element oxidation state
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An oxide (ËˆÉ’ksaÉªd) is an anion of oxygen in the oxidation state of -2 or a chemical compound formally containing an oxygen in this state. Most of the Earth's crust consists of oxides. Oxides result when elements are oxidized by oxygen in air. Combustion of hydrocarbon s affords the
Most elements have more than one possible oxidation state with carbon having at least twelve, ...
In coordination chemistry, the oxidation number of a central atom in a coordination compound is the charge that it would have if all the ligands were removed along with the electron pairs that were shared with the central atom. Oxidation numbers are often confused with oxidation states.
The oxidation number is used in the nomenclature of inorganic compounds. It is represented by a Roman numeral. The oxidation number is placed either as a right superscript to the element symbol, for example FeIII, or in parentheses after the name of the element, iron(III): in the latter case, there is no space between the element name and the oxidation number.
Oxidation number versus oxidation state
The oxidation number is usually numerically equal to the oxidation state and so the terms oxidation state and oxidation number are often used interchangeably. However, oxidation number is used in coordination chemistry with a slightly different meaning since the rules used for counting electrons are different: every electron belongs to the ligand, regardless of electronegativity. Also, oxidation numbers are conventionally represented with Roman numerals while oxidation states are given in Arabic numerals. The oxidation number of a central atom may be part of the compounds name (for example iron(II,III) oxide); the oxidation state of atoms is not included in compound names.
The oxidation state can differ from the oxidation number in a few cases where the ligand atom is less electronegative than the central atom (e.g., in iridiumphosphine complexes), resulting in a formal oxidation state that is different from the oxidation number.
Spectroscopic oxidation states
Although formal oxidation numbers can be helpful for classifying compounds, they are unmeasurable and their physical meaning can be ambiguous. Formal oxidation numbers require particular caution for molecules where the bonding is covalent, since the formal oxidation numbers require the heterolytic removal of ligands, which essentially denies covalency. Spectroscopic oxidation states, as defined by Jorgenson and reiterated by Wieghardt, are measurables that are bench-marked using spectroscopic and crystallographic data.
oxidation and reduction complementary chemical reactions characterized by the loss or gain, respectively, of one or more electrons by an atom or molecule. Originally the term oxidation was used to refer to a reaction in which oxygen combined with an element or compound, e.g., the reaction of magnesium with oxygen to form magnesium oxide or the combination of carbon monoxide with oxygen to form carbon dioxide. Similarly, reduction referred to a decrease in the amount of oxygen in a substance or its complete removal, e.g., the reaction of cupric oxide and hydrogen to form copper and water. When an atom or molecule combines with oxygen, it tends to give up electrons to the oxygen in forming a chemical bond . Similarly, when it loses oxygen, it tends to gain electrons. Such changes are now described in terms of changes in the oxidation number, or oxidation state, of the atom or molecule (see valence ). Thus oxidation has come to be defined as a loss of electrons or an increase in oxidation number, while reduction is defined as a gain of electrons or a decrease in oxidation number, whether or not oxygen itself is actually involved in the reaction. In the formation of magnesium oxide from magnesium and oxygen, the magnesium atoms have lost two electrons, or the oxidation number has increased from zero to +2. This is also true when magnesium reacts with chlorine to form magnesium chloride. In solution, ferrous iron (oxidation number +2) may be oxidized to ferric iron (oxidation number +3) by the loss of an electron. In the reduction of cupric oxide the oxidation number of copper has changed from +2 to zero by the gain of two electrons. The two processes, oxidation and reduction, occur simultaneously and in chemically equivalent quantities. In the formation of magnesium chloride, for every magnesium atom oxidized by a loss of two electrons, two chlorine atoms are reduced by a gain of one electron each. Oxidation-reduction reactions, called also redox reactions, are most simply balanced in the form of chemical equations by arranging the quantities of the substances involved so that the number of electrons lost by one substance is equaled by the number gained by another substance. In such reactions, the substance losing electrons (undergoing oxidation) is said to be an electron donor, or reductant, since its lost electrons are given to and reduce the other substance. Conversely, the substance that is gaining electrons (undergoing reduction) is said to be an electron acceptor, or oxidant. Common reductants (substances readily oxidized) are the active metals, hydrogen, hydrogen sulfide, carbon, carbon monoxide, and sulfurous acid. Common oxidants (substances readily reduced) include the halogens (especially fluorine and chlorine), oxygen, ozone, potassium permanganate, potassium dichromate, nitric acid, and concentrated sulfuric acid. Some substances are capable of acting either as reductants or as oxidants, e.g., hydrogen peroxide and nitrous acid. The corrosion of metals is a naturally occurring redox reaction. Industrially, many redox reactions are of great importance: combustion of fuels; electrolysis (oxidation occurs at the anode and reduction at the cathode); and metallurgical processes in which free metals are obtained from their ores.
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Answers:In a compound, always assign 2- to oxygen. Almost always assign 1+ to hydrogen (unless it's bound to a metal or boron) Always give group I metals a charge of 1+ and group 2 metals a charge of 2+ NH4 is a polyatomic ion with a charge of 1+, so the N there has a charge of 3-. With that, you can work out the states of the others there. Let me know if you have issues with that.
Answers:Group 1 metals, +1 Group 2 metals, +2 Group 13 elements, +3 toward the top, +1 toward the bottom Group 14 elements, +4 for silicon and germanium, +4 and +2 for tin, primarily +2 for lead, +4 to -4 for carbon Group 15 elements large range, -3 to +5 Group 16 elements large range -2 to +6 Group 17 elements large range -1 to +7 except fluorine, only 0 and -1, He, Ne, Ar only 0 Kr only 0 and 2 Xe 0 through 8
Answers:The easiest way to work with these problems is to start with what you know. Na is +1 and each O atom = -2, so to balance everything out in a molecule with a +1 and a -8, we need Bi to = +7 Next, Mg is in group 2, so it is +2, O still = -2 so we have +4 from Mg and -14 from O leaving us with -10 worth of charge to deal with using P. Since there are two P atoms, we can divide the +10 by 2 to see that each P atom is +5. It works the same for the others. Generally, group 1 or column 1 gives up 1 electron, group 2, gives up 2 electrons. When it comes to the negatively charged stuff it helps to be familiar with the anions since there isn't always an easy way to know what's up with them. Transitional metals are a little different so be careful. They will often be +1, +2, +3 depending on how easy it is for them to lower their energy level by losing electrons from d or s orbitals. Good Luck.
Answers:You work on the assumption that H is always +1 in a compound and oxygen is -2 in a compound. Metals will have an oxidation state equal to to the "ionic" charge it would have if it were in solution as an ion. The oxidation number of polyatomic ions are well known. (i.e. sulfate is -2) Lastly, the oxidation number of any element in the elemental state is zero. The sum of the oxidation numbers of all of the elements in a comopund is zero. The sum of the oxidation numbers of all the elements in a polyatomic ion is equal to the charge on the ion. With all of that in mind we can easily assign oxidation numbers. +2 +6 -2 ...........0...........+1+6-2 ............. +3..+6-2 ....... +1-1 2FeSO4 (aq) + Cl2 (g) + H2SO4 (aq) ---> Fe2(SO4)3 + 2HCl