Objectives Theories of acids and Bases Define acids and bases according to the Brønsted – Lowry and Lewis theories Deduce whether or not a species could act as a Brønsted – Lowry and / or a Lewis acid or base Deduce the formula of the conjugate acid (or base) of any Brønsted – Lowry base (or acid)
The history of acid-base theory • There have been many attempts to define acids and bases. • The sour taste and the effect on vegetable colourings, such as litmus, characterized acids. • The soapy feel and detergent power characterized alkalis. • Acids seem to react with alkalis and also with some other compounds to give salts. • The term base replaced the term alkali as meaning the opposite of an acid.
A base was defined as a substance which would react with an acid to form a salt and water. • An acid was defined by Liebig in 1838 as a compound that contains hydrogen that can be displaced by a metal. • A big advance was the Ostwald-Arrhenius theory of electrolytic dissociation in 1880. • They defined acids as substances that produce hydrogen ions in solution. • Bases, they said, produce hydroxide ions in solution and neutralize acid by the reaction: H+ + OH- H2O
Before long this theory ran into difficulties. It was noted that, while hydrochloric acid conducts electricity, pure hydrogen chloride does not. • What is the significance of conducting electricity? • Should hydrogen chloride be classified as an acid or does it only become an acid on contact with water? • Also bases such as ammonia neutralize acids by picking up hydrogen ions, rather than by providing hydroxide ions. NH3 + H+ NH4+
As work proceeded, doubt was cast on the very existence of the hydrogen ion, H+, in solution. • The proton H+, is very small (10-15m diameter) compared with other cations (around 10-10m diameter). • The electric field around it is so intense that it attracts any molecule with unshared electrons, such as H2O • The reaction H+ + H2O H3O+ was shown to liberate 1300kJ mol-1. • As the reaction was so exothermic, it gave evidence that unhydrated protons can not exist in solution. • The hydrated proton, H3O+, is called the oxonium ion, and is also referred to as the hydrogen ion.
The Brønsted-Lowry definition of an acid The Brønsted – Lowry definition of acids and bases are about competition for protons, H+ This defines an acidas a substance that can donate a proton (a H+ ion ) and a base as a substance that can accept a proton. Acids and bases can only react in pairs – one acid and one base.
H H H : : O H N : : H N+ H : O- H : H H H Example 1 • Ammonia acts as a base in water by accepting a proton from a molecule of water, using its lone pair to form the bond. • In this case the water is acting as an acid + +
HCl + H2O H3O+ Cl- + H : O : H Example 2 • Hydrogen chloride is a covalently bonded gas, but when it dissolves in water, the water acts as a base and hydrogen chloride donates a proton to a molecule of water, acting as an acid. A proton has no electrons at all, so to form a covalent bond with another species, that species must have a lone pair of electrons. This is the case with water:
Conjugate acid-base pairs • The Cl- ion left after HCl has donated a proton is itself able to accept a proton (to go back to HCl) and is therefore a base. • It is called the conjugate base of HCl. • In the same way, the NH4+ ion is an acid as it can donate a proton to return to being NH3. • It is the conjugate acid of ammonia. • So we have two acid-base pairs. HCl + NH3 Cl- + NH4+ acid base conjugate base conjugate acid
Water as an acid and a base • HCl can donate a proton to water, so that water acts as a base. • Write the equation to show this. HCl + H2O H3O+ + Cl- Water may also act as an acid. For example Here OH- is the conjugate base and H2O the acid NH4+ + NH3 OH- + H2O Water is described as an amphoteric or amphiprotic solvent
Quick questions • Hydrogen bromide, HBr, is acidic. What is its conjugate base? • OH- is a base. What is its conjugate acid? • Identify the conjugate acid/base pairs in: HNO3 + OH- → NO3- + H2O • Magnesium oxide is described as basic whereas sodium hydroxide is an alkali. Explain the difference? • What species are formed when the following bases accept a proton? • OH- • NH3 • H2O • Cl-
The Lewis definition of acids and Bases • There are reactions which appear to us, on common-sense grounds, to be acid-base reactions, and which do not come within the scope of the Brønsted-Lowry definition. • Such reactions are: CaO + SO3 CaSO4 NH3 + BF3 NH3BF3
H F F H H N H N : + B B F F F F H H • To accommodate reactions of this type, G N Lewis(during 1930 to 1940) proposed a fresh definition of acids and bases. • He described acid-base reactions as reactions in which an unshared electron pair in the base molecule is accepted by the acid molecule, with the formation of a covalent bond. • For example, in this reaction Ammonia is the base, and boron trifluoride is the acid
The Lewis definition of a base includes the Brønsted-Lowry bases because a species with a lone pair of electrons will accept a proton from a Brønsted-Lowry acid. • Lewis acids, such as BF3 and SO3, are not acids in the Brønsted-Lowry sense, and acids such as HCl, H2SO4 and CH3CO2H are not acids according to the Lewis definition • The Brønsted-Lowry description of acids and bases lends itself readily to a quantitative treatment of the strengths of acids and bases. • No such quantitative treatment is possible for Lewis acids and bases.
