Humboldt State University ® Department of Chemistry

Richard A. Paselk

Chem 110	General Chemistry	Fall 2003
Lecture Notes::Lec 24_27 October	© R. Paselk 2003
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The Chemistry of the Elements

The Representative Elements

Hydrogen:

• Most abundant element in the Universe (over 90% of the atoms in the Universe, 75% by mass, about 25% helium, everything else is trace).
  • Rare in Earth's atmosphere because it is too light to be held by Earth's gravity.
  • Common in the Earth's crust (includes oceans etc.): 0.9% by mass, but 15% by atom, making it the third most abundant atom after oxygen and silicon.
• Three isotopes: H (¹H, 99.985 atom%), D (²H, 0.015 atom%), T (³H, 10^-17 atom%, t_1/2 = 12y).
  • Large isotope effects due to large mass differences.
    • fp of deuterium oxide is 3.8°C.
    • pD of pure deuterium oxide is 7.35.
    • rates of reactions differ significantly between H and D versions of compounds.
• Usually prepared by single displacement redox reactions.

  Most commonly an acid is used to react with a metal. e.g.:

• Note that any metal with a negative reduction potential can be used in principle.
• It can also be prepared by the reaction of water with calcium:

• Of course the alkali metals also react with water to give hydrogen, but too reactive for practical purposes, igniting the hydrogen gas produced.
• Hydrogen gas is generally rather unreactive due to the strong H₂ bond (DH° = 436kJ/mol for H₂ Æ 2 H). That is must have a high temperature to break the bond or a catalyst such as Pt or Pd.
  • Of course a spark or other source of heat can ignite hydrogen oxygen mixtures, releasing enough energy to continue the reaction (2 H₂ + O₂ Æ 2 H₂O; DG° = -229kJ/mol
• More compounds contain H than any other element.
  • Oxidation states of +1 and -1.
  • Major industrial use is ammonia synthesis.
  • Alkali and alkaline earths react directly with hydrogen to produce metal hydrides (e.g. NaH, KH, etc.). A variety of transition metals also form metal hydrides, many of which are non-stoichiometric, that is there is not a fixed small whole number ratios. This occurs when the hydrogens fit into cavities in the metal crystal.

Group I - The Alkali Metals

• Very reactive metals with electronic configurations of: [noble gas]ns¹.
  • Easily lose outermost electron to attain a noble gas configurations.
  • Outermost electron is also easily excited to a higher energy state, giving them intense flame colors.
    • s Æ p transition, with very slightly different p energies gives a doublet for the line.
• Group I elements are very metallic, held by metallic bonding with little or no covalent bond character
  • As a result very soft metals.
    • Li is exception, get some covalency due to small size.
  • Note properties in table below, discuss mp, density, ionization energies, reduction potentials and chemistry etc.

Properties of Group I - the Alkali Metals

Property		Li	Na	K	Rb	Cs	Fr
Valence-shell electron configuration		2s1	3s1	4s1	5s1	6s1	7s1
Melting point (°C)		180	98	64	39	29	-
Density (g/cm3)		0.54	0.97	0.86	1.5	1.9	-
Ionization energies - 1st & 2nd (kJ/mol) M(s) Æ M+(aq) + e-		520 7296	496 4563	419 3069	403 2650	376 2420	370 2170
Standard Reduction Potentials (V, 25°C) M+(aq) + e- Æ M(s)		-3.04	-2.71	-2.92	-2.92	-2.93	-
Electronegativity		1.0	0.9	0.8	0.8	0.7	0.7

Chemistry of the Alkali Metals

• Preparation: All of the alkali metals can be prepared by the electrolytic reduction.
  • Chemical reduction can be used for production of K, Rb, & Cs
    • Rb₂CO_3(s) + 2C_(s) Æ 2Rb_(g) + 3CO_(g)
• Reactions: All of the alkali metals give ionic salts with the exception of a few lithium compounds.
  • Water: all react to give hydrogen gas and base (hence the name alkali metals), e.g.:
    • 2Na_(s) + 2H₂O_(l) Æ 2Na⁺_(aq) + 2OH^-_(aq) + H_2(g)
  • Oxygen: all react with gaseous oxygen, but give different products:
    • Lithium give an oxide: 4Li_(s) + O_2(g) Æ 2Li₂O_(s)
    • Sodium gives a peroxide: 2Na_(s) + O_2(g) Æ Na₂O_2(s)
    • Potassium, Rubidium and Cesium give superoxides: K_(s) + O_2(g) Æ KO_2(s)
  • Lithium is the only element on the Periodic chart which reacts with nitrogen at room temperature (other alkali metals react at higher temps), giving lithium nitride:
    • 6Li_(s) + 2N_2(g) Æ 2Li₃N_(s)
• Compounds: The alkali metals form a wide range of compounds.
  • The salts are generally soluble. Sodium and potassium salts are so soluble they generally cannot be readily analyzed by wet chemical means (they are usually analyzed by flame tests or electrochemical means).
  • The oxides and hydroxides are very basic.
  • Only lithium salts typically crystallize as hydrates.
  • Alkali metals occur along with other metals in a wide range of minerals.

Group II - the Alkaline Earth Metals

Let's begin our look at the Group II elements with some physical properties:

• Valence shell electron configurations: all ns².
• Ionization energies and oxidation states: If you look at the values in the table below one might expect to see +1 as well as +2 as an oxidation state, or, even +3. Why don't these values show up?
  • If we were to look at the atoms we would see that, like Na the atom has a diffuse valence cloud of electrons surrounding a tight core. A dot density image of a beryllium atom is shown below:

  

  • If the Be atom loses a single electron it does not reduce in size all that much, since there is still an electron in the valence shell, as a result the ion cannot pack as tightly in compounds nor approach a negative counter ion as closely as the +2 core.
  • The difference in energy due to being able to bring the ion much closer is greater than the energy to remove the second electron, so +1 ions are not stable. (The only place +1 ions of Group II are observed as "stable" species is in high vacuum situations, where ion concentrations are very low.)
  • A similar argument holds in solutions, where the hydration energy of the +2 ion is much higher (the waters can pack closer), so the energy difference drives the formation of the +2 ion and the +1 ion is not stable.

Properties of Group II - the Alkaline Earth Metals

Property		Be	Mg	Ca	Sr	Ba	Ra
Valence-shell electron configuration		2s2	3s2	4s2	5s2	6s2	7s2
Melting point (°C)		1280	651	851	800	850	960
Density (g/cm3)		1.86	1.75	1.55	2.6	3.6	5.0
Ionization energies - 1st ,2nd & 3rd (kJ/mol) M(s) Æ M2+(aq) + 2 e-		899 1757 14849	738 1450 7730	590 1145 4941	549 1064 4207	503 965 3420	509 978 -
Standard Reduction Potentials (V, 25°C) M2+(aq) + 2 e- Æ M(s)		-1.69	-2.37	-2.87	-2.89	-2.91	-2.92
Electronegativity		1.5	1.2	1.0	1.0	0.9	0.9

© R A Paselk

Last modified 28 October 2003