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ChemWiki: The Dynamic Chemistry E-textbook > Inorganic Chemistry > Descriptive Chemistry > Main Group Elements > Group 1: The Alkali Metals

Group 1: The Alkali Metals

Alkali metals are the chemical elements found in Group 1 of the periodic table. The alkali metals include: Lithium (Li), Sodium (Na),  Potassium (K), Rubidium (RB), Cesium (Cs), and Francium (Fr). Hydrogen, while it appears to be listed within Group 1, is not included in the alkali metals since it rarely exhibits similar behavior. The word "alkali" received its name from the Arabic word "al qali," meaning "from ashes". These particular elements were given the name "Alkali" because they react with water to form hydroxide ions, creating a basic solution (pH>7), which are also called alkaline solutions.

Properties and Facts About Alkali Metals

Alkali metals are known for being some of the most reactive metals. This is due in part to their larger atomic radii and low ionization energies. They tend to donate their electrons in reactions and often have an oxidation state of +1. These metals are characterized as being extremely soft and silvery in color. They also have low boiling and melting points and are less dense than most elements. Li, Na, and K have the ability to float on water because of their low density. All of these characteristics can be attributed to the large atomic radii and weak metallic bonding these elements possess. Group 1 elements have a valence electron configuration is ns1 and are good reducing agents (meaning they are easily oxidized). All of the alkali metals are found naturally in nature, but not in their pure forms. Most combine with oxygen and silica to form minerals in the Earth and are readily mined as they are of relatively low densitys and thus do not sink.

Alkali Metal Reactions

1) With Hydrogen: all alkali metals react with hydrogen to form hydrides 

2K(l) + H2(g) → 2KH(s)

2) With Water: Alkali metals and water react violently to form strong bases and hydrogen gas.

General Reaction:

2M(s) + 2H2O → MOH(aq) + H2(g)

where M=alkali metal

example: 2Na(s) + 2H2O → 2NaOH(aq) + H2(g)

note: Reactivity with water increases as you go down the group. The explosive reaction of sodium with water. In this case, the exothermic reaction is enough to ignite the hydrogen gas that

3) With Halogens: Alkali metals and halogens combine to form ionic salts

General Reaction: M(s) + X(g)→ MX(s)

where M=alkali metal and X=halogen

example: Na+(s) + Cl-(g) → NaCl (s) 

4) With Nitrogen: only Lithium reacts with Nitrogen at room temperature

6Li(s) + N2(g) → 2Li3N(s)

5) With OxygenAlkali metals form multiple types of oxides, peroxides and superoxides when combined with oxygen:

  • Oxide ion= O2-
    •  compounds generally look like M2O
      • ex. Li2O
  • Sodium forms Peroxides
    • Peroxide Ion= O22-
      • compounds generally look like M2O2
        • ex. Na2O2
  • Potassium, Cesium, and Rubidium form superoxides
    • Superoxide ion=O2-
      •  compounds generally look like MO2
        • ex. KO2

Trends

  • Electronegativity and Ionization energy increase from LEFT TO RIGHT and BOTTOM TO TOP
    • Alkali metals have the lowest electronegativity and ionization energy
    • Francium is the least electronegative element.
  • Atomic radius increases from RIGHT TO LEFT and TOP TO BOTTOM
    • Francium is the largest element
  • Boiling points and melting points increase going BOTTOM TO TOP
    • Lithium has the highest boiling point and Francium has the lowest boiling point in Group 1. 

Uses

  • Sodium Vapor Lamps
  • Atomic Clocks
  • Table Salt

Flame Colors

All alkali metals have their own specific flame color. The colors are caused by the difference in energy among the valence shell of s and p orbitals, which corresponds to wavelengths of visible light. When the element is introduced into the flame, its outer electrons are excited and jump to a higher electron orbital. The electrons then fall and emit energy in the form of light. The different colors of light depend on how much energy or how far the electron falls back to a lower energy level. For this reason, they are often used in fireworks. Each alkali metal has a unique color and is easily identifiable

 

Group 1 Element Flame Color
Lithium Crimson
Sodium Golden Yellow
Potassium Red/Violet
Rubidium Blue/Violet
Cesium Blue/Violet

 

                              

 Lithium                       Sodium                 Potassium            

Elements of the Alkali Metal Group

Lithium

  • named after the Greek word for stone (lithos)
  • discovered in Sweden in 1817 
  • Atomic number: 3
  • Atomic weight: 6.941
  • the lightest and least dense of all alkali metals
  • highly reactive
  • a soft metal
  • has a low ionization energy
  • Electron configuration: [He]2s1
  • Often used in rechargeable batteries.
    • include those used in cell phones, camcorders, laptop computers, and cardiac pacemakers.

