ACIDS_&_BASES_MAH_(NXPowerLite) Opp (1)

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Unformatted text preview: •100’s of free ppt’s from www.pptpoint.com library Compiled by MAH 1 Topic 8: Acids and Bases (6 hours) • 8.1 Theories of Acids and Bases • 8.1.1 Define acids and bases according to the • • • • Brønsted–Lowry and Lewis theories. 8.1.2 Deduce whether or not a species could act as a Brønsted–Lowry and/or a Lewis acid or base. 8.1.3 Deduce the formula of the conjugate acid (or base) of any Bronsted­Lowry acid (or base) The members of a conjugate acid–base pair always differ by a single proton (H+). Students should make clear the location of the proton transferred, e.g. CH3COOH/CH3COO­ rather than C2H4O2/C2H3O2 ­ 2 Naming acids 3 Acid & Bases Before you start it would be helpful to… • know the simple properties of acids, bases and alkalis 4 ACIDS AND BASES ACIDS BRØNSTED-LOWRY THEORY ACID proton donor HCl ——> H+(aq) + Cl¯(aq) BASE proton acceptor NH3 (aq) + H+(aq) Conjugate systems Acids are related to bases ACID Bases are related to acids BASE ——> PROTON + PROTON NH4+(aq) + CONJUGATE BASE CONJUGATE ACID 5 ACIDS AND BASES BRØNSTED-LOWRY THEORY ACID proton donor HCl ——> H+(aq) + Cl¯(aq) BASE proton acceptor NH3 (aq) + H+(aq) ——> NH4+(aq) 6 ACIDS AND BASES ACIDS BRØNSTED-LOWRY THEORY ACID proton donor HCl ——> H+(aq) + Cl¯(aq) BASE proton acceptor NH3 (aq) + H+(aq) Conjugate systems Acids are related to bases ACID Bases are related to acids BASE ——> NH4+(aq) PROTON + + CONJUGATE BASE PROTON CONJUGATE ACID For an acid to behave as an acid, it must have a base present to accept a proton... HA acid example + CH3COO¯ + base B base H 2O acid BH+ + A¯ conjugate conjugate acid base CH3COOH acid + OH¯ base 7 The Bronsted­Lowry Theory of acids and bases • • Read the following passage carefully and try to answer the questions at the end. The theory An acid is a proton (hydrogen ion) donor. A base is a proton (hydrogen ion) acceptor. When hydrogen chloride gas dissolves in water to produce hydrochloric acid, the hydrogen chloride molecule gives a proton (a hydrogen ion) to a water molecule. A co­ordinate (dative covalent) bond is formed between one of the lone pairs on the oxygen and the hydrogen from the HCl. Hydroxonium ions, H3O+, are produced. 8 9 It is important to realise that whenever you talk about hydrogen ions in solution, H+(aq), what you are actually talking about are hydroxonium ions/oxonium ions 10 10 The hydrogen chloride / ammonia reaction: • Whether you are talking about the reaction in solution or in the gas state, ammonia is a base because it accepts a proton (a hydrogen ion). The hydrogen becomes attached to the lone pair on the nitrogen of the ammonia via a co­ordinate bond. 11 11 • If it is in solution, the ammonia accepts a proton • • from a hydroxonium ion: If the reaction is happening in the gas state, the ammonia accepts a proton directly from the hydrogen chloride: Either way, the ammonia acts as a base by accepting a hydrogen ion from an acid. 12 12 Conjugate pairs • When hydrogen chloride dissolves in water, almost 100% of it reacts with the water to produce hydroxonium ions and chloride ions. Hydrogen chloride is a strong acid, and we tend to write this as a one­way reaction: • But for a weak acid HA you have to think of the reaction with water as being reversible. 13 13 Conjugate pairs continued • Thinking about the forward reaction: • The HA is an acid because it is donating a • • • • proton (hydrogen ion) to the water. The water is a base because it is accepting a proton from the HA. But there is also a back reaction between the hydroxonium ion and the A­ ion: The H3O+ is an acid because it is donating a proton (hydrogen ion) to the A­ ion. The A­ ion is a base because it is accepting a proton from the H3O+ 14 14 The reversible reaction contains two acids and two bases. We think of them in pairs, called conjugate pairs. 