Acid, Bases and Salt

                                      Bases

                  In chemistry, there are three definitions in common use of the word base, known as Arrhenius bases, Brønsted bases, and Lewis bases. All definitions agree that bases are substances which react with acids as originally proposed by G.-F. Rouelle in the mid-18th century.

Svante Arrhenius proposed in 1884 that a base is a substance which dissociates in aqueous solution to form hydroxide ions OH⁻. These ions can react with hydrogen ions (H⁺ according to Arrhenius) from the dissociation of acids to form water in an acid–base reaction. A base was therefore a metal hydroxide such as NaOH or Ca(OH)₂. Such aqueous hydroxide solutions were also described by certain characteristic properties. They are slippery to the touch, can taste bitter^[1] and change the color of pH indicators (e.g., turn red litmus paper blue).

In water, by altering the autoionization equilibrium, bases yield solutions in which the hydrogen ion activity is lower than it is in pure water, i.e., the water has a pH higher than 7.0 at standard conditions. A soluble base is called an alkali if it contains and releases OH⁻ ions quantitatively. Metal oxides, hydroxides, and especially alkoxides are basic, and conjugate bases of weak acids are weak bases.

Bases and acids are seen as chemical opposites because the effect of an acid is to increase the hydronium (H₃O⁺) concentration in water, whereas bases reduce this concentration. A reaction between aqueous solutions of an acid and a base is called neutralization, producing a solution of water and a salt in which the salt separates into its component ions. If the aqueous solution is saturated with a given salt solute, any additional such salt precipitates out of the solution.

In the more general Brønsted–Lowry acid–base theory (1923), a base is a substance that can accept hydrogen cations (H⁺)—otherwise known as protons. This does include aqueous hydroxides since OH⁻ does react with H⁺ to form water, so that Arrhenius bases are a subset of Brønsted bases. However, there are also other Brønsted bases which accept protons, such as aqueous solutions of ammonia (NH₃) or its organic derivatives (amines).^[2] These bases do not contain a hydroxide ion but nevertheless react with water, resulting in an increase in the concentration of hydroxide ion.^[3] Also, some non-aqueous solvents contain Brønsted bases which react with solvated protons. For example in liquid ammonia, NH₂⁻ is the basic ion species which accepts protons from NH₄⁺, the acidic species in this solvent.

G. N. Lewis realized that water, ammonia, and other bases can form a bond with a proton due to the unshared pair of electrons that the bases possess.^[3] In the Lewis theory, a base is an electron pair donor which can share a pair of electrons with an electron acceptor which is described as a Lewis acid.^[4] The Lewis theory is more general than the Brønsted model because the Lewis acid is not necessarily a proton, but can be another molecule (or ion) with a vacant low-lying orbital which can accept a pair of electrons. One notable example is boron trifluoride (BF₃).

Some other definitions of both bases and acids have been proposed in the past, but are not commonly used today.

General properties of bases include:

• Concentrated or strong bases are caustic on organic matter and react violently with acidic substances.
• Aqueous solutions or molten bases dissociate in ions and conduct electricity.
• Reactions with indicators: bases turn red litmus paper blue, phenolphthalein pink, keep bromothymol blue in its natural colour of blue, and turn methyl orange-yellow.
• The pH of a basic solution at standard conditions is greater than seven.
• Bases are bitter.^[5]

Reactions between bases and water

The following reaction represents the general reaction between a base (B) and water to produce a conjugate acid (BH⁺) and a conjugate base (OH⁻):^[3]

B_(aq) + H₂O_(l) ⇌ BH⁺_(aq) + OH⁻_(aq)

The equilibrium constant, K_b, for this reaction can be found using the following general equation:^[3]

K_b = [BH⁺][OH⁻]/[B]

In this equation, the base (B) and the extremely strong base (the conjugate base OH⁻) compete for the proton.^[6] As a result, bases that react with water have relatively small equilibrium constant values.^[6] The base is weaker when it has a lower equilibrium constant value.^[3]

Neutralization of acids

Ammonia fumes from aqueous ammonium hydroxide (in test tube) reacting with hydrochloric acid (in beaker) to produce ammonium chloride (white smoke).

