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IB CHEMISTRY Topic 8 Acids and bases Higher level
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8.1 Theories of acids and bases
OBJECTIVES • A Brønsted–Lowry acid is a proton/H+ donor and a Brønsted–Lowry base is a proton/H+ acceptor. • Amphiprotic species can act as both Brønsted–Lowry acids and bases. • A pair of species differing by a single proton is called a conjugate acid-base pair. • Deduction of the Brønsted–Lowry acid and base in a chemical reaction. • Deduction of the conjugate acid or conjugate base in a chemical reaction.
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NOS: Acid-Base Hypothesis
Greeks classifying substances based on taste – sweet, salty, sour.... Alkalis were known to neutralize the sour taste 1777 Antoine Lavoiser burnt phosphorus and sulfur with oxygen. These compounds were acidic when dissolved in water. Therefore acids contain oxygen. Created the word oxygen: Oxys = sour, genes = born in Greek. 1809 Humphry Davy questioned this as HCl, H2S, H2Te were acidic with no oxygen.
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Key – Best hypothesis is the one with the most predictive power
Arrhenius proposed that salts form ions in solution even without an electric current. Hence acids were possibly substances that released H+ ions, and bases released OH- ions. But this did not explain outside of aqueous solutions or for bases that did not have an –OH group. Bronsted and Lowry – acid is a H+ donor, base is a H+ acceptor. NH3 can now be explained as a base. Lewis using his dot structures defined acids as an electron pair acceptor, and a base as an electron pair donor. BF3 + NH3 can now be explained.
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Learning check - match Phosphorus and sulfur oxides make acids Lewis
HCl does not contain oxygen Acids give off H+ ions Acid is a H+ donor, base is a H+ acceptor Acids are electron pair acceptors and and bases are electron pair donors Lewis Lavosier Davy Bronsted and Lowry Arrhenius
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1 of 3. Arrhenius Acids and Bases
Arrhenius Acid – produce H+ ions (or hydronium ions H3O+) eg. HCl Arrhenius Base – produce OH- ions eg. NaOH (Problem: some bases don’t have hydroxide ions!)
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2 of 3. Bronsted-Lowry Acids and Bases
Bonsted-Lowry Acid – proton donor eg. H2O Bronsted-Lowry Base – proton acceptor eg. NH3 A “proton” is really just a hydrogen atom that has lost it’s electron!
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A Brønsted-Lowry acid is a proton donor
A Brønsted-Lowry base is a proton acceptor Base – conjugate acid pair conjugate acid conjugate base base acid Acid – conjugate base pair
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The Brønsted definition means NH3 is a BASE in water — and water is itself an ACID
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Conjugate Pairs
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Amphoteric – can act as either an acid or a base
eg. Al2O3 as a base: Al2O3 + 6 HCl —> 2 AlCl3 + 3 H2O Al2O3 as an acid: Al2O3 + 2NaOH + 3 H2O —> NaAl(OH)4 Amphiprotic – a type of amphoteric substance that specifically can act as either an acid or base by accepting or donating hydrogen ions eg. water, diprotic acids
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Learning Check! Label the acid, base, conjugate acid, and conjugate base in each reaction: HCl + OH- Cl- + H2O H2O + H2SO4 HSO4- + H3O+
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8.2 Properties of acids and bases
OBJECTIVES • Most acids have observable characteristic chemical reactions with reactive metals, metal oxides, metal hydroxides, hydrogen carbonates and carbonates. • Salt and water are produced in exothermic neutralization reactions. • Balancing chemical equations for the reaction of acids. • Identification of the acid and base needed to make different salts. • Candidates should have experience of acid-base titrations with different indicators.
