Acids and Bases pH and Titrations - Geneva High School

Acids and Bases pH and Titrations - Geneva High School

Acids and Bases pH and Titrations Overview Acid Base Concepts Arrhenius Brnsted Lowry Lewis Acid and Base Strengths Relative Strengths of Acids and Bases

Molecular Structure and Acid Strength Self Ionization of Water and pH Self Ionization of Water Solutions of a Strong Acid or Base The pH of a Solution II. Theories of acid and bases A. Arrhenius theory 1. Arrhenius Acid as known as a traditional acid a. a chemical compound that contains

hydrogens and ionizes in aqueous solution to form hydrogen ions 2. Arrhenius base a. any chemical that produces hydroxide ions B. Bronsted -Lowry theory 1923 (worked independently) 1. Bronsted Acid a. an ion or molecule that is a proton donor 2. Bronsted base a. anything which will accept a proton

3. Conjugate acid - base pairs HF + HOH Acid1Base2 ===> H3O+ + Acid2 Base1 F-

a. Strong acids have weak conjugate bases b. Strong bases and weak conjugate acids 4. Autoionization - self-ionization HOH <===> H+ + OH - Conjugate Acid-Base Pairs Fig. 17-1, p. 507 Conjugate Acid-Base Pairs

p. 507 Brnsted-Lowry Theory of Acids & Bases Conjugate Acid-Base Pairs General Equation p. 507 Brnsted-Lowry Theory of Acids & Bases

p. 504 5. Amphiprotic ability to act as an acid or base C. Lewis theory 1. Lewis acid a. anything which can accept a pair of electrons 2. Lewis Base a. anything which will donate a pair of

electrons Lewis Theory of Acids & Bases p. 506 Objectives Properties of acids and bases The pH scale Distinguish between strong and weak acids and list the clinical uses of these

acids Distinguish between strong and weak bases and list the clinical uses of these acids Understand neutralisation and the clinical applications of neutralisation On page 355 Do questions 4,6,7,10,14, 17,18,22,23,25,26,28 Properties of Acids

Aqueous solutions of acids have a sour taste. Acids change the color of acid-base indicators. Some acids react with active metals to release hydrogen gas. Acids react with bases to produce salts and water. Some acids conduct electric current. PROPERTIES OF ACIDS & BASES Acids Produce hydrogen ions (H+) in H2O Taste sour

Act as electrolytes in solution Neutralise solutions containing hydroxide ions (OH -) React with several metals releasing H2(g) corrosion React with carbonates releasing CO2(g) Destroy body tissue How many foods can you think of that are sour? Chances are, almost all of the foods that you

associate with being sour, owe their sour taste to an acid: Lemons citric acid Grapefruit citric acid Apples malic acid Sour milk lactic acid Vinegar acetic acid Grapes tartaric acid Strength of acids A strong acid is one that ionizes completely

in aqueous solutions. The strength of an acid depends on the polarity of the bond between hydrogen and the element to which it is bonded and the ease with which that bond can be broken. Acid strength increases with increasing polarity and decreasing bond energy. Acids that are weak electrolytes are weak acids. Strong acids are: Strong electrolytes

~ 100% ionisation good conductors Severe burns to body tissue *** Stomach lining protected against HCl by mucus STRENGTHS OF ACIDS Strong Acids (very few) Eg HCl Hydrochloric Acid ~ Stomach acid

Marieb, Fig 26.11 HNO3 Nitric Acid ~ May be used to cauterise warts ~Drugs, explosives, fertilisers, dyes H2SO4 Sulphuric Acid ~ conc. to treat stomach hypoacidity ~ Fertilisers, dyes, glues Factors Affecting Acid Strength Binary acids:

Bond strength is directly related to the acid strength (bond size). HI and HBr have larger bonds lengths and are more acidic than HF and HCl, even though fluorine is most electronegative. For bonds of similar size the acid strength is related to electronegativity difference. Oxyacids are acidic substances that contain oxygen and some other nonmetal, e.g. HNO3 An increase in the electronegativity of an

atom bound to oxygen increases in polarity of the bond and makes it more acidic. More oxygen = more polar. Weak Acids (most acids in nature) CH3COOH Acetic Acid ~ Vaginal jellies, antimicrobial solution ears, plastics, dyes, insecticides H2CO3 Carbonic Acid ~Bicarbonate buffer system, carbonated drinks H3PO4 Phosphoric Acid