15 15 • When the acid, HA, loses a proton it forms a • • • base, A­. When the base, A­, accepts a proton back again, it obviously refoms the acid, HA. These two are a conjugate pair. Members of a conjugate pair differ from each other by the presence or absence of the transferable hydrogen ion. If you are thinking about HA as the acid, then A­ is its conjugate base. If you are thinking about A­ as the base, then HA is its conjugate acid. 16 16 • The water and the hydroxonium ion are also a • conjugate pair. Thinking of the water as a base, the hydroxonium ion is its conjugate acid because it has the extra hydrogen ion which it can give away again. Thinking about the hydroxonium ion as an acid, then water is its conjugate base. The water can accept a hydrogen ion back again to reform the hydroxonium ion. 17 17 A second example of conjugate pairs 18 18 • Think first about the forward reaction. Ammonia is a • • base because it is accepting hydrogen ions from the water. The ammonium ion is its conjugate acid ­ it can release that hydrogen ion again to reform the ammonia. The water is acting as an acid, and its conjugate base is the hydroxide ion. The hydroxide ion can accept a hydrogen ion to reform the water. Looking at it from the other side, the ammonium ion is an acid, and ammonia is its conjugate base. The hydroxide ion is a base and water is its conjugate acid. 19 19 Amphoteric substances • You may possibly have noticed (although probably not!) • that in one of the last two examples, water was acting as a base, whereas in the other one it was acting as an acid. A substance which can act as either an acid or a base is described as being amphoteric. 20 20 8.1 Theories of Acids and Bases • • • • • At the end of this topic you should be able to Define what a Brønsted–Lowry acid is and give a chemical equation showing a Bronsted­Lowry acid reacting: Define what a Brønsted–Lowry base is and give a chemical equation showing a Bronsted­Lowry base reacting: Define what a Lewis acid is and give a suitable chemical equation to illustrate your defiinition: Define what a Lewis base is giving chemical equations to support your defiinition: 21 21 The basic definitions 22 22 The Lewis Theory of acids and bases The theory • An acid is an electron pair acceptor. • A base is an electron pair donor. 23 23 ACIDS AND BASES LEWIS THEORY ACID lone pair acceptor BF3 H+ BASE lone pair donor NH3 H2O LONE PAIR DONOR LONE PAIR DONOR AlCl3 LONE PAIR ACCEPTOR LONE PAIR ACCEPTOR 24 24 The relationship between the Lewis theory and the Bronsted­Lowry theory • Lewis bases • It is easiest to see the relationship by looking at exactly what Bronsted­Lowry bases do when they accept hydrogen ions. Three Bronsted­Lowry bases we've looked at are hydroxide ions, ammonia and water, and they are typical of all the rest. • The Bronsted­Lowry theory says that they are acting as bases because they are combining with hydrogen ions. The reason they are combining with hydrogen ions is that they have lone pairs of electrons ­ which is what the Lewis theory says. The two are entirely consistent. • So how does this extend the concept of a base? At the moment it doesn't ­ it just looks at it from a different angle. 25 25 • But what about other similar reactions of ammonia or water, for example? On the Lewis theory, any reaction in which the ammonia or water used their lone pairs of electrons to form a co­ordinate bond would be counted as them acting as a base. • Here is a reaction which you will find talked about on the page dealing with co­ordinate bonding. Ammonia reacts with BF3 by using its lone pair to form a co­ordinate bond with the empty orbital on the boron. 