A. Biologically important elements
There are 81 stable elements in nature. Fifteen of these are present in all living things, and a further 8 –10 are only found in particular organisms. The illustration shows the first half of the periodic table ,containing all of the biologically important elements. In addition to physical and chemical data, it also provides information about the distribution of the elements in the living world and their abundance in the human body. The laws of atomic structure underlying the periodic table are discussed in chemistry textbooks.
More than 99%of the atoms in animals’ bodies are accounted for by just four elements —hydrogen (H),oxygen (O),carbon © and nitrogen (N).Hydrogen and oxygen are the constituents of water, which alone makes up 60 –70%of cell mass (see p.196).Together
with carbon and nitrogen, hydrogen and oxygen are also the major constituents of the organic compounds on which most living processes depend. Many biomolecules also contain sulfur (S)or phosphorus §.The above macroelements are essential for all organisms.
A second biologically important group of elements, which together represent only about0 .5%oft he body mass, are present almost exclusively in the form of inorganic ions .
This group includes the alkali metals sodium (Na)and potassium (K),and the alkaline earth metals magnesium(Mg)and calcium(Ca). The halogen chlorine (Cl)is also always ionized in the cell. All other elements important for life are present in such small quantities that they are referred to as trace elements .These include transition metals such as iron(Fe),zinc (Zn),copper (Cu),cobalt (Co)and manganese (Mn).A few nonmetals, such as iodine (I)and selenium (Se),can also be classed as essential trace elements.
B. Electron configurations: examples
The chemical properties of atoms and the types of bond they form with each other are determined by their electron shells. The electron configurations of the elements are therefore also shown in Fig. A .Fig. B explains the symbols and abbreviations used. More detailed discussions of the subject are available in chemistry textbooks.
The possible states of electrons are called orbitals .These are indicated by what is known as the principal quantum number and by a letter —s,p ,o rd .The orbitals are filled one by one as the number of electrons increases. Each orbital can hold a maximum of two electrons, which must have oppositely directed “spins.” Fig. A shows the distribution of the electrons among the orbitals for each of the elements. For example, the six electrons of carbon (B1 )occupy the 1s orbital, the 2s orbital, and two 2p orbitals. A filled 1s orbital has the same electron configuration as the noble gas helium (He).This region of the electron shell of carbon is therefore abbreviated as “He” in Fig. A .Below this, the numbers of electrons in each of the other filled orbitals (2s and 2p in the case of carbon) ares hown on the right margin. For example, the electron shell of chlorine (B2 ) consists of that of neon (Ne) and seven additional electrons in 3s and 3p orbitals. In iron(B3 ),a transition metal of the first series, electrons occupy the 4s orbital even though the 3d orbitals are still partly empty. Many reactions of the transition metals involve empty d orbitals —e.g., redox reactions or the formation of complexes with bases.
Particularly stable electron arrangements arise when the outermost shell is fully occupied with eight electrons (the “octet rule ”).
This applies, for example, to the noble gases, as well as to ions such as Cl –(3s 2 3p 6 ) and Na + (2s 2 2p 6 ).It is only in the cases of hydrogen and helium that two electrons are already sufcient to fill the outermost 1s orbital.