~ Drugs, fertilisers, soaps, detergents, animal feed Bases Produce or cause an increase in hydroxide ions (OH-) in H2O Taste bitter Turn red litmus blue Act as electrolytes in solution Neutralise solutions containing hydrogen ions (H +) Have a slippery, soapy feel

Destroy body tissue/ dissolve fatty (lipid) material How many bases can you think of? Ammonia

Sodium hydroxide lye drain cleaner Milk of magnesia Mg(OH)2 antacid Aluminum hydroxide antacid Baking soda sodium hydrogen carbonate 4. STRENGTHS OF BASES Strong Bases NaOH Sodium Hydroxide ~ Removes grease drains, ovens Mg(OH)2 Magnesium hydroxide ~ Antacid ~ Laxative

Al(OH)3 Aluminium hydroxide ~ Antacid ~ Absorbs toxins, gases, ~ Causes constipation Strength of Bases The strength of a base depends on the extent to which the base dissociates. Strong bases are strong electrolytes. Strong bases: calcium hydroxide, barium hydroxide, sodium hydroxide, etc

Weak bases: ammonia, aniline Table 15-4 page 461 Relative Strengths of Acids and Bases and Extent of Reaction The table of relative strengths of acids and conjugate bases can be used to predict if a reaction

will produce product. E.g. Which will NO3 produce product? HNO3 + CN or HCN + Strong bases are: Strong electrolytes ~ 100% dissociation in water good conductors

Severe damage to skin & eyes (Group 1A elements) Weak Bases Eg NH3 Ammonia ~ Waste product of protein break down in body. CO3 2- In antacids

HCO3 In antacids, buffers HPO4 2- In buffers Weak bases are: Weak electrolytes Do not contain OH but react with H2O small numbers of OH Reaction with Water : Weak bases NH3(g) + H2O

NH4 + (aq) + OH (aq) HCO3 (aq) + H2O H2CO3 (aq) + OH-(aq)

Acids & Bases STRONG vs WEAK _ completely ionized

_ partially ionized _ strong electrolyte _ weak electrolyte _ ionic/very polar bonds _ some covalent bonds Strong Acids: HClO4 H2SO4 HI HBr

HCl HNO3 Strong Bases: LiOH NaOH KOH Ca(OH)2 Sr(OH)2 Ba(OH)2

Dissociation in Water : Strong bases Metal hydroxides ions H2O NaOH(s) Mg(OH)2(s) Al(OH)3(s) Na+(aq) + H2O

H2O Mg 2 + (aq) + Al 3+(aq) + OH-(aq) OH- (aq) OH- (aq) 5. ACID-BASE NEUTRALISATION

Neutralisation Reaction Acid + Base Salt + Water HCl + NaOH NaCl + H2O H+ + OH H2O

Neutralise each other Must be equal concentrations Antacids clinical applications (Check for side effects!!) ~ Neutralise excess stomach acid ~ Raise stomach pH > 4 \Pepsin inactive ~ Assist with ulcer treatment ~ solubility in H2O but still produce high % of ions

CaCO 3 2HCl + CaCO 3 CaCl 2 + H2O + CO 2 (g) ~Also a Ca 2 + supplement

Long term overuse Ca 2 + levels risk kidney stones (renal calculi) Eg Mg(OH) 2 Milk of Magnesia & in Mylanta 2HCl + Mg(OH) 2 MgCl 2 + 2H2O Al(OH) 3

3HCl In Mylanta + Al(OH) 3 AlCl 3 + 3H2O NaHCO 3 Baking Soda Not recommended!! HCl + NaHCO 3 NaCl + H2O + CO2 (g)

~ Elderly tend to OD Stomach can explode Weak acids are: Weak electrolytes Small % ionisation weak conductors Dissociation in Water : Weak acids Polar covalent molecules Mainly stay as molecules