Bases react with acids to neutralize each other at a fast rate both in water and in alcohol.^[7] When dissolved in water, the strong base sodium hydroxide ionizes into hydroxide and sodium ions:

NaOH → Na+
 + OH−

and similarly, in water the acid hydrogen chloride forms hydronium and chloride ions:

HCl + H
2O → H
3O+
 + Cl−

When the two solutions are mixed, the H
3O+
 and OH−
 ions combine to form water molecules:

H
3O+
 + OH−
 → 2 H
2O

If equal quantities of NaOH and HCl are dissolved, the base and the acid neutralize exactly, leaving only NaCl, effectively table salt, in solution.

Weak bases, such as baking soda or egg white, should be used to neutralize any acid spills. Neutralizing acid spills with strong bases, such as sodium hydroxide or potassium hydroxide, can cause a violent exothermic reaction, and the base itself can cause just as much damage as the original acid spill.

Alkalinity of non-hydroxides

Bases are generally compounds that can neutralize an amount of acids. Both sodium carbonate and ammonia are bases, although neither of these substances contains OH−
 groups. Both compounds accept H⁺ when dissolved in protic solvents such as water:

Na₂CO₃ + H₂O → 2 Na⁺ + HCO₃⁻ + OH⁻

NH₃ + H₂O → NH₄⁺ + OH⁻

From this, a pH, or acidity, can be calculated for aqueous solutions of bases. Bases also directly act as electron-pair donors themselves:

CO₃²⁻ + H⁺ → HCO₃⁻

NH₃ + H⁺ → NH₄⁺

A base is also defined as a molecule that has the ability to accept an electron pair bond by entering another atom's valence shell through its possession of one electron pair.^[7] There are a limited number of elements that have atoms with the ability to provide a molecule with basic properties.^[7] Carbon can act as a base as well as nitrogen and oxygen. Fluorine and sometimes rare gases possess this ability as well.^[7] This occurs typically in compounds such as butyl lithium, alkoxides, and metal amides such as sodium amide. Bases of carbon, nitrogen and oxygen without resonance stabilization are usually very strong, or superbases, which cannot exist in a water solution due to the acidity of water. Resonance stabilization, however, enables weaker bases such as carboxylates; for example, sodium acetate is a weak base.

Strong bases

A strong base is a basic chemical compound that can remove a proton (H⁺) from (or deprotonate) a molecule of even a very weak acid (such as water) in an acid–base reaction. Common examples of strong bases include hydroxides of alkali metals and alkaline earth metals, like NaOH and Ca(OH)
2, respectively. Due to their low solubility, some bases, such as alkaline earth hydroxides, can be used when the solubility factor is not taken into account.^[8] One advantage of this low solubility is that "many antacids were suspensions of metal hydroxides such as aluminium hydroxide and magnesium hydroxide."^[9] These compounds have low solubility and have the ability to stop an increase in the concentration of the hydroxide ion, preventing the harm of the tissues in the mouth, oesophagus, and stomach.^[9] As the reaction continues and the salts dissolve, the stomach acid reacts with the hydroxide produced by the suspensions.^[9] Strong bases hydrolyze in water almost completely, resulting in the leveling effect."^[7] In this process, the water molecule combines with a strong base, due to the water's amphoteric ability; and, a hydroxide ion is released.^[7] Very strong bases can even deprotonate very weakly acidic C–H groups in the absence of water. Here is a list of several strong bases:

Lithium hydroxide	LiOH
Sodium hydroxide	NaOH
Potassium hydroxide	KOH
Rubidium hydroxide	RbOH
Cesium hydroxide	CsOH
Magnesium hydroxide	Mg(OH) 2
Calcium hydroxide	Ca(OH) 2
Strontium hydroxide	Sr(OH) 2
Barium hydroxide	Ba(OH) 2
Tetramethylammonium hydroxide	N(CH 3) 4OH
Guanidine	HNC(NH 2) 2

The cations of these strong bases appear in the first and second groups of the periodic table (alkali and earth alkali metals). Tetraalkylated ammonium hydroxides are also strong bases since they dissociate completely in water. Guanidine is a special case of a species that is exceptionally stable when protonated, analogously to the reason that makes perchloric acid and sulfuric acid very strong acids.

Acids with a p K_a of more than about 13 are considered very weak, and their conjugate bases are strong bases.

Super bases

Main article: Superbase

Group 1 salts of carbanions, amides, and hydrides tend to be even stronger bases due to the extreme weakness of their conjugate acids, which are stable hydrocarbons, amines, and dihydrogen. Usually, these bases are created by adding pure alkali metals such as sodium into the conjugate acid. They are called superbases, and it is impossible to keep them in water solution because they are stronger bases than the hydroxide ion. As such, they deprotonate conjugate acid water. For example, the ethoxide ion (the conjugate base of ethanol) in the presence of water undergoes this reaction.

