CHM 101 Notes PDF

ACIDS AND BASES The terms acid and base have been defined in different ways, depending on the particular way of looking at the properties of acidity and basicity. Arrhenius first defined acids as compounds which ionize to produce hydrogen ions, and bases as compounds which ionize to produce hydroxi...

ACIDS AND BASES The terms acid and base have been defined in different ways, depending on the particular way of looking at the properties of acidity and basicity. Arrhenius first defined acids as compounds which ionize to produce hydrogen ions, and bases as compounds which ionize to produce hydroxide ions. According to the Lowry-Bronsted definition, an acid is a proton donor, and a base is a proton acceptor. According to the Lewis definition, acids are molecules or ions capable of coordinating with unshared electron pairs, and bases are molecules or ions having unshared electron pairs available for sharing with acids. To be acidic in the Lewis sense, a molecule must be electron deficient. This is the most general acid base concept. All Lowery Bronsted acids are Lewis acids but, in addition, the Lewis definition includes many other reagents such as boron trifluoride, aluminum chloride, etc. Theories of Acids and Bases Three different theories have been put forth in order to define acids and bases. These theories include the Arrhenius theory, the Bronsted-Lowry theory, and the Lewis theory of acids and bases. Acids and bases can be defined via three different theories. Arrhenius Concept of Acids and Bases The Arrhenius concept of an acid and a base stated as follows: An acid is a substance that, when dissolved in water, increases the concentration of hydronium ion, H3O +(aq). A base is a substance that, when dissolved in water, increases the concentration of hydroxide ion, OH −(aq). 1 Arrhenius proposed that acid-base reactions are characterized by acids if they dissociate in aqueous solution to form hydrogen ions (H+) and bases if they form hydroxide (OH – ) ions in aqueous solution. Limitations of Arrhenius Concept The presence of water is absolutely necessary for acids and bases. Dry HCl can’t act as an acid. HCl acts as an acid in water only and not any other solvent. The concept does not explain the acidic and basic character of substances in non-aqueous solvents. The neutralization process is only possible for reactions which can occur in aqueous solutions, although reactions involving salt formation can occur in the absence of a solvent. The acidic character of some salts such as AlCl3 in aqueous solution can’t be explained. An extended as well as artificial explanation is needed to define the basic nature of NH 3. Brønsted–Lowry Concept of Acids and Bases In 1923 the Danish chemist Johannes N. Brønsted (1879–1947) and, independently, the British chemist Thomas M. Lowry (1874–1936) pointed out that acid–base reactions can be seen as proton-transfer reactions and that acids and bases can be defined in terms of this proton (H +) transfer. According to the Brønsted–Lowry concept, an acid is the species donating a proton in a proton-transfer reaction. A base is the species accepting the proton in a proton-transfer reaction. Consider, for example, the reaction of hydrochloric acid with ammonia, writing it as an ionic equation, you have After cancelling Cl− you obtain the net ionic equation 2 In this reaction in aqueous solution, a proton, H +, is transferred from the H 3 O+ ion to the NH3 molecule, giving H 2 O and NH 4 +. Here H 3 O+ is the proton donor, or acid, and NH 3 is the proton acceptor, or base. Note that in the Brønsted–Lowry concept, acids (and bases) can be ions as well as molecular substances. You can also apply the Brønsted–Lowry concept to the reaction of HCl and NH3 dissolved in benzene, C6 H 6 ,. In benzene, HCl and NH 3 are not ionized. The equation is Here the HCl molecule is the proton donor, or acid, and the NH 3 molecule is the proton acceptor, or base. In any reversible acid–base reaction, both forward and reverse reactions involve proton transfers. Consider the reaction of NH 3 with H 2 O. In the forward reaction, NH 3 accepts a proton from H 2 O. Thus, NH 3 is a base and H2 O is an acid. In the reverse reaction, NH 4 + donates a proton to OH −. NH 4 + ion is the acid and OH − is the base. Note that NH 3 and NH 4 + differ by a proton. That is, NH 3 becomes the NH 4 + ion by gaining a proton, whereas the NH 4 + ion becomes the NH 3 molecule