Chapter 9 Slides 2017 - Calvin

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• A hydrogen atom contains a proton and an electron, thus a hydrogen ion (H+) is a proton:

Proton (H+) transfer between molecules isthe basis of acid/base chemistry

• Acids:Ø Are proton donorsØ Produce protons (H+) when dissolved in water

• Bases:Ø Are proton acceptorsØ Produce hydroxide ions (OH-) when dissolved in water

ACIDS AND BASES, DEFINED

CHAPTER 9: Acids, Bases, pH, and Buffers

• Molecules that can act as an acid or a base are called amphoteric

Solutions containing amphoteric molecules have spontaneous exchange of protons between their molecules, some acting as acids and some as bases

• Pure water is an important example of an amphoteric molecule:

H2O + H2O H3O+ + OH-

Ø This happens spontaneously, but very rarely in pure water!

WATER IS AMPHOTERIC


Acts as base Acts as acid

• 9.1 Properties of Acids and Bases

• 9.2 pH

• 9.3 Buffers

OUTLINE


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• In aqueous solution, acids donate a proton to water to produce a hydronium ion, H3O+:

• Example: Hydrochloric acid (“gastric juice”)

hydrochloric acid water hydronium ion chloride ion

ACIDS AND HYDRONIUM IONS


H+ + H2O

à

• The carboxylic acid group is called an “acid” because it can “lose” a proton in aqueous solutionØ This forms a carboxylate ionØ The name of the ion ends in –ate

CHAPTER 7: Organic Functional Groups

IONIZATION OF CARBOXYLIC ACIDS

A carboxylate ion is an example of a polyatomic ionH3C H

Acetic acid

H3C

+ H+

Acetate

The hydrogen atoms highlighted in pink are released as protons (H+) in solution

Note the relevant functional groups

COMMON BIOLOGICAL ACIDS


carbonic acid acetic acid phosphoric acid

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• In aqueous solution, bases produce hydroxide ion

• Hydroxides can form in two ways in water:1. The base may “steal” a proton from H2O:

2. The base may dissociate, releasing a hydroxide ion:

BASES AND HYDROXIDE IONS


Base + H2O

AMINES AS COMMON BIOLOGICAL BASES


The nitrogen atoms highlighted in blue can all accept a hydrogen atom

Note the relevant functional group

A base must contain at least one nonbonding pair of electrons to accept an H+ from water

H H

ammonia ammonium ion

dopamine

• Amino acids contain both the amino group and the carboxylic acid group:Ø In neutral (pH 7) solution, both groups are charged

• Amino acids are the “building blocks” of proteins.

CHAPTER 7: Organic Functional Groups

AMINO ACIDS: AN AMINE + CARBOXYLIC ACID

Carboxylates are (-) charged

Amines are (+) charged

Would you expect amino acids to be soluble in aqueous solution?

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• Acids or bases may be classified as strong or weak, depending on how they behave in water:

• A strong acid or base is one that fully dissociates in water:Ø Similar to an ionic compoundØ Produces electrolytes in water: H+ + anion

• A weak acid or base is one that only partially dissociates in water:Ø The resulting solution contains a mixture of both the

acid and its conjugate base

STRENGTH OF ACIDS AND BASES


• HCl is an example of a strong acid as it completely dissociates when dissolved in water:

• Strong acids are uncommon in biological systems

STRONG ACIDS


• Sodium hydroxide (NaOH) is a good example of a strong base that completely dissociates in water:

NaOH (s) + H2O à Na+ (aq) + OH- (aq)

• Strong bases are also uncommon in biological systems

STRONG BASES


• Strong bases that utilize a metal cation are poorly soluble in water

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• In an acid-base neutralization reaction, a proton transfers from an acid to a base, producing a salt(ionic compound) and water.

• The basis of all neutralization reactions is: H+ + OH- à H2O

acid base neutral

• The type of salt formed depends on what specific acid & base are used in the reaction:

HI (aq) + KOH (aq) à H2O (l) + KI (aq)acid base water salt

ACID-BASE NEUTRALIZATION REACTIONS


• Antacids are able to neutralize stomach acids because they are bases

• All of the antacids below are poorlysoluble strong bases

Ø They are poorly soluble so are safe due to limited release of their hydroxide ions.

