Metals

The Alkali Metals

Alkali metals are chemical elements from the s-block of the periodic table. They have homologous physical and chemical properties.

Learning Objectives

Recall the periodic trends observed in the alkali metals.

Key Takeaways

Key Points

  • The alkali metals are a group of chemical elements from the s-block of the periodic table with similar properties: they appear silvery and can be cut with a plastic knife.
  • Alkali metals are highly reactive at standard temperature and pressure and readily lose their outermost electron to form cations with charge +1.
  • All the discovered alkali metals occur in nature.
  • Most alkali metals have many different applications, such as rubidium and caesium atomic clocks, sodium-vapor lamps, and table salt.

Key Terms

  • lye: A strong caustic alkaline solution of potassium or sodium salts, obtained by leaching wood ashes. It is much used in making soap as well as in biodiesel.
  • alkali metal: Any of the soft, light, reactive metals of Group 1 of the periodic table; lithium, sodium, potassium, rubidium, cesium, and francium.
  • caesium atomic clock: A primary frequency standard in which electronic transitions between the two hyperfine ground states of caesium-133 atoms are used to control the output frequency.

The alkali metals are a group of chemical elements in the periodic table with the following physical and chemical properties:

  • shiny
  • soft
  • silvery
  • highly reactive at standard temperature and pressure
  • readily lose their outermost electron to form cations with a charge of +1

They can all be cut easily with a plastic knife due to their softness, and their shiny surface tarnishes rapidly in air due to oxidation. Because of their high reactivity, alkali metals must be stored under oil to prevent reaction with air. In the modern IUPAC nomenclature, the alkali metals comprise the group 1 elements, excluding hydrogen. All the alkali metals react with water, with the heavier alkali metals reacting more vigorously than the lighter ones.

image

Alkali Metals: Lithium is stored in oil because of its high reactivity.

Periodic Trends of Alkali Metals

The alkali metals are lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs),and francium (Fr). This group lies in the s-block of the periodic table, as all alkali metals have their outermost electron in an s-orbital. The alkali metals provide the best example of group trends in properties in the periodic table, with elements exhibiting similar properties. For instance, when moving down the table, all known alkali metals show:

  • increasing atomic radius,
  • decreasing electronegativity
  • increasing reactivity
  • decreasing melting and boiling points

In general, their densities increase when moving down the table, with the exception of potassium, which is less dense than sodium.

Reactions of Alkali Metals

Alkali metals react violently with water, halogens, and acids. The reactions release surprising amounts of heat and light. In a chemical equation, alkali metals are represented with an M. Here are some sample reaction equations:

  • Alkali metals react with oxygen to form oxides, which have a duller appearance and lower reactivity. The oxides are much less reactive than the pure metals.

[latex]4{ \text{M} }_{ (\text{s}) }+{ \text{O} }_{ 2(\text{g}) }\rightarrow 2{ \text{M} }_{ 2 }\text{O}[/latex]

  • The oxides react vigorously with water to form a hydroxide. The resulting hydroxides of these elements dissociate completely in water to form some of the strongest bases known. Sodium hydroxide (NaOH), also called lye, is an industrial-strength base.

[latex]{ \text{M} }_{ 2 }\text{O}_{ (\text{s}) }+\text{H}_{ 2 }{ \text{O} }\rightarrow 2\text{MOH}_{ (\text{aq}) }[/latex]

  • The pure alkali metal can also react directly with water. In this case, the metal is a basic anhydride. Gaseous hydrogen is released, which is flammable.

[latex]2{ \text{M} }_{ (\text{s}) }+2{ \text{H} }_{ 2 }\text{O}\rightarrow 2\text{MOH}_{ (\text{aq}) }+{ \text{H} }_{ 2(\text{g}) }[/latex]

  • Exposing an alkali metal to a halogen will cause an extremely exothermic reaction that results in an ionic salt. Almost every salt of an alkali metal is highly soluble in water. They form conducting solutions, proving their ionic nature.

[latex]2{ \text{M} }_{ (\text{s}) }+{ \text{Cl} }_{ 2(\text{g}) }\rightarrow 2\text{MCl}_{ (\text{s}) }[/latex]

Occurrence in Nature

All the discovered alkali metals occur in nature. Experiments have been conducted to attempt the synthesis of ununennium (Uue), which is likely to be the next member of the group if the attempt is successful. It is predicted that the next alkali metal after ununennium would be unhexpentium (Uhp), an element that has not yet received even attempts at synthesis due to its extremely high atomic number.

Applications of Alkali Metals

Most alkali metals have many different applications. Two of the most well-known applications of the pure elements are rubidium and cesium atomic clocks, of which cesium atomic clocks are the most accurate representation of time known as of 2012. A common application of the compounds of sodium is the sodium-vapor lamp, which emits very efficient light. Table salt, or sodium chloride, on the other hand, has been used since antiquity.

The Alkaline Earth Metals

The alkaline earth metals are chemical elements in the s-block of the periodic table with very similar physical and chemical properties.

Learning Objectives

Predict which oxidation state an alkaline earth metal will adopt.