CH
3CH
2O−
 + H
2O → CH
3CH
2OH + OH−

Examples of common superbases are:

• Butyl lithium (n-C₄H₉Li)
• Lithium diisopropylamide (LDA) [(CH₃)₂CH]₂NLi
• Lithium diethylamide (LDEA) (C
  2H
  5)
  2NLi
• Sodium amide (NaNH₂)
• Sodium hydride (NaH)
• Lithium bis(trimethylsilyl)amide [(CH
  3)
  3Si]
  2NLi

Strongest superbases were only synthesised in gas phase:

• Ortho-diethynylbenzene dianion (C₆H₄(C₂)₂)²⁻ (This is the strongest superbase ever synthesized)
• Meta-diethynylbenzene dianion (C₆H₄(C₂)₂)²⁻ (second strongest superbase)
• Para-diethynylbenzene dianion (C₆H₄(C₂)₂)²⁻ (third strongest)
• Lithium monoxide anion (LiO⁻) was considered the strongest superbase before diethynylbenzene dianions were created.

Weak bases

Main article: weak base

A weak base is one which does not fully ionize in an aqueous solution, or in which protonation is incomplete. For example, ammonia transfers a proton to water according to the equation^[10]

${\displaystyle NH_{3}(aq)+H_{2}O(l)\rightleftharpoons NH_{4}^{+}(aq)+OH^{-}(aq)}$

The equilibrium constant for this reaction at 25 °C is 1.8 x 10⁻⁵,^[11] so that the extent of reaction or degree of ionization is quite small.

Lewis bases

A Lewis base or electron-pair donor is a molecule with a high-energy pair of electrons which can be shared with a low-energy vacant orbital in an acceptor molecule to form an adduct. In addition to H⁺, possible acceptors (Lewis acids) include neutral molecules such as BF₃ and metal ions such as Ag⁺ or Fe³⁺. Adducts involving metal ions are usually described as coordination complexes.^[12]

According to the original formulation of Lewis, when a neutral base forms a bond with a neutral acid, a condition of electric stress occurs.^[7] The acid and the base share the electron pair that formerly only belonged to the base.^[7] As a result, a high dipole moment is created, which can only be destroyed by rearranging the molecules.^[7]

Solid bases

Examples of solid bases include:

• Oxide mixtures: SiO₂, Al₂O₃; MgO, SiO₂; CaO, SiO₂^[13]
• Mounted bases: LiCO₃ on silica; NR₃, NH₃, KNH₂ on alumina; NaOH, KOH mounted on silica on alumina^[13]
• Inorganic chemicals: BaO, KNaCO₃, BeO, MgO, CaO, KCN^[13]
• Anion exchange resins^[13]
• Charcoal that has been treated at 900 degrees Celsius or activates with N₂O, NH₃, ZnCl₂-NH₄Cl-CO₂^[13]

Depending on a solid surface's ability to successfully form a conjugate base by absorbing an electrically neutral acid, the basic strength of the surface is determined.^[14] "The number of basic sites per unit surface area of the solid" is used to express how much base is found on a solid base catalyst.^[14] Scientists have developed two methods to measure the amount of basic sites: titration with benzoic acid using indicators and gaseous acid adsorption.^[14] A solid with enough basic strength will absorb an electrically neutral acid indicator and cause the acid indicator's color to change to the color of its conjugate base.^[14] When performing the gaseous acid adsorption method, nitric oxide is used.^[14] The basic sites are then determined using the amount o

Bases as catalysts

f carbon dioxide than is absorbed.^[14]

Basic substances can be used as insoluble heterogeneous catalysts for chemical reactions. Some examples are metal oxides such as magnesium oxide, calcium oxide, and barium oxide as well as potassium fluoride on alumina and some zeolites. Many transition metals make good catalysts, many of which form basic substances. Basic catalysts have been used for hydrogenations, the migration of double bonds, in the Meerwein-Ponndorf-Verley reduction, the Michael reaction, and many other reactions. Both CaO and BaO can be highly active catalysts if they are treated with high temperature heat.^[14]

