equal, but some animals are more equal than others", all alkali metal hydroxides are strong bases, but some of them are stronger bases than others. The order of basicity, LiOH < NaOH < KOH < RbOH < CsOH, parallels that of the size of the metal ions, M+ (Table 1). A rationale for the higher basicity of the MOH with a bulkier alkali metal cation is provided in Figure 3-5 using LiOH vs. KOH as an example. Being larger in size than Li+, K+ experiences a weaker electrostatic attraction from the OH- anion, which makes electrolytic dissociation of KOH more facile.
Figure 3-5. The order of basicity of alkali metal hydroxides, explained by weaker electrostatic interactions for the larger alkali metal cation.


3.1.5. Preparation and Reactions of NaOH and KOH. The two by far most important alkali metal hydroxides are NaOH and KOH. Lithium hydroxide is much less common, and both RbOH and CsOH are exotic compounds.

Both NaOH and KOH are manufactured and used on a very large scale. Like other alkali metal hydroxides, NaOH and KOH are formed on treatment of the corresponding metal or its oxide with water (Figure 3-6), but these methods are far from being economical, let alone safe. As we touched on previously (Volume 2), sodium and potassium hydroxides are manufactured by the chloralkali process that involves electrolysis of NaCl and KCl, respectively, in water (Figure 3-6, bottom equation).
Figure 3-6. Methods to make NaOH. KOH is formed similarly.


As very strong bases, NaOH and KOH readily react with compounds exhibiting acidic and even weakly acidic properties. Examples of reactions of NaOH with a strong acid, a weak acid, an acidic oxide, an amphoteric oxide, and an amphoteric hydroxide are presented in Figure 3-7. As already mentioned in Volume 1, the formulas of the products of the reactions with amphoteric ZnO and Al(OH)3 are simplified for our introductory course.
Figure 3-7. Examples of reactions of NaOH. KOH reacts similarly.


3.1.6. Selected Examples of Alkali Metal Salts. Sodium and potassium salts are ubiquitous in nature and life. Salts of lithium are much less frequently encountered, and those of Rb and Cs are very rare. All alkali metal salts of simple inorganic acids are white in color and soluble in water.

Sodium salts are found in virtually every household. The most common sodium salt that everyone has seen and used is sodium chloride (NaCl), table salt. Nearly as common is baking soda (NaHCO3, sodium bicarbonate). The active ingredient of laundry bleach is sodium hypochlorite, NaOCl.

Potassium is critical for maintaining the sodium-potassium balance in the human body. A lower than normal potassium level in the blood (hypokalemia) results in high blood pressure and an increased risk of cardiovascular diseases. Potassium chloride (KCl) is used to provide potassium enrichment in such foods as baby formulas, potato chips and other snacks, cereals, and sports drinks. Potassium is one of the three elements that are essential for plant nutrition (the other two are nitrogen and phosphorus). Potassium salts used as fertilizers include potassium chloride (KCl), potassium nitrate (KNO3), potassium carbonate (K2CO3), and potassium sulfate (K2SO4).

Lithium has been used for over 70 years in treatment of bipolar and schizophrenic disorders. The most widely used lithium drug is lithium carbonate (Li2CO3), which is on the World Health Organization's list of essential medicines.

3.1.7. Exercises.

1. What are the two most widespread alkali metals? Answer

2. All reactions of the alkali metals as simple substances are (a) redox reactions where an alkali metal serves as the oxidant; (b) ion exchange reactions; (c) redox reactions in which an alkali metal is oxidized; (d) redox reactions in which an alkali metal reduces another substance; (e) neutralization reactions. Answer

3. Write the electron configuration of Na and K atoms and explain why the alkali metals are reducing agents.

4. The oxidation state of any alkali metal is always +1. True or false? Answer

5. Lithium is the only alkali metal that is less dense than water. True or false? Answer

6. Write a balanced chemical equation for the reaction of Na with N2 to give sodium nitride, Na3N. Is N2 an oxidant or reductant in this reaction? Answer

7. All alkali metals react with O2 to give M2O (M = Li, Na, K, Rb, Cs). True or false? Answer

8. All alkali metal hydroxides are equally strong bases. True or false? Answer

9. Provide an explanation for the order of basicity LiOH < NaOH < KOH < RbOH < CsOH. [Answer: See 3.1.4]

10. 1 g of Li metal was treated with water in excess and the volume of H2 produced was measured. The experiment was then repeated with 1 g of Na metal in place of the Li. Which of the two reactions produced a larger volume of hydrogen gas? Solve the problem mentally. Answer

11. The reaction of a 23-g chunk of sodium metal with water in excess produced (a) 22.4 L of H2; (b) 11.2 L of H2; (c) 40 g of NaOH and 22.4 L of H2; (d) 20 g of NaOH and 11.2 L of H2; (e) 40 g of NaOH and 11.2 L of H2. Answer

12. Alkali metal salts of simple inorganic acids are all white in color, except for cesium salts that are yellow. True or false? Answer

13. Relative to sodium, potassium is (a) a weaker reductant; (b) a stronger reductant; (c) both are equally strong reductants; (d) neither one is a reductant, as both are strong oxidants. Answer

14. In the reaction of sodium with water (a) sodium is reduced and hydrogen is oxidized; (b) oxygen is reduced and sodium is oxidized; (c) sodium is oxidized and hydrogen is reduced; (d) sodium is oxidized and both oxygen and hydrogen are reduced. Answer