ACID BASE AND SALT

1. ACID BASE AND SALT

2. ACID An acid (often represented by the generic formula HA [H+A−]) is traditionally considered any chemical compound that, when dissolved in water, gives a solution with a hydrogenion activity greater than in pure water, i.e. a pH less than 7.0. That approximates the modern definition of Johannes Nicolaus Brønsted and Martin Lowry, who independently defined an acid as a compound which donates a hydrogen ion (H+) to another compound (called abase). Common examples include acetic acid (in vinegar) and sulfuric acid (used in car batteries). Acid/base systems are different from redox reactions in that there is no change in oxidation state.

3. DEFINITIONS The word "acid" comes from the Latinacidus meaning "sour," but in chemistry the term acid has a more specific meaning. There are four common ways to define an acid. • ARRHENIUS:- According to this definition developed by the Swedish chemist Svante Arrhenius, an acid is a substance that increases the concentration of hydrogen ions (H+), which are carried as hydronium ions (H3O+) when dissolved in water, while bases are substances that increase the concentration of hydroxide ions (OH-). This definition limits acids and bases to substances that can dissolve in water. Around 1800, many French chemists, including Antoine Lavoisier, incorrectly believed that all acids contained oxygen. Indeed the modern German word for oxygen is Sauerstoff (lit. sour substance), as are the Afrikaans and Dutch words for oxygen suurstof and zuurstof respectively, with the same meaning. English chemists, including Sir Humphry Davy, at the same time believed all acids contained hydrogen. Arrhenius used this belief to develop this definition of acid.

4. 2) BRONSTED-LOWRY :- According to this definition, an acid is a proton (hydrogen nucleus) donor and a base is a proton acceptor. The acid is said to be dissociated after the proton is donated. An acid and the corresponding base are referred to as conjugate acid- base pairs. Brønsted and Lowry independently formulated this definition, which includes water-insoluble substances not in the Arrhenius definition. Acids according to this definition are variously referred to as Brønsted acids, Brønsted- Lowry acids, proton acids, protic acids, or protonic acid. 3) SOLVENT-SYSTEM DESINITION :- According to this definition, an acid is a substance that, when dissolved in an autodissociating solvent, increases the concentration of the solvonium cations, such as H3O+ in water, NH4+ in liquid ammonia, NO+ in liquid N2O4, SbCl2+ in SbCl3, etc. Base is defined as the substance that increases the concentration of the solvate anions, respectively OH-, NH2-, NO3-, or SbCl4-. This definition extends acid-base reactions to non-aqueous systems and even some aprotic systems, where no hydrogen nuclei are involved in the reactions. This definition is not absolute, a compound acting as acid in one solvent may act as a base in another.

5. 4) LEWIS :- According to this definition developed by Gilbert N. Lewis, an acid is an electron- pair acceptor and a base is an electron-pair donor. (These are frequently referred to as "Lewis acids" and "Lewis bases," and are electrophiles and nucleophiles, respectively, in organic chemistry; Lewis bases are also ligands in coordination chemistry.) Lewis acids include substances with no transferable protons (ie H+hydrogen ions), such as iron(III) chloride, and hence the Lewis definition of an acid has wider application than the Brønsted-Lowry definition. In fact, the term Lewis acid is often used to exclude protic (Brønsted-Lowry) acids. The Lewis definition can also be explained with molecular orbital theory. In general, an acid can receive an electron pair in its lowest unoccupied orbital (LUMO) from the highest occupied orbital (HOMO) of a base. That is, the HOMO from the base and LUMO from the acid combine to a bonding molecular orbital.

6. Although not the most general theory, the Brønsted-Lowry definition is the most widely used definition. The strength of an acid may be understood by this definition by the stability of hydronium and the solvated conjugate base upon dissociation. Increasing or decreasing stability of the conjugate base will increase or decrease the acidity of a compound. This concept of acidity is used frequently for organic acids such as carboxylic acid. The molecular orbital description, where the unfilled proton orbital overlaps with a lone pair, is connected to the Lewis definition

7. PROPERTIES Brønsted–Lowry acids: Are generally sour in taste. (For example, the sour taste of lemon juice is due to citric acid). • Strong or concentrated acids or their fumes often produce a stinging feeling on mucous membranes. • Change the color of pH indicators as follows: turn blue litmus and methyl orange red, turn phenolphthalein colorless. • React with metals to produce a metal salt and hydrogen. • React with metal carbonates to produce water, CO2 and a salt. • React with metal hydroxides and metal oxides to produce water and a salt. • Conduct electricity, depending on the degree of dissociation in aqueous solution. • Produce solvonium ions, such as oxonium (H3O+) ions in water

8. Acids can be gases, liquids, or solids. Respective examples (at 20 °C and 1 atm) are hydrogen chloride, sulfuric acid and citric acid. Solutions of acids in water are liquids, such as hydrochloric acid - an aqueous solution of hydrogen chloride. At 20 °C and 1 atm, linear carboxylic acids are liquids up to nonanoic acid (nine carbon atoms) and solids beginning from decanoic acid (ten carbon atoms). Aromatic carboxylic acids, the simplest being benzoic acid, are solids. Strong acids and some concentrated weak acids are corrosive and can cause severe burns even after short contact. Generally, acid burns on the skin are treated by rinsing the affected area abundantly with running water, followed up with immediate medical attention. Particular acids may also be dangerous for reasons not related to their acidity. Material Safety Data Sheets (MSDS) can be consulted for detailed information on dangers and handling instructions.