Uses of bases

• Sodium hydroxide is used in the manufacture of soap, paper, and the synthetic fiber rayon.
• Calcium hydroxide (slaked lime) is used in the manufacture of bleaching powder.
• Calcium hydroxide is also used to clean the sulfur dioxide, which is caused by the exhaust, that is found in power plants and factories.^[9]
• Magnesium hydroxide is used as an 'antacid' to neutralize excess acid in the stomach and cure indigestion.
• Sodium carbonate is used as washing soda and for softening hard water.
• Sodium bicarbonate (or sodium hydrogen carbonate) is used as baking soda in cooking food, for making baking powders, as an antacid to cure indigestion and in soda acid fire extinguisher.
• Ammonium hydroxide is used to remove grease stains from clothes

Acidity of bases

The number of ionizable hydroxide (OH-) ions present in one molecule of base is called the acidity of bases.^[15] On the basis of acidity bases can be classified into three types: monoacidic, diacidic and triacidic.

Monoacidic bases

Sodium hydroxide

When one molecule of a base via complete ionization produces one hydroxide ion, the base is said to be a monoacidic base. Examples of monoacidic bases are:

Sodium hydroxide, potassium hydroxide, silver hydroxide, ammonium hydroxide, etc

Diacidic bases

When one molecule of base via complete ionization produces two hydroxide ions, the base is said to be diacidic. Examples of diacidic bases are:

Barium hydroxide, magnesium hydroxide, calcium hydroxide, zinc hydroxide, iron(II) hydroxide, tin(II) hydroxide, lead(II) hydroxide, copper(II) hydroxide, etc.

Triacidic bases

When one molecule of base via complete ionization produces three hydroxide ions, the base is said to be triacidic.^[16] Examples of triacidic bases are:

Aluminium hydroxide, ferrous hydroxide, Gold Trihydroxide,^[17]

Etymology of the term

The concept of base stems from an older alchemical notion of "the matrix":

The term "base" appears to have been first used in 1717 by the French chemist, Louis Lémery, as a synonym for the older Paracelsian term "matrix." In keeping with 16th-century animism, Paracelsus had postulated that naturally occurring salts grew within the earth as a result of a universal acid or seminal principle having impregnated an earthy matrix or womb. ... Its modern meaning and general introduction into the chemical vocabulary, however, is usually attributed to the French chemist, Guillaume-François Rouelle. ... In 1754 Rouelle explicitly defined a neutral salt as the product formed by the union of an acid with any substance, be it a water-soluble alkali, a volatile alkali, an absorbent earth, a metal, or an oil, capable of serving as "a base" for the salt "by giving it a concrete or solid form." Most acids known in the 18th century were volatile liquids or "spirits" capable of distillation, whereas salts, by their very nature, were crystalline solids. Hence it was the substance that neutralized the acid which supposedly destroyed the volatility or spirit of the acid and which imparted the property of solidity (i.e., gave a concrete base) to the resulting salt.

Acid, Base and Salt

                                                Acid

 This article is about acids in chemistry. For other uses, see Acid (disambiguation).

"Acidity" redirects here. For the novelette, see Acidit

"Acidic" redirects here. For the band, see Acidic (band).y (novelette).

An acid is a molecule or ion capable of either donating a proton (i.e., hydrogen ion, H⁺), known as a Brønsted–Lowry acid, or, capable of forming a covalent bond with an electron pair, known as a Lewis acid.^[1]

The first category of acids are the proton donors, or Brønsted–Lowry acids. In the special case of aqueous solutions, proton donors form the hydronium ion H₃O⁺ and are known as Arrhenius acids. Brønsted and Lowry generalized the Arrhenius theory to include non-aqueous solvents. A Brønsted or Arrhenius acid usually contains a hydrogen atom bonded to a chemical structure that is still energetically favorable after loss of H⁺.

                                            

Aqueous Arrhenius acids have characteristic properties which provide a practical description of an acid.^[2] Acids form aqueous solutions with a sour taste, can turn blue litmus red, and react with bases and certain metals (like calcium) to form salts. The word acid is derived from the Latin acidus/acēre, meaning 'sour'.^[3] An aqueous solution of an acid has a pH less than 7 and is colloquially also referred to as "acid" (as in "dissolved in acid"), while the strict definition refers only to the solute.^[1] A lower pH means a higher acidity, and thus a higher concentration of positive hydrogen ions in the solution. Chemicals or substances having the property of an acid are said to be acidic.

