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The Periodic Table: Crash Course Chemistry #4

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-Stored in oil
-Cesium/Rubidium are used to make atomic clocks
-Cesium can be spelled like "caesium"

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Chemistry for Kids: Elements - Alkali Metals

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Kids learn about the alkali metals of the periodic table. Which elements are in this group. Properties, similarities, and other facts.

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-Heat
-Metals
-Storage/keeping fresh
-Strong reducing agents

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In IUPAC (International Union of Pure and Applied Chemistry) nomenclature, alkali metals or Group 1A elements belong to group 1A in the periodic table. Scroll down to learn about the uses of alkali metals.

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In IUPAC (International Union of Pure and Applied Chemistry) nomenclature, alkali metals or Group 1A elements belong to group 1A in the periodic table. Scroll down to learn about the uses of alkali metals.

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-When they are heated; lithium, sodium, potassium, rubidium, and cesium ignite through combustion reactions with oxygen.

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Oxygen is a highly reactive element that is very abundant on earth and in the human body. It is found in many compounds that are used to sustain basic life forms and modern civilization. Compounds containing oxygen are of great interest in the field of chemistry.

Background

Oxygen is ubiquitous; it comprises approximately 46% of the crust, 21% of the atmosphere, and 61% of the human body. Because of oxygen's high reactivity, it is most often found in compounds. Oxygen's high reactivity is due to its biradical electron configuration. As shown in a molecular orbital drawing of O ₂, the two unpaired electrons make the molecule highly susceptible to bond formation.

Oxygen has two allotropes (dioxygen, O ₂, and ozone, O ₃), both excellent oxidizing agents ( Table P2). Oxygen is typically observed observed in the -2 oxidation state, in the form O ²⁻, but it can also form other ions such as peroxide, O ₂ ²⁻, and superoxide, O ₂⁻. With different possible oxidation states, many possible molecular compounds can be formed when an element reacts with oxygen. Many reactions involving oxygen occur in biological processes, including cellular respiration and photosynthesis.

Oxides are chemical compounds that contain at least one oxygen atom and at least one atom of another element. There are four principle oxidation states of oxygen: -2, -1, -1/2, and 0. The oxide ion, O ²⁻, has a oxidation state of -2; the peroxide ion, O ₂ ²⁻, has a oxidation state of -1; and the superoxide ion, O ₂⁻, has a oxidation state of -1/2. With metals, oxygen forms oxides that are largely ionic in character.

There are general trends in the reactions between main group elements and oxygen:

Most nonmetals form the oxide with the highest possible oxidation state, with the halides excepted. Most metals form oxides with the oxygen in a -2 oxidation state.

As a general rule, metal oxides are basic and nonmetal oxides are acidic. Basicity of an oxide increases with increasing ionic (metallic) character. Metal oxides, peroxides, and superoxides dissolve in water actually react with water to form basic solutions. Oxygen also forms covalent oxides with non-metals, that react with water to form acidic solutions.

Oxygen does not react with fluorine or noble gases.

Exceptions to all of these trends are discussed below.

Reactions with Hydrogen

Oxygen reacts with hydrogen to produce two compounds: water (\(H_2O\)) and hydrogen peroxide (\(H_2O_2\)). Water is a versatile compound and participates in acid-base equilibrium and oxidation-reduction reactions. It can act as an acid, base, reducing agent, or oxidizing agent. Water's multifaceted abilities make it one of the most important compounds on earth. The reaction between hydrogen and oxygen to form water is given below:

\[2H_{2 (g)} + O_{2 (g)} \rightarrow 2H_2O_{(l)} \label{1}\]

Hydrogen peroxide's potent oxidizing abilities give it great industrial potential. The following equation shows the reaction of hydrogen and oxygen to form hydrogen peroxide:

\[H_2 + O_2 \rightarrow H_2O_2 \label{2}\]

The product of this reaction is called a peroxide because oxygen is in the \(O_2^{2-}\) form (hydrogen has a +1 oxidation state). This concept is further explained regarding lithium below.

Reactions with Group 1 Elements

Oxygen reacts rapidly with Group 1 elements. All alkali metal oxides form basic solutions when dissolved in water. The principal combustion product is the most stable product with respect to the reactants. For example, with careful control of oxygen, the oxide M ₂ O (where M represents any alkali metal) can be formed with any of the alkali metals. When heated, lithium, sodium, potassium, rubidium, and cesium ignite through combustion reactions with oxygen.
Lithium
Lithium, the first metal in Group 1, reacts with oxygen to form Li ​2O and burns with a red flame. The oxygen in this compound is an oxide (O ²⁻). The formation of Li ₂ O, the principal combustion product, is illustrated by the equation below:

\[4 Li(s) + O_2(g) \rightarrow 2 Li_2O(s)\label{3}\]

However, if there is excess oxygen present, it is possible that a small amount of the compound Li ₂O ₂ can be formed. Because alkali metals always have a +1 oxidation state, oxygen is in the O ₂ ²⁻ form. When oxygen is in this state, the compound is called a peroxide. The formation of this peroxide, the less-likely non-principal combustion product, under excess oxygen is illustrated by the equation below:

\[2 Li(s) + O_2(g) \rightarrow Li_2O_2(s) \label{4}\]
Sodium
Sodium burns in air with often little more than an orange glow. Using larger amounts of sodium or burning it in pure oxygen produces a strong orange flame. A white solid mixture of sodium oxide and sodium peroxide is formed. The equation for the formation of the simple oxide is analogous to that for lithium:

\[ 4Na(s) + O_2 (g) \rightarrow 2Na_2O (s) \label{5}\]