by losing a proton. The species NH4+ and NH 3 are a conjugate acid–base pair. A conjugate acid–base pair consists of two species in an acid–base reaction, one acid and one base, that differ by the loss or gain of a proton. The acid in such a pair is called the conjugate acid of the base, whereas the base is the conjugate base of the acid. Here NH 4 + is the conjugate acid of NH 3 , and NH 3 is the conjugate base of NH 4 +. Limitations of Bronsted Lowry Concept Bronsted Lowry could not explain the reaction occurring in the non-protonic solvent like SO 2 , N 2 O4 , etc. 3  It cannot explain the reactions between acidic oxides like etc and the basic oxides like etc which can easily take place in the absence of solvent as well e.g. (No proton transfer) Substances like BF3 , AlCl3 etc, do not contain hydrogen which means they can’t donate a proton, still they behave as acids. Exercises In the following equations, label each species as an acid or a base. Show the conjugate acid–base pair The Brønsted–Lowry concept defines a species as an acid or a base according to its function in the acid–base, or proton-transfer, reaction. some species can act as either an acid or a base. An amphiprotic species is a species that can act as either an acid or a base (it can lose or gain a proton), depending on the other reactant. For example, HCO 3 − acts as an acid in the presence of OH− but as a base in the presence of HF. Anions with ionizable hydrogens, such as HCO 3 −, and certain solvents, such as water, are amphiprotic. The amphiprotic characteristic of water is important in the acid–base properties of aqueous solutions. Consider, for example, the reactions of water with the base NH 3 and with the acid HC2 H3 O 2 (acetic acid). In the first case, water reacts as an acid with the base NH3. In the second case, water reacts as a base with the acid HC 2 H 3 O2. we have now seen several ways in which the Brønsted–Lowry concept of acids and bases has greater scope than the Arrhenius concept. In the Brønsted–Lowry concept: 4 1. A base is a species that accepts protons; OH − is only one example of a base. 2. Acids and bases can be ions as well as molecular substances. 3. Acid–base reactions are not restricted to aqueous solution. 4. Some species can act as either acids or bases, depending on what the other reactant is. Lewis Concept of Acids and Bases Certain reactions have the characteristics of acid-base responses but do not fit the Brønsted–Lowry concept. An example is the reaction of the basic oxide Na 2 O with the acidic oxide SO 3 to give the salt Na2 SO 4. G. N. Lewis, who proposed the electron-pair theory of covalent bonding, realized that the concept of acids and bases could be generalized to include reactions of acidic and basic oxides and many other reactions, as well as proton-transfer reactions. According to this concept, a Lewis acid is a species that can form a covalent bond by accepting an electron pair from another species; a Lewis base is a species that can form a covalent bond by donating an electron pair to another species. The Lewis and the Brønsted–Lowry concepts are simply different ways of looking at certain chemical reactions. Such different views are often helpful in understanding reactions Consider again the neutralization of NH 3 by HCl in aqueous solution, mentioned earlier. It consists of the reaction of a proton from H 3 O + with: NH 3. 5 The proton is an electron-pair acceptor, so it is a Lewis acid. Ammonia, NH3, which has a lone pair of electrons, is an electron-pair donor and therefore a Lewis base. Now let us look at the reaction of Na2 O with SO 3. It involves the reaction of the oxide ion, O 2−, from the ionic solid, Na2 O, with SO 3. Here SO 3 accepts the electron pair from the O 2− ion. Thus, O 2− is the Lewis base and SO 3 is the Lewis acid. The Lewis concept embraces many reactions that we might not think of as acid–base reactions. The reaction of boron trifluoride with ammonia is an example In this reaction, the NH 3 molecule donates the lone pair of electrons on the nitrogen atom to the boron atom of BF3. Boron trifluoride accepts the electron pair and so is a Lewis acid. Ammonia donates the electron pair and so is a Lewis base. Limitations of Lewis Concept Lewis concept gave a generalized idea including all coordination reactions and compounds. This is not true always. An idea about the relative strength of acids and bases is not provided by Lewis concept. Lewis concept is not in line with the acid-base reaction concept. 