Key Takeaways

Key Points

  • The alkaline earth metals are shiny, silvery-white, and somewhat reactive metals at standard temperature and pressure.
  • All the alkaline earth metals readily lose their two outermost electrons to form cations with a 2+ charge.
  • All of the alkaline earth metals except magnesium and strontium have at least one naturally occurring radioisotope.
  • Magnesium and calcium are ubiquitous and essential to all known living organisms.

Key Terms

  • Alkaline earth metals: A group of chemical elements in the periodic table with similar properties: shiny, silvery-white, somewhat reactive at standard temperature and pressure. They readily lose their two outermost electrons to form cations with charge +2.

Properties of Alkaline Earth Metals

The alkaline earth metals (beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra)) are a group of chemical elements in the s-block of the periodic table with very similar properties:

  • shiny
  • silvery-white
  • somewhat reactive metals at standard temperature and pressure
  • readily lose their two outermost electrons to form cations with a 2+ charge
  • low densities
  • low melting points
  • low boiling points

The alkaline earth metals comprise the group 2 elements. All the discovered alkaline earth metals occur in nature.

Reactions of Alkaline Earth Metals

All the alkaline earth metals have two electrons in their valence shell, so they lose two electrons to form cations with a 2+ charge. Most of the chemistry has been observed only for the first five members of the group; the chemistry of radium is not well established due to its radioactivity.

In chemical terms, all of the alkaline metals react with the halogens to form ionic alkaline earth metal halides. All the alkaline earth metals except beryllium also react with water to form strongly alkaline hydroxides which should be handled with great care. The heavier alkaline earth metals react more vigorously than the lighter ones.

The alkaline metals have the second-lowest first ionization energies in their respective periods of the periodic table. This is due to their low effective nuclear charges and the ability to attain a full outer shell configuration by losing just two electrons. The second ionization energy of all of the alkaline metals is also somewhat low.

Beryllium is an exception. It does not react with water or steam, and its halides are covalent. All compounds that include beryllium have a covalent bond. Even beryllium fluoride, which is the most ionic beryllium compound, has a low melting point and a low electrical conductivity when melted.

Here is the list of some of the common reactions of alkaline earth metals, where E = elements that act as reducing agents:

  • The metals reduce halogens to form ionic halides: [latex]\text{E}_{ (\text{s}) }+\text{X}_{ 2 }\rightarrow \text{EX}_{ 2(\text{s}) }[/latex]where X = F, Cl, Br or I
  • The metals reduce O2 to form the oxides:

[latex]2\text{E}_{ (\text{s}) }+\text{O}_{ 2 }\rightarrow 2\text{EO}_{ (\text{s}) }[/latex]

  • The larger metals react with water to produce hydrogen gas: [latex]\text{E}_{ (\text{s}) }+2\text{H}_{ 2 }\text{O}_{ (\text{l}) }\rightarrow \text{E}_{ (\text{aq}) }^{ 2+ }+2\text{OH}_{ (\text{aq}) }^{ - }+\text{H}_{ 2(\text{g}) }[/latex] where E = Ca, Sr or Ba
  • The metals undergo transmetallation reactions to exchange ligands: [latex]\text{Ae}+\text{Hg}{ \{ \text{N}(\text{SiMe}_{ 3 })_{ 2 }\} _{ 2 } }\rightarrow [\text{Ae}\{ { \text{N}(\text{SiMe}_{ 3 })_{ 2 }\} _{ 2 } }(\text{THF})_{ 2 }][/latex] where Ae = Ca, Sr, or Ba.

Alkaline Earth Metal Compounds

  • Alkylmagnesium halides (RMgX where R = hydrocarbon group and X = halogen) are used to synthetise organic compounds. Here’s an example: [latex]3\text{RMgCl}+\text{SnCl}_{ 4 }\rightarrow 3\text{MgCl}_{ 2 }+\text{R}_{ 3 }\text{SnCl}[/latex]
  • Magnesium oxide (MgO) is used as a material to refract furnace brick and wire insulation (melting point of 2852°C).
  • Calcium carbonate (CaCO3) is mainly used in the construction industry and for making limestone, marble, chalk, and coral.

Radioactivity

All of the alkaline earth metals, except magnesium and strontium, have at least one naturally occurring radioisotope: beryllium-7, beryllium-10, and calcium-41 are trace radioisotopes. Calcium-48 and barium-130 have very long half-lives and thus occur naturally. All isotopes of radium are radioactive.

Occurrence in Nature

Emerald is a naturally occurring compound of beryllium. Calcium and magnesium are abundant in the earth’s crust, making up several important rock forming minerals such as dolomite (dolostone) and calcite (limestone). The other non-radioactive members of the group are only present in smaller quantities. Deposits of each of these minerals are mined to extract the elements for further use. Radium, with a maximum half-life of 1,601 years, is only present in nature when it is resupplied by a decay chain from the radioactive decay of heavier elements.

image

Emerald: Emerald is a variety of beryl, a mineral that contains the alkaline earth metal beryllium. Beryllium only occurs naturally in combination with other elements in minerals.

