Acids and Bases - Essay Example

Chemistry Research Paper 7 April Acids and Bases 0 Introduction Since the 17th century, sourness was used asthe defining property of acids while bitterness was used for bases (Silberg 591). An acid was any substance that reacted with metals such as magnesium, zinc and aluminum producing hydrogen gas in the process. A chemical reaction concerning a base and an acid was known to produce substances called salts, which were ionic in nature (Ebbin and Gammon 624). Although these definitions of acids and bases described distinctive features, they gave definitions based on the molecular behavior of acids and base. Discovery of new phenomena in the field of science often leads to the replacement of limited definitions with the broader ones. A correct definition of acids and bases is important to our daily applications: the human bloodstream, cleaning materials, environment and industry. 1.1 Acid-Base Theories Just like other chemical theories, acids and bases have undergone many changes in the recent past. The changes implemented have been done to ensure a more general theory. The three main theories that are applicable in the science field are: the Arrhenius Theory the Bronsted‐Lowry Theory the Lewis Theory 1.1.1 Arrhenius Theory In 1987, the first theory describing bases and acids was proposed by a Swedish chemist Svante Arrhenius (Ebbin and Gammon 624). His definition of acids and bases was based on the effects these substances had on water. An acid was defined as a substance, which when dissolved in water, increases the concentration of hydronium ion, H3O+ (Ebbin and Gammon 625). He defined a base as a substance, which when dissolved in water, increases the concentration of hydroxide ion, OH-. Frequently chemists use hydrogen ion, H+(aq) for H3O+. Hydrogen ions are sometimes referred as protons due to an absence of electrons in their valence shell. HCl(aq) + H2O(l) H3O+(aq) + Cl-(aq) NaOH(aq) + H2O(l) Na+(aq) + OH-(aq) These hydrogen and hydroxide ions may combine to form a water molecule. OH-(aq) + H+(aq) H2O(l) Arrhenius theory was limited to substances, which could produce hydrogen or hydroxide ions in water and therefore could not explain why acids and bases, which were not in aqueous state, did not produce hydrogen and hydroxide ions (Silberg 592). An example of such substance is ammonia. NH3(g) + HCl(g) NH4Cl(s) 1.1.2 Bronsted‐Lowry Theory Danish chemist Johannes N. Bronsted and British chemist Thomas M. Lowry proposed an extension to Arrhenius definition of an acid and a base in 1923. According to them, a reaction between a base and an acid involves the transfer of protons from one species to the other. They defined an acid as a species, which donates a proton (H+) and a base a species that accepts a proton (Silberg 600). Therefore, an acid must contain H in its formula while a base must contain a lone pair of electrons, which binds with proton example is ammonia. In the forward reaction, H3O+ is the acid because it donates H+ to NH3 while NH3 is the base as it receives a proton from H3O+ in this proton-transfer reaction. In the reverse reaction, NH4+ is the acid because it donates H+ to H2O to form H3O+ while H2O is the base as it accepts a proton from NH4+.This theory focuses on both the reactants and products in its definition. The species H3O+ and H2O are a conjugate acid-base pair (Silberg 601). H3O+ is a conjugate acid of the base H2O while NH3 is a conjugate base of acid NH4+. Generally, a weak acid forms a strong conjugate base and vice versa. A monoprotic acid or base is an acid or base that donates or accepts a single proton while polyprotic acid or base is an acid or base that donates or accepts more than a single proton. 1.1.3 Lewis Theory Lewis G.N proposed a theory, which defined acids and bases in terms of electron transfer. His definitions of acids and bases did not rely on the presence of any particular solvent. He defined an acid as a species, which is electron-deficient or a species and seeks a molecular species that contain an available pair of electrons (Ebbin and Gammon 628). A base was defined as a species that contain a pair of an electron that can be donated to another species. An example is the formation of ammonium ion. Ammonia donates a pair of an electron to hydrogen ion; ammonia is the base while hydrogen ion is the acid. Acids and bases are, therefore, acceptor and donor molecules, respectively. 