Common aqueous acids include hydrochloric acid (a solution of hydrogen chloride which is found in gastric acid in the stomach and activates digestive enzymes), acetic acid (vinegar is a dilute aqueous solution of this liquid), sulfuric acid (used in car batteries), and citric acid (found in citrus fruits). As these examples show, acids (in the colloquial sense) can be solutions or pure substances, and can be derived from acids (in the strict^[1] sense) that are solids, liquids, or gases. Strong acids and some concentrated weak acids are corrosive, but there are exceptions such as carboranes and boric acid.

The second category of acids are Lewis acids, which form a covalent bond with an electron pair. An example is boron trifluoride (BF₃), whose boron atom has a vacant orbital which can form a covalent bond by sharing a lone pair of electrons on an atom in a base, for example the nitrogen atom in ammonia (NH₃). Lewis considered this as a generalization of the Brønsted definition, so that an acid is a chemical species that accepts electron pairs either directly or by releasing protons (H⁺) into the solution, which then accept electron pairs. However, hydrogen chloride, acetic acid, and most other Brønsted–Lowry acids cannot form a covalent bond with an electron pair and are therefore not Lewis acids.^[4] Conversely, many Lewis acids are not Arrhenius or Brønsted–Lowry acids. 

                                       Types Of Acids

   There are two basic types of acids organic and inorganic acids. Inorganic acids are sometimes referred to as mineral acids. As a group, organic acids are generally not as strong as inorganic acids. The main difference between the two is the presence of carbon in the compound; inorganic acids do not contain carbon.

• Inorganic acids – Inorganic acids are often termed mineral acids. The anhydrous form may be gaseous or solid. An inorganic anhydride is an oxide of metalloid which can combine with water to form an inorganic acid.

Example:

1. Sulphuric acid (H2SO4)
2.  Phosphoric acid (H3PO4)
3. Nitric acid (HNO3)

• Organic acids – Organic acids are corrosive and toxic. Corrosivity is a form of toxicity to the tissues that the acid contacts. Organic acids and their derivatives cover a wide range of substances. They are used in nearly every type of chemical manufacture. Because of the variety in the chemical structure of the members of the organic acid group.

Example:

1. Acetic acid
2. Citric acid
3. Formic acid

                                   Properties Of Acids

• (i)               Aqueous solutions of acids are electrolytes, meaning that they conduct electrical current. ...
• (ii)              Acids have a sour taste. ...
• (iii)             Acids change the color of certain acid-base indicates. ...
• (iv)             Acids react with active metals to yield hydrogen gas. ...
• (v)              Acids react with bases to produce a salt compound and water.
•                                             Uses Of Acids
•     

  1. Hydrochloric Acid (HCl)

  • Dilute hydrochloric acid is used in various industries that use heating applications. It is applied to remove deposits from the inside of the boilers.
  • Hydrochloric acid is also used for cleaning sinks and sanitary ware.

  2. Sulphuric Acid (H2SO4)
  Sulphuric acid is such an important industrial chemical that it is called the king of chemicals. Some of its major uses are as follows:

  • Sulphuric acid is used in car batteries.
  • It is used in the manufacture of paints, drugs, dyes, and to produce fertilizers.

  3. Nitric Acid (HNO3)

  • It is used by goldsmiths for cleaning gold and silver ornaments.
  • It is also used for the production of fertilizers such as ammonium nitrate.

  4. Acetic Acid (CH3COOH)

  • Acetic acid is used directly to enhance the flavour of food. In fact, we commonly know acetic acid as vinegar.
  • It is also used as a cleansing agent in products meant for cleaning windows, floors, utensils, etc.
  • It also helps to remove stains on woodwork such as furniture and carpets.
  • Acetic acid is used as a preservative in pickles, etc. Most microorganisms cannot live in an acidic environment. An acidic environment either slows down their activities or can also kill them. This is why you will find vinegar in many commonly packaged food items such as pickles, sauce, ketchups, etc.

Nitrogen

                                                Nitrogen

         Nitrogen is the chemical element with the symbol N and atomic number 7. It was first discovered and isolated by Scottish physician Daniel Rutherford in 1772. Although Carl Wilhelm Scheele and Henry Cavendish had independently done so at about the same time, Rutherford is generally accorded the credit because his work was published first. The name nitrogène was suggested by French chemist Jean-Antoine-Claude Chaptal in 1790 when it was found that nitrogen was present in nitric acid and nitrates. Antoine Lavoisier suggested instead the name azote, from the Ancient Greek: ἀζωτικός "no life", as it is an asphyxiant gas; this name is used instead in many languages, such as French, Italian, Russian, Romanian, Portuguese and Turkish, and appears in the English names of some nitrogen compounds such as hydrazine, azides and azo compounds.

