Group 1 alkali metals: structure and reactions
Atomic structure and the periodic table • The periodic table
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Outer shell electron and ion formation
Group 1 atoms have one electron in their highest energy level. The single outer electron experiences weaker attraction to the nucleus than inner electrons because of shielding by inner shells. The weak attraction causes the outer electron to be lost easily, producing a stable +1 ion. Formation of a +1 ion explains chemical behaviour: ionic bonding with non-metals, basic oxides or hydroxides, and strong reducing character. The tendency to lose the outer electron increases as atomic radius increases down the group because the outer electron is further from the nucleus and more shielded.
Metallic structure and physical properties
Group 1 elements form metallic lattices in which positive ions sit in a sea of delocalised electrons. The single valence electron per atom gives relatively weak metallic bonding compared with metals that supply more valence electrons for bonding. Weaker metallic bonding causes low melting points, softness, and low density in the group. Trends in metallic bonding depend on atomic size and number of delocalised electrons. Larger atoms produce fewer delocalised electrons per unit volume and weaker bonding, so melting points and hardness tend to decrease down the group while atomic mass increases.
Predicting trends down the group
Reactivity with non-metals and water increases down the group. Cause: outer electron is further from the nucleus and more shielded, so nuclear attraction falls and electron loss becomes easier. Effect: lithium is least reactive of the three, potassium is most reactive. Melting and boiling points fall down the group because metallic bonding weakens as atomic size increases. Density generally increases down the group because atomic mass increases faster than atomic volume, but atomic packing and crystal structure can cause anomalies; predictions should allow for small exceptions.
Reactions with oxygen
Lithium forms a simple oxide when it reacts with oxygen: 4Li + O2 -> 2Li2O. Sodium reacts with oxygen to form a peroxide under typical conditions: 2Na + O2 -> Na2O2. Potassium reacts with oxygen to form a superoxide: K + O2 -> KO2. Observations on combustion differ: lithium burns with a red flame to give white oxide, sodium forms a white peroxide powder often with orange flame, and potassium burns with a lilac flame producing a yellow-white superoxide. The type of oxide reflects increasing ease of electron loss and stabilisation of different oxygen anions down the group.
Reactions with chlorine
Group 1 metals react vigorously with chlorine to form ionic chlorides. Lithium, sodium and potassium form halide salts with formula MCl where M is the metal: 2Li + Cl2 -> 2LiCl; 2Na + Cl2 -> 2NaCl; 2K + Cl2 -> 2KCl. Reactions are exothermic and produce white crystalline solids. The metal loses one electron; chlorine atoms gain one electron to form chloride ions. The ionic lattice stabilises the +1 and -1 ions formed.
Reactions with water
Group 1 metals react with water to give a metal hydroxide and hydrogen gas. Reaction general equation: 2M + 2H2O -> 2M(OH) + H2 (M = Li, Na, K). Lithium reacts steadily to form lithium hydroxide and hydrogen. Sodium reacts more vigorously, often melting to form a mobile, silvery ball and producing sodium hydroxide and hydrogen. Potassium reacts very vigorously and often ignites the hydrogen produced. Reactivity with water increases down the group because the outer electron is held less tightly, so electron transfer to water (producing H2) occurs more readily for larger atoms.
Key notes
Important points to keep in mind