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Properties of Acids Often produce H+ in water (also written as the hydronium ion H3O+ which is a hydrogen ion attached to a water molecule) Taste sour Corrode metals Electrolytes React with bases (metal hydroxides) to form a salt, water React with carbonates to form salt, water, CO2 React with metals/metal oxides to form salt, H2 pH is less than 7 Turns blue litmus paper red, phenolphthalein colourless, methyl orange red
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Key reactions 1. Acid + base salt + water Eg. HCl + NaOH NaCl + H2O *Base can be a metal hydroxide, or a metal oxide. 2. Acid + metal salt + hydrogen Eg. HCl + Mg MgCl2 + H2 (unbalanced) 3. Acid + metal carbonate salt + water + carbon dioxide Eg. HCl + Na2CO3 NaCl + H2O + CO2 (unbalanced)
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Acid Nomenclature Flowchart
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Learning check – nomenclature:
HBr (aq) H2CO3 H2SO3 hydrobromic acid carbonic acid sulfurous acid
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Properties of Bases Often produce OH- ions in water Taste bitter
Are electrolytes Feel soapy, slippery React with acids to form salts and water pH greater than 7 Turns red litmus paper blue, phenolphthalein pink, methyl orange yellow An alkali is a base that dissolves in water
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8.3 The pH scale OBJECTIVES • pH = − log[H+(aq)] and [H+] = 10−pH. • A change of one pH unit represents a 10-fold change in the hydrogen ion concentration [H+]. • pH values distinguish between acidic, neutral and alkaline solutions. • The ionic product constant, Kw = [H+] [OH-] = 10−14 at 298 K. • Solving problems involving pH, [H+] and [OH-]. • Students should be familiar with the use of a pH meter and universal indicator.
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The pH scale is a way of expressing the strength of acids and bases
The pH scale is a way of expressing the strength of acids and bases. Instead of using very small numbers, we just use the NEGATIVE power of 10 on the Molarity of the H+ (or OH-) ion. pH 1 is 10 times stronger than pH 2 pH 1 is 100 times stronger than pH 3 etc… Under 7 = acid = neutral Over 7 = base
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Calculating pH pH = - log10 [H+(aq)] [H+] = 10-pH
where [H+] is the concentration of hydrogen ions in mol dm-3 Working backwards: [H+ (aq)] = antilog (-pH) Or [H+] = 10-pH
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Ionic product constant of water - Kw
Water undergoes very slight ionisation ... H2O(l) ⇌ H+(aq) + OH¯(aq) Applying the equilibrium law: Kc = [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). At 25°C: Kw = [H+(aq)] [OH¯(aq)] mol2 dm-6 = 1 x mol2 dm-6 Because the constant is based on an equilibrium, Kw VARIES WITH TEMPERATURE Also note: pKw = 14 = pH + pOH
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Kw variation with temperature
• 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
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Calculate the pH of 0.02M HCl?
Problem 1: Calculate the pH of 0.02M HCl? HCl H+ + Cl¯ [H+] = 0.02M = 2x10-2 mol dm-3 pH = - log [H+] = 1.7
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Calculate the pH of 0.1M NaOH?
Problem 2: Calculate the pH of 0.1M NaOH? NaOH Na+ + OH¯ [OH¯] = 0.1M = 1x 10-1 mol dm-3 Method1: Kw = [H+][OH¯] = 1 x mol2dm-6 [H+] = Kw / [OH¯] = 1 x moldm-3 pH = - log[H+] = 13 Method 2: pOH = 1 pOH + pH = 14 1 + pH = 14 pH = 13
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8.4 Strong and weak acids and bases
OBJECTIVES • Strong and weak acids and bases differ in the extent of ionization. • Strong acids and bases of equal concentrations have higher conductivities than weak acids and bases. • A strong acid is a good proton donor and has a weak conjugate base. • A strong base is a good proton acceptor and has a weak conjugate acid. • Distinction between strong and weak acids and bases in terms of the rates of their reactions with metals, metal oxides, metal hydroxides, metal hydrogen carbonates and metal carbonates and their electrical conductivities for solutions of equal concentrations.