Objectives Strong and weak acids and bases Distinguish between strong and weak acids and bases in terms of the extent of dissociation, reaction with water and electrical conductivity State whether a given acid or base is strong or weak Distinguish between strong and weak acids and bases, and determine the relative strengths of acids and bases using experimental data
Strong acids • The acidity of a substance depends on the concentration of H+(aq). This is what pH measures. • When acids such as hydrochloric and nitric acid dissolve into water they break up completely into ions. • This is called complete dissociation. • For example: HCl(g) H+(aq) + Cl-(aq) HNO3(aq) H+(aq) + NO3-(aq)
Acids that completely dissociate into ions in aqueous solutions are called strong acids. • The word strong refers only to how much the acid dissociates and not in any way to how concentrated it is. • Can you think of a reason why hydrogen chloride dissolved into an organic liquid does not have acidic qualities?
Weak acids • Weak acids are not fully dissociated when dissolved in water. • Ethanoic acid (the acid in vinegar, also called acetic acid) is a good example. • In a 1 mol dm-3 solution of ethanoic acid, only about 4 in every thousand ethanoic acid molecules are dissociated into ions. • We say the degree of dissociation is 4/1000; the rest remain dissolved as covalently bonded molecules.
CH3CO2H(aq) H+(aq) + CH3CO2-(aq) Before dissociation 0 0 1000 996 4 4 at equilibrium • In fact an equilibrium is set up: Acids like this are called weak acids – weak refers only to the degree of dissociation. We can have a very dilute solution of a strong acid, or a very concentrated solution of a weak acid. Strength and concentration are independent.
Bases • In aqueous solutions, bases dissociate to produce the OH- ion (which can react with a H+ ion by forming a bond through the electrons of one of its lone pairs). • Bases can also be classified as weak or strong. • Strong bases are completely dissociated into ions in aqueous solutions, while weak bases are only partially dissociated.
Dissociation of strong and weak bases • Sodium hydroxide is a strong base; NaOH(aq) Na+(aq) + OH-(aq) • A solution of ammonia in water (NH3(aq)) , also called ammonium hydroxide, NH4OH, is a weak base: NH3(aq) + H2O(l) NH4+(aq) +OH-(aq) Or NH4OH(aq) NH4+(aq) + OH-(aq)
Examples of strong acids and bases • Examples of strong acids and bases include: Note: because one mole of HCl produces one mole of H+ ions it is known as monoprotic (monobasic). Sulphuric acid is known as a diprotic (dibasic) acid, as one mole of sulphuric acid produces two moles of hydrogen ions.
