The alkaline earth metals or alkaline earths are a set of six elements found in the second group 2 of the periodic table. Each of these elements has two electrons in its outer electron shell, which gives them similar chemical properties and makes them distinct from other metal groups. These elements are all shiny, silvery-white, relatively soft metals that are somewhat reactive and typically form compounds rather than existing in their pure elemental form in nature.
Key Takeaways: Alkaline Earth Metals
- Alkaline earth metals are the elements in Group 2 of the periodic table.
- Each atom has two valence electrons in an outer s orbital.
- These elements form divalent cations (2⁺) and have a +2 oxidation state.
- They are reactive metals, though less reactive than alkali metals.
- They occur in compound form in nature, not in their pure state.
- Reactivity increases down the group, with radium being the most reactive.
List of the Alkaline Earth Metals
There are six alkaline earths. In order of increasing atomic number, they are:
⚛️ Note: Scientists predict that Element 120, once synthesized, will also be an alkaline earth metal based on its predicted electron configuration.
Element Profiles
Here are short, informative snapshots of each of the six alkaline earth metals:
Beryllium (Be)
- Atomic Number: 4
- Appearance: Hard, brittle, silver-gray metal
- Uses: Aerospace components, X-ray windows, nuclear reactors
- Biological Role: None
- Toxicity: Highly toxic when inhaled; causes lung disease
Magnesium (Mg)
- Atomic Number: 12
- Appearance: Silvery-white, lightweight metal
- Uses: Fireworks, aerospace alloys, dietary supplements
- Biological Role: Essential nutrient; supports enzyme function, muscle and nerve activity
- Toxicity: Generally low; excessive intake can cause gastrointestinal issues
Calcium (Ca)
- Atomic Number: 20
- Appearance: Dull gray or silver metal (tarnishes quickly)
- Uses: Cement, steelmaking, dietary supplements, biological signaling
- Biological Role: Critical for bone and teeth structure, blood clotting, and muscle contraction
- Toxicity: Essential; high intake may cause kidney stones or interfere with absorption of other nutrients
Strontium (Sr)
- Atomic Number: 38
- Appearance: Soft, silvery metal; tarnishes quickly
- Uses: Fireworks (red color), ferrite magnets, formerly in CRT glass
- Biological Role: Limited; strontium mimics calcium in bone tissue
- Toxicity: Strontium is relatively safe in small amounts; radioactive isotopes (e.g., Sr-90) are dangerous
Barium (Ba)
- Atomic Number: 56
- Appearance: Silvery-white, soft metal
- Uses: Barium sulfate for medical imaging, drilling fluids, fireworks (green color)
- Biological Role: None
- Toxicity: Soluble compounds are toxic; insoluble barium sulfate is considered safe for medical use
Radium (Ra)
- Atomic Number: 88
- Appearance: Silvery-white metal (glows faintly due to radioactivity)
- Uses: Formerly in luminous paint; now limited to scientific research and cancer radiotherapy
- Biological Role: None
- Toxicity: Extremely radioactive; accumulates in bone and poses a severe cancer risk
Alkaline Earth Metal Properties
Because all of these elements have two valence electrons (ns² configuration), they exhibit many shared physical and chemical properties:
Physical Properties
- Luster: All are shiny and silver-colored when freshly cut.
- Malleability and ductility: They can be shaped and stretched like other metals.
- Density: Generally low for metals (though higher than alkali metals).
- Hardness: Softer than many transition metals, but harder than alkali metals.
- Melting and boiling points: Moderate compared to other metals—higher than alkali metals but lower than transition metals.
Chemical Properties
- Valence electrons: Each atom has a full outer s-orbital (ns²).
- Charge and oxidation state: Commonly form 2⁺ cations with an oxidation state of +2.
- Electronegativity and electron affinity: Both are low, meaning they do not easily gain electrons.
- Reactivity:
- Reactivity increases down the group (Be < Mg < Ca < Sr < Ba < Ra).