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In a nutshell: A strong acid or base is one that completely dissociates in water. Eg. HCl H+ + Cl- <1% >99%
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Strong acids/bases Note the arrow indicates COMPLETE DISSOCIATION
Strong acids completely dissociate (split up) into ions in aqueous solution Eg. HCl H+(aq) Cl¯(aq) MONOPROTIC HNO3 H+(aq) NO3¯(aq) H2SO4 2H+(aq) SO42-(aq) DIPROTIC Strong bases completely dissociate into ions in aqueous solution Eg. NaOH Na+(aq) + OH¯(aq) Note the arrow indicates COMPLETE DISSOCIATION
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Weak acids/bases ⇌ Note the arrow indicates PARTIAL DISSOCIATION
Weak acids partially dissociate into ions in aqueous solution Eg. ethanoic acid CH3COOH(aq) ⇌ CH3COO¯(aq) + H+(aq) Can also be written as: HA(aq) + H2O(l) ⇌ A¯(aq) + H3O+(aq) or HA(aq) ⇌ A¯(aq) + H+(aq) Weak bases partially dissociate into ions in aqueous solution Eg. NH3(aq) + H2O(l) ⇌ NH4+(aq) + OH¯(aq) ⇌ Note the arrow indicates PARTIAL DISSOCIATION
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Determining weak/strong, acids/bases experimentally
A weak acid has a higher pH than a strong acid of equal concentration A weak acid does not conduct electricity as well as a strong acid A weak acid reacts more slowly than a strong acid eg. acid + metal H2
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Examples to be learnt - weak/strong, acids/bases
Strong acids: hydrochloric acid, nitric acid, sulfuric acid Weak acids: carboxylic acids (eg. ethanoic acid is vinegar), carbonic acid Strong bases: group 1 hydroxides (eg. NaOH), barium hydroxide Weak bases: ammonia and amines 18M glacial acetic acid pH of %
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Learning check: Weak or strong, acid or base?
Weak base Strong acid Strong base Ammonia NH4+ Hydrochloric acid HCl Potassium hydroxide KOH Ethanoic acid CH3COOH
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8.5 Acid deposition OBJECTIVES • Rain is naturally acidic because of dissolved CO2 and has a pH of 5.6. Acid deposition has a lower pH, usually below 5.0. • Acid deposition is formed when nitrogen or sulfur oxides dissolve in water to form HNO3, HNO2, H2SO4 and H2SO3. • Sources of the oxides of sulfur and nitrogen and the effects of acid deposition should be covered. • Balancing the equations that describe the combustion of sulfur and nitrogen to their oxides and the subsequent formation of HNO3, HNO2, H2SO4 and H2SO3. • Distinction between the pre-combustion and post-combustion methods of reducing sulfur oxides emissions. • Deduction of acid deposition equations for acid deposition with reactive metals and carbonates.
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Acid deposition Carbonic acid from CO2 in the atmosphere makes rain a little acidic naturally: CO2(g) + H2O(l) ⇌ H2CO3(aq) H2CO3(aq) ⇌ H+(aq) + HCO3-(aq) HCO3-(aq) ⇌ H+(aq) + CO32-(aq) Precipitation that is below pH 5.6 is considered acid deposition.
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Acid deposition – nitrogen
At high temperatures such as in engines: N2(g) + O2(g) 2NO(g) (primary reaction) This further reacts with oxygen: 2NO(g) + O2(g) 2NO2(g) This then reacts with water: 2NO2(g) + H2O(l) HNO3(aq) + HNO2(aq) 1. Keep adding oxygen to the radicals, 2. add water to form acid ions
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Acid deposition – sulphur
Sulphur dioxide is released by burning coal (up to 3% sulphur): S(s) + O2(g) SO2(g) (primary reaction) This reacts with water to form sulfurous acid: SO2(g) + H2O(l) ⇌ H2SO3(aq) This can also react further with air, then water to form sulphuric acid: 2SO2(g) + O2(g) ⇌ 2SO3(g) SO3(g) + H2O(l) H2SO4(aq)
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Prevention of acid deposition
Pre-combustion methods involve crushing and washing coal with water to remove up to 90% of solid sulphur. Post-combustion methods such as calcium oxide (lime) can remove gaseous sulphur (and other pollutants): CaO(s) + SO2(g) CaSO3(s)
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18.1 Lewis acids and bases Higher level OBJECTIVES • A Lewis acid is a lone pair acceptor and a Lewis base is a lone pair donor. • When a Lewis base reacts with a Lewis acid a coordinate bond is formed. • A nucleophile is a Lewis base and an electrophile is a Lewis acid. • Application of Lewis’ acid–base theory to inorganic and organic chemistry to identify the role of the reacting species.