Dissociation in Water : Strong acids Polar covalent molecules ions Eg. HCl(l) H2O H+(aq) + Cl-(aq) H2 O

HNO3(l) H+ (aq) + NO3- (aq) H2O H2SO4(l) 2H+ (aq) + SO42- (aq) Dissociation in water : Weak acids (cont) H2O

CH3COOH (l) H2CO3 (l) H3PO4 (l) H+ (aq) + CH3COO- (aq) H2O

H2O H+ (aq) + HCO3-(aq) H+ (aq) + H2PO4- (aq) 2. THE pH SCALE Ion Product of Water Pure H2O at 25C Some molecules ionise H2O H+ + OH[H+ ] = 1 x 10-7 M = [OH- ]

On the pH scale, values below 7 are acidic, a value of 7 is neutral, and values above 7 are basic. The pH scale The pH scale ranges from 1x 100 to 1 x 10-14 mol/L or from 1 to 14. pH = - log [H+] 1 2 3 4 5 6 7 8 9 10 11 12 13 14 acid

neutral base Stomach juice: pH = 1.0 3.0 Human blood: pH = 7.3 7.5 Lemon juice: pH = 2.2 2.4 Seawater: pH = 7.8 8.3 Vinegar: pH = 2.4 3.4 Ammonia: pH = 10.5 11.5 Carbonated drinks: pH = 2.0 4.0 0.1M Na2CO3: pH = 11.7

Orange juice: pH = 3.0 4.0 1.0M NaOH: pH = 14.0 pH describes [H+ ] & [OH- ] \ Indicates if a fluid is : 0 Acidic [H+ ] = 100

7 Neutral [H+ ] = 10-7 [OH- ] =10-7 14 Basic [OH- ] = 100

[H+ ] = 10-14 [OH- ] =10-14 Using the pH scale Exponential values for [H+ ] & [OH- ] inconvenient in a clinical workplace \Simplify pH scale acid-base concentration

p potential or Power H Hydrogen Acidic solution [H+ ] > [OH- ] Neutral solution [H+ ] = [OH- ] Basic solution [H+ ] < [OH- ]

Water Equilibrium Kw = [H+] [OH-] = 1.0 x 10-14 Equilibrium constant for water Water or water solutions in which [H+] = [OH-] = 10-7 M are neutral solutions. A solution in which [H+] > [OH-] is acidic A solution in which [H+] < [OH-] is basic Autoionization of Water Water can act as both an acid and a

base equilibrium is: H2O + H2O H3O+ + OH. Kw = [OH][H3O+] = 1.00 x 1014M2. Since [OH] = [H3O+] [H3O+] = 1 x 107 M (called a neutral solution) \ Ion Product of H2O: [H+ ] x [OH- ] = [1 x 10-7 ] x [1 x 10-7 ]

* Add exponents = 1 x 10-14 Acidic [H3O+] > 1.00x107M Neutral [H3O+] = 1.00x107M Basic [H3O+] < 1.00x107M All acids/bases dissolved in water must obey equation for the ionization of water. They either add H3O+ or OH to water. Most of the acids in this chapter will be stronger

than water and add significantly to the hydronium ion concentration. E.g. the hydronium ion concentration of an acidic solution was 1.00x105 M. What was the [OH]? E.g. what is the hydronium ion concentration if the hydroxide concentration was 2.50x103 M? The pH loop Fig. 17-3, p. 516

The pH of a Solution pH = log[H3O+] and [H3O+] = 10pH Acidic pH < 7.00 Neutral pH = 7.00 Basic pH > 7.00 E.g. determine the pH of a solution in which [H3O+] = 5.40x106 M E.g.2 determine the pH of a solution in which the [OH] = 3.33x103 M

E.g.3 determine the pOH of a solution in which the [OH] = 3.33x103 M E.g.4 Determine the [H3O+] if the pH of the solution is 7.35. The term pX is defined in exactly the same way as pH Eg.5 What is the pCa if [Ca2+] = 6.44x10-4 E.g. 6 what is pH and [OH] of 0.125 M Ba(OH)2. [H3O+] of water is small compared to added [H3O+] from the acid and ignored in the calculation.