                                       

Nitrogen is the lightest member of group 15 of the periodic table, often called the pnictogens. It is a common element in the universe, estimated at about seventh in total abundance in the Milky Way and the Solar System. At standard temperature and pressure, two atoms of the element bind to form dinitrogen, a colourless and odorless diatomic gas with the formula N₂. Dinitrogen forms about 78% of Earth's atmosphere, making it the most abundant uncombined element. Nitrogen occurs in all organisms, primarily in amino acids (and thus proteins), in the nucleic acids (DNA and RNA) and in the energy transfer molecule adenosine triphosphate. The human body contains about 3% nitrogen by mass, the fourth most abundant element in the body after oxygen, carbon, and hydrogen. The nitrogen cycle describes movement of the element from the air, into the biosphere and organic compounds, then back into the atmosphere.

Many industrially important compounds, such as ammonia, nitric acid, organic nitrates (propellants and explosives), and cyanides, contain nitrogen. The extremely strong triple bond in elemental nitrogen (N≡N), the second strongest bond in any diatomic molecule after carbon monoxide (CO),^[3] dominates nitrogen chemistry. This causes difficulty for both organisms and industry in converting N₂ into useful compounds, but at the same time means that burning, exploding, or decomposing nitrogen compounds to form nitrogen gas releases large amounts of often useful energy. Synthetically produced ammonia and nitrates are key industrial fertilisers, and fertiliser nitrates are key pollutants in the eutrophication of water systems.

Apart from its use in fertilisers and energy-stores, nitrogen is a constituent of organic compounds as diverse as Kevlar used in high-strength fabric and cyanoacrylate used in superglue. Nitrogen is a constituent of every major pharmacological drug class, including antibiotics. Many drugs are mimics or prodrugs of natural nitrogen-containing signal molecules: for example, the organic nitrates nitroglycerin and nitroprusside control blood pressure by metabolizing into nitric oxide. Many notable nitrogen-containing drugs, such as the natural caffeine and morphine or the synthetic amphetamines, act on receptors of animal neurotransmitters.

                                            Properties of Nitrogen

   Nitrogen makes up the bulk of earth's atmosphere: 78.1 percent by volume. It is so inert at standard temperature and pressure that it was termed "azote" (meaning "without life") in Antoine Lavoisier's Method of Chemical Nomenclature. Nevertheless, nitrogen is a vital part of food and fertilizer production and a constituent of the DNA of all living things.

Characteristics

Nitrogen gas (chemical symbol N) is generally inert, nonmetallic, colorless, odorless and tasteless. Its atomic number is 7, and it has an atomic weight of 14.0067. Nitrogen has a density of 1.251 grams/liter at 0 C and a specific gravity of 0.96737, making it slightly lighter than air. At a temperature of -210.0 C (63K) and a ressure of 12.6 kilopascals, nitrogen reaches its triple point (the point an element can exist in gaseous, liquid and solid forms simultaneously).

Other States

At temperatures below nitrogen's boiling point of -195.79 C (77K), gaseous nitrogen condenses into liquid nitrogen, a fluid that resembles water and remains odorless and colorless. Nitrogen solidifies at a melting point of -210.0 C (63K) into a fluffy solid resembling snow.

Molecular Bonding

Nitrogen forms trivalent bonds in most compounds. In fact, molecular nitrogen exhibits the strongest possible natural triple bond due to the five electrons in the outer shell of the atom. This strong triple bond, along with nitrogen's high electronegativity (3.04 on the Pauling scale), explains its nonreactivity.

Uses

Nitrogen gas is useful in industrial and production settings due to its abundance and nonreactivity. In food production, nitrogen gas suppression systems can extinguish fires without fear of contamination. Iron, steel and electronic components, which are sensitive to oxygen or moisture, are produced in a nitrogen atmosphere. Nitrogen gas is commonly combined with hydrogen gas to produce ammonia.

