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4.4 Properties of the s-block elements
References: K&T 22.1 - 22.4; SAL 8.1 - 8.5; 9.1 - 9.8
This is an appropriate time in the course of our discussion to introduce the first section of the periodic table in detail. In
particular, our discussion of Brønsted acidity has meant a focus on the properities of the proton, the cationic form of the first
element, hydrogen. It is therefore natural to discus the chemistry of this element in a little bit more detail. The similarities
and differences between hydrogen and the other Group 1 elements will be considered next, and then after briefly considering
the closely-related Group 2 elements, we will go on to look at the structures of ionic compounds, which are of such great
importance for the elements of these groups.
4.4.1 Hydrogen, the unique element
Hydrogen is located on most periodic tables in Group 1, and this certainly fits with the electron configuration 1s1. But
there is a lot of evidence from chemical behavior that hydrogen has properties that bridge both Groups 1 and 17. Now if you
think about this, it is hard to think of two more disparate groups in the periodic table than these two: one is the most
electropositive, the other the most electronegative group in the periodic table. One readily gives up an electron to produce a
monocation, the other readily accepts and electron to form a monoanion. These considerations might imply that hydrogen has
chameleon-like variability of properties, but the reality is rather that hydrogen has distinctly intermediate properties between
the two groups.
The first indication that this is so is provided by its Pauling electronegativity of 2.20 (compare Cs, 0.79 and F, 3.98).
This puts hydrogen very close to elements such as boron, carbon, phosphorus, platinum, and in general the whole middle
section of the periodic table. This argument is emphasized by the periodic table graphic shown above. The reasons for this
placement are summarized in the following table.
Unique placement of hydrogen in the periodic table
Group 1 (alkali metals) Forms monopostive ion, H+ Is not a metal
Has a single s electron Does not react with water
Group 17 (halogens) Is a non-metal H– is unstable and reactive
Forms a diatomic molecule, H2 H2 is relatively unreactive, with a strong bond
Forms a mononegative anion, H
Isotopes of Hydrogen
Hydrogen has three isotopes, which are unique among the elements in having been given their own names and symbols.
The most abundant ("normal") form is sometimes called protium, 1 H . This isotope is 99.985% abundant, high enough that
for most purposes the presence of the other isotopes in natural sources of the element can be ignored. The other stable isotope
of hydrogen is deuterium, 1 H or 1 D , which is 0.015 % abundant.
The third is a radioactive form called tritium, 1 H or 1T , which is only 10–15% abundant, and has a short half-life of
only about 12 years. The fact that there is always a very small amount of tritium in natural sources of hydrogen (primarily in
water) is a consequence of the production of this isotope by cosmic bombardment of atmospheric gases. For example,
neutrons react with nitrogen atoms in the upper atmosphere to produce carbon and tritium:
7 N+ n→
6 C + 3T
The isotopes of hydrogen are unique in being so different in their mass (i.e. 1, 2 and 3 amu) that their compounds differ
significantly in physical properties, and even in chemical behavior, despite the common adage that isotopes have no effect on
chemical behavior. Thus, for example, D2O, heavy water, melts at 3.8° and boils at 101.4° It is in fact possible to
concentrate heavy water simply by repeated boiling of large quantities of pure water. Its density is about 10% higher than
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H2O, hence the name "heavy". As another example of this isotope effect, consider the following table depicting the properties
of the elemental forms of the three isotopes:
Physical properties of the isotopes of molecular hydrogen
Isotope Molar mass, g/mol Boiling point, K Bond energy, kJ/mol
H2 2.02 20.6 436
D2 4.03 23.9 443
T2 6.03 25.2 446
Tritium is used in a variety of tracer applications in science and technology. It is a low energy β-emitter which does
relatively little damage to tissue for medical tracer applications. However, the main market for tritium is as the fuel for the so-
called "Hydrogen bomb", which is really a tritium bomb, as the material of choice for the production of thermonuclear
warheads. The isolation of tritium from sea-water is prohibitively expensive, so the weapons industry produces its tritium
from the less-abundant lithium isotope 6Li by the reaction:
3 Li + n→
0 He + 3T
Production of tritium has been so aggressive that lithium that is commercially available has actually been artificially
depleted of its 6Li content, and the natural average atomic mass for lithium cannot be used in high-accuracy work with
compounds of this element. Tritium decays to produce the rare helium isotope 3He, which is of inestimable value for
cryogenic physics because of its use in extreme low-temperature refrigeration apparatus.
