These minerals represent dissolved salts that have been removed from sea water by evaporation and are mined for table salt and other uses. Interaction with pore water can affect the chemistry of the ocean. Pore water is sea water that has been trapped between sediment grains. The chemistry of the pore water is susceptible to change by biological processes.
For example, bacteria in the sediment consume organic tissue, at the same time using up much or even all the oxygen in the pore water. These processes are slow but continuous, and in the long term affect the chemistry of sea water. Many questions concerning the chemistry of sea water remain unanswered. In recent years, work on the composition of tiny sea-water bubbles included in salt crystals from ancient evaporites suggests that the composition of sea water may have changed over the past million years. The changes amount to about a factor of two for several of the major elements; these observations will lead to much future research on the processes that establish ocean chemistry and their changes over time.
Broecker, Wallace. Tracers in the Sea. Chester, Roy Marine Geochemistry, 2nd ed. London, U. Dasch, E.
Elements as Building Blocks
Julius, ed. Encyclopedia of Earth Sciences. Libes, Susan. An Introduction to Marine Biogeochemistry. Pilson, Michael E. An Introduction to the Chemistry of the Sea. Toggle navigation. Addition—Removal Processes and Considerations The composition of sea water is controlled by many different processes, all acting at the same time, and adding and removing substances at different rates see figure.
Residence Time. Hydrothermal Processes. Biological Processes. Pore-Water Interactions. Bibliography Broecker, Wallace. User Contributions: 1. Clain Jones.
According to their shared physical and chemical properties, the elements can be classified into the major categories of metals , metalloids and nonmetals. Metals are generally shiny, highly conducting solids that form alloys with one another and salt-like ionic compounds with nonmetals other than noble gases. A majority of nonmetals are coloured or colourless insulating gases; nonmetals that form compounds with other nonmetals feature covalent bonding.
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In between metals and nonmetals are metalloids, which have intermediate or mixed properties. Metal and nonmetals can be further classified into subcategories that show a gradation from metallic to non-metallic properties, when going left to right in the rows. The metals may be subdivided into the highly reactive alkali metals, through the less reactive alkaline earth metals, lanthanides and actinides, via the archetypal transition metals, and ending in the physically and chemically weak post-transition metals.
Nonmetals may be simply subdivided into the polyatomic nonmetals , being nearer to the metalloids and show some incipient metallic character; the essentially nonmetallic diatomic nonmetals , nonmetallic and the almost completely inert, monatomic noble gases. Specialized groupings such as refractory metals and noble metals , are examples of subsets of transition metals, also known  and occasionally denoted.
Placing elements into categories and subcategories based just on shared properties is imperfect. There is a large disparity of properties within each category with notable overlaps at the boundaries, as is the case with most classification schemes. Radon is classified as a nonmetallic noble gas yet has some cationic chemistry that is characteristic of metals. Other classification schemes are possible such as the division of the elements into mineralogical occurrence categories , or crystalline structures. Categorizing the elements in this fashion dates back to at least when Hinrichs  wrote that simple boundary lines could be placed on the periodic table to show elements having shared properties, such as metals, nonmetals, or gaseous elements.
The electron configuration or organisation of electrons orbiting neutral atoms shows a recurring pattern or periodicity. The electrons occupy a series of electron shells numbered 1, 2, and so on. Each shell consists of one or more subshells named s, p, d, f and g. As atomic number increases, electrons progressively fill these shells and subshells more or less according to the Madelung rule or energy ordering rule, as shown in the diagram. The electron configuration for neon , for example, is 1s 2 2s 2 2p 6.
With an atomic number of ten, neon has two electrons in the first shell, and eight electrons in the second shell; there are two electrons in the s subshell and six in the p subshell. In periodic table terms, the first time an electron occupies a new shell corresponds to the start of each new period, these positions being occupied by hydrogen and the alkali metals.
Since the properties of an element are mostly determined by its electron configuration, the properties of the elements likewise show recurring patterns or periodic behaviour, some examples of which are shown in the diagrams below for atomic radii, ionization energy and electron affinity. It is this periodicity of properties, manifestations of which were noticed well before the underlying theory was developed , that led to the establishment of the periodic law the properties of the elements recur at varying intervals and the formulation of the first periodic tables.
Atomic radii vary in a predictable and explainable manner across the periodic table. For instance, the radii generally decrease along each period of the table, from the alkali metals to the noble gases ; and increase down each group. The radius increases sharply between the noble gas at the end of each period and the alkali metal at the beginning of the next period.
These trends of the atomic radii and of various other chemical and physical properties of the elements can be explained by the electron shell theory of the atom; they provided important evidence for the development and confirmation of quantum theory. The electrons in the 4f-subshell, which is progressively filled across the lanthanide series, are not particularly effective at shielding the increasing nuclear charge from the sub-shells further out.
The elements immediately following the lanthanides have atomic radii that are smaller than would be expected and that are almost identical to the atomic radii of the elements immediately above them. This is known as the lanthanide contraction. The effect of the lanthanide contraction is noticeable up to platinum element 78 , after which it is masked by a relativistic effect known as the inert pair effect.
The first ionization energy is the energy it takes to remove one electron from an atom, the second ionization energy is the energy it takes to remove a second electron from the atom, and so on. For a given atom, successive ionization energies increase with the degree of ionization. Electrons in the closer orbitals experience greater forces of electrostatic attraction; thus, their removal requires increasingly more energy. Ionization energy becomes greater up and to the right of the periodic table.
Large jumps in the successive molar ionization energies occur when removing an electron from a noble gas complete electron shell configuration. Similar jumps occur in the ionization energies of other third-row atoms. Electronegativity is the tendency of an atom to attract a shared pair of electrons. The higher its electronegativity, the more an element attracts electrons. It was first proposed by Linus Pauling in Hence, fluorine is the most electronegative of the elements, [n 5] while caesium is the least, at least of those elements for which substantial data is available.
There are some exceptions to this general rule. Gallium and germanium have higher electronegativities than aluminium and silicon respectively because of the d-block contraction. Elements of the fourth period immediately after the first row of the transition metals have unusually small atomic radii because the 3d-electrons are not effective at shielding the increased nuclear charge, and smaller atomic size correlates with higher electronegativity.
The electron affinity of an atom is the amount of energy released when an electron is added to a neutral atom to form a negative ion. Although electron affinity varies greatly, some patterns emerge. Generally, nonmetals have more positive electron affinity values than metals.
Chlorine most strongly attracts an extra electron. The electron affinities of the noble gases have not been measured conclusively, so they may or may not have slightly negative values. Electron affinity generally increases across a period. This is caused by the filling of the valence shell of the atom; a group 17 atom releases more energy than a group 1 atom on gaining an electron because it obtains a filled valence shell and is therefore more stable.
A trend of decreasing electron affinity going down groups would be expected.
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The additional electron will be entering an orbital farther away from the nucleus. As such this electron would be less attracted to the nucleus and would release less energy when added. In going down a group, around one-third of elements are anomalous, with heavier elements having higher electron affinities than their next lighter congenors. Largely, this is due to the poor shielding by d and f electrons.
A uniform decrease in electron affinity only applies to group 1 atoms. The lower the values of ionization energy, electronegativity and electron affinity, the more metallic character the element has. Conversely, nonmetallic character increases with higher values of these properties.