6  Lewis concept has not discussed the behaviour of protonic acids like HCl. Exercises In the following reactions, identify the Lewis acid and the Lewis base. Strong and Weak Acids We classify acids as strong or weak depending on the extent to which they donate H + ions. Strong acids are completely ionized in water; weak acids are not. The strong acids are of two types: those formed by binary molecular compound of the generic formula HX, where X is nonmetal, and those called oxoacids with the generic formula H mXO n , where X is a nonmetal. The strength of an acid depends on its ability to donate a proton, which depends in turn on the strength of the bond to the acidic proton. In this section, we apply trends in atomic and bond properties to determine the trends in acid strength of nonmetal hydrides and oxoacids. Acid Strength of Nonmetal Hydrides (HX) Two factors determine how easily a proton is released from a nonmetal hydride: The electronegativity of the central nonmetal (X) The strength of the X—H bond 1. Across a period, acid strength increases. The electronegativity of the nonmetal X determines the trend. From left to right, as X becomes more electronegative, it withdraws electron density from H, and the X—H bond becomes more polar. As a result, H+ is pulled away more easily by an O atom of a water molecule. In aqueous solution, the hydrides of Groups 3A(13) to 5A(15) do not behave as acids, but an increase in acid strength is seen from Group 6A(16) to 7A(17). 2. Down a group, acid strength increases. The X—H bond strength determines the trend. As X becomes larger, the X—H bond becomes longer and weaker, so H + comes off more easily. For example, the hydrohalic acid strength increases down the group: 7  HF HOClO 3 > HOClO 2 > HOClO > HOCl Similarly, HNO 3 is stronger than HNO 2 , H 2 SO 4 is stronger than H 2 SO3 , and so forth. Strong and Weak Bases A strong base is one that completely dissociates in solution. NaOH, for example, is a strong base. Some strong bases, such as Sr (OH)2 , contain two O H − ions. These bases completely dissociate, producing 2 mol of OH − per mole of the base. For example, Sr (OH)2 dissociates as follows: 8 Neutralization When an acid and a base are mixed, the H +(aq) from the acid combines with the O H −(aq) from the base to form H 2 O(l). Consider the reaction between hydrochloric acid and sodium hydroxide mentioned earlier Acid–base reactions (also called neutralization reactions ) generally form water and an ionic compound called a salt that usually remains dissolved in the solution. The net ionic equation for many acid–base reactions is: H+(aq) + OH−(aq) →H2O(l) Another acid–base reaction is the reaction that occurs between sulfuric acid and potassium hydroxide. H 2 SO 4 (aq) + 2 KOH → 2 H O(l) + K SO (aq) Notice the pattern of acid and base reacting to form water and a salt. Acid + Base →Water + Salt When writing equations for acid–base reactions, write the formula of the salt using the procedure for writing formulas of ionic compounds presented in Section What is Salt in Chemistry? Salt is an ionic compound that contains a cation (base) and an anion (acid). It is present in large quantities in seawater, where it is the main mineral constituent. Salt is essential for animal life and saltiness is one of the basic human tastes. Salt is an ionic compound that has a cation other than H + and an anion other than OH – and is obtained along with water in the neutralization reaction between acids and bases. Eg:- NaCl, CuCl2 etc. Acid + Base → Salt + water 9 Sodium chloride is one of the best-known salt. One salt is known to almost everyone because of its widespread use in every day. Types of Salt 1. Acidic salt – The salt formed by partial neutralization of a diprotic or a polyprotic acid is known as an acidic salt. These salts have ionizable H + ion along with another cation. Mostly the ionizable H + is a part of the anion. Some acid salts are used in baking. For eg:- NaHSO 4 , KH 2 PO 4 , NH 4 Cl, CuSO 4 etc. H2 SO4 + NaOH → NaHSO4 + H2 O 2. Basic or Alkali Salt – The salt formed by the partial neutralization of a strong base by a weak acid is known as a basic salt. They hydrolyze to form a basic solution. It is because when hydrolysis of basic salt takes place, the conjugate base of the weak acid is formed in the solution. For example: - Na2 S, CH 3 COONa 3. Double salt – The salts that contain more than one cation or anion are known as double salt. They are obtained by the combination of two different salts crystallized in the same ionic lattice. For example: - Potassium sodium tartrate (KNaC4 H4 O 6.4H2 O) also known as Rochelle salt K 2 SO 4.Al2 (SO 4 )3. 