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3 of 3. Lewis Acids and Bases
Higher level Lewis acid - a substance that accepts an electron pair eg. BF3 Lewis base - a substance that donates an electron pair eg. NH3
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Formation of hydronium ion is also an excellent example.
Higher level Formation of hydronium ion is also an excellent example. Electron pair of the new O-H bond originates on the Lewis base. Bond formed is a coordinate bond Water acts as the nucleophile H+ acts as the electrophile
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Higher level Coordinate bond Coordinate bond
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Lewis Acid-Base interactions in Biology
Higher level Lewis Acid-Base interactions in Biology The heme group in hemoglobin can interact with O2 and CO. The Fe ion in hemoglobin is a Lewis acid O2 and CO can act as Lewis bases Heme group
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18.2 Calculations involving acids and bases
Higher level OBJECTIVES • The expression for the dissociation constant of a weak acid (Ka) and a weak base (Kb). • For a conjugate acid base pair, Ka × Kb = Kw. • The relationship between Ka and pK a is (pKa = -log Ka), and between Kb and pKb is (pKb = -log Kb). • Solution of problems involving [H+(aq)], [OH–(aq)], pH, pOH, Ka, pKa, Kb and pKb. • Discussion of the relative strengths of acids and bases using values of Ka, pKa, Kb and pKb.
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Acid dissociation constant Ka and pH
Higher level A weak acid, HA, only partially dissociates: HA(aq) ⇌ H+(aq)+ A¯(aq) Applying the Equilibrium Law: 𝐾 𝑎 = 𝐴 − [ 𝐻 + ] [𝐻𝐴] 𝐾 𝑎 = 𝐻 + [ 𝐻 + ] [𝐻𝐴] (as A- and H+ are in equal amounts) Therefore [H+] = Ka[HA] From this pH can be determined.
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Other calculations Higher level CH3CHOOH ⇌ CH3COO- + H+ 𝐾 𝑎 = 𝐶𝐻3𝐶𝑂𝑂 − [ 𝐻 + ] 𝐶𝐻3𝐶𝑂𝑂𝐻 CH3COO- + H20 ⇌ CH3CHOOH + OH- 𝐾 𝑏 = 𝐶𝐻3𝐶𝑂𝑂𝐻 [ 𝑂𝐻 − ] 𝐶𝐻3𝐶𝑂𝑂 − Therefore KaKb = [H+][OH-] = Kw By mathematical rule: pKa + pKb = pKw Addition
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KEY formula Higher level pH = -log[H+] [H+] = 10-pH K = [products]/[reactants] Kw = Ka x Kb = pKw = pH + pOH = 14
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Problem 1: Calculate the pH of a weak acid HX of concentration 0.1M (Ka = 4x10-5 mol dm-3). HA(aq) ⇌ H+(aq)+ A¯(aq) 𝐾 𝑎 = 𝐴 − [ 𝐻 + ] [𝐻𝐴] (H+ and A- are in equal amounts) [H+] = [HA]Ka Key assumption: Equilibrium and original concentration of weak acid is the same due to little dissociation: [H+(aq)] = x 4x10-5 mol dm-3 = x10-6 mol dm-3 = x10-3 mol dm-3 pH = - log [H+(aq)] = 2.699
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Size of Ka Kb pKa pKb values
Higher level Size of Ka Kb pKa pKb values Similarly to pH, a large [H+] ion concentration gives a small pH value. A large Ka value means the acid is strong, and hence small pKa value also means the acid is strong. A strong acid ‘completely dissociates’ HCl Ka = 1.3 x 106, H2SO4 Ka = 1.0 x 103, HNO3 Ka = 24 A weak acid ‘only partially dissociates’ Ethanoic acid 1.8 x 10-5, Carbonic acid 4.4 x 10-7 Addition
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18.3 pH curves Higher level OBJECTIVES • The characteristics of the pH curves produced by the different combinations of strong and weak acids and bases. • An acid–base indicator is a weak acid or a weak base where the components of the conjugate acid–base pair have different colours. • The relationship between the pH range of an acid–base indicator, which is a weak acid, and its pKa value. • The buffer region on the pH curve represents the region where small additions of acid or base result in little or no change in pH. • The composition and action of a buffer solution. • The general shapes of graphs of pH against volume for titrations involving strong and weak acids and bases with an explanation of their important features. • Selection of an appropriate indicator for a titration, given the equivalence point of the titration and the end point of the indicator. • While the nature of the acid–base buffer always remains the same, buffer solutions can be prepared by either mixing a weak acid/base with a solution of a salt containing its conjugate, or by partial neutralization of a weak acid/base with a strong acid/base. • Prediction of the relative pH of aqueous salt solutions formed by the different combinations of strong and weak acid and base.