Biological Role and Toxicity of Alkaline Earth Metals

Magnesium and calcium are essential to all known living organisms. They are involved in more than one role. For example, magnesium or calcium ion pumps play a role in some cellular processes. Magnesium functions as the active center in some enzymes, and calcium salts take a structural role in bones.

Strontium plays an important role in marine aquatic life, especially hard corals, which use strontium to build their exoskeletons. Strontium and barium have some uses in medicine. For example “barium meals” are used in radiographic imaging, while strontium compounds are employed in some toothpastes.

However, beryllium and radium are toxic. Beryllium’s low aqueous solubility means it is rarely available to biological systems. It has no known role in living organisms, and, when encountered by them, is usually highly toxic. Radium has a low availability and is highly radioactive, making it toxic to life.

Aluminum

Aluminum is a soft, silvery metal in the boron group of the periodic table.

Learning Objectives

Describe the properties of aluminum.

Key Takeaways

Key Points

  • Aluminum is a soft, lightweight, and malleable silvery metal that is not soluble in water.
  • The vast majority of compounds feature aluminum in the oxidation state 3+, but compounds with +1 and +2 oxidation states are known.
  • Aluminum has many known isotopes, whose mass numbers range from 21 to 42.
  • Aluminum is the most widely used non-ferrous metal and is mostly alloyed, which improves its mechanical properties.

Key Terms

  • aluminum: A metallic chemical element (symbol Al) with an atomic number of 13.
  • passivation: Refers to a material becoming “passive,” that is, being less affected by environmental factors such as air or water.

Physical Properties of Aluminum

Aluminum is:

  • relatively soft
  • durable
  • lightweight
  • ductile
  • malleable
  • appearance ranging from silvery to dull gray
  • not soluble in water under normal circumstances
  • nonmagnetic
  • does not ignite easily
  • capable of being a superconductor

Chemical Properties

Aluminum is resistant to corrosion due to the phenomenon of passivation. A thin surface layer of aluminum oxide is formed when the metal is exposed to air. This oxide layer protects the aluminum beneath the surface from further oxidation. Like many other metals, aluminum can also be oxidized by water to produce hydrogen and heat:

[latex]2\text{Al}\quad +\quad 3{ \text{H} }_{ 2 }\text{O}\quad \longrightarrow \quad { \text{Al} }_{ 2 }{ \text{O} }_{ 3 }+3{ \text{H} }_{ 2 }[/latex]

Although aluminum is extremely easily oxidized, it is possible to remove the oxide layer from a sample without it immediately reforming. The simplest and safest way is to connect a battery to the sample and perform electrolysis under either an inert atmosphere (like argon gas) or vacuum conditions.

The vast majority of aluminum compounds feature the metal in the oxidation state 3+. The coordination number of aluminum can vary, but generally Al3+ is tetra- or hexacoordinate. This means it will have 4 or 6 ligands.

Aluminum Halides: Use as Lewis Acids

Aluminum is a very reactive metal that readily reacts to product trivalent compounds. Its halides (AlF3, AlCl3, AlBr3 and AlI3) are common examples. Trivalent aluminum is electron-deficient and therefore exceptionally useful as a Lewis acid, particularly in organic synthesis.

Aluminum Hydrides and Organoaluminum Compounds

A variety of compounds of empirical formula AlR3 and AlR1.5Cl1.5 exist. These species usually feature tetrahedral Al centers. With large organic groups, triorganoaluminum exist as three-coordinate monomers, such as triisobutylaluminum.

The important aluminum hydride is lithium aluminum hydride (LiAlH4), which is used as a reducing agent in organic chemistry. It can be produced from lithium hydride and aluminum trichloride:

[latex]4\text{LiH}\quad +\quad \text{Al}{ \text{Cl} }_{ 3 }\quad \longrightarrow \quad \text{LiAl}{ \text{H} }_{ 4 }\quad +\quad 3\text{LiCl}[/latex]

General Use of Aluminum

Aluminum is the most widely used non-ferrous metal. Aluminum is almost always alloyed, which markedly improves its mechanical properties, especially when tempered. For example, the common aluminum foils and beverage cans are alloys of 92% to 99% aluminum. Some of the many uses for aluminum metal are in:

  • Transportation as sheet, tube, castings, etc
  • Packaging (cans, foil, etc. )
  • Construction (windows, doors, siding, building wire, etc. )
  • A wide range of household items, from cooking utensils to baseball bats, and watches
  • Street lighting poles, sailing ship masts, walking poles, etc.
  • Outer shells of consumer electronics, also cases for equipment (e.g. photographic equipment)
  • Electrical transmission lines for power distribution
  • Super purity aluminum, used in electronics and CDs
  • Heat sinks for electronic appliances, such as transistors and CPUs
  • Substrate material of metal-core copper clad laminates used in high brightness LED lighting
  • Powdered aluminum used in paint and in pyrotechnics
  • A variety of countries, including France, Italy, Poland, Finland, Romania, Israel, and the former Yugoslavia, have issued coins struck in aluminum or aluminum-copper alloys
image

Use of aluminum in transportation: Aluminum-bodied Austin “A40 Sports” (c. 1951).