Likewise, the reaction for peroxide formation takes the same form for both metals:

\[ 2Na(s) + O_2 (g) \rightarrow Na_2O_2 (s) \label{6}\]
Figure 1: Sodium explosion. Image used with permission from Wikipedia. Potassium
Small pieces of potassium heated in air tend to melt instantly into a mixture of potassium peroxide and potassium superoxide with no visible flame. Larger pieces of potassium burn with a lilac-colored flame. The equation for the formation of the peroxide is identical to that for sodium (click here for more information):

\[ 2K(s) + O_2 (g) \rightarrow K_2O_2 (s) \label{7}\]

The superoxide generating reaction is given below:

\[ K(s) + O_2 (g) \rightarrow KO_2 (s) \label{8}\]
Other alkali metals
The other alkali metals (Rb, Cs, Fr) form superoxide compounds (in which oxygen takes the form O ₂⁻) as the principal combustion products. The following equation shows the formation of superoxide, where M represents K, Rb, Cs, or Fr:

\[M(s) + O_2(g) \rightarrow MO_2(s) \label{9}\]

These compounds tend to be effective oxidizing agents due to the fact that O ₂⁻ is one electron short of a complete octet and thus has a strong affinity for another electron. It is easily reduced, and therefore act as an effective oxidizing agent.

Reactions with Group 2 Elements

The elements of Group 2 are beryllium, magnesium, calcium, strontium, barium, and radioactive radium. Alkaline earth metals also react with oxygen, though not as rapidly as Group 1 metals; these reactions also require heating. Similarly to Group 1 oxides, most group 2 oxides and hydroxides are only slightly soluble in water and form basic, or alkaline solutions.

All Group 2 metals all react similarly, burning to form oxides (compounds containing the O ²⁻ ion) as shown:

\[2 M(s) + O_2(g) \rightarrow 2 MO(s) \label{10}\]

Once initiated, the reactions with oxygen are vigorous. The only peroxides (compounds containing the O ₂ ²⁻ ion) that can be formed from alkaline metals are strontium peroxide and barium peroxide. Both reactions require heat and excess oxygen. The general reaction is given below:

\[M(s) + O_2(g) \rightarrow MO_2(s)\label{11}\]

where M represents Sr or Ba.
Beryllium
Beryllium is unreactive with air and water. The chemical behavior of beryllium is best attributed to its small size and high ionization energy of its atoms.
All other group 2 metals
Except beryllium, the other alkaline earth metals form oxides in air at room temperature.

\[2 M(s) + O_2(g) \rightarrow 2 MO(s) \label{12}\]

where M represents Be, Mg, Ca, Sr, Ba, or Ra.

Peroxides, of the form MO ₂, are formed for all these elements except beryllium as shown:

\[M(s) + O_2(g) \rightarrow MO_2(s) \label{13}\]

Magnesium, calcium, strontium and barium oxides react with water to form hydroxides:

\[ MO(s) + H_2O(l) \rightarrow M(OH)_2(s) \label{14}\]

All the oxides and hydroxides of the group 2 metals, except of those of beryllium, are bases:

\[ M(OH)_2(s) \rightarrow M^{2+}(aq) + 2OH^-(aq) \label{15}\]

Reactions with Group 13 Elements

Group 13 consists of the following elements: boron, aluminum, gallium, indium, and thallium. Boron is the only element in this group that possesses no metallic properties. These elements vary in their reactions with oxygen. Recall that oxides of metals are basic and oxides or nonmetals are acidic; this is true for all elements in Group 13, except Al and Ga.

All other Group 13 elements also produce compounds of the form of M ​2O ₃, but adhere to the acid-base rules of metal and nonmetal oxides. Here is the equation of the reaction of oxygen and a Group 13 element:

\[4M(s) + 3O_2(g) \rightarrow 2M_2O_3(s) \label{16}\]

where M is any Group 13 element. At high temperatures, thallium also reacts with oxygen to produce Tl ​2 O:

\[4Tl(s) + O_2(g) \rightarrow 2Tl_2O \label{17}\]

Boron:

The most common oxide form of boron, B ₂O ₃ or boron trioxide, is obtained by heating boric acid:

\[2B(OH)_3 \xrightarrow{\Delta} B_2O_3 + 3H_2O \label{18}\]

Aluminum:

Aluminum occurs almost exclusively in the +3 oxidation state. It rapidly reacts with oxygen in air to give a water-insoluble coating of Al ₂O ₃. This oxide layer protects the metal beneath from further corrosion. The reaction is shown below:

\[4Al(s) + 3O_2(g) \rightarrow 2Al_2O_3 \label{19}\]

Aluminum trioxide, Al ​2O ₃, is amphoteric (acts both as an acid and a base):

\[Al_2O_3(s) + 6HCl(aq) \rightarrow 2AlCl_3(aq) + 3H_2O(l) \label{20}\]

\[Al_2O_3(s) + 2NaOH(aq) + 3H_2O(l) \rightarrow 2Na[Al(OH)_4](aq) \label{21}\]

Except for thallium in which the +1 oxidation state is more stable than the +3 state, aluminum, gallium, and indium favor +3 oxidation states

All of group 13 metal elements are known to form a trivalent oxide.

\[ 4M(s) + 3O_2(g) \rightarrow 2M_2O_3(s) \label{22}\]

with M represents Al, Ga, In, or Tl

Thallium is the only element in this group favors the formation of oxide over trioxide.

\[2M(s) + O_2(g) \rightarrow 2MO(s) \label{23}\]

Reactions with Group 14 Elements

Group 14 is made up of both metals (toward the bottom of the group), metalloids, and nonmetals (at the top of the group). The oxides of the top of Group 4 elements are slightly acidic, and the acidity of the oxides decreases down the group.