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Salt hydrolysis Higher level
Salts of weak acids and bases react with water in such a way that the pH of these salt solution is not neutral. Eg. FeCl2 salt: Fe2+ + H2O ⇌ Fe(OH)2 + 2H+ Hence a salt of a weak base will make the solution slightly acidic. CH3COONa salt: CH3COO- + H2O ⇌ CH3COOH + OH- Hence a salt of a weak acid will make the solution slightly alkaline. Na+ will not convert to NaOH, Cl- will not covert to HCl, as they are strong acids and bases. Think – weak acids and bases, really are strong! Strong at holding onto their hydrogen or hydroxide ions!
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Buffers Higher level small quantities of acid or alkali are added.”
Definition “Solutions which resist changes in pH when small quantities of acid or alkali are added.” Acidic Buffer (pH < 7) made from a weak acid its sodium or potassium salt ethanoic acid sodium ethanoate Alkaline Buffer (pH > 7) made from a weak base its chloride ammonia ammonium chloride Uses Standardising pH meters Buffering biological systems (eg in blood) Maintaining the pH of shampoos
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Buffer solutions - ideal concentration
Higher level The concentration of a buffer solution is also important If the concentration is too low, there won’t be enough CH3COOH and CH3COO¯ to cope with the ions added. Summary For an acidic buffer solution one needs ... large [CH3COOH(aq)] - for dissociating into H+(aq) when alkali is added large [CH3COO¯(aq)] - for removing H+(aq) as it is added This situation can’t exist if only acid is present; a mixture of the acid and salt is used. The weak acid provides the equilibrium and the large CH3COOH(aq) concentration. The sodium salt provides the large CH3COO¯(aq) concentration. One uses a WEAK ACID + its SODIUM OR POTASSIUM SALT
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How buffers work CH3CHOOH ⇌ CH3COO- + H+ In equilibrium
Higher level CH3CHOOH ⇌ CH3COO- + H+ In equilibrium CH3CHOONa CH3COO- + Na+ Complete dissociation If an acid is added: HCl H+ + Cl- The excess H+ will cause equilibrium to be pushed to the left and so the pH is remains the same: CH3COO- + H+ CH3CHOOH If a base is added: NaOH Na+ + OH- The H+ will neutralize the OH-: H+ + OH- H2O The equilibrium will be pushed to the right due to loss of H+ and so the pH remains the same: CH3CHOOH CH3COO- + H+ (note: state symbols must also be written with these equations in the exam)
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Indicator chemical reactions
Higher level Indicator chemical reactions Indicators are weak acids which have a different colour to their conjugate base HIn H+ + In- colour colour 2 low pH: equilibrium pushed left = colour 1 high pH: equilibrium pushed right = colour 2
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Higher level Methyl orange pH < 3.2 pH > 4.4
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Higher level acid end-point alkali
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Higher level Phenolphthalein pH < 8.2 pH > 10.0
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Higher level phenolphthalein acid alkali
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Relationship of pH to pKa value
Higher level Relationship of pH to pKa value
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Relationship of pH to pKa value
Higher level Relationship of pH to pKa value When pH = pKa the colour will be midway. If you increase [H+] it pushes the indicator to the acid colour. It takes about a 10 times difference in indicator colour species to see the colour change therefore most colour changes take about 2 pH units.