Examples of weak acids and bases • Weak acids and bases are only slightly dissociated (ionized) into their ions in aqueous solution:
Experiments to distinguish between strong and weak acids and bases • It is important for us to be able to distinguish between strong and weak acids and bases, ways in which we can do this are: • pH measurement • Conductivity measurement • Using universal indicator • Observation from vigour of reaction
pH measurement • Because a strong acid produces a higher concentration of hydrogen ions than a weak acid, with the same concentration, the pH of a strong acid will be lower than a weak acid. • Similarly a strong base will have a higher pH than a weak base, with the same concentration. • The most accurate way to determine pH is to use a pH meter 0.1 moldm-3 HCl(aq) pH = 1.0 0.1 moldm-3 CH3COOH(aq) pH = 2.9
Conductivity measurement • Strong acids and strong bases in solution will give much higher readings on a conductivity meter than equimolar (equal concentration) solutions of weak acids or bases, because they contain more ions in solution • Using Universal indicator • Strong acids produce a red colourwith universal indicator weak acids produce orange/yellow colours • Strong bases produce a purple colourwith U.I. weak bases produce a blue colour
Observation from reaction • Reacting an acid with a carbonate or a hydrogen carbonate will produce carbon dioxide gas . • The reaction will be more vigorous with a stronger acid compared to a weaker acid. HCl + Na2CO3 2NaCl + H2O + CO2 will be more vigorous than 2CH3COOH + Na2CO3 2CH3COONa + H2O + CO2
Acid – base reactions and Le Chatelier’s principle CH3COOH + NaOH → CH3COONa + H2O HCl + NaOH → NaCl + H2O • The volume of alkali in both cases will be 25 cm3 • Although there are fewer hydrogen ions in the ethanoic acid solution, once they react with hydroxide ions then, according to Le Chatelier’s principle, more of the acid dissociates to replace them until eventually all of the acid has been neutralized. x cm3 25cm3 0.10mol dm-3 0.10mol dm-3 x cm3 25cm3 0.10mol dm-3 0.10mol dm-3
Titrations • The only difference observed in titrations involves temperature rather than volume. • Strong acids and bases in aqueous solution always release the same amount of heat per mole of hydrogen ions neutralized, because the only reaction taking place is the neutralisation of hydrogen ions by hydroxide ions. • Titrations involving weak acids or bases will release a different amount of heat, because heat energy is required to dissociate the molecules to form ions, and heat energy is released when these ions become hydrated, in addition to the heat evolved during neutralisation.
Objectives Outline the characteristic properties of acids and bases in aqueous solution With reactive metals With neutralization reactions with bases With carbonates With hydrogen carbonates With indicators
Metal + acid salt + hydrogen MgCl2(aq) + H2 (aq) Mg (s) + 2HCl (aq) Reactions of acids • All acids react to form salts and the reactions may be written as balanced or ionic equations. Metals above hydrogen in the reactivity series react with acids to produce hydrogen gas and the metal salt Write the balanced symbol equation for the reaction between magnesium and hydrochloric acid
Metal + acid salt + hydrogen Reactions of acids • All acids react to form salts and the reactions may be written as balanced or ionic equations. Metals above hydrogen in the reactivity series react with acids to produce hydrogen gas and the metal salt Write the balanced symbol equation for the reaction between magnesium and hydrochloric acid
water Metal oxide + acid salt + MgO (s) + 2HCl (aq) MgCl2(aq) + H2O (l) Metal hydroxide + acid salt + water + H2O (l) NaOH(aq) HCl (aq) NaCl (aq) + Acids react with bases such as metal oxides and hydroxides to form a salt and water. For example: Can you write a balanced symbol equation showing an example of each of these reactions?
If both the acid and the base are strong,– that is, they both completely dissociate in water – for example, the reaction between nitric acid and sodium hydroxide can be written: HNO3(aq) + NaOH(aq) NaNO3(aq) + H2O(aq) • This can also be written as: H+(aq)+ NO3-(aq) + Na+(aq) + OH-(aq) Na+(aq) + NO3-(aq) + H20(l) • We can remove the spectator ions to leave us with: H20(l) H+(aq) + OH-(aq)
The nitrate ions and the sodium ions are ‘spectator ions’ , and so the only reaction that is actually taking place is the neutralization of hydrogen ions by hydroxide ions to form water molecules. • The enthalpy change for the reaction is -57.3kJ mol-1, and it will therefore always be the same when a strong acid is neutralized by a strong base in aqueous solution. • A base can be defined as a substance that neutralizes an acid to form a salt and water, so other acid-base reactions include the reactions of acids with most metal oxides.
