H 2 o. 2na(s) + cl 2 (g) →10. thus, redox reactions and redox

The chalcogens are the first group in the p block to have no stable metallic elements. All isotopes of polonium Pothe only metal in group 16, are radioactive, and only one element in the group, tellurium Tecan even be described as a semimetal. As in groups 14 and 15, the lightest element of group 16, oxygen, is found in nature as the free element.

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Of the group 16 elements, only sulfur was known in ancient times; the others were not discovered until the late 18th and 19th centuries. Sulfur is frequently found as yellow crystalline deposits of essentially pure S 8 in areas of intense volcanic activity or around hot springs. As early as the 15th century BC, sulfur was used as a fumigant in Homeric Greece because, when burned, it produces SO 2 fumes that are toxic to most organisms, including vermin hiding in the walls and under the floors of houses.

Hence references to sulfur are common in ancient literature, frequently in the context of religious purification. Though used primarily in the production of sulfuric acid, sulfur is also used to manufacture gunpowder and as a cross-linking agent for rubber, which enables rubber to hold its shape but retain its flexibility.

Oxygen was not discovered untilwhen the Swedish pharmacist Carl Wilhelm Scheele found that heating compounds such as KNO 3Ag 2 CO 3and HgO produced a colorless, odorless gas that supported combustion better than air.

Chapter 11.6: Types of Chemical Reactions

Oxygen is used primarily in the steel industry during the conversion of crude iron to steel using the Bessemer process. Another important industrial use of oxygen is in the production of TiO 2which is commonly used as a white pigment in paints, paper, and plastics.

The silvery-white metal had the same density as antimony but very different properties. Tellurium is used to color glass and ceramics, in the manufacture of blasting caps, and in thermoelectric devices. Berzelius was born into a well-educated Swedish family, but both parents died when he was young. He studied medicine at the University of Uppsala, where his experiments with electroshock therapy caused his interests to turn to electrochemistry.

Berzelius devised the system of chemical notation that we use today. In addition, he discovered six elements cerium, thorium, selenium, silicon, titanium, and zirconium. The heaviest chalcogen, polonium, was isolated after an extraordinary effort by Marie Curie.

Although she was never able to obtain macroscopic quantities of the element, which she named for her native country of Poland, she demonstrated that its chemistry required it to be assigned to group Marie Curie was awarded a second Nobel Prize in Chemistry in for the discovery of radium and polonium.

The same process that is used to obtain nitrogen from the atmosphere produces pure oxygen. Oxygen can also be obtained by the electrolysis of water, the decomposition of alkali metal or alkaline earth peroxides or superoxides, or the thermal decomposition of simple inorganic salts, such as potassium chlorate in the presence of a catalytic amount of MnO 2 :.

Sulfur is also recovered from H 2 S and organosulfur compounds in crude oil and coal and from metal sulfide ores such as pyrite FeS 2. Because selenium and tellurium are chemically similar to sulfur, they are usually found as minor contaminants in metal sulfide ores and are typically recovered as by-products.

Both of these elements are notorious for the vile odors of many of their compounds. With their ns 2 np 4 electron configurations, the chalcogens are two electrons short of a filled valence shell.

This tendency is greatest for oxygen, the chalcogen with the highest electronegativity. As with the other groups, the lightest member in the group, in this case oxygen, differs greatly from the others in size, ionization energy, electronegativity, and electron affinity, so its chemistry is unique.

Also as in the other groups, the second and third members sulfur and selenium have similar properties because of shielding effects.The chemical reactions we have described are only a tiny sampling of the infinite number of chemical reactions possible.

How do chemists cope with this overwhelming diversity? How do they predict which compounds will react with one another and what products will be formed? The key to success is to find useful ways to categorize reactions. Familiarity with a few basic types of reactions will help you to predict the products that form when certain kinds of compounds or elements come in contact.

Most chemical reactions can be classified into one or more of five basic types:. The general forms of these five kinds of reactions are summarized in Table It is important to note, however, that many reactions can be assigned to more than one classification, as you will see in our discussion.

The classification scheme is only for convenience; the same reaction can be classified in different ways, depending on which of its characteristics is most important. Oxidation—reduction reactions, in which there is a net transfer of electrons from one atom to another, and condensation reactions are discussed in this section.

Acid—base reactions and one kind of exchange reaction—the formation of an insoluble salt such as barium sulfate when solutions of two soluble salts are mixed together—will be discussed in Section Table The term oxidation The loss of one or more electrons in a chemical reaction.

The substance that loses electrons is said to be oxidized. When iron is exposed to air in the presence of water, for example, the iron turns to rust—an iron oxide. When exposed to air, aluminum metal develops a continuous, coherent, transparent layer of aluminum oxide on its surface.