26 26 27 27 28 28 • As far as the ammonia is concerned, it is behaving exactly the same as when it reacts with a hydrogen ion ­ it is using its lone pair to form a co­ordinate bond. If you are going to describe it as a base in one case, it makes sense to describe it as one in the other case as well. 29 29 Lewis acids • Lewis acids are electron pair acceptors. In the • • above example, the BF3 is acting as the Lewis acid by accepting the nitrogen's lone pair. On the Bronsted­Lowry theory, the BF3 has nothing remotely acidic about it. This is an extension of the term acid well beyond any common use. What about more obviously acid­base reactions ­ like, for example, the reaction between ammonia and hydrogen chloride gas? 30 30 • The lone pair on the nitrogen of an ammonia molecule is • • attracted to the slightly positive hydrogen atom in the HCl. As it approaches it, the electrons in the hydrogen­ chlorine bond are repelled still further towards the chlorine. Eventually, a co­ordinate bond is formed between the nitrogen and the hydrogen, and the chlorine breaks away as a chloride ion. This is best shown using the "curly arrow" notation commonly used in organic reaction mechanisms. The whole HCl molecule is acting as a Lewis acid. It is accepting a pair of electrons from the ammonia, and in the process it breaks up. Lewis acids don't necessarily have to have an existing empty orbital. 31 31 Now it’s question time! 32 32 33 33 34 34 Lewis acid type questions 35 35 36 36 3. Choosing suitable examples from the following: • NH3, O2− , Cu2+, OH–, NH2­, H2O explain, using a different equation in each case, the meaning of the terms below. • (i) Brønsted­Lowry acid (2) • (ii) Lewis acid (2) • (iii) conjugate acid­base pair (Identify each member of both acid­base pairs.) (3) 37 37 8.2 Properties of Acids and Bases (1h) • 8.2.1 • • • Outline the characteristic properties of acids and bases in aqueous solution. Bases that are not hydroxides, such as ammonia, soluble carbonates and hydrogencarbonates, should be included. Alkalis are bases that dissolve in water. Students should consider the effects on indicators and reactions of acids with bases, metals and carbonates 38 38 Things acids do 39 39 Things alkalis do 40 40 Less useful acid reactions 41 41 Other not so good acid reactions 42 42 8.3 Strong and Weak Acids and Bases (1h) • 8.3.1 Distinguish between strong and weak acids and bases in • • • • • • terms of the extent of dissociation, reaction with water and electrical conductivity. The term ionization can be used instead of dissociation. Solutions of equal concentration can be compared by pH and/or conductivity Aim 8: Although weakly acidic solutions are relatively safe, they can still cause damage over long periods of time. Students could consider the effects of acid deposition on limestone buildings and living things. 8.3.2 State whether a given acid or base is strong or weak. Students should consider hydrochloric acid, nitric acid and sulfuric acid as examples of strong acids and ethanoic acid and carbonic acid (aqueous carbon dioxide) as weak acids. Students should consider all group 1 hydroxides and barium hydroxide as strong bases and ammonia and amines as weak bases. 8.3.3 Distinguish between strong and weak acids and bases, and 43 43 determine the relative strengths of acids and bases, using STRONG ACIDS AND BASES STRONG ACIDS completely dissociate (split up) into ions in aqueous solution e.g. HCl ——> H+(aq) + Cl¯(aq) MONOPROTIC 1 replaceable H HNO3 ——> H+(aq) + NO3¯(aq) H2SO4 ——> 2H+(aq) + SO42-(aq)DIPROTIC 2 replaceable H’s 44 44 STRONG ACIDS AND BASES STRONG ACIDS e.g. completely dissociate (split up) into ions in aqueous solution HCl ——> H+(aq) + Cl¯(aq) MONOPROTIC 1 replaceable H DIPROTIC 2 replaceable H’s HNO3 ——> H+(aq) + NO3¯(aq) H2SO4 ——> 2H+(aq) + SO42-(aq) STRONG BASES e.g. completely dissociate into ions in aqueous solution NaOH ——> Na+(aq) + OH¯(aq) 45 45 WEAK ACIDS Weak acids partially dissociate into ions in aqueous solution e.g. ethanoic acid When a weak acid dissolves in water an equilibrium is set up CH3COOH(aq) HA(aq) + H2O(l) CH3COO¯(aq) + H+(aq) A¯(aq) + H3O+(aq) The water stabilises the ions To make calculations