Non-metal: The non-metal carbon of Group 14 (and its compounds) burn to form \(CO_{2\;(g)}\) and, in smaller amounts, \(CO_{(g)}\); both are acidic under different conditions. Carbon monoxide is only slightly soluble in water and does not react with it. Click here for more Information.

Metalloid: The metalloid silicon reacts with oxygen to form only one stable compound, \(SiO_{​2}\), which dissolves slightly in water and is weakly acidic(Figure 2).

The three metals in this group have many different oxide compounds due to their extended octets. All of these oxides are amphoteric (exhibit both basic and acidic properties). For example:

Germanium: \(GeO\), \(GeO_{​2}\)

Tin: \(SnO\), \(SnO_{​2}\)

Lead: \(PbO\), \(PbO_​{2}\), \(Pb_​{3}O_4\)

Figure 2: \(SiO_{​2}\): There are various different structures for silicon dioxide. The easiest to remember and draw is based on a diamond structure with each of the silicon atoms being bridged to its other four neighbors via an oxygen atom.

Reactions with Group 15 Elements

The nitrogen family, Group 15, is capable of reacting with oxygen in many different ways. Nitrogen and phosphorus are nonmetallic, arsenic and antimony are metalloids, and bismuth is metallic.
Nitrogen
Nitrogen reacts with oxygen to form many oxides ranging in oxidation states from +1 to +5: All these oxides are gases at room temperature except for N ​2O ₅, which is solid. The nitrogen oxides are given below:

NO, N ₂O, N ₂O ₃, NO ₂, N ₂O ₅

All of these reactions are endothermic, requiring energy for oxygen to react directly with N ₂(g). The oxides of nitrogen are acidic (because they are nonmetal oxides). N ​2O ₃ and N ​2O ₅ react with water to give acidic solutions of oxoacids. These reactions are shown below:

Nitrous acid: \[N_2O_3(s) + H_2O(l) \rightarrow 2HNO_2(aq) \label{24}\]

Nitric acid: \[N_2O_5(s) + H_2O(l) \rightarrow 2HNO_3(aq) \label{25}\]
Phosphorus
There are two forms of allotropes of phosphorus, white phosphorus and red phosphorus. Red phosphorus is less reactive than white phosphorus. Phosphorus reacts with oxygen, usually forming two oxides depending on the amount available oxygen: P ₄O6 when reacted with a limited supply of oxygen, and P ​4O 10 when reacted with excess oxygen; the latter is shown below.

\[P_4O_{10} + 6 H_2O \rightarrow 4 H_3PO \label{26}\]

On rare occasions, P ​4O ₇, P ​4O ₈, and P ​4O ₉ are also formed, but in small amounts.

Both P ​4O ₄ and P ​4O ₁₀ react with water to generate oxoacids. Reactions are shown below.

Phosphorous acid: \[P_4O_6(l) + 6H_2O(l) \rightarrow 4H_3PO_3(aq) \label{27}\]

Phosphoric acid: \[P_4O_{10}(s) + 6H_2O(l) \rightarrow 4H_3PO_4(aq) \label{28}\]
Other Group 15 Elements
Arsenic, antimony and bismuth react with oxygen when burned. The common oxidation states for arsenic, antimony, and bismuth are +3 and +5. There are two main types of oxides for each element:

Arsenic: As​2O₃, As​2O₅

Antimony: Sb​2O₃, Sb​2O₅

Bismuth: Bi​2O₃, Bi​2O₅

There are other oxides, such as Sb ​4O ₁₀, that are not formed directly through reaction with oxygen. Arsenic(III) oxide and antimony(III) oxide are amphoteric, whereas bismuth(III) oxide acts only as a base (this is because it is the most metallic element in the group).

Reactions with Group 16 Elements

The elements in Group 16 include oxygen, sulfur, selenium, tellurium, and polonium. Oxygen reacts with the elements in its own group to form various oxides, mostly in the form of AO ₂ and AO ₃.
Oxygen
Although oxygen is located in Group 16, it is unique in its extreme electronegativity; this allows it to readily gain electrons and create hydrogen bonds. Because it is the smallest element in its group, it is capable of forming double bonds. It has no d-orbitals, and cannot expand its valence shell.

Oxygen is capable of reacting with itself, forming allotropes. One of oxygen's allotropes, ozone (O ₃), is formed when oxygen gas, O ₂, is subjected to ultraviolet light.
Sulfur
Sulfur dioxide, SO ₂, and sulfur trioxide, SO ₃, are the only common sulfur oxides.

\[S(s) + O_2(g) \xrightarrow{\Delta} SO_2(g) \label{29}\]

Sulfur's reaction with oxygen produces the oxides mentioned above as well as oxoacids. All are powerful oxidizing agents. SO ₂ is mainly used to make SO ₃, which reacts with water to produce sulfuric acid (recall that nonmetals form acidic oxides). These sequential reactions are shown below:

\[ 2SO_2(g) + O_2(g) \rightleftharpoons 2SO_3(g) \label{30}\]

\[2SO_3(g) + H_2O(l) \rightarrow H_2SO_4(aq) \label{31}\]
Selenium and tellurium
Selenium and tellurium adopt compounds of the forms AO ₂, AO ₃, and AO.