Potential

In 2001, "Nature" reported that Carnegie Institution of Washington scientists were able to transform gaseous nitrogen into a solid state by subjecting the gaseous form to intense pressure. The researchers pressed a sample of nitrogen between two pieces of diamond with a force equivalent to 1.7 million times that of atmospheric air pressure, transforming the sample into a clear solid resembling ice, but with a crystal structure like that of diamond. At temperatures below -173.15°C (100K) the sample remained a solid when pressure was removed. When it reverts back to gaseous state nitrogen releases great amounts of energy, leading physics professor Dr. Richard M. Martin to speculate on its use as a rocket fuel.

Oxygen

                                       Oxygen

                  All living orgaisms oxygen to survive. All animals and plants used oxygen in respiration. Therefore it is also called Pranavayu. It was names oxygen by the scientist Levoshe.

                 Oxygen is formed in the largest quantity of elements fund on Earth. About half of the top surface of the land is oxygen. It is 21% in air and 89% in water.

                  The structure of an oxygen atom consists of eight protons and eight neutrons in the nucleus.

                  The two oxygen atoms 'O' and 'O' interact chemically to form one molecule of oxygen gas. It is spelled O2 as a symbol.

                   Oxygen has atomic number 8, atomic weight and mass number 16 and valency 2. Trees make oxygen gas by photosynthesis.

                       Properties Of Oxygen

(i)                This gas is colorless and odorless.

(ii)               This gas is less soluble in water, somewhat heavier then air.

(iii)              In the presence of oxygen, the match stick burns more rapidly with red light. In this process, the carbon present in the matchlet burns with oxygen to form carbon dioxide gas.                        C + O2 -------->  CO2

                      

Oxygen Uses

(i)              Oxygen is helpful in burning things.

(ii)               Oxygen is necessary for respiration in animals and plants.

(iii)            Aquatic animals and plants use oxygen dissolved in water in respiration.

                                      

Phosphorus

                                    Phosphorus

                     Atomic number 15, atomic weight 31, electronic structure 1s²2s²2p⁶3s²3p³  

^{formula 31/15 P, valency +3, +5.}

              ^{Scientist Brand of germany discovered Phosphorus in 1669. He first obtained this element by distillation of urine, sand and coal, As this element shines in the darkness. Hence, it was named Phosphorus.}

               ^{It is not found in free state in nature, as it is highly functional. It is found in the form of compounds in nature in the joint state.}

                               

                              ^{Phosphorus impurities}

               ^{These are several allotropes of phosphorus of which white, red and black phosphorus are the main ones.}

^{(i.)        White Phosphorus-}

                ^{White phosphorus is white in pure state but gradually becomes yellow, it has a garlic like odor and is poisonous. It is a molecular solid consisting of quadratically arranged P4 units. The value of P-P-P bond angle is 60' and there is more tension in the structure. This is why white phosphorus is more active. In the presence of air, it catches fire, so it is kept in cold water. It is soft and can be easily cut with a knife.}

 ^{(ii.)        Red Phosphorus-}

                  ^{The structure of red phosphorus is a complex chain structure. It is odorless and has a high threshold. The red phosphorus has a P-P-P bond angle of 100' .}

 ^{(iii.)       Black Phosphorus-}

                   ^{Black phosphorus is more stable due to its flaky structure. There are many layers of phosphorus atoms attached to it. It is also odorless like red phosphorus.}

                   ^{The formation of black phosphorus and red phosphorus from white phosphorus can be explained by it.}

                           ^{Use Of Phosphorus}

 ^{(1.)      White phosphorus is used in making smoke balloons, fireballs, fire sports demonstrations and colorful flashing matches.}

 ^{(2.)       Red phosphorus is used in making matches.}

 ^(3.)       Red phosphorus is used in making an alloy called phosphorus bronzes. This alloy contains copper, tin and phosphorus.

 (4.)       Phosphorus compounds, zinc phosphorus and calcium phosphide are used as rat medicine.

Non-Crystalline Sulfur

                                 Non-Crystalline Sulfur

                  This  Sulfur is found in three allotropes.

(i.)              Plastic Sulfur-

                           When boiling sulfur is poured into cold water, a soft material similar to rubber is obtained which is called plastic sulfur. It is a temporary alteration of sulfur which gradually converts to rhombic sulfur.

(ii.)             Δ (Delta) Sulfur or Milk Sulfur-

                                 It is a white colored non-crystalling sulfur. It is used in making medicines.

(iii.)            Colloid Sulfur-

                              This form of colloidal sulfur is obtained by flowing H2S gas in dilute nitric acid.