1 T→ He +
This process occurs for the material in all thermonuclear warheads, which means that the warheads need to be periodically
"topped up" with fresh tritium. At the same time, the valuable 3He is collected and so-far has been made available to scientists
at reasonable cost. This may be said to be the only good thing that has come out of the thermonuclear weapons industry; in
general the benefits from the nuclear industry, created originally for weapons production, have been substantial. We can
mention the generation of electricity (which despite the risks may have a significantly lower environmental load than
continued burning of fossil fuels and hydro-electric dam construction), the production of isotopes for scientific and medical
use, and the general understanding of the nuclear processes that are a significant part of the energetic processes occurring in
All three isotopes of hydrogen are important NMR nuclei. Nuclear magnetic resonance is the single most important
structural tool in modern chemistry, and is also used in a number of industrial and medical applications (for example, in
magnetic resonance imaging, MRI). 1H is the most sensitive nucleus known for NMR work, and is of extreme importance for
NMR properties of the isotopes of hydrogen
Isotope Abundance Nuclear spin Relative sensitivity
H 99.985% ½ 1.00
D 0.015%, enriched 1 1.45 × 10–6
T always enriched ½ 1.21
all applications of NMR. Deuterium is used as a substitute for hydrogen in 1H NMR precisely because its nuclei do not absorb
radio waves at the same frequency as 1H. Thus solvents enriched to 99% or better in 2D are typically used in 1H NMR
spectroscopy. However, deuterium is an NMR nucleus in its own right. It finds a variety of specialist applications where it is
desireable to have it present in lower abundances. Thus for example, Prof. Siminovitch in our Physics department studies
biopolymer membranes by deuterium NMR using samples that have significant, but still low, percentages of deuterium
substituted for some of the hydrogen atoms in the membranes. He obtains information about the mobility of the membranes
from the signals that he measures. Deuterium NMR is also used by chemists studying subsitution sites on molecules, where
specific hydrogen atoms are replaced quantitatively by 2D, and then located by 2D NMR spectroscopy.
Properties of Hydrogen
Hydrogen is a non-metal element, and exists as a diatomic molecule in the elemental form. In this it formally resembles
the Group 17 elements, all of which are molecular diatomic elements. However, the H–H bond is much stronger (436 kJ/mol)
than that of any of the halogens. This is the main reason that hydrogen is relatively unreactive. This is an important point.
We are so used to thinking of hydrogen as a dangerous explosive gas (e.g. the Hindenburg disaster, in which an airship filled
with volatile hydrogen to produce lift caught fire, and passengers were killed) that it is easy to lose sight of its relative low
reactivity with the majority of the elements. In fact, the reaction with O2 and F2 are some of the few spontaneous reactions of
hydrogen with the elements. Thus the reaction with nitrogen is much more difficult, and indeed the production of ammonia is
a very difficult process. Many thermodynamically favored hydrogen reactions, such as the addition to C=C double bonds
require the presence of catalysts which promote the breaking of the H–H bond.
Under standard conditions, hydrogen is a colorless gas. Its very low boiling point, 20.7 K, reflects its nonpolar character
and low molecular mass. It is, of course, the least dense gas known. It freezes at 14 K. Hydrogen is ninth in abundance in the
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earth’ crust, 0.9% by mass, where it occurs primarily in water and in fossil fuels. In its earliest history, H2 was mainly used
as a fuel. In the middle of the 19th century it was found that heating soft coal (in the absence of air) gave a gas that could be
used for cooking and lighting. This gas, called coal gas, contains about 20% H2 along with several lightweight hydrocarbons.
Because coal gas was useful, new methods were sought for its production. It was found that injecting water into a bed of red
hot coke produces a mixture of H2 and CO. This mixture is known as water gas or synthesis gas (syngas).
C(s) + H2O(g) à H2(g) + CO(g)
Water gas burns cleanly and can be handled readily. The amount of heat produced, however, is only about half that from the
combustion of coal gas, and its flame is nearly invisible. Furthermore, carbon monoxide is highly toxic and has no odor.
Despite its hazards, water gas was used as a cooking gas to some extent until about 1950, but only after adding another
material to make the flame luminous and a malodorous compound so that leaks would be detected. Although no longer used
as a fuel, there is renewed interest in syngas because recent chemical research has shown that it can be used to manufacture
Preparation of Hydrogen
About 300 billion L (STP) of hydrogen gas is produced worldwide in a year, and virtually all is used immediately in other
processes. The largest quantity of hydrogen is produced by the catalytic steam reformation of hydrocarbons. This process uses
methane, CH4, as the primary starting material. Methane reacts with steam at high temperature to give CO and H2.