Reactions with Group 17 Elements

The elements in Group 17 include fluorine, chlorine, bromine, and iodine. These elements are called halogens, from Greek roots translating to "salt formers." The halogens react with oxygen, but many of the resulting compounds are unstable, lasting for only moments at a time. They range in structure from X ​2O to X ​2O ₇, where X represents a halogen. Their extended octets allow them to bond with many oxygen atoms at a time.
Fluorine:
The most electronegative element adopts the -1 oxidation state. Fluorine and oxygen form OF ₂, which is known as oxygen fluoride.

\[2 F_2 + 2NaOH \rightarrow OF_2 + 2 NaF + H_2O \label{32}\]
Other Halogens
The other halogens form oxoacids instead of oxides. For example:

Oxidation state of halogen

Chlorine

Bromine

Iodine

+1

HOCl

HOBr

HOI

+3

HClO ₂

--

--

+5

HClO ₃

HBrO ₃

HIO HIO ₃

+7

HClO ₄

HBrO ₄

HIO ₄; H ₅IO ₆

Reactions with Group 18 Elements

The Group 18 noble gases include helium, neon, krypton, xenon, and radon. Noble gases are chemically inert with the exception of xenon, which reacts with oxygen to form XeO ​3 and XeO ​4 at low temperatures and high pressures. The ionization energy of xenon is low enough for the electronegative oxygen atom to capture electrons. XeO ​3 is highly unstable, and is known to spontaneously detonate in a clean, dry environment.

References

1. Reactions of Selenium and Oxygen. Matrix Infrared Spectra and Density Functional Calculations of Novel SexOy Molecule

2. G. Dana Brabson and, Lester Andrews, Colin J. Marsden. The Journal of Physical Chemistry 1996 100 (41), 16487-16494

3. Petrucci, Ralph H. General Chemistry: Principles and Modern Applica tions - 9th ed. Upper Saddle River: Prentice Hall, 2007. Print.

4. Zumdahl, Steven S. Chemical Principles - 5th ed. Houghton Mifflin Company, 2005. Print.

5. Madsen, Dorte. Class Lecture Notes. Main Group Elements. University of California, Davis, California. Spring 2011.

Problems

1. If there are 4.00 g of potassium, how much K₂O is created when during combustion (in excess oxygen)?

2. Which noble gas(es), if any, react with oxygen? Why?

3. Complete and balance the following reaction: Al(s) + O₂(g) →

4. What is the best environment for nitrogen and oxygen to react in?

5. What are the principal combustion products of each Group 1 element with oxygen?

Contributors

Kathleen M. (UCD Spring 2010), Antoinette Mursa (UC Davis Spring 2011)

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Reactions of Main Group Elements with Oxygen

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Oxygen is a highly reactive element that is very abundant on earth and in the human body. It is found in many compounds that are used to sustain basic life forms and modern civilization. Compounds containing oxygen are of great interest in the field of chemistry.

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-Sodium and Potassium are vital to life
-Rubidium and Cesium are not very important

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ALKALI METALS CONCEPT Group 1 of the periodic table [1] of elements consists of hydrogen, and below it the six alkali metals [2]: lithium, sodium, potassium, rubidium, cesium, and francium.

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-Metals are usually more of an importance compounds

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In IUPAC (International Union of Pure and Applied Chemistry) nomenclature, alkali metals or Group 1A elements belong to group 1A in the periodic table. Scroll down to learn about the uses of alkali metals.

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In IUPAC (International Union of Pure and Applied Chemistry) nomenclature, alkali metals or Group 1A elements belong to group 1A in the periodic table. Scroll down to learn about the uses of alkali metals.

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-They can react well with nonmetals and water
-Basic
-Highly reactive

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Structural Biochemistry/Alkali Metals - Wikibooks, open books for an open world

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The alkali metals are a series of chemical elements forming Group 1 (IUPAC style) of the periodic table: lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs), and francium (Fr). (Hydrogen, although nominally also a member of Group 1, very rarely exhibits behavior comparable to the alkali metals).

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All the elements are metals

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As you go down on the table, the elements get more reactive

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Li: Lithium
Na: Sodium
K: Potassium
Rb: Rubidium
Cs: Cesium
Fr: Francium

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1 valence electron

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-Hydrogen: s1
-Lithium: s1
-Sodium: s1
-Potassium: s1
-Rubidium: s1
-Cesium: s1
-Francium: s1
Source: Ck-12

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- larger atomic radii and low ionization energies
-  reaction with water forms alkalies
-  alkali metals have low melting points
-react readily with atmospheric oxygen and water vapor
Source:
https://chem.libretexts.org

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- high thermal and electrical conductivity
- high lustre
- high ductility
- high malleability
Source: Britannica

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Alkali metal , any of the six chemical elements that make up Group 1 (Ia) of the periodic table -namely, lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), and francium (Fr). The alkali metals are so called because reaction with water forms alkalies (i.e., strong bases capable of neutralizing acids ). Sodium and potassium are the sixth and seventh most abundant of the elements, constituting, respectively, 2.6 and 2.4 percent of Earth's crust . The other alkali metals are considerably more rare, with rubidium, lithium, and cesium, respectively, forming 0.03, 0.007, and 0.0007 percent of Earth's crust. Francium, a natural radioactive isotope , is very rare and was not discovered until 1939.

