Acids and BasespH and Titrations

Overview Acid – Base Concepts

Arrhenius Brønsted – 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 solutionto 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 ===> H3O+ + F-

Acid1Base2 Acid2 Base1

a. Strong acids have weak conjugate bases

b. Strong bases and weak conjugate acids

4. Autoionization - self-ionization

HOH <===> H+ + OH -

Fig. 17-1, p. 507

Conjugate Acid-Base Pairs

p. 507

Conjugate Acid-Base Pairs

p. 507

Brønsted-Lowry Theory of Acids & BasesConjugate Acid-Base Pairs

General Equation

p. 504

Brønsted-Lowry Theory of Acids & Bases

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

p. 506

Lewis Theory of Acids & Bases

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

Some acids conduct electric current.

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.

PROPERTIES OF ACIDS & BASES Acids

Produce hydrogen ions (H+) in H2OTaste sourAct 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 youassociate with being sour, owe their sour taste to an acid:Lemons – citric acid

Grapefruit – citric acidApples – malic acid

Sour milk – lactic acid

Vinegar – acetic acid

Grapes – tartaric acid

Strength of acids

Acids that are weak electrolytes are weak acids.

A strong acid is one that ionizes completely in aqueous solutions.

The strength of an acid depends on the polarity ofthe 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.

Strong acids are:

Strong electrolytes

~ 100% ionisation good conductors

Severe burns to body tissue *** Stomach lining protected against HCl

by mucus

STRENGTHS OF ACIDSStrong Acids (very few)EgHCl Hydrochloric Acid~ Stomach acid

HNO3 Nitric Acid~ May be used to cauterise warts~Drugs, explosives, fertilisers, dyes

H2SO4 Sulphuric Acid~ conc. to treat stomach hypoacidity~ Fertilisers, dyes, glues

Marieb, Fig 26.11

Factors Affecting Acid Strength

For bonds of similar size the acid strength isrelated to electronegativity difference.

Binary acids:

Bond strength is directly related to the acid strength (bond size).

HI and HBr have larger bonds lengths and aremore acidic than HF and HCl, even though fluorine is most electronegative.

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

Destroy body tissue/ dissolve fatty (lipid) material

Have a slippery, ‘soapy’ feel

Neutralise solutions containing hydrogen ions (H +)

Act as electrolytes in solution

Turn red litmus blue

Taste bitter

Produce or cause an increase in hydroxide ions (OH-) in H2O

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

~ Causes constipation~ Absorbs toxins, gases,~ Antacid

Al(OH)3 Aluminium hydroxide

~ Antacid ~ Laxative

Mg(OH)2 Magnesium hydroxide

~ Removes grease – drains, ovens

NaOH Sodium Hydroxide

Strong Bases

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 produce product?

HNO3 + CN or HCN +

3NO

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: Strong Bases:HClO4 LiOHH2SO4 NaOHHI KOHHBr Ca(OH)2

HCl Sr(OH)2

HNO3 Ba(OH)2

Dissociation in Water : Strong basesMetal hydroxides ions

H2O

H2O

H2O

NaOH(s) Na+(aq) + OH-

(aq)

Mg(OH)2(s) Mg 2 + (aq) + OH-

(aq)

Al(OH)3(s) Al 3+(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 + levelsÞ risk kidney stones (renal calculi)

Eg

Mg(OH) 2 Milk of Magnesia & in Mylanta

2HCl + Mg(OH) 2 MgCl 2 + 2H2O

Al(OH) 3 In Mylanta

3HCl + 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 acidsPolar covalent molecules ions

Eg.

HCl(l) H+(aq) + Cl-(aq)

 

HNO3(l) H+ (aq) + NO3

- (aq)

 

H2SO4(l) 2H+ (aq) + SO4

2- (aq)

H2O

H2O

H2O

Dissociation in water : Weak acids (cont)

CH3COOH (l) H+ (aq) + CH3COO- (aq)

H2O

H2O

H2O

H3PO4 (l) H+ (aq) + H2PO4

- (aq)

H2CO3 (l) H+ (aq) + HCO3

-(aq)

2. THE pH SCALE

Ion Product of Water

Pure H2O at 25°CSome 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 scaleThe 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 [OH- ] =10-14

7 Neutral [H+ ] = 10-7 [OH- ] =10-7

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

Using the pH scale

Exponential values for [H+ ] & [OH- ] inconvenient in a clinical workplace

\ Simplify pH scale acid-base concentration

p potential or PowerH 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

They either add H3O+ or OH to water.

All acids/bases dissolved in water must obey equation for the ionization of water.

Acidic [H3O+] > 1.00x107MNeutral [H3O+] = 1.00x107MBasic [H3O+] < 1.00x107M

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?

Most of the acids in this chapter will be stronger than water and add significantly to the hydronium ion concentration.

Fig. 17-3, p. 516

The pH loop

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.2 determine the pH of a solution in which the [OH] = 3.33x103 M

E.g. determine the pH of a solution in which [H3O+] = 5.40x106 M

Eg.5 What is the pCa if [Ca2+] = 6.44x10-4

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

E.g.3 determine the pOH of a solution in which the [OH] = 3.33x103 M

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 scaleThe 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 pHAlgebraic 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 con’t

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}

{ reactants}

The { } are read as concentration

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 /literit 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 con’t

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.

TITRATIONTitration 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

TITRATIONMAVA = 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

TITRATION1. 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.

Fig. 17-2, p. 511

Predicting Acid-Base Reactions

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

p. 504

Brønsted-Lowry Theory of Acids & Bases

Notice that water is both an acid & a base = amphoteric

Reversible reaction

p. 514

Water Equilibrium