ANTACID NEUTRALIZATION REACTIONS


• Weak acids only partially dissociate in water:Ø A large portion of undissociated acid molecules remain

intact in aqueous solutionØ A small fraction of molecules donate a hydrogen atom

• Carboxylic acids are generally weak acids:Ø A common example is acetic acid (vinegar)

WEAK ACIDS


Strong = full dissociation Weak = partial dissociation

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• Weak bases in solution also undergo reversible reactions with water: Ø Nearly all of the base form is present, along with a small

amount of its conjugate acid as well as hydroxide ion.

:NH3 + H2O NH4+ + OH-

ammonia water ammonium ion hydroxide ion

• Amines are generally weak bases:Ø The nitrogen on amines has a

lone electron pair capable ofbinding a proton (hydrogen atom)

WEAK BASES


• When an acid dissolves in water, hydronium ions form, and the dissociated acid remains as an anion.

Ø The anion portion of the acid is its “conjugate base”

• An acid and its conjugate base are together called a conjugate acid-base pair

CONJUGATE ACID-BASE PAIRS


For each of the following equations, label the conjugate acid-base pairs. Indicate whether water acts as an acid or a base in the reaction.

a. NH4+ + H2O H3O+ + NH3

b. HCO3- + H2O OH- + H2CO3

c. CH3COO- + H2O OH- + CH3COOH

PRACTICE PROBLEM


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• Reactions involving a weak acid or base contain a bidirectional arrow, indicating it is reversible:

In reversible reactions, both the forward reaction and the reverse reaction are going on at the same time.

Ø All chemical “species”—the hydronium ion, the acetate ion, and acetic acid—are present in solution

REVERSIBLE REACTIONS


acetic acid acetate ion hydronium ionwater

• Reversible reactions have two separate reaction rates—both a forward AND a reverse rate.

A reaction is at equilibrium when the forward and reverse reaction rates are equivalent

Ø If both rates are the same, then there is no net change in reactant or product concentration over time.

• Chemical equilibrium is a dynamic equilibrium, and not a “static equilibrium”

ACID-BASE EQUILIBRIA



CH3COOH + H2O CH3COO- + H3O+

• An equilibrium may be disturbed by changes in conditions, such as concentration or temperature

Le Châtelier’s principle holds that when an equilibrium is disturbed, the reaction responds by shifting in the direction needed to restore equilibrium

LE CHÂTELIER’S PRINCIPLE


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• The extent to which an acid or base react with water at equilibrium is defined by its equilibrium constant (K)

Ø Equilibrium constants from reactions in aqueous solution do NOT account for the “concentration” of water

EQUILIBRIUM CONSTANTS


K = = concentration of products [products]concentration of reactants [reactants]


CH3COOH + H2O CH3COO- + H3O+

K =[CH3COO-] [H3O+]

[CH3COOH]

• In pure water, two H2O molecules can ionize to produce exactly one hydronium ion & one hydroxide ion:

H2O + H2O H3O+ + OH-

Ø The product of their concentration is a constant value, termed the ion-product constant, or KW.

Ø In pure water, concentration of each is equal at 10-7 M

Kw = [H3O+] [OH-] = (1.0 ´ 10-7 M) ´ (1.0 ´ 10-7 M) = 1.0 ´ 10-14

ION-PRODUCT CONSTANT, KW


Because this relationship is constant, it allows us to calculate the amount of hydronium ion or hydroxide ion present, even if the equilibrium is disturbed

• Because the relationship between H3O+ and OH- is constant, we can simply measure one of them to determine the acidity and basicity of a solution:Ø In practice, we only measure the concentration of hydronium

ions (H3O+) to determine the pH balance of a solution.

• The pH of a solution is defined as the negative log of the hydronium ion concentration:

pH = -log [H3O+]• The inverse it true as well:

[H3O+] = 10-pH

MEASURING ACIDITY WITH pH


Use of brackets around a chemical formula indicate its molar concentration(M = moles/liter)

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• Aqueous solutions are measured by the pH scale in units that range between 0 and 14:

• This is a logarithmic scale!Ø Each change in one pH unit corresponds to a 10-fold difference

in hydronium ion concentration.

Ø If the pH of a solution changes from 4 to 2, it becomes 100 timesmore acidic (not twice as acidic)

THE pH SCALE


Ø Acidic solutions are at the lower end of this pH range.

Ø Basic solutions are at the higher end of this pH range.