The alkali metals are so reactive that they are generally found in nature combined with other elements. Simple minerals , such as halite (sodium chloride, NaCl), sylvite (potassium chloride, KCl), and carnallite (a potassium-magnesium chloride, KCl · MgCl ₂· 6H ₂O), are soluble in water and therefore are easily extracted and purified. More complex, water-insoluble minerals are, however, far more abundant in Earth's crust. A very dilute gas of atomic sodium (about 1,000 atoms per cubic cm [about 16,000 atoms per cubic inch]) is produced in Earth's mesosphere (altitude about 90 km [60 miles]) by ablation of meteors . Subsequent reaction of sodium with ozone and atomic oxygen produces excited sodium atoms that emit the light we see as the "tail" of a meteor as well as the more diffuse atmospheric nightglow. Smaller amounts of lithium and potassium are also present.

The alkali metals have the silver-like lustre, high ductility, and excellent conductivity of electricity and heat generally associated with metals. Lithium is the lightest metallic element. The alkali metals have low melting points, ranging from a high of 179 °C (354 °F) for lithium to a low of 28.5 °C (83.3 °F) for cesium. Alloys of alkali metals exist that melt as low as −78 °C (−109 °F).

The alkali metals react readily with atmospheric oxygen and water vapour. (Lithium also reacts with nitrogen .) They react vigorously, and often violently, with water to release hydrogen and form strong caustic solutions. Most common nonmetallic substances such as halogens , halogen acids, sulfur , and phosphorus react with the alkali metals. The alkali metals themselves react with many organic compounds, particularly those containing a halogen or a readily replaceable hydrogen atom.

Sodium is by far the most important alkali metal in terms of industrial use. The metal is employed in the reduction of organic compounds and in the preparation of many commercial compounds. As a free metal, it is used as a heat-transfer fluid in some nuclear reactors . Hundreds of thousands of tons of commercial compounds that contain sodium are used annually, including common salt (NaCl), baking soda (NaHCO ₃), sodium carbonate (Na ₂CO ₃), and caustic soda (NaOH). Potassium has considerably less use than sodium as a free metal. Potassium salts, however, are consumed in considerable tonnages in the manufacture of fertilizers . Lithium metal is used in certain light-metal alloys and as a reactant in organic syntheses. An important use of lithium is in the construction of lightweight batteries . Primary lithium batteries (not rechargeable) are widely used in many devices such as cameras, cellular telephones, and pacemakers. Rechargeable lithium storage batteries that could be suitable for vehicle propulsion or energy storage are the subject of intensive research. Rubidium and cesium and their compounds have limited use, but cesium metal vapour is used in atomic clocks , which are so accurate that they are used as time standards.

History

Alkali metal salts were known to the ancients. The Old Testament refers to a salt called neter (sodium carbonate), which was extracted from the ash of vegetable matter. Saltpetre (potassium nitrate) was used in gunpowder , which was invented in China about the 9th century ad and had been introduced into Europe by the 13th century.

In October 1807 the English chemist Sir Humphry Davy isolated potassium and then sodium . The name sodium is derived from the Italian soda, a term applied in the Middle Ages to all alkalies; potassium comes from the French potasse, a name used for the residue left in the evaporation of aqueous solutions derived from wood ashes.

Lithium was discovered by the Swedish chemist Johan August Arfwedson in 1817 while analyzing the mineral petalite. The name lithium is derived from lithos, the Greek word for "stony." The element was not isolated in pure form until Davy produced a minute quantity by the electrolysis of lithium chloride.

While the German chemists Robert Bunsen and Gustav Kirchhoff were investigating the mineral waters in the Palatinate in 1860, they obtained a filtrate that was characterized by two lines in the blue region of its spectrum (the light emitted when the sample was inserted into a flame ). They suggested the presence of a new alkali element and called it cesium , derived from the Latin caesius, used to designate the blue of the sky. The same researchers, on extracting the alkalies from the mineral lepidolite , separated another solution, which yielded two spectral lines of red colour. They proposed the name rubidium for the element in this solution from the Latin rubidus, which was used for the darkest red colour. Francium was not discovered until 1939 by Marguerite Perey of the Radium Institute in Paris.

In the 19th century the only use for the alkali metals was the employment of sodium as a reagent in the manufacture of aluminum . When the electrolytic process for aluminum purification was established, it appeared that large-scale use of sodium would cease. Subsequent improvements in the electrolytic production of sodium, however, reduced the cost of this element to such an extent that it can be employed economically to manufacture gasoline additives, reagents for chemical industry , herbicides , insecticides , nylon , pharmaceuticals, and reagents for metal refining. The continuous electrolysis of sodium hydroxide, a technique called the Castner process, was replaced in 1926 by the Downs cell process. This process, in which a molten sodium chloride-calcium chloride mixture (to reduce the melting point ) is electrolyzed, produces both sodium metal and chlorine .