pH scale The scale for measuring the hydronium ion concentration [H+] in any solution must be able to cover a large range. A logarithmic scale covers factors of 10. A solution with a pH of 1 is 10 times stronger than a solution with a pH of 2 A solution with a pH of 1 has [H+] of 0.1 mol/L or 10-1 A solution with a pH of 3 has [H+] of 0.001 mol/L or 10-3

A solution with a pH of 7 has [H+] of 0.0000001 mol/L or 10-7 Manipulating pH Algebraic manipulation of: pH = - log [H+] allows for: [H+] = 10-pH If pH is a measure of the hydronium ion

concentration then the same equations could be used to describe the hydroxide (base) concentration. [OH-] = 10-pOH pOH = - log [OH-] thus: pH + pOH = 14 ; the entire pH range!

Table 17-3, p. 518 Methods of Measuring pH pH paper is used that has compounds in it which are change to different colors for different pH ranges. An colored indicator can be placed in the solution and its color correlated with pH. HIn(aq) + H2O(l) H3O+(aq) + In(aq). E.g. phenolphthalein is colorless in acid form but

pink in basic form. The pH at which they change color depends on their equilibrium constant. More accurate and precise measurements are made with a pH meter. A combination of voltmeter and electrodes. Indicators and pH meters

Acid-Base indicators are compounds whose color is sensitive to pH. The color of an indicator changes as the pH of a solution changes. Indicators come in many colors. The pH range over which an indicator changes color also varies. This pH range is also called the transition interval. Indicators and pH meters cont

A universal indicator is made by mixing several different indicators. If an exact value for the pH of a solution is needed, a pH meter should be used. A pH meter determines the pH of a solution be measuring the voltage between the two electrodes that are placed in the solution. Sample Problems Q1: Calculate the pH of a solution if

[H+] = 2.7 x 10-4 M pH = -log[H+] pH = -log(2.7 x 10-4) = 3.57 Q2: Find the hydrogen ion concentration of a solution if its pH is 11.62. [H+] = 10-pH [H+] = 10-11.62 = 2.4 x 10-12M Q3: Find the pOH and the pH of a solution if its hydroxide ion concentration is 7.9 x 10-5M

pOH = -log[OH-] pOH = -log(7.9 x 10-5) = 4.10 pH + pOH = 14 pH = 14 - 4.10 pH = 9.9 A brief introduction to equilibrium constants and Ionization constants Equilibrium is defined as a state where concentrations of products and reactants are unchanging The equilibrium constant is a mathematical expression of the

ratio of the products compared to the reactants K = { products} The { } are read as concentration { reactants} All species are considered to be at equilibrium Equilibrium constant and Ionization constants can be determine for many reactions For acids

Ka = { H3O +} { A-} {HA} {H2O} Since the concentration of H2O is roughly 55 moles /liter it can be deleted from the above equation therefore the equation becomes Ka = { H3O +} { A-} {HA} Titration A neutralization reaction occurs between an acid and a base.

Because acids and bases react, the progressive addition of an acid to a base or vice versa can be used to compare the concentrations of the acid and the base. Titrations cont Titration is the controlled addition and measurement of the amount of a solution of known concentration required to react completely with a measure amount of a solution of unknown concentration

Titration provides a sensitive means of determining the chemically equivalent volumes of acidic and basic solutions. TITRATION Titration of a strong acid with a strong base ENDPOINT = POINT OF NEUTRALIZATION = EQUIVALENCE POINT At the end point for the titration of a strong acid with a strong base, the moles of acid (H+) equals the moles

of base (OH-) to produce the neutral species water (H2O). If the mole ratio in the balanced chemical equation is NOT 1:1 then you must rely on the mole relationship and handle the problem like any other stoichiometry problem. MOLES OF ACID = MOLES OF BASE nacid = nbase TITRATION MAVA = MBVB