Ø A pH of 7 is considered neutral

THE pH SCALE


What is the fold change in [H3O+] between your blood and a can of Coke?

pH ~7.2

pH ~3.0

1. What is the pH of a urine sample with an [H3O+] of 5.6 ´ 10-7 M?

2. Calculate the [H3O+] of grapefruit juice, which has a pH of 3.22.

PRACTICE PROBLEMS


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• pH values can vary widely in different tissue compartments of the body:

• Many medical conditions canalter blood pH & cause health complications.

Ø Diabetes Ø Drug abuse

pH IN CELLS & LIVING TISSUES


Ø The physiological pH of blood (and inside cells) has a range of pH of 7.35 - 7.45

Ø pH values for blood are critical for good health

Ø Kidney diseaseØ Many other conditions

• The form at physiological pH represents the biological useful structure.

• The change in chemical structure at other pH values results in loss of biological functionality.

FUNCTIONAL GROUPS AT PHYSIOLOGICAL pH


-1

-2

+1

Notice that certain functional groups are ionized ….theycarry a charge at physiological pH

• A buffer is a solution that resists changes in pH upon addition of small amounts of an acid or base:Ø Buffers ensure that the pH environment of an aqueous

solution remains relatively constant

Ø Our bodies use several different buffering systems to maintain constant pH in our blood, cells & tissues

• A buffer contains a weak acid and its conjugate base in roughly equal concentrations.

BUFFERS


H3C H

Acetic acid

H3C

Acetate

Weak acid Conjugate base

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• Sodium acetate is an ionic compound that completely dissociates into sodium & acetate ions:Ø Acetate (CH3COO-) is the conjugate base for acetic acidØ In aqueous solution, some of the acetate ions will

reform the acid by reacting with water

• Adding acid (H+) or base to the solution will simply shift the equilibria between acetate:acetic acid:Ø Remember Le Châtelier’s principle!

THE ACETATE BUFFER SYSTEM


H3C

Acetate

+ H3C H

Acetic acid

H2O OH-+

• At some point, a buffer will lose its effectiveness:Ø pH will begin to change

Ø One of the components will eventually be used up.

• The buffering capacity of a buffer depends upon: 1. Concentration of the weak

acid and conjugate base2. Specific properties of the

acid/base pair

LIMITS OF BUFFERING


A titration curve demonstrates the effective buffering range of acetic acid & the acetate ion

When the weak base ammonia (NH3) is mixed with ammonium chloride (NH4Cl) in water, it forms a “buffered” solution:

a. Write the equilibrium equation representing this buffer.

b. What is the purpose of NH4Cl in the buffer?

c. How would this buffer react if H3O+ were added? Show the equation.

d. How would this buffer react if OH- were added? Show the equation.

PRACTICE PROBLEM


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• PhysiologicalbufferingsystemsmaintainaconstantpHinlivingtissuesandfluids:1.Phosphate (PO4

-3),concentrationinthemM rangeinsidecells

2.Histidine,efficientbufferinsidecells atneutralpH

3.Bicarbonate (HCO3-),importantforpHbufferingoutsidecells

inbloodplasma

• Thebicarbonatebuffer helpsmaintainbloodpH:Ø Theweakacidiscarbonicacid(H2CO3).Ø Itsconjugatebaseisbicarbonateion(HCO3

-)

H2CO3 + H2O HCO3- + H3O+

weak acid conjugate base

BLOOD BUFFER


• The source of bicarbonate in the bloodstream is carbon dioxide (CO2), a metabolic waste product:

• Breathing (respiration) regulates pH by changing the concentration of CO2 in the blood:Ø Respiration removes CO2, shifting equilibrium to the left

Ø Not breathing keeps CO2 in the blood, shifting equilibrium to the right

BREATHING & BLOOD pH BALANCE


CO2 + H2O H2CO3

Buffering system

+ H2O HCO3- + H3O+

Respiration (increases pH)

Not breathing (decreases pH)

• Acidosis results when blood pH drops below acceptable limits (too acidic)Ø Respiratory acidosis occurs when the lungs are not

able to remove CO2 efficiently Ø Metabolic acidosis occurs when there is an increase in

acidity in the blood

• Alkalosis occurs when blood pH rises above acceptable limits (too basic = “alkaline”)Ø Respiratory alkalosis results from too-rapid gas

exchange (altitude sickness or hyperventilation)Ø Metabolic alkalosis may come about through vomiting

or consuming too much antacid

MEDICAL CHANGES IN BLOOD pH


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Chapter 9 Slides 2017 - Calvin