General properties of the group

Physical properties

The alkali metals have the high thermal and electrical conductivity, lustre, ductility, and malleability that are characteristic of metals. Each alkali metal atom has a single electron in its outermost shell. This valence electron is much more weakly bound than those in inner shells. As a result, the alkali metals tend to form singly charged positive ions ( cations ) when they react with nonmetals. The compounds that result have high melting points and are hard crystals that are held together by ionic bonds (resulting from mutually attractive forces that exist between positive and negative electrical charges). In the metallic state, either pure or in alloys with other alkali metals, the valence electrons become delocalized and mobile as they interact to form a half-filled valence band. As with other metals, such a partially filled valence band is a conduction band and is responsible for the valence properties typical of metals . In passing from lithium to francium , the single electron tends to be less strongly held. Generally, the energy necessary to remove the outermost electron from the atoms of an element, the ionization energy , decreases in the periodic table toward the left and downward in each vertical file, with the result that the most easily ionizable element in the entire table is francium , followed closely by cesium . The alkali metals, which make up the extreme left-hand file, have ionization energies ranging from 124.3 kilocalories per mole (kcal/mole) in lithium to 89.7 kcal/mole in cesium (omitting the rare radioactive element francium). The alkaline-earth metals , the next group to the right, have higher ionization energies ranging from 214.9 in beryllium to 120.1 kcal/mole in barium .

The electronegativity scale of the elements compares the ability of the atoms of the various elements to attract electrons to themselves. In the periodic table the electronegativities range from 0.7 for cesium, the least electronegative of the elements, to 4.0 for fluorine , the most electronegative. Metals are ordinarily considered to be those elements having values less than 2.0 on the electronegativity scale. As a group the alkali metals are the least electronegative of the elements, ranging from 0.7 to 1.0 on the scale, while the alkaline earths, the next group on the table, have electronegativities ranging from about 0.9 to 1.5.

The table summarizes the important physical and thermodynamic properties of the alkali metals. At atmospheric pressure these metals are all characterized by a body-centred cubic crystallographic arrangement (a standard pattern of atoms in their crystals ), with eight nearest neighbours to each atom. The closest distance between atoms, a characteristic property of crystals, increases with increasing atomic weight of the alkali metal atoms. As a group, the alkali metals have a looser crystallographic arrangement than any of the other metallic crystals, and cesium-because of its greater atomic weight -has an interatomic distance that is greater than that of any other metal.

Vapour-pressure data for the alkali metals and for two alloys formed between elements of the group show that the vapour pressures increase in regular fashion with increasing atomic weight. Cesium is the most volatile of the alkali metals, with a boiling point of 671 °C (1,240 °F). The boiling points of the alkali metals decrease in regular fashion as the atomic numbers increase, with the highest, 1,317 °C (2,403 °F), being that of lithium.

The melting points of the alkali metals as a group are lower than those of any other nongaseous group of the periodic table, ranging between 179 °C (354 °F) for lithium and 28.5 °C (83.3 °F) for cesium. Among the metallic elements, only mercury has a lower melting point (−38.9 °C, or −38.02 °F) than cesium. The low melting points of the alkali metals are a direct result of the large interatomic distances in their crystals and the weak bond energies associated with such loose arrays. These same factors are responsible for the low densities, low heats of fusion, and small changes in volume upon fusion of the metals. Lithium, sodium , and potassium are less dense than water.

The large size of an alkali metal atom (and the resulting low density of the metal) results from the presence of only one, weakly bound electron in the large outer s-type orbital . Upon going from the noble-gas configuration of argon (atomic number 18) to potassium (atomic number 19), the added electron goes into the large 4 s orbital rather than the smaller 3 p orbital. When, however, potassium, rubidium , or cesium metals are subjected to increasing pressure (up to one-half million atmospheres or more), a number of phase transitions occur. Ultimately, occupation of a d-type orbital becomes preferred over that of the s-orbital, with the result that these alkali metals resemble transition metals . Under such circumstances, alloys with transition metals (such as iron ) can form, a result that does not occur at low pressures. It has been proposed that the lower-than-expected density of Earth's core may be the result of the formation of a potassium-iron alloy under the extreme pressures that occur there.

The alkali metals have played an important role in quantum physics . Some alkali metal isotopes, such as rubidium-87, are bosons . Dilute atomic gases of such alkali metal isotopes, confined by magnetic fields or " laser mirrors" and cooled to temperatures near absolute zero , form Bose-Einstein condensates. In this state, the cluster of atoms is in a single quantum state and exhibits macroscopic behaviours normally seen only with atomic-sized particles. These include interference effects and coherent motion of the entire "cloud" of atoms.

Chemical properties

Since the alkali metals are the most electropositive (the least electronegative) of elements, they react with a great variety of nonmetals. In its chemical reactivity, lithium more closely resembles Group 2 (IIa) of the periodic table than it does the other metals of its own group. It is less reactive than the other alkali metals with water , oxygen , and halogens and more reactive with nitrogen , carbon , and hydrogen .

