Redox reactions oxidation numbers

Oxidation-reduction (redox) reactions involve the loss of electrons and increase in oxidation number (oxidation) by one substance or system, with an associated gain of electrons and decrease in oxidation number (reduction) by another substance or system. Thus for every oxidation reaction there must be a reduction reaction. The oxidation number of an atom represents the hypothetical charge an atom would have if the ion or molecule were to dissociate.46-47... 

Recall from Chapter 4 that an oxidation-reduction (redox) reaction involves a transfer of electrons from the reducing agent to the oxidizing agent, and that oxidation involves a loss of electrons (an increase in oxidation number) and reduction involves a gain of electrons (a decrease in oxidation number). 

As you can see, the oxidation number of bromine changed from —1 to 0, an increase of 1. At the same time, the oxidation number of chlorine changed from 0 to -1, a decrease of 1. Therefore, chlorine is reduced and bromine is oxidized. All redox reactions follow the same pattern. When an atom is oxidized, its oxidation number increases. When an atom is reduced, its oxidation number decreases. Note that there is no change in the oxidation number of potassium. The potassium ion takes no part in the reaction and is called a spectator ion. How would the reaction differ if you used zinc bromide (ZnBr2) instead of potassium bromide ... 

You can always recognize a redox reaction by analyzing oxidation numbers. First determine the oxidation number of each element wherever it appears in the reaction. If no elements change in oxidation numbers, the reaction is not an oxidation-reduction reaction. If changes do occur, the reaction is an oxidation-reduction reaction. Remember that oxidation and reduction must always occur together if some atoms increase in oxidation numbers, then others must decrease. 

Is the reaction redox If any atoms change their oxidation number, the reaction is redox. 

However, now there are changes in oxidation numbers H changes from +1 in HI to 0 in H2, and I changes from —1 in HI to 0 in fy. So this is also an oxidation-reduction (redox) reaction. 

Oxidation is defined as electron loss, and reduction as electron gain. In an oxidation-reduction (redox) reaction, electrons move from one reactant to the other the reducing agent is oxidized (loses electrons), and the oxidizing agent is reduced (gains electrons). Chemists use oxidation numbers, the numbers of electrons owned by the atoms in reactants and products, to follow these changes. 

Section 4.1 polar molecule (109) solvated (110) electrolyte (110) nonelectrolyte (112) Section 4.2 molecular equation (113) total ionic equation (114) spectator ion (114) net ionic equation (114) Section 4.3 precipitation reaction (115) precipitate (115) metathesis reaction (116) Section 4.4 acid-base reaction (117) neutralization reaction (117) acid (117) base (118) salt (119) titration (11 9) equivalence point (120) end point (120) Section 4.5 oxidation-reduction (redox) reaction (123) oxidation (124) reduction (124) oxidizing agent (124) reducing agent (124) oxidation number (O.N.) (or oxidation state) (124) Section 4.6 activity series of the metals (130)... 

The number of electrons lost or gained by an atom in forming an ionic bond is equal to its VALENCE. Atoms that readily lose electrons are said to be ELECTROPOSITIVE, those that readily gain electrons to be ELECTRONEGATIVE. The loss of electrons is called OXIDATION, and the gain of electrons is called reduction. By definition, the formation of an ionic bond from elements must involve an oxidation-reduction (REDOX) reaction. The more electropositive element is oxidized, and the more electronegative one is reduced. 

Olefin synthesis starts usually from carbonyl compounds and carbanions with relatively electropositive, redox-active substituents mostly containing phosphorus, sulfur, or silicon. The carbanions add to the carbonyl group and the oxy anion attacks the oxidizable atom Y in-tramolecularly. The oxide Y—O" is then eliminated and a new C—C bond is formed. Such reactions take place because the formation of a Y—0 bond is thermodynamically favored and because Y is able to expand its coordination sphere and to raise its oxidation number. 

In earlier sections of this chapter, we showed how to write and balance equations for precipitation reactions (Section 4.2) and acid-base reactions (Section 4.3). In this section we will concentrate on balancing redox equations, given the identity of reactants and products. To do that, it is convenient to introduce a new concept, oxidation number. 

The concept of oxidation number is used to simplify the electron bookkeeping in redox reactions. For a monatomic ion (e.g., Na+, S2 ), the oxidation number is, quite simply, the charge of the ion (+1, —2). In a molecule or polyatomic ion, the oxidation number of an element is a pseudo-charge obtained in a rather arbitrary way, assigning bonding electrons to the atom with the greater attraction for electrons. 

Nitrogen cannot have an oxidation number lower than —3, which means that when NH3 takes part in a redox reaction, it always acts as a reducing agent. Ammonia may be oxidized to elementary nitrogen or to a compound of nitrogen. An important redox reaction of ammonia is that with hypochlorite ion ... 

A redox reaction, therefore, is any reaction in which there are changes in oxidation numbers. 

K.9 For each of the following redox reactions, identify the substance oxidized and the substance reduced by the change in oxidation numbers. 

Balancing the chemical equation for a redox reaction by inspection can be a real challenge, especially for one taking place in aqueous solution, when water may participate and we must include HzO and either H+ or OH. In such cases, it is easier to simplify the equation by separating it into its reduction and oxidation half-reactions, balance the half-reactions separately, and then add them together to obtain the balanced equation for the overall reaction. When adding the equations for half-reactions, we match the number of electrons released by oxidation with the number used in reduction, because electrons are neither created nor destroyed in chemical reactions. The procedure is outlined in Toolbox 12.1 and illustrated in Examples 12.1 and 12.2. 

The sulfur-rich oxides S 0 and S 02 belong to the group of so-called lower oxides of sulfur named after the low oxidation state of the sulfur atom(s) compared to the best known oxide SO2 in which the sulfur is in the oxidation state +4. Sulfur monoxide SO is also a member of this class but is not subject of this review. The blue-green material of composition S2O3 described in the older literature has long been shown to be a mixture of salts with the cations S4 and Ss and polysulfate anions rather than a sulfur oxide [1,2]. Reliable reviews on the complex chemistry of the lower sulfur oxides have been published before [1, 3-6]. The present review deals with those sulfur oxides which contain at least one sulfur-sulfur bond and not more than two oxygen atoms. These species are important intermediates in a number of redox reactions of elemental sulfur and other sulfur compounds. 

Magnesium cations and oxide anions attract each other strongly, forming the ionic solid, MgO. Notice that in the balanced redox reaction there is no net change in the number of electrons two Mg atoms lose four electrons, and one O2 molecule gains four electrons. 

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