CH4(g) + H2O(g) à 3 H2(g) + CO(g) ∆H rxn = +206 kJ
The reaction is rapid in the 900° to 1000° range and goes nearly to completion. More hydrogen can be obtained in a
second step in which the CO that is formed reacts with more water. This so-called water gas shift reaction is run at 400° to
500° and is slightly exothermic.
H2O(g) + CO(g) à H2(g) + CO2(g) ∆H rxn = –41 kJ
The CO2 formed in the process is removed by reaction with CaO (to give CaCO3), thus leaving fairly pure hydrogen.
Electrolysis of water is the cleanest method of H2 production (at least 99.9 % pure), and it provides a valuable byproduct,
high-purity O2. Electric energy is quite expensive, however, so this method is not generally used.
A number of reactions can be used in the laboratory to form hydrogen, one of the simplest of which is the reaction of a
metal with an acid. In 1783, Charles (of Charles’ law) used the reaction of sulfuric acid with iron to produce the hydrogen for
a lighter-than-air balloon. The reaction of aluminum with NaOH also generates hydrogen as one product. During World War
II, this method was used to obtain hydrogen to inflate small balloons for weather observation and to raise radio antennas.
Metallic aluminum was plentiful because it came from damaged aircraft. Finally, the third method shown below is perhaps the
most efficient way to synthesize H2 in the laboratory. It is also useful for removing traces of water from liquid compounds that
do not have a reactive OH group.
Methods for Preparing H2 in the Laboratory
1. Metal + Acid gives metal salt + H2
Mg(s) + 2 HCl(aq) à MgCl2(aq) + H2(g)
2. Metal + H2O or base gives metal hydroxide or oxide + H2
2 Na(s) + 2 H2O(l) à 2 NaOH(aq) + H2(g)
2 Fe(s) + 3 H2O(l) à Fe2O3 (s) + 3 H2(g)
2 Al(s) + 2 KOH(aq) + 6 H2O(l) à 2 KAl(OH)4(aq) + 3 H2(g)
3. Metal hydride + H2O gives metal hydroxide + H2
CaH2(s) + 2 H2O(l) à Ca(OH)2(s) + 2 H2(g)
Hydrogen combines chemically with virtually every other element except the noble gases. Three different types of
hydrogen-containing binary compounds (compounds made up from only two kinds of elements) are known: ionic hydrides,
covalent hydrides and metallic or interstitial hydrides. The later are an unusual type of compound in that they are frequently
non-stoichiometric. The distribution of the three major types of hydrides is shown in the periodic table below. The formulas
and structures of the element hydrides of the main group elements are often model systems for a wide variety of other
derivatives of these elements, including halogen compounds and hydrocarbon-derivatives (known generally as organometallic
compounds). Formally, the ionic hydrides are compounds of H– with metal cations. They are sometimes called the saline
hydrides, emphasizing the analogy to the salts of the halogen anions. The covalent hydrides are compounds where hydrogen
has H+ character, although of course these are defined to be covalent compounds. The metallic hydrides belong to an
intermediate category which is neither fully ionic nor covalent. In the most extreme cases, e.g. palladium, the metal seems to
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be able to absorb a large and variable range of hydrogen with little apparent change to the properties of the metal or obvious
change of appearance.
Classification of Element Hydrides Ionic hydrides
H Covalent hydrides He
2 13 14 15 16 17
Li Be Metallic hydrides B C N O F Ne
Na Mg Al Si P S Cl Ar
3 4 5 6 7 8 9 10 11 12
K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe
Cs Ba Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po As Rn
Lanthanides La Ce Pr Nd Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
Actinides Ac Th Pa U Np Pu Am Cm
Ionic metal hydrides form when H2 reacts with Group 1 and 2 metals.
2 Na(s) + H2(g) à 2 NaH(s)
Ca(s) + H2(g) à CaH2(s)
These ionic compounds contain the hydride ion, H– , in which hydrogen is in the –1 oxidation state. All ionic hydrides are
extremely reactive towards water, and are moisture sensitive. They are excellent drying agents for non-halogenated aprotic
organic solvents. They are powerful reducing agents capable of converting most element halides to element hydrides.
Covalent hydrides are formed with electronegative elements, such as the non-metals carbon, nitrogen, oxygen, and
fluorine. Here the formal oxidation number of the hydrogen atom is +1.