Reactions with oxygen

The alkali metals tend to form ionic solids in which the alkali metal has an oxidation number of +1. Therefore, neutral compounds with oxygen can be readily classified according to the nature of the oxygen species involved. Ionic oxygen species include the oxide, O ²⁻, peroxide , O ₂ ²⁻, superoxide, O ₂⁻, and ozonide O ₃⁻. Compounds that can be prepared that contain an alkali metal, M, and oxygen are therefore the monoxide, M ₂O, peroxide, M ₂O ₂, superoxide, MO ₂, and ozonide, MO ₃. Rubidium and cesium and, possibly, potassium also form the sesquioxide, M ₄O ₆, which contains two peroxide anions and one superoxide anion per formula unit. Lithium forms only the monoxide and the peroxide.

All the alkali metals react directly with oxygen; lithium and sodium form monoxides, Li ₂O and Na ₂O, and the heavier alkali metals form superoxides, MO ₂. The rate of reaction with oxygen, or with air, depends upon whether the metals are in the solid or liquid state , as well as upon the degree of mixing of the metals with the oxygen or air. In the liquid state, alkali metals can be ignited in air with ease, generating copious quantities of heat and a dense choking smoke of the oxide .

The free energy of formation (a measure of stability) of the alkali metal oxides at 25 °C (77 °F) varies widely from a high of −133 kcal/mole for lithium oxide to −63 kcal/mole for cesium oxide. The close approach of the small lithium ion to the oxygen atom results in the unusually high free energy of formation of the oxide. The peroxides (Li ₂O ₂and Na ₂O ₂) can be made by passing oxygen through a liquid- ammonia solution of the alkali metal, although sodium peroxide is made commercially by oxidation of sodium monoxide with oxygen. Sodium superoxide (NaO ₂) can be prepared with high oxygen pressures, whereas the superoxides of rubidium, potassium, and cesium can be prepared directly by combustion in air. By contrast, no superoxides have been isolated in pure form in the case of lithium or the alkaline-earth metals, although the heavier members of that group can be oxidized to the peroxide state. The cyanides of potassium, rubidium, and cesium, which are less stable than the lower oxides, can be prepared by the reaction of the superoxides with ozone.

Reactions with water

The alkali metals all react violently with water according to M + H ₂O → MOH + 1/2 H ₂. The rate of the reaction depends on the degree of metal surface presented to the liquid. With small metal droplets or thin films of alkali metal, the reaction can be explosive. The rate of the reaction of water with the alkali metals increases with increasing atomic weight of the metal. With the heavier alkali metals, the hydroxides are highly soluble; thus, they are removed readily from the reacting surface, and the reaction can proceed with unabated vigour. The reaction involves equimolar mixtures (that is, equal numbers of atoms or molecules ) of the alkali metal and water to form a mole (an amount equal to that of the reactants) of alkali metal hydroxide and half a mole of hydrogen gas. These reactions are highly exothermic (give off heat), and the hydrogen that is generated can react with oxygen to increase further the heat that is generated.

Reactions with nonmetals

Of the alkali metals, only lithium reacts with nitrogen, and it forms a nitride (Li ₃N). In this respect it is more similar to the alkaline-earth metals than to the Group 1 metals. Lithium also forms a relatively stable hydride , whereas the other alkali metals form hydrides that are more reactive. Lithium forms a carbide (Li ₂C ₂) similar to that of calcium . The other alkali metals do not form stable carbides, although they do react with the graphite form of carbon to give intercalation compounds (substances in which the metal atoms are inserted between layers of carbon atoms in the graphite structure).

The alkali metals can be burned in atmospheres of the various halogens to form the corresponding halides . The reactions are highly exothermic , producing up to 235 kcal/mole for lithium fluoride. The alkali metals react with nonmetals in Groups 15 and 16 (Va and VIa) of the periodic table. Sulfides can be formed by the direct reaction of the alkali metals with elemental sulfur , furnishing a variety of sulfides. Phosphorus combines with the alkali metals to form phosphides with the general formula M ₃ P.

Formation of alloys

The characteristics of alloy behaviour in alkali metals can be evaluated in terms of the similarity of the elements participating in the alloy. Elements with similar atomic volumes form solid solutions (that is, mix completely in all proportions); some dissimilarity in atomic volumes results in eutectic-type systems (solutions formed over limited concentration ranges), and further dissimilarity results in totally immiscible systems. The high-pressure transition in potassium, rubidium, and cesium that converts these s-type metals to more transition metal-like d-type metals yields atomic volumes that are similar to those of many transition metals at the same pressure. This permits alloys or compounds to form between these alkali metals and such transition metals as nickel or iron.

The elements potassium, rubidium, and cesium, which have rather similar atomic volumes and ionization energies, form complete solid solutions and mixed crystals. Sodium, which is a significantly smaller atom than potassium and has a higher ionization energy, tends to form eutectic systems with potassium, rubidium, and cesium. Even greater dissimilarity exists in the atomic volumes of sodium and lithium, resulting in insolubilities of the liquid phases. The consolute temperature (the temperature at which the two liquids become completely miscible) increases on going from the lithium-sodium alloy system to the lithium-cesium system. Lithium and cesium can coexist as two separate liquid phases at temperatures up to at least 1,100 °C (2,000 °F).

There is only one example of solid miscibility in alkali-alkaline-earth-metal binaries-the lithium-magnesium system, in which the two elements are very similar. Sodium forms compounds only with barium in the alkaline-earth-metal series. The heavier alkali metals all tend to form immiscible liquid phases with the alkaline earth metals.

