4.76.
Measurements of the pH of blood and urine are commonly used in medical diagnoses.
Experimentally.
Between approximately 10% and 90% neutralization of the proton-donor species.
Buffering.
Tetany, which involves muscle contractions in the hands, arms, feet, and larynx.
To determine the amount of an acid in a given solution.
A mixture of equal concentrations of acetic acid and acetate ion.
It expresses the extent of ionization of a weak acid or base.
7.4
pH = -log[H+]
Stronger acids have larger ionization constants, while weaker acids have smaller ionization constants.
By the pKa value.