Metal carbonate + acid salt + water + carbon dioxide Metal hydrogen carbonate + acid salt + water + carbon dioxide • All carbonates and hydrogen carbonates will neutralize acids to produce carbon dioxide andwater. For example: CaCO3(s) + 2HCl(aq) CaCl2(aq) + H2O(l) + CO2(g) 2NaHCO3(aq) + H2SO4(aq) NaSO4(aq) 2H2O(l) + 2CO2(g) • The carbonate ion (the conjugate base of the hydrogen carbonate ion, HCO3-) and the hydrogen carbonate ion (the conjugate base of carbonic acid, H2CO3) both behave as Brønsted-Lowry bases and accept protons from the acid
With indicators • Acid- base indicators can be used to determine whether or not a solution is acidic. • Common indicators include:
Characteristic reactions of bases • The simplest definition of a base is that it is asubstance that can neutralize an acid. • A base that is soluble in water is known as an alkali. • Typical reactions of bases in solution include the following. • Neutralization of acids • Displacement of ammonia from ammonium salts • With indicators
Displacement of ammonia from ammonium salts NH4Cl(s) + NaOH(aq)NaCl(aq) + NH3(g) + H2O(l) • Ammonia is very soluble in water, so it may be formed in solution, or, if small amounts of sodium hydroxide are added to solid ammonium chloride, it can be collected as a gas. • In this reaction the chloride ions and sodium ions are spectator ions, so the actual reaction taking place in solution is: NH4+(aq) + OH-(aq) NH3(g) + H2O(l)
The equilibrium lies very much on the right, so a single arrow has been shown, but technically the reaction should be written NH4+(aq) + OH-(aq) NH3(g) + H2O(l) • Under the Brønsted-Lowry definition, ammonium ions (NH4+) and water molecules (H2O) are acting as acids, and hydroxide ions (OH-) and ammonia molecules (NH3) are acting as bases. • As water is a weak acid, its conjugate bases is strong, so the position of equilibrium lies very much to the right.
With indicators • If a base is insoluble in water, for example copper (II) oxide (CuO), then there will be no reaction with indicators. • Indicators actually measure the concentration of hydrogen ion in aqueous solution, H+(aq), so only soluble bases (alkalis) that produce hydroxide ions in water will change the colour of indicators. • Increasing the concentration of hydroxide ions will lower the concentration of hydrogen ions H+(aq) + OH-(aq) H2O(l)
Objectives The ionisation of water Understand how water ionises and give the ionic product of water at 298K AHL State the expression for the ionic product of water (Kw). Deduce [H+(aq)] and [OH-(aq)] for water at different temperatures given Kw values.
The ionisation of water • Water is slightly ionised: H2O(l) H+(aq) + OH-(aq) • Or this may be written: H2O(l) + H2O(l) H3O+(aq) + OH-(aq) • This equilibrium is established in water and all aqueous solutions, write the equilibrium expression for this reaction: [H+(aq)]eqm [OH-(aq)]eqm Kc = [H2O(l)]eqm
The equilibrium lies very much over to the left hand side and so we say the concentration of H2O stays constant. • To avoid having two constants (Kc and the concentration of water) in the same expression, a new equilibrium constant is defined. This is called the ionic product for water and is given the symbol Kw Kw = [H+(aq)]eqm [OH-(aq)]eqm • The value of Kw is usually taken to be 1x10-14 mol2 dm-6 at 298K
Each H2O that dissociates (splits up) gives rise to one H+ and one OH- so, in pure water at 298K [H+(aq)]eqm = [OH-(aq)]eqm So, 1x10-14 = [H+(aq)]2eqm [H+(aq)] = 1x10-7mol dm-3 = [OH-(aq)]eqm • Thismeans that only 2 in every 1000,000,000 water molecules is split up into hydrogen and hydroxide ions!!!
Kw and temperature • If Kw is 5.13 x 10-13 mol2 dm-6 what is the concentration of [H+] Kw = [H+]2 [H+]2 = 5.13 x 10-13 [H+] = 5.13 x 10-13 [H+] =
Objectives The pH scale Distinguish between aqueous solutions that are acidic, neutral or alkaline using the pH scale Identify which of two or more aqueous solutions is more acidic or alkaline using pH scales. State that each change of one pH unit represents a 10-fold change in the hydrogen ion concentration [H+(aq)] Deduce changes in [H+(aq)] when the pH of a solution changes by more than one pH unit
The pH Scale • The acidity of a solution depends on the concentration of H+(aq) and is measured on the pH scale. pH is defined as –log10 [H+(aq)] Remember that square brackets, [ ], mean the concentration in mol dm-3. This expression is more complicated than simply stating the concentration of H+(aq), but it does away with awkward numbers like 10-13, which occur because the concentration of H+(aq) solutions is so small. The minus sign makes almost all pH values positive (because the logs of numbers less than 1 are negative).
On the pH scale: • The smaller the pH, the greater the concentration of H+ (aq) • A difference of one pH number means a tenfold difference in acidity, so that pH 2 is ten times as acidic as pH 3 • pH measures alkalinity as well as acidity, because as [H+(aq)] goes up, [OH-(aq)] goes down. • If a solution contains more H+(aq) than OH-(aq), its pH will be less than 7 and we call it acidic. • If a solution contains more OH-(aq) than H+(aq), its pH will be greater than 7 and we call it alkaline