1. Suppose 75.00 mL of hydrochloric acid was required to neutralize 22.50 mL of 0.52 M NaOH. What is the molarity of the acid? HCl + NaOH H2O + NaCl Ma Va = Mb Vb rearranges to Ma = Mb Vb / Va so Ma = (0.52 M) (22.50 mL) / (75.00 mL) = 0.16 M Now you try: 2. If 37.12 mL of 0.843 M HNO3 neutralized 40.50 mL of KOH, what is the molarity of the

base? Mb = 0.773 mol/L TITRATION 1. If 37.12 mL of 0.543 M LiOH neutralized 40.50 mL of H2SO4, what is the molarity of the acid? 2 LiOH + H2SO4 Li2SO4 + 2 H2O First calculate the moles of base: 0.03712 L LiOH (0.543 mol/1 L) = 0.0202 mol LiOH Next calculate the moles of acid:

0.0202 mol LiOH (1 mol H2SO4 / 2 mol LiOH)= 0.0101 mol H2SO4 Last calculate the Molarity: Ma = n/V = 0.010 mol H2SO4 / 0.4050 L = 0.248 M 2. If 20.42 mL of Ba(OH)2 solution was used to titrate 29.26 mL of 0.430 M HCl, what is the molarity of the barium hydroxide solution? Mb = 0.308 mol/L TITRATION

Titration of a strong acid with a strong base ENDPOINT = POINT OF NEUTRALIZATION = EQUIVALENCE POINT At the end point for the titration of a strong acid with a strong base, the moles of acid (H+) equals the moles of base (OH-) to produce the neutral species water (H2O). If the mole ratio in the balanced chemical equation is 1:1 then the following equation can be used. MOLES OF ACID = MOLES OF BASE

nacid = nbase Since M=n/V MAVA = MBVB Equivalence Point (E.P.) The point at which the two solutions used in a titrations are present in chemically equivalent amounts is the equivalence point (e.p.). Indicators and pH meters cane be used to

determine the e.p. A pH meter will show a large voltage change occurring at the e.p. An indicator changes color over a range that includes the pH of the e.p. End point The point in a titration at which an indicator changes color is called the end point.

Which indicator should you use? Indicators that undergo transition at about pH 7 are used to determine the equivalence point of strong acid/ strong base titrations because the neutralization of strong acids with strong bases produces a salt solution with a pH of approximately 7. Which indicator should

you use? Indicators that change color at pH lower than 7 are useful in determining the equivalence point of strong acid/ weak base titrations. The equivalence point of this titration is acidic because the salt formed is a weak acid. Thus the salt solution has a pH lower than 7. Which indicator should you use?

Indicators that change color at pH higher than 7 are useful in determining the equivalence point of weak acid/ strong base titrations. These reactions produce salt solutions whose pH is greater than 7. This occurs because the salt solution formed is a weak base. Which indicator should you use? To determine the equivalence point for a weak

acid/ weak base, we use no indicator. The pH of the equivalence point of weak acids and weak bases may be almost any value, depending on the relative strengths of the reactants. The color transition of an indicator helps very little in determining whether reactions between weak acids and bases are complete. Molarity and titration To calculate the molarity of a substance

using titration: 1. Start with the balanced equation and determine the chemically equivalent amounts of the acid and base. 2. Determine the moles of acid or base from the known solution used in the titration. Molarity and titration 3. Determine the moles of solute of the unknown solution used during the

titration. 4. Determine the molarity of the unknown solution. Predicting Acid-Base Reactions Fig. 17-2, p. 511 Review questions (cont) Distinguish between strong & weak bases; List clinical uses of these

bases & write equations for their dissociation in water. Complete simple equations for the neutralisation reaction of an acid & a base; Discuss clinical applications of acid-base neutralisation. Review questions List the properties of acids & bases. Discuss the pH scale: *Define the ion product of water &

indicate how this is determines the pH scale. * Use the pH scale to determine if a given solution is acidic, neutral or basic. Distinguish between strong acids & weak acids: List clinical uses of these acids & write equations for their dissociation in water Brnsted-Lowry Theory of Acids & Bases

Notice that water is both an acid & a base = amphoteric Reversible reaction p. 504 Water Equilibrium p. 514

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