1.2 Strengths of Acids and Bases 1.2.1 Relative Strength of Acids An acid strength depends entirely on its ability to transfer H+ to a base in order to form its conjugate base (Arun, Bahl and Tuli 940). Dissolving a monoprotic acid (HA) in water, results in the transfer of a proton from the acid to water to form a conjugate base and H3O+. taking H3O+ = H+, the above equation can be written as if we apply the law of mass action to the equilibrium reaction, Ka is the acid dissociation constant (Arun, Bahl and Tuli 941). Water is excluded from the equilibrium expression because its concentration remains constant. An acid strength is, therefore, defined as the concentration of hydrogen ions (H+) in its aqueous solution at a specific temperature (Arun, Bahl and Tuli 941). Since the concentration of H+ depends on the numerical value of Ka, then the Ka is the measure of the acidic strength of any acid (Cox 7638). From the dissociation constant expression, the larger the Ka value, the higher the concentration of [H+] hence the stronger the acid. For a strong acid, approximately all the HA dissociate in water giving a large Ka value. Similarly, a weak acid dissociates partially in water giving a small Ka value. 1.2.2 Relative Strength of Bases Based on the Arrhenius definition of a base, a substance that produces hydroxide ions (OH-) in an aqueous state, the basic properties of a base are due to the presence of hydroxyl ions. The dissociation of a base BOH in water can be represented as If we apply the law of mass action to this equilibrium, then the equilibrium expression can be written as Kb is the base dissociation constant defined as the concentration of hydroxide ions in its aqueous state at a specific temperature (Arun, Bahl and Tuli 943). The larger the Kb value, the higher the concentration of [OH-] and hence the stronger the base. 1.3 Relationship between Ka and Kb Water molecule undergoes self-ionization to produce H3O+ and OH-; a proton is transferred from one water molecule to another (Ebbin and Gammon 636). The concentration of water remains constant since the concentration of the ions formed is small. Kc[H2O]2 is constant at 250C denoted as Kw. Kw is the water dissociation constant and has a numerical value of 1.0 × 10-14. Therefore, . For the equilibrium reactions Therefore, Ka × Kb = Kw (Silberg 613). 1.4 The pH Scale The [OH-] and [H+] in aqueous solution can assume numerous orders of magnitude in terms of their values. There is, therefore, need to represent them on a logarithmic scale. pH = - log [H+], pOH = - log[OH-] and pKw = - log Kw. From the three equations, it is evident that pKw = pOH + pH (the value = 14.0 for pure water at 250C) (Ebbin and Gammon 640). The pH of neutral solution at 250C is 7.0; higher pH values correspond to alkaline solutions while lower pH values correspond to acidic solutions. 2.0 Conclusion The first successful theory in definitions of acids and bases was the Arrhenius concept. Later in 1923, Bronsted and Lowry described acid-base reactions as proton transfer reactions. Their concept of acids and bases was superior to that of Arrhenius because it could explain basicity of substances, which did not contain OH- in their original structures. The latest theory of acids and bases to be developed was that of Lewis whose definition of an acid is an electron-pair acceptor and a base as an electron-pair donor. The strength of acids and bases depend on the strength and polarity of chemical bond in the acidic hydrogen bond. Water molecule ionizes to produce OH- and H3O+ ions whose concentration relate to the dissociation constant of water Kw. The acidity or basicity of a solution is determined by the concentration of hydronium ion. Works Cited Arun, Bahl, Bahl, B.S and Tuli, G.D. Essentials of Physical Chemistry. 2000. Print. Cox, R. A. "Acids and Bases: Solvent Effects on Acid-Base strength." Angewandte Chemie International Edition 52.30 (2013): 7638. Web. 7 Sep. 2015. Ebbin D. Darrell and Gammon, D. Steven. General Chemistry. 9th ed. New York: Houghton Mifflin Company, 2009. Print. Silberg, S. Martin. Principles of General Chemistry. 2nd ed. New York: McGraw-Hill, 2010. Print. Read More