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# 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, though listed in Group 1 due to its electronic configuration, 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 basic solutions (pH>7), which are also called alkaline solutions.

### Properties and Facts About Alkali Metals

Alkali metals are among 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 have an oxidation state of +1. These metals are characterized by their soft texture and silvery color. They also have low boiling and melting points and are less dense than most elements. Lithium, sodium, and potassium float on water because of their low density. All these characteristics can be attributed to these elements' large atomic radii and weak metallic bonding. Group 1 elements have a valence electron configuration of ns1 and are good reducing agents (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; they are readily mined because they have relatively low densities and do not sink.

### Alkali Metal Reactions

Reactions with hydrogen

All alkali metals react with hydrogen to form hydrides as follows:

$2K(l) + H_2(g) \rightarrow 2KH(s)$

Reactions with water

Alkali metals and water react violently to form strong bases and hydrogen gas according to the following general reaction:

$2M(s) + 2H_2O \rightarrow 2MOH(aq) + H_2(g)$

where M denotes an alkali metal. An example involving sodium is given below:

$2Na(s) + 2H_2O \rightarrow 2NaOH(aq) + H_2(g)$

The 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, resulting in an explosion like the one shown below:

Reactions with halogens

Alkali metals and halogens combine to form ionic salts with the general reaction:
$M(s) + X_2(g) \rightarrow M_2X(s)$

where M represents an alkali metal and X represents a halogen. For example

$2Na(s) + Cl_2(g) \rightarrow 2NaCl (s)$

Reactions with nitrogen

Only lithium reacts with nitrogen at room temperature as follows:

$6Li(s) + N_2(g) \rightarrow 2Li_3N(s)$

Reactions with oxygen

Alkali metals form multiple types of oxides, peroxides and superoxides when combined with oxygen:
• Oxide ion: O2-
•  compounds generally take the form M​2O: e.g. Li​2O
• Sodium forms Peroxides
• Peroxide Ion:  O22-
• compounds generally take the form M​2O2: e.g. Na​2O2
• Potassium, Cesium, and Rubidium form superoxides
• Superoxide ion: O2-
•  compounds generally take the form MO2: e.g. KO2

### Trends

• Electronegativity and Ionization energy increase across a period and up the group.
• Alkali metals have the lowest electronegativity and ionization energy
• Francium is the least electronegative element.
• Atomic radius increases from right to left across a period and down the group.
• Francium is the largest element
• Boiling points and melting points increase from the bottom of the group to the top of the group.
• Lithium and francium have the highest and lowest boiling points, respectively, in Group 1.

### Flame Colors

Each alkali metal has a specific, characteristic flame color. The colors are caused by the difference in energy among the valence shells of s and p orbitals (not d-d transitions), which corresponds to wavelengths of visible light. When the element is introduced into the flame, its outer electrons are excited to higher-energy orbitals. The electrons then relax back to lower-energy orbitals, emitting energy in the form of light. The different colors of light depend on how much energy or how far the electron falls to a lower energy level. The alkali metals' bright flame colors make them useful in firework manufacturing. Each has a unique color and is easily identifiable

Lithium                       Sodium                 Potassium

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

### 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
• Soft metal
• Low ionization energy
• Electron configuration: [He]2s1
• Often used in rechargeable batteries.
• Including 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
• 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
• Also 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 long 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
• Hardly any occurs 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

### 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: $$Li_2O_2 + H_2O \rightarrow ?$$
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. Complete and balance the following equation: $$Li(s)+N_2(g) \rightarrow ?$$
10. Compounds of the form  M​2O2 are formed from an alkali metal and what kind of oxygen ion?.

1. The melting point and boiling point generally decrease down Group 1 Alkali Metals, so sodium has a higher boiling point (97.72°C) than francium (27°C). Atomic size increases as you move down the group and the larger the atom, the greater the distance between neighboring atoms in a metal lattice. This results in a weaker lattice energy and lower thermal energy to separate the atoms (lower melting point)
2. False. Group 1 and 2 form ionic hydrides. P block elements forms molecular hydride and nitrogen is in P block.
3. [Kr] 5s1
4. K, Rb and Cs.
5. $$Li_2O_2 + H_2O \rightarrow 2LiOH + H_2O_2$$
6. Lithium
7. False. Rubidium is a stable element (or at least has stable isotopes)
8. False. Only lithium reacts directly with nitrogen
9. $$6Li(s)+N_2(g) \rightarrow 2Li_3N(s)$$
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

10:33, 2 Sep 2014

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