Several elements in Group 12 (IIb) of the periodic table ( zinc , cadmium , and mercury ) react with the alkali metals to form compounds. Mercury forms at least six compounds, commonly termed amalgams, with each of the five alkali metals, and with the exception of the amalgam with lithium, the highest melting point compound in each series has the formula MHg ₂. Lithium and sodium also form compounds with cadmium and zinc.

Formation of complexes

Until the late 1960s there were few complexes of the alkali metal cations with organic molecules. Specialized biological molecules such as valinomycin were known to complex selectively the potassium cation K⁺ for transport across cell membranes , but synthetic ionophores (molecules that can form complexes with ions) were rare. All the alkali cations have a charge of +1 and, except for lithium , are chemically similar and rather inert. The only significant difference between one alkali cation and another is the size.

The synthesis of crown ethers by American chemist Charles J. Pedersen in 1967 provided size-selective cyclic molecules consisting of ether oxygens forming a ring or "crown" that could complex a cation of the right size to fit into the hole in the centre of the molecule. In some cases two crown ether molecules can encapsulate a cation in a "sandwich" fashion. For example, K⁺ just fits into the centre of an 18-crown-6 ring (18 atoms in the ring, 12 of which are carbon atoms and 6 are ether oxygen atoms) to form a 1:1 complex (that is, 1 cation:1 crown ether), K⁺(18C6). Cs⁺ is too large to fit into the ring but can be complexed on one side to form the Cs⁺(18C6) complex or can be sandwiched between two 18-crown-6 molecules to form the 1:2 complex, Cs⁺(18C6) ₂. Thus, the selectivity of a crown ether for a particular cation depends on the ring size. Common crown ethers are 12-crown-4, 15-crown-5, and 18-crown-6. These molecules are selective for Li⁺, Na⁺, and K⁺, respectively.

Even greater affinity for alkali cations was achieved by the synthesis of cryptands by French chemist Jean-Marie Lehn in 1968 and spherands by American chemist Donald Cram in 1979. These are three-dimensional molecules with an internal cavity or crypt that can completely encapsulate the alkali cation. By synthesizing molecules with different cavity sizes, the selectivity for particular cations over those of the "wrong" size to fit in the cavity can be controlled. It should be noted, however, that these molecules are not rigid and that flexibility of the framework can alter the cavity size to accommodate alkali cations of different sizes, although with differences in the strength of complexation.

Since the initial syntheses of crown ethers and cryptands, thousands of complexants for cations of various sizes, charges, and geometries have been synthesized. This has led to an entirely new branch of chemistry called supramolecular chemistry.

Analytical chemistry of the alkali metals

Classical methods of separation and analysis of alkali metals are rather difficult and time consuming. For lithium they include such procedures as selective extraction of lithium chloride into organic solvents and the detection of lithium with azo dyes that give highly sensitive colour reactions in alkaline solutions. A modification of the uranyl acetate test (the precipitation of an insoluble sodium salt with uranyl acetate) has been used as a standard test for the presence of sodium. The use of a cobaltinitrite solution permits separation of potassium from sodium by precipitation of the insoluble potassium salt. There are essentially no satisfactory analytical methods for rubidium and cesium based on the use of reagents in solution.

Classical methods of separation of the alkali metals have been largely supplanted by chromatographic elution. Strongly acidic cation -exchange resins and aqueous acidic solutions are used. Generally the affinity increases with atomic weight so that the ions are eluted in the order Fr⁺ > Cs⁺ > Rb⁺ > K⁺ > Na⁺ > Li⁺, which is the order of decreasing size of the hydrated ions. Ion-exchange resins that are specific for lithium have been developed. Macrocyclic compounds such as crown ethers and cryptands that are selective for particular alkali metal ions have been synthesized. They form cationic complexes that can be dissolved in organic solvents such as chloroform (CHCl ₃) with counterions such as picrate (C ₆H ₂[NO ₂] ₃O⁻).

The characteristic flame colours of the alkali metals (red, yellow, violet, red, and blue for Li, Na, K, Rb, and Cs, respectively) are qualitative indicators of the modern analytical methods used to determine the concentrations of alkali-metal salts in aqueous solution. The intensities of the characteristic spectral lines in emission after excitation by a flame or ICP (inductively coupled plasma) give quantitative measures of the individual alkali metal concentration in the parts per million range or lower. Determination of one alkali metal in the presence of another, however, can result in interference, which can be reduced by using specially prepared standard solutions that contain known amounts of the interfering metals.

The analysis of the alkali-metal samples for the presence of nonmetallic elements, such as oxygen , carbon , hydrogen , and nitrogen , requires specialized techniques. The oxygen content of sodium and potassium samples can be determined by extraction of the free alkali metal with mercury , leaving behind mercury-insoluble oxides and carbonates , which can subsequently be analyzed by means of solution methods. The oxygen content of rubidium and cesium can be accurately determined by precise measurement of the freezing point of these two elements.

The carbon content of alkali metals can be analyzed by oxidation of the alkali metal in pure oxygen, followed by infrared measurement of the carbon dioxide generated during combustion. For the analysis of nitride in lithium, the nitride commonly is converted to ammonia , and the ammonia is measured by colorimetric analysis.
Frederick Tepper James L. Dye

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alkali metal | Definition, Properties, & Facts

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Alkali metal , any of the six chemical elements that make up Group 1 (Ia) of the periodic table -namely, lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), and francium (Fr). The alkali metals are so called because reaction with water forms alkalies (i.e., strong bases capable of neutralizing acids ).