In both cases, the metal acquires a positive charge by transferring electrons to the neutral oxygen atoms of an oxygen molecule. Because the metals have lost electrons to oxygen, they have been oxidized ; oxidation is therefore the loss of electrons. Conversely, because the oxygen atoms have gained electrons, they have been reducedso reduction is the gain of electrons. For every oxidation, there must be an associated reduction.Post a comment.

Where there is oxidation, there is always reduction --Chemistry is essentially a study of redox systems. Oxidation reaction and reduction reaction—. Addition of oxygen. Removal of oxygen. Redox reaction—.

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Reactions in which both oxidation and reduction reactions take place simultaneously. Reactions which involve change in oxidation number of the interacting species. Reactant that oxidize the other reactant and itself that get reduce in given reaction is known as oxidizing agent or oxidant. Reactant that reduce the other reactant and itself that get oxidized in given reaction is known as reducing agent or reductant.

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Oxidation number Oxidation number denotes the oxidation state of an element in a compound ascertained according to a set of rules formulated on the basis that electron in a covalent bond belongs entirely to more electronegative element.

Difference between oxidation no.

h 2 o. 2na(s) + cl 2 (g) →10. thus, redox reactions and redox

Rules to determine the oxidation number In elements, in the free or the uncombined state, each atom bears an oxidation number of zero. For ions composed of only one atom, the oxidation number is equal to the chargeon the ion. The oxidation number of oxygen in most compounds is —2. However, we come acrosstwo kinds of exceptions here. One arises in the case of peroxides and superoxides, the compounds of oxygen in which oxygen atoms are directly linked to each other.

While in peroxides e. The second exception appears rarely, i. In such compounds e. The number assigned to oxygen will depend upon the bonding state of oxygen but this number would now be a positive figure only. In all its compounds, fluorine has an oxidation number of —1. Other halogens Cl, Br, and I also have an oxidation number of —1, when they occur as halide ions in their compounds. Chlorine, bromine and iodine when combined with oxygen, for example in oxoacids and oxoanions, have positive oxidation numbers.

The algebraic sum of the oxidation number of all the atoms in a compound must be zero. In polyatomic ion, the algebraic sum of all the oxidation numbers of atoms of the ion must equal the charge on the ion.

Thus, the sum of oxidation number of three oxygen atoms and one carbon atom in the carbonate ion, CO 3 2— must equal. Trend of ox. In a period it increases 1to 7 from left to right.Magnesium is oxidised due to addition of an electronegative element, fluorine.

HOW TO PICK OUT REDOX REACTIONS?

Potassium ferricyanide is oxidised due to removal of the electropositive element, Potassium. Removal of electronegative element, chlorine from ferric chloride reduces Fe. Process explaining Redox Reaction.

There is also a concept of half reaction; oxidation and reduction reactions when separately mentioned are known as half reactions and together they form a redox reaction.

A redox reaction, can hence, be easily identified by breaking the reaction into two half reactions or accounting for the change of electron or change in oxidation number. This reaction can be redox if either A and B or both A and B are in the elemental form. Decomposition Reactions are the opposite of combination reactions. During a decomposition, a compound breaks into two or more components. If any one of the components is in its elemental form, the said decomposition reaction is a redox reaction.

The further classification includes:. Reaction in which an element shows both reduction and oxidation then that reaction is called Disproportionation Reaction. A requirement for this type of reaction is that reacting substance can exist in at least three oxidation states. Thus, the reacting substance has intermediate oxidation number and the products formed have both higher and lower oxidation states. In above example, focusing on oxygen, we can see that oxygen has its intermediate oxidation number of -1 in reacting substance and a higher and lower oxidation state of 0 and -1 respectively in products formed.

These reactions are precisely opposite of Disproportionation reaction. Here two recanting substances, each containing the same element but with a different oxidation number, will form a product in which the mentioned element is in another oxidation state.

This method uses half reactions involving oxidation and reduction and then balancing them individually only to add them together again to form the complete reaction. The following steps are involved in this task:. Step 1 : Write a skeleton equation for the reaction in ionic form to form a basic idea:. Step 3 : The key is to balance the atoms other than O and H in each of half reaction individually, here Fe and Cr.

In our example, oxidation half takes care of itself while for balancing the reduction half we multiply it by 2 thus balancing the Cr atoms.

Step 4 : As water is neutral, to balance our oxygen atoms we use H 2 O. Thus, we get:. Step 5 : For a reaction to be stable, it must be electrically neutral. Thus, we have to balance charges in both of our half reactions and the overall reaction.

h 2 o. 2na(s) + cl 2 (g) →10. thus, redox reactions and redox

The oxidation half reaction is thus rewritten to balance the charge:. The total positive charge on the left hand side of the equation is twelve while right hand side only have six, thus creating an unbalance in the charge. To nullify extra positive charge we add six electrons to left hand side with excess positive charge. Adding six electrons to left hand side causes imbalance of negative charge when both the half reactions are compared. Thus, we balance charges across both the equations by multiplying the oxidation half by six:.