- React with oxygen to form oxides (e.g., CaO, MgO).
- React with halogens (e.g., Cl₂, Br₂) to form halides (e.g., CaCl₂, MgCl₂).
- Water reaction:
- Beryllium and magnesium react slowly, if at all.
- Calcium, strontium, barium, and radium react readily with water to produce metal hydroxides and hydrogen gas.
Biological Role and Toxicity
While the alkaline earth metals are chemically similar, they vary widely in their importance to biological systems and their toxicity:
- Essential Elements:
- Calcium (Ca) and magnesium (Mg) play critical biological roles in living organisms. Calcium is essential for bones, nerve function, and blood clotting, while magnesium participates in over 300 enzymatic reactions.
- Conditionally Essential or Supplementary:
- Strontium (Sr) has some physiological effects similar to calcium. Certain strontium salts treat osteoporosis, though their use is limited and controversial.
- Toxic Elements:
- Beryllium (Be) is highly toxic if inhaled as dust or fumes, leading to chronic beryllium disease (a lung condition). It has no known biological role.
- Barium (Ba) compounds that are soluble (such as barium chloride or nitrate) are toxic and affect muscle and nerve function by disrupting potassium ion channels.
- Radium (Ra) is radioactive and highly hazardous. It accumulates in bones like calcium, increasing cancer risk through radiation exposure.
Why Are They Called “Alkaline Earth” Metals?
The name “alkaline earth metals” comes from the alkaline oxides they form. These oxides are:
- Basic (alkaline): They react with water to form strong bases (alkaline hydroxides).
- Earthy: In alchemical terms, “earth” referred to insoluble, non-metallic substances remaining after heating minerals.
These oxides include:
- Beryllia (BeO)
- Magnesia (MgO)
- Lime (CaO)
- Strontia (SrO)
- Baryta (BaO)
Their strong basic nature and mineral origins led to the group name.
Real-World Examples and Uses
- Magnesium: Used in fireworks, lightweight alloys, and in the human body for enzyme activity.
- Calcium: Essential for bones and teeth, and used in construction (e.g., cement and plaster).
- Strontium: Gives red color to fireworks and is used in some ceramics and magnets.
- Barium: Used in X-ray imaging of the digestive system (as barium sulfate).
- Radium: Formerly used in self-luminous paints; now of interest in cancer therapy.
- Beryllium: Used in aerospace materials and X-ray windows, though toxic.
FAQs About Alkaline Earth Metals
Q: Why are alkaline earth metals less reactive than alkali metals?
A: Alkaline earth metals have two valence electrons, so it takes more energy to remove them compared to the single valence electron in alkali metals.
Q: Which alkaline earth metal is the most reactive?
A: Radium is the most reactive, though it is radioactive and rarely studied directly.
Q: Why don’t we find these metals in their pure form in nature?
A: They readily react with oxygen, water, and other substances, forming stable compounds instead.
Q: Are alkaline earth metals safe to handle?
A: Beryllium and radium are toxic, and others can be reactive. Always follow lab safety guidelines.
Q: What is the trend in atomic and ionic size?
A: Both atomic and ionic radii increase down the group due to the addition of electron shells.
References
- Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
- Krebs, Robert E. (2006). The History and Use of our Earth’s Chemical Elements: A Reference Guide. Greenwood Publishing Group. pp. 65–81. ISBN 0-313-33438-2.
- Lide, David R. (2004). Handbook of Chemistry and Physics (84th ed.). CRC Press. ISBN 978-0-8493-0566-5.
- Miledi, R. (1966). “Strontium as a Substitute for Calcium in the Process of Transmitter Release at the Neuromuscular Junction”. Nature. 212 (5067): 1233–4. doi:10.1038/2121233a0
- Weeks, Mary Elvira (1932). “The discovery of the elements. X. The alkaline earth metals and magnesium and cadmium”. Journal of Chemical Education. 9 (6): 1046. doi:10.1021/ed009p1046