9. NOMENCLATURE In the classical naming system, acids are named according to their anions. That ionic suffix is dropped and replaced with a new suffix (and sometimes prefix), according to the table below. For example, HCl has chloride as its anion, so the- ide suffix makes it take the Form hydrochloric acid. In the IUPAC naming system, "aqueous" is simply added to the name of the ionic compound. Thus, for hydrogen chloride, the IUPAC name would be aqueous hydrogen chloride. The prefix "hydro-" is added only if the acid is made up of just hydrogen and one other element.

10. Classical naming system:

11. NEUTRALIZATION Neutralization is the reaction between an acid and a base, producing a salt and neutralized base; for example, hydrochloric acid and sodium hydroxide form sodium chloride and water: HCl(aq) + NaOH(aq) → H2O(l) + NaCl(aq)

12. Neutralization is the basis of titration, where a pH indicator shows equivalence point when the equivalent number of moles of a base have been added to an acid. It is often wrongly assumed that neutralization should result in a solution with pH 7.0, which is only the case with similar acid and base strengths during a reaction. Neutralization with a base weaker than the acid results in a weakly acidic salt. An example is the weakly acidic ammonium chloride, which is produced from the strong acid hydrogen chloride and the weak base ammonia. Conversely, neutralizing a weak acid with a strong base gives a weakly basic salt, e.g. sodium fluoride from hydrogen fluoride and sodium hydroxide.

13. BASE In chemistry, a base is most commonly thought of as an aqueous substance that can accept protons. A base is also often referred to as an alkali if OH− ions are involved. This refers to the Brønsted-Lowry theory of acids and bases. Alternate definitions of bases include electron pair donors (Lewis), as sources of hydroxide anions (Arrhenius). In addition to this, bases can commonly be thought of as any chemical compound that, when dissolved in water, gives a solution with a pH higher than 7.0. Examples of simple bases are sodium hydroxide and ammonia. Bases can be thought of as the chemical opposite of acids. A reaction between an acid and base is called neutralization. Bases and acids are seen as opposites because the effect of an acid is to increase the hydronium ion (H3O+) concentration in water, whereas bases reduce this concentration. Bases react with acids to produce water and salts (or their solutions).

14. DEFINITIONS A strong base is a base which hydrolyzes completely, raising the pH of the solution towards 14. Strong bases, like strong acids, attack living tissue and cause serious burns. They react differently to skin than acids do, so while strong acids are corrosive, we say that strong bases are caustic. Superbases are a class of especially basic compounds and non- nucleophilic bases are a special class of strong bases with poor nucleophilicity. Bases may also be weak bases such as ammonia, which is used for cleaning. Arrhenius bases are water-soluble and these solutions always have a pH greater than 7. An alkali is a special example of a base, where in an aqueous environment, hydroxide ions (also viewed as OH−) are donated. There are other more generalized and advanced definitions of acids and bases.

15. The notion of a base as a concept in chemistry was first introduced by the French chemist Guillaume François Rouelle in 1754. He noted that acids, which in those days were mostly volatile liquids (like acetic acid), turned into solid salts only when combined with specific substances. These substances form a concrete base for the salt and hence the name.

16. PROPERTIES Some general properties of bases include: • Slimy or soapy feel on fingers, due to saponification of the lipids in human skin. • Concentrated or strong bases are caustic (corrosive) 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 and phenolphthalein pink

17. BASE AND PH The pH of (impure) water is a measure of its acidity. In pure water, about one in ten million molecules dissociate into hydronium ions (H3O+) And hydroxide ions (OH−), according to the following equation: 2H2O(l) → H+ (aq) + OH− (aq) The concentration, measured in molarity (M or moles per dm³), of the ions is indicated as [H3O+] and [OH−]; their product is the dissociation constant of water with and has the value 10−7 M. The pH is defined as −log [H3O+]; thus, pure water has a pH of 7. (These numbers are correct at 23 °C and slightly different at other temperatures.)

18. A base accepts (removes) hydronium ions(H3O+) from the solution, or donates hydroxide ions(OH−) to the solution. Both actions will lower the concentration of hydronium ions, and thus raise pH. By contrast, an acid donates H3O+ ions to the solution or accepts OH−, thus lowering pH. For example, if 1 mole of sodium hydroxide (40 g) is dissolved in water to make 1 litre of solution, the concentration of hydroxide ions becomes [OH−] = 1 mol/L. Therefore [H+] = 10−14 mol/L, and pH = −log 10−14 = 14. Note that in this calculation, it is assumed that the activity is equivalent to the concentration, which is not realistic at concentrations over 0.1 mol dm−3. The base dissociation constant or Kb is a measure of basicity. pKb is the negative log of Kb and related to the pKa by the simple relationship pKa + pKb = 14 . Alkalinity is a measure of the ability of a solution to neutralize acids to the equivalence points of carbonates or bicarbonates.

19. NEUTRALISATION WITH ACIDS When dissolved in water, the strong base sodium hydroxide decomposes into hydroxide and sodium ions: NaOH → Na+ + OH− and similarly, in water hydrogen chloride forms hydronium and chloride ions: HCl + H2O → H3O+  + Cl− When the two solutions are mixed, the H3O+ and OH− ions combine to form water molecules: H3O+  + OH−  → 2 H2O

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