easier the dissociation can be written... HA(aq) A¯(aq) + H+(aq) 46 46 WEAK ACIDS Weak acids partially dissociate into ions in aqueous solution e.g. ethanoic acid CH3COOH(aq) When a weak acid dissolves in water an equilibrium is set up CH3COO¯(aq) HA(aq) + H2O(l) + H+(aq) A¯(aq) + H3O+(aq) The water stabilises the ions To make calculations easier the dissociation can be written... The weaker the acid HA(aq) A¯(aq) + H+(aq) the less it dissociates the more the equilibrium lies to the left. The relative strengths of acids can be expressed as Ka or pKa values The dissociation constant for the weak acid HA is Ka = [H+(aq)] [A¯(aq)] mol dm-3 [HA(aq)] 47 47 WEAK BASES Partially react with water to give ions in aqueous solution e.g. ammonia When a weak base dissolves in water an equilibrium is set up NH3 (aq) + H2O (l) NH4+ (aq) + OH¯ (aq) as in the case of acids it is more simply written NH3 (aq) + H+ (aq) NH4+ (aq) 48 48 WEAK BASES Partially react with water to give ions in aqueous solution e.g. ammonia When a weak base dissolves in water an equilibrium is set up NH3 (aq) + H2O (l) NH4+ (aq) + OH¯ (aq) as in the case of acids it is more simply written NH3 (aq) + The weaker the base H+ (aq) NH4+ (aq) the less it dissociates the more the equilibrium lies to the left The relative strengths of bases can be expressed as Kb or pKb values. 49 49 STRONG & WEAK • This is concerned with how the acid dissociates into ions • Strong acids dissociate almost completely • These includes hydrochloric acid, sulfuric acid and nitric acid 50 50 51 51 relationship of acid strength and conjugate base strength for the dissociation reaction. 52 52 WEAK ACIDS • These do not dissociate completely and the weaker they are, the less they dissociate Ethanoic acid (acetic acid) is weak and is the main ingredient in vinegar. 53 53 8.4 The pH Scale (1h) • 8.4.1 Distinguish between aqueous solutions that are acidic, neutral or • alkalin using the pH scale. 8.4.2 Identify which of two or more aqueous solutions is more acidic or alkaline, using pH values. • Students should be familiar with the use of a pH meter or pH paper. • 8.4.3 State that each change of one pH unit represents a tenfold change in the hydrogen ion concentration [H+(aq)] • Relate integral values of pH to [H+(aq)] expressed as powers of ten. • Calculation of pH from [H+(aq)] is not required • 8.4.4 Deduce changes in [H+(aq)] when the pH of a solution changes by more than one pH unit. 54 54 55 55 pH differences between acids and bases Notice every 1 unit change in pH changes the concentration X10!!! 56 56 57 57 Hydrogen ion concentration [H+(aq)] Introduction hydrogen ion concentration determines the acidity of a solution hydroxide ion concentration determines the alkalinity for strong acids and bases the concentration of ions is very much larger than their weaker counterparts which only partially dissociate. 58 58 Hydrogen ion concentration [H+(aq)] pH hydrogen ion concentration can be converted to pH to convert pH into hydrogen ion concentration [H+(aq)] = antilog (-pH) An equivalent calculation for bases converts the hydroxide ion concentration to pOH pOH pH = - log10 [H+(aq)] pOH = - log10 [OH¯(aq)] in both the above, [ ] represents the concentration in mol dm-3 [H+] 100 OH¯ 10-14 10-13 10-12 10-11 10-10 10-9 10-8 10-7 10-6 10-5 10-4 10-3 10-2 10-1 10-0 pH 0 10-1 10-2 10-3 10-4 10-5 10-6 10-7 10-8 10-9 10-10 10-11 10-12 10-13 10-14 1 2 STRONGLY ACIDIC 3 4 5 WEAKLY ACIDIC 6 7 8 NEUTRAL 9 10 WEAKLY ALKALINE 11 12 13 14 STRONGLY ALKALINE 59 59 Ionic product of water - Kw Despite being covalent, water conducts electricity to a very small extent. This is due to the slight ionisation ... H2O(l) + H2O(l) or, more simply H2O(l) H3O+(aq) + OH¯(aq) H+(aq) + OH¯(aq) 60 60 Ionic product of water - Kw Despite being covalent, water conducts electricity to a very small extent. This is due to the slight ionisation ... H2O(l) + H2O(l) or, more