Step 6 : Adding to half reactions, after all the above steps, we can get fully balanced ionic equation:. Step 7 : Now the last step involves manual verification; check whether both sides have the same number of a single type of atoms and the charges cancel out each other thus concluding the fully balanced equation.

The above example is discussed for reaction in acidic medium but for the reaction in a basic medium; first balance the atoms as is done in the acidic medium. Step 2 : Note down the oxidation numbers for Cr and S on the top of reaction thus giving an idea for the oxidising and reducing agent.Introduction The reaction which involve both oxidation and reduction reactions is called Redox reaction.

Examples of oxidation:. Redox Reactions In Terms Of Electron Transfer Reactions According to electronic concept, the processes which involves loss of electrons are called oxidation reactions. Similarly, processes which involve gain of electrons are called reduction reactions.

The atom which reduced, act as oxidising agent and the atom which oxidised act as reducing agent. Competitive Electron Transfer Reaction Place a strip of metallic zinc in an aqueous solution of copper nitrate. You may notice that the strip becomes coated with reddish metallic copper and the blue colour of the solution disappears.

Copper ion is reduced by gaining electrons from the zinc. Oxidation Number Oxidation number of an element may be defined as the charge which an atom of the element has or appears so have when present in the combined state in a compound.

The rules for calculation of oxidation number are: 1. In elements, in the free or the uncombined state, each atom bears an oxidation number of zero. Evidently each atom in H2 has the oxidation number zero. For ions composed of only one atom, the oxidation number is equal to the charge on the ion. The oxidation number of oxygen in most compounds is However, we come across two kinds of exceptions here.

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The second exception appears rarely, i. In such compounds e. The number assigned to oxygen will depend upon the bonding state of oxygen but this number would now be a positive figure only. In all its compounds, fluorine has an oxidation number of Other halogens Cl, Br, and I also have an oxidation number of-1, when they occur as halide ions in their compounds.

Chlorine, bromine and iodine when combined with oxygen, for example in oxoacids and oxoanions, have positive oxidation numbers. The algebraic sum of the oxidation number of all the atoms in a compound must be zero. In polyatomic ion, the algebraic sum of all the oxidation numbers of atoms of the ion must equal the charge on the ion. Thus, the sum of oxidation number of three oxygen atoms and one carbon atom in the carbonate ion, CO 3 2- must equal A term that is often used interchangeably with the oxidation number is the oxidation state.

Oxidation state of a metal is a compound is sometimes represented by Stock notation. According to this, the oxidation number is written as Roman numeral in parenthesis after the symbol of the metal in the molecular formula. In terms of oxidation number, oxidation may be defined as a chemical change in which there occurs an increase in the oxidation number of an atom or atoms.

h 2 o. 2na(s) + cl 2 (g) →10. thus, redox reactions and redox

Reduction may be defined as a chemical change in which there occurs a decrease in the oxidation number of an atom or atoms. Thus, a redox reaction may be defined as a reaction in which the oxidation number of atoms undergoes a change. Types Of Redox Reactions 1. Decomposition Reaction: Decomposition reactions are the opposite of combination reactions.An acid—base reaction is a chemical reaction that occurs between an acid and a base.

It can be used to determine pH. Their importance becomes apparent in analyzing acid—base reactions for gaseous or liquid species, or when acid or base character may be somewhat less apparent.

The first of these concepts was provided by the French chemist Antoine Lavoisieraround It is important to think of the acid-base reaction models as theories that complement each other. The first scientific concept of acids and bases was provided by Lavoisier in around The Lavoisier definition held for over 30 years, until the article and subsequent lectures by Sir Humphry Davy in which he proved the lack of oxygen in H 2 SH 2 Teand the hydrohalic acids.

However, Davy failed to develop a new theory, concluding that "acidity does not depend upon any particular elementary substance, but upon peculiar arrangement of various substances". InJustus von Liebig proposed that an acid is a hydrogen-containing compound whose hydrogen can be replaced by a metal.

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Liebig's definition, while completely empirical, remained in use for almost 50 years until the adoption of the Arrhenius definition. The first modern definition of acids and bases in molecular terms was devised by Svante Arrhenius.

The Arrhenius definitions of acidity and alkalinity are restricted to aqueous solutions, and refer to the concentration of the solvent ions. Under this definition, pure H 2 SO 4 and HCl dissolved in toluene are not acidic, and molten NaOH and solutions of calcium amide in liquid ammonia are not alkaline.

Overall, to qualify as an Arrhenius acid, upon the introduction to water, the chemical must either cause, directly or otherwise:. Conversely, to qualify as an Arrhenius base, upon the introduction to water, the chemical must either cause, directly or otherwise:. The reaction of an acid with a base is called a neutralization reaction.

The products of this reaction are a salt and water. In this traditional representation an acid—base neutralization reaction is formulated as a double-replacement reaction. For example, the reaction of hydrochloric acidHCl, with sodium hydroxideNaOH, solutions produces a solution of sodium chlorideNaCl, and some additional water molecules.

The modifier aq in this equation was implied by Arrhenius, rather than included explicitly. It indicates that the substances are dissolved in water. An acid—base reaction is, thus, the removal of a hydrogen ion from the acid and its addition to the base. The reception of a proton by a base produces its conjugate acidwhich is the base with a hydrogen ion added. An acid and a base react not to produce a salt and a solvent, but to form a new acid and a new base. The concept of neutralization is thus absent.

The calculation of pH under the Arrhenius model depended on alkalis bases dissolving in water aqueous solution.

Science – Oxidation and Reduction (Redox) Reaction – Chemical Reaction Types Part 6 – English

Water is amphoteric —that is, it can act as both an acid and a base.The seventh member of the group, francium Fr is radioactive and so rare that only 20 atoms of Fr may exist on Earth at any given moment [1]. All the alkali metals are soft and, except for Cs which is yellow, are silvery-gray in color. Lithium, sodium, potassium, rubidium, and cesium have a great many other properties in common.

Each has metallic properties such as good conduction of heat and electricity, malleability the ability to be hammered into sheetsand ductility the ability to be drawn into wires. The high thermal heat conductivity and the relatively low melting point for a metal of sodium make it an ideal heat-transfer fluid. It is used to cool certain types of nuclear reactors liquid-metal fast breeder reactors, LMFBRs and to cool the valves of high-powered automobile engines for this reason.

h 2 o. 2na(s) + cl 2 (g) →10. thus, redox reactions and redox

Some general properties of the alkali metals are summarized in the table below. All these metal atoms contain a singles electron outside a noble-gas configuration, and so the valence electron is-well shielded from nuclear charge and the atomic radii are relatively large. The large volume of each atom results in a low density—small enough that Li, Na, and K float on water as they react with it.

Because they readily donate electrons in this way, all the alkali metals are strong reducing agents. They are quite reactive, even reducing water. Weak attraction for the valence electron also results in weak metallic bonding, because it is attraction among nuclei and numerous valence electrons that holds metal atoms together.

22.6: The Other Group 16 Elements: S, Se, Te, and Po

Weak metallic bonding results in low melting points, especially for the larger atoms toward the bottom of the group. Cs, for example, melts just above room temperature.

Weak metallic bonding also accounts for the fact that all these metals are rather soft. Two other elements are found in group IA. Hydrogen, although many of its compounds have formulas similar to the alkali metals, is a nonmetal and is almost unique in its chemical behavior.

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Therefore it is not usually included in this group. Francium Fr is quite radioactive, and only small quantities are available for study; so it too is usually omitted.

Its properties, however, appear to be similar to those of Cs and the other alkali metals. Hydrogen gas is given off, which propels the the lithium metal across the water as it reacts. If the excess water is evaporated, the compound lithium hydroxide LiOH remains behind.

Thus the equation for this reaction is. The elements sodium, potassium, rubidium, and cesium also combine violently with water to form hydroxides. The equations for their reactions are. Since the alkali metals all react with water in the same way, a general equation may be written:.

In addition to their behavior when added to water, the alkali metals react directly with many elements. All combine swiftly with oxygen in air to form white oxide:. Potassium, rubidium, and cesium are sufficiently reactive that yellow superoxides whose general formula is MO 2 can be formed:. Unless the surface of a sample of an alkali metal is scraped clean, it will appear white or gray instead of having a silvery metallic luster.

This is due to the oxide, peroxide, or superoxide coating that forms after a few seconds of exposure to air. The following movie shows how a freshly cut piece of lithium is shiny, but dulls to gray when exposed to oxygen in the air. The video also focuses on another important property of alkali metals: they are soft, and easy to cut, compared to other metals. A dull gray oxidized cylinder of lithium metal is cut, revealing a shiny silvery surface.

After 1 minute, the surface has dulled, and after 10 minutes, the cut surface has returned to the dull gray of the rest of the lithium metal. Since the alkali metal is lithium, the only reaction with oxygen that occurs is:. These oxides, hydrides, hydroxides, and sulfides all dissolve in water to give basic solutions, and these compounds are among the strong bases. The peroxides and superoxides formed when the heavier alkali metals react with O 2 also dissolve to give basic solutions:.

Both of the latter equations describe redox as well as acid-base processes, as you can confirm by assigning oxidation numbers. Therefore disproportionation simultaneous oxidation and reduction of O 2 2— or O 2 — to the more common oxidation states of 0 in O 2 and —2 in OH — is possible.


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