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Indicators Higher level For an indicator to work in a titration, the pH range of the indicator’s colour change must be within the range of the pH change at the end-point. Indicator colour at low pH pH range of colour change high pH methyl orange red 3.2 – 4.4 orange phenolphthalein colourless 8.2 – 10.0 purple
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pH curves weak acid (CH3COOH) vs strong alkali (NaOH)
Higher level Types There are four types of acid-base titration; each has a characteristic curve. strong acid (HCl) vs strong base (NaOH) weak acid (CH3COOH) vs strong alkali (NaOH) strong acid (HCl) vs weak base (NH3) weak acid (CH3COOH) vs weak base (NH3) In the following examples, alkali (0.1M) is added to 25cm3 of acid (0.1M) End points need not be “neutral‘ due to the phenomenon of salt hydrolysis
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Terminology Higher level
Equivalence point (stoichiometric point) – exact point at which the solution is neutralized for the given acid (solution may not actually be neutral), given as a volume Endpoint – point at which an indicator changes colour, given as a pH Point of inflection – place in the pH curve where a large change in pH occurs Half-equivalence – point in the graph where half of the acid has been neutralized and turned into a salt Buffer region – point on the graph that is most resistant to changes in pH
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Strong acid (HCl) vs strong base (NaOH)
Higher level Strong acid (HCl) vs strong base (NaOH) Curve levels off at pH 13 due to excess 0.1M NaOH (a strong alkali) Very sharp change in pH over the addition of less than half a drop of NaOH Very little pH change during the initial 20cm3 pH 1 at the start due to 0.1M HCl (strong monoprotic acid)
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Higher level Any of the indicators listed will be suitable - they all change in the ‘vertical’ portion PHENOLPHTHALEIN LITMUS METHYL ORANGE
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Strong acid (HCl) vs weak base (NH3)
Higher level Curve levels off at pH 10 due to excess 0.1M NH3 (a weak alkali) Sharp change in pH over the addition of less than half a drop of NH3 Very little pH change during the initial 20cm3 pH 1 at the start due to 0.1M HCl
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Higher level Only methyl orange is suitable - it is the only one to change in the ‘vertical’ portion PHENOLPHTHALEIN LITMUS METHYL ORANGE
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Weak acid (CH3COOH) vs strong base (NaOH)
Higher level Curve levels off at pH 13 due to excess 0.1M NaOH (a strong alkali) Sharp change in pH over the addition of less than half a drop of NaOH Steady pH change pH 4 due to 0.1M CH3COOH (weak monoprotic acid)
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Higher level PHENOLPHTHALEIN LITMUS METHYL ORANGE
Only phenolphthalein is suitable - it is the only one to change in the ‘vertical’ portion
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Weak acid (CH3COOH) vs weak base (NH3)
Higher level Curve levels off at pH 10 due to excess 0.1M NH3 (a weak alkali) NO SHARP CHANGE IN pH Steady pH change pH 4 due to 0.1M CH3COOH (weak monoprotic acid)
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Higher level There is no suitable indicator- none change in the ‘vertical’ portion. The end point can be detected by plotting a curve using a pH meter. PHENOLPHTHALEIN LITMUS METHYL ORANGE NOTHING SUITABLE
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SUMMARY for Phenolphthalein
Higher level SUMMARY for Phenolphthalein
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SUMMARY for methyl orange
Higher level SUMMARY for methyl orange
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