simply Applying the equilibrium law to the second equation gives [ ] is the equilibrium concentration in mol dm-3 H3O+(aq) + OH¯(aq) H2O(l) Kc H+(aq) + OH¯(aq) = [H+(aq)] [OH¯(aq)] [H2O(l)] 61 61 Ionic product of water - Kw Despite being covalent, water conducts electricity to a very small extent. This is due to the slight ionisation ... H2O(l) + H2O(l) or, more simply Applying the equilibrium law to the second equation gives H3O+(aq) + OH¯(aq) H2O(l) Kc [ ] is the equilibrium concentration in mol dm-3 H+(aq) + OH¯(aq) = [H+(aq)] [OH¯(aq)] [H2O(l)] As the dissociation is small, the water concentration is very large compared with the dissociated ions and any changes to its value are insignificant; its concentration can be regarded as constant. This “constant” is combined with (Kc) to get a new constant (Kw). Kw = [H+(aq)] [OH¯(aq)] mol2 dm-6 = 1 x 10-14 mol2 dm-6 (at 25°C) Because the constant is based on an equilibrium, Kw VARIES WITH TEMPERATURE 62 62 Ionic product of water - Kw The value of Kw varies with temperature because it is based on an equilibrium. Temperature / °C 0 20 25 30 60 Kw / 1 x 10-14 mol2 dm-6 0.11 0.68 1.0 1.47 5.6 H+ / x 10-7 mol dm-3 pH 0.33 7.48 0.82 7.08 1.0 7 1.27 6.92 2.37 6.63 What does this tell you about the equation H2O(l) H+(aq) + OH¯(aq) ? 63 63 Ionic product of water - Kw The value of Kw varies with temperature because it is based on an equilibrium. Temperature / °C 0 20 25 30 60 Kw / 1 x 10-14 mol2 dm-6 0.11 0.68 1.0 1.47 5.6 H+ / x 10-7 mol dm-3 pH 0.33 7.48 0.82 7.08 1.0 7 1.27 6.92 2.37 6.63 What does this tell you about the equation H2O(l) H+(aq) + OH¯(aq) ? • Kw gets larger as the temperature increases • this means the concentration of H+ and OH¯ ions gets greater • this means the equilibrium has moved to the right • if the concentration of H+ increases then the pH decreases • pH decreases as the temperature increases 64 64 Ionic product of water - Kw The value of Kw varies with temperature because it is based on an equilibrium. Temperature / °C 0 20 25 30 60 Kw / 1 x 10-14 mol2 dm-6 0.11 0.68 1.0 1.47 5.6 H+ / x 10-7 mol dm-3 pH 0.33 7.48 0.82 7.08 1.0 7 1.27 6.92 2.37 6.63 What does this tell you about the equation H2O(l) H+(aq) + OH¯(aq) ? • Kw gets larger as the temperature increases • this means the concentration of H+ and OH¯ ions gets greater • this means the equilibrium has moved to the right • if the concentration of H+ increases then the pH decreases • pH decreases as the temperature increases Because the equation moves to the right as the temperature goes up, it must be an ENDOTHERMIC process 65 65 Relationship between pH and pOH Because H+ and OH¯ ions are produced in equal amounts when water dissociates their concentrations will be the same. [H+] = [OH¯] = 1 x 10-7 mol dm-3 Kw = [H+] [OH¯] = 1 x 10-14 mol2 dm-6 take logs of both sides log[H+] + log[OH¯] = -14 multiply by minus - log[H+] - log[OH¯] = 14 change to pH and pOH pH + pOH = 14 (at 25°C) 66 66 Relationship between pH and pOH Because H+ and OH¯ ions are produced in equal amounts when water dissociates their concentrations will be the same. [H+] = [OH¯] = 1 x 10-7 mol dm-3 Kw = [H+] [OH¯] = 1 x 10-14 mol2 dm-6 take logs of both sides multiply by minus - log[H+] - log[OH¯] = 14 change to pH and pOH N.B. log[H+] + log[OH¯] = -14 pH + pOH = 14 (at 25°C) As they are based on the position of equilibrium and that varies with temperature, the above values are only true if the temperature is 25°C (298K) Neutral solutions may be regarded as those where [H+] = [OH¯]. Therefore a neutral solution is pH 7 only at a temperature of 25°C (298K) Kw is constant for any aqueous solution at the stated temperature 67 67 THE END 68 68 Indicators 69 69 ...
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This note was uploaded on 11/17/2011 for the course PHYS 121 taught by Professor Burgeson during the Fall '11 term at BYU.

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