Sodium

  • named after the Latin word for soda, Natria
  • discovered in 1807 
  • Atomic number: 11
  • Atomic weight: 22.9897
  • soft silvery metal.
  • extremely reactive metal
  • Electron configuration: [Ne]3s1
  • used in nuclear reactors because of its low boiling point.
  • Sodium is reacted with chlorine to produce the ionic halide, NaCl
    • Sodium chloride is an important part of human diet
      • It is used during winter months to control the ice on the road.

Potassium

  • named after the word Potash
    • Potash: means that Potassium is an element contained in the compound
  • discovered in 1807
  • Atomic number: 19
  • Atomic Weight: 39.0983
  • one of the most abundant elements in the earth's crust
  • oxidizes easily
  • lavender flame color
  • Electron configuration: [Ar]4s1
  • used mostly to produce chemicals, such as fertilizers for use in agriculture.
    • Potassium is an important nutrient needed for plant growth.

Rubidium

  • named after the latin word for red, rubidius
  • Atomic number: 37
  • soft metal
  • reddish flame color
  • Electron configuration: [Kr] 5s1.
  • discovered in 1861
  • known to have about 26 isotopes
  • very large half life at an estimated 49 billion years 

Cesium

  • Atomic number: 55
  • forms a strong base with water
  • Atomic Weight: 132.91
  • discovered in 1860
  • often used as a catalyst in various hydrogenation organic reactions
  • low melting point
  • Electron configuration: [Xe]6s1.

Francium 

  • discovered in 1939
  • very radioactive
  • hardly any Francium occurring naturally in the earth's crust
  • Atomic number: 87
  • Electron configuration: [Rn]6s1
  • heaviest and most electropositive metal
  • has the lowest boiling point
    • melts at low temperatures.
  • most reactive of the alkali metals group

Problems

  1. Which alkali metal has a higher melting point, Sodium (Na) or Francium (Fr)? Explain .
  2. True or False. NH3 is an ionic hydride.
  3. What is the electron configuration of Rubidium?
  4. Which alkali metals form superoxides?
  5. Complete and balance the following equation:   Li2O2 + H2O → ? 
  6. Which element is the most electronegative: Francium, Potassium, or Lithium?
  7. True or False: Rubidium has a very short half-life and decays quickly.
  8. True or False: All alkali metals react with Nitrogen.
  9. Balance the following equation: Li(s)+N2(g)→ ?
  10. Compounds that generally look like M2O2 are formed with a metal and what kind of oxygen ion?.

Answers

Highlight to see the answers. 

  1. Sodium has a higher boiling point because it has a larger atomic radius that Francium. Greater atomic radius means a bigger molecule thus having a higher boiling point.
  2. False. Group 1 and 2 form ionic hydrides. P block forms molecular hydride and Nitrogen is in P block.
  3. [Kr] 5s1
  4. K, Rb and Cs.
  5. Li2O2 + H2O 2LiOH + H2O2
  6. Lithium
  7. False
  8. False, only Lithium reacts with Nitrogen
  9. 6Li(s)+N2(g)→2Li3N(s)
  10. Peroxide Ion= O22-

References

  1. Fresenius, C. Remigius. Manual Qualitative Chemical Analysis. J. Wiley & Sons, 1897. Page 430
  2. Massey, A. G. Main Group Chemistry. London: Ellis Horwodd, 1990. Print.
  3. Petrucci et al., General Chemistry, Principles & Modern Applications, by Macmillan Publishing Company,Ninth Edition,Page 877.

Contributors

  • Kellie Berman, Nilshita Devi

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Last Modified
11:27, 27 Mar 2014

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