24H2 O (Potash Alum). 4. Mixed Salts – The salt that consists of a fixed proportion of two salts, often sharing either a common cation or common anion is known as mixed salt. For example :- CaOCl2 Properties of Salt Colour: colourless, transparent, or white. Taste: different salts can depict all the five basic tastes, sour, sweet, salty, bitter, and even a hint of savoury or umami. Odour: Odourless 10  Melting Point/Boiling Point: The melting point of sodium chloride is 800 degrees Celsius or 1,072 degrees Kelvin. The boiling point of sodium chloride or salt is 1,737 degrees Kelvin or 1,464 degrees Celsius. (roughly) Saltwater contains ions and is a fairly good conductor of electricity. Hydrolysis of a Salt Hydrolysis of salt refers to the reaction of salt with water. It is the reverse of a neutralization reaction. In this reaction, when salt undergoes reaction with water, the constituent acid and base are formed as products. In hydrolysis, the salt dissociates to form ions, completely or partially depending upon the solubility product of that salt. Salt is a compound formed by the neutralisation reaction between an acid and a base. They generally ionise in water furnishing cations and anions. The cations or anions formed during the ionisation of salts either exist as hydrated ions in aqueous solutions or interact with water to regenerate the acids and bases. The process of interaction between cations or anions of salts and water is known as hydrolysis of salts. Based on hydrolysis, salts are divided into three categories: Acidic salts Basic salts Neutral salts Let us discuss hydrolysis of salts of the following types: Salts of strong acid and strong base: Salts formed by the neutralisation of strong acid and strong base are neutral in nature as the bonds in the salt solution will not break apart. They generally get hydrated but do not hydrolyse. Therefore, such salts are generally known as neutral salts. For example: NaCl Salts of weak acid and strong base: Salts formed by the neutralisation of weak acid and strong base are basic in nature. For example: CH3 COONa CH3 COONa (aq) → CH3 COO − (aq) + Na+ (aq) 11  Acetate ion formed undergoes hydrolysis to form acetic acid and OH – ions. CH3 COO−(aq) + H2 O ⇋ CH3 COOH(aq) + OH−(aq) As we know acetic acid is a weak acid, it remains unionised in the solution. This results in an increase in the concentration of OH – ions which makes the solution alkaline. The pH of the solution is greater than 7. Salts of strong acid and weak base: Salts formed by the neutralisation of strong acid and weak base are acidic in nature. For example: NH4 Cl NH4 Cl(aq) → Cl−(aq) + NH4 +(aq) Ammonium ion formed undergoes hydrolysis to form ammonium hydroxide and H+ ions. NH4 +(aq) + H2 O ⇋ NH4 OH(aq) + H+(aq) As we know ammonium hydroxide is a weak base, it remains unionised in the solution. This results in an increase in the concentration of H+ ions which make the solution acidic. the pH of such solutions is less than 7. Salts of a weak acid and weak base: Salts formed by the neutralisation of weak acid and weak base are acidic, basic or neutral, depending on the nature of acids and bases involved. For example: CH3 COONH4. A general mechanism for the hydrolysis of ions formed from these salts: CH3 COO− + NH4 + + H2 O ⇋ CH3 COOH + NH4 OH It is a bit more complex and will require the K a and K b to be taken into account. Whichever is the stronger acid will be the dominate factor in determining whether it is acidic or basic. The cation will be the acid, and the anion will be the base and will form a hydronium ion or a hydroxide ion depending on which ion reacts more readily with the water. The degree of hydrolysis in such cases is independent of the concentration of the solution. 12 Solubility of salts Different types of salt Solubility Nitrate salts Soluble Ammonium salts Soluble Chloride salts Soluble excluding AgCl, Hg2Cl2, PbCl2 Sulphate salts Soluble excluding PbSO 4, CaSO 4, BaSO 4, and Ag2SO 4 Carbonate salts Insoluble excluding Na2CO3, K 2CO 3, (NH 4)2CO 3 Lead(ll) salts Insoluble excluding Pb(NO3)2, Pb(CH 3COO)2 Ethanoate salts Soluble Sodium and potassium salts Soluble 13