N2(g) + 3 H2(g) à 2 NH3(g)
F2(g) + H2(g) à 2 HF(g)
Metallic (interstitial) hydrides
Hydrogen is absorbed by many metals forming interstitial hydrides, in which hydrogen atoms reside in the spaces between
metal atoms (called interstices) in the crystal lattice. For example, when a piece of palladium metal is used as an electrode for
the electrolysis of water, the metal can soak up a thousand times its volume of hydrogen (at STP). Most interstitial hydrides
are non-stoichiometric, that is, the ratio of metal and hydrogen does not involve whole numbers. When interstitial hydrides
are heated, H2 can be driven out. Thus, these materials can be used to store H2, the same as a sponge can store water. This is
one way to store hydrogen for use as a fuel in automobiles
Intermolecular contacts are normally of much lower energy than bonds within molecules. But one
15 16 17
such contact is sufficiently strong and has sufficiently distinct properties to have been named a "bond".
N O F
That is the hydrogen bond. This refers to any intermolecular contact in which a proton is shared between 7 8 9
two atoms drawn from the most electronegative elements: F, O, N, and to a lesser extent S, Cl and Br. Its P S Cl
properties are a direct consequence of the unique electron configuration of hydrogen, which is 1s1. When 15 16 17
H is bonded to an electronegative element, electron density is drawn away exposing the nucleus, and this As Se Br
33 34 35
makes the hydrogen strongly susceptible to electron pair donors.
Examples of hydrogen bonding are the structure of solid HF and a comparison of the properties of the isomers dimethyl
ether and ethanol. The latter, for which hydrogen bonding is possible, has a boiling point over 100 degrees higher. Above all
else, hydrogen bonding affects the properties of water, giving is a whole range of unique properties. The open cage structure
of ice is also due to hydrogen bonding, and this makes the so-called ice-I structure less dense than the liquid at 0°C. This
makes ice float on water, a very, very important factor in making this planet habitable. In fact, the density of water
maximizes at about 4°C. Thus hydrogen bonding is behind most of the weather-controlling factors that water and ice cause
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on global and local weather phenomena. Ice-I is the form of ice that exists at the freezing point under 1 atm of applied
pressure. There perhaps as many as seven different crystal structures for ice which form at colder temperatures, and
particularly under high-pressures (such as exist at the bottoms of glaciers, etc.) The open-network crystal structure of ice-I is
depicted in the picture below. The hydrogen bonds are linear but asymmetrically distributed so that there remain long and
short O–H bonds in the structure.
One of the strongest known hydrogen bonds (169 kJ/mol) is found in the [F–H–F]– ion, a linear ion in which the proton
is centred between the to fluoride ions, and each F–H bond is 2.27 Å long. Most other hydrogen bonds are assymetrical, with
short distances of about 1.1 Å, and long contacts of 2.5 to 3.5 Å. These are often non-linear, but bond angles are usually large
and easily deformed.
Hydrogen is a key element in living organisms. In fact, the existence of life depends on two particular properties of
hydrogen: the closeness of the electronegativities of carbon and hydrogen and the ability of hydrogen to form hydrogen bonds
when covalently bonded to nitrogen or oxygen. The low polarity of the carbon-hydrogen bond contributes to the stability of
organic compounds in our chemically reactive world. Biological processes also rely on both polar and nonpolar surfaces, the
best example of the latter being the lipids. It is important to realize that nonpolar sections of biological molecules, usually
containing just carbon and hydrogen atoms, are just as significant as their polar regions.
Hydrogen bonding is a vital part of all
biomolecules. Proteins are held in shape by
hydrogen bonds that form cross-links between
chains. The strands of DNA and RNA, the genetic
material, are held together by hydrogen bonds as
well. But more than that, the hydrogen bonds in the
double helices are not random; they form between
specific pairs of organic bases. These pairs are
preferentially hydrogen bonded because the two
components fit together to give particularly close
approaches of the hydrogen atoms involved in the
hydrogen bonding. This bonding is illustrated at
right for the interaction between two particular base
units, thymine and adenine. It is the specific
matching that results in the precise ordering of the
components in the DNA and RNA chains, a system
that allows those molecules to reproduce themselves almost completely error-free.
All proteins depend on hydrogen bonding for their function as well. Proteins consist mainly of one or more strands of linked
amino acids. But to function, most proteins must form a compact shape. To do this, the protein strand loops and intertwines
with itself, being held in place by hydrogen bonds cross-linking one part of the strand to another.