                        H2S + 2HNO3(meager)----------> 2NO2 + 2H2O + S

                              This sulfur is soluble in carbon disulfide and its insoluble in water. On heating or after some time it is converted into rhombic sulfur. It is also used in making medicines.

                            Uses Of Sulfur

   (i.)            Sulfur is used extensively in the industrial production of sulfuric acid.

    (ii.)           It is used in making gunpowder and matchmak.

     (iii.)         Also useful as a pesticide.

                                 

Some Important Metals

                                Some Important Metals

 Nonmetal Sulfur (Sulfur)

                        Atomic number 16, Atomic weight 32, Electronic structure 1s²2s²2p⁶3s²3p⁴. valency +2, +4 and +6 and 2 in some compounds.

                        Sulfur has been known since ancient time. Sulfur come out with lava when the volcano bursts. Sulfur compounds are found in many springs and wells / stepwells.

                         Ancient Indian physicians used sulfur make various medicines. The properties of copper are destroyed when heated with sulfur.

                         A scientist named Lavoisier called is an element based on the study of its properties. Sulfur is found in free from in nature and various compounds.

                             Sulfur Allotropy

       Allotropism-

                             Two or more forms of an element that exhibit differences in structure and other physical properties from each other but possess similar chemical properties are called as allotropes of that element. This properties of elements is called allotroopism. 

                                     

                                       Crystalline sulfur

(i)          Rhombic Sulfur-

                              This form of sulfur is also known as alpha sulfur. This form sulfur is very stable at normal temperatures. It is insoluble in water but soluble in carbon disulfide. On heating it at 368.6 K (95.6' C), it is converted into another unconformable, monoclinic sulfur.

 (ii.)      Mono-clinic Sulfur-

                               This is also called beta sulfur. It occurs as a needle-shaped crystal. Therefore, it is also called cone sulfur, it is insoluble in the water and soluble in carbon disulfide. This alkali of sulfur is stable at temperatures above 368.6 K. At low temperatures, it converts to (alpha) sulfur. Both forms of sulfur coexist at 368.6 K. This temperatures is called transition temperature.

α

Non- Metals

                                       Non- Metals

                  There are many elements that are different from metals in physical and chemical properties. called non- metals. For example carbon, oxygen, hydrogen, nitrogen, sulfur, etc. Nonmetals are also called electrically negative elements. Nonmetal make anion.

                                           

                                 

           

                                        Properties of Non- Metals

    (i.)                  Generally these are brittle, So its sheets or wires cannot be made.

    (ii.)                 They do not have any special shine, but iodine is a shiny nonmetal.

    (iii.)                Its melting point are low but the melting point of diamond and graphite it very high, about 3000' C.

    (iv.)                 They are often non-conductors of heat and electricity but graphite is a good conductor of electricity and heat.

     (v.)                  It dose not contain free electron. The outermost orbit of their atom often has 5, 6 or 7 electrons.

     (vi.)                 Non-metal oxides are acidic.

     (vii.)               Those substance which are made up of the some element, but their structure and composition are different, are called allotropic.

                                                     

                                    Type of Non- Metals             

     (i.)        Solid Non- Metals-

                                            Carbon, Sulfur, Phosphorus, Iodine etc.

      (ii.)      Fluid Non- Metals-

                                            Bromine. 

      (iii.)     Gas Non- Metals-

                                            Hydrogen, Oxygen, Chlorine, Ozone etc.

Refining

                                                          Refining

                          The method of extraction of pure metal from impure or raw metal proper conditions is called refining of the metal. Several methods are used to obtain metals of high purity. These are as follows.

                                                          

       Distillation-

                           Zn, Cd, Hg metals are treated by this method because they are low boiling point metals. 

       Electrolytic Refining-

                            Copper, silver, gold, aluminum and lead metal are treated by this method, anode of impure metal, and cathode of metal are made.

       Zone Refining-

                          This method is based in the principle that the solubility of impurities is higher in the molten state than in the solid state of the metal. This metal is useful in obtaining very high purity semiconductors and other highly pure metals such as germanium, silic, boron, gallium and indium.

        Vapour Phase Refining-

                             In this method, the metal is converted in to a voltatile compound. Thereafter, it is decomposed to obtain pure metal. Example  Mode process It used for nickel.

         Von-Arkell Method-

                              It is used for zirconium and titanium.

           Chromatographic Method-

                              This method is based on the principle that absorption of different components of a mixture is different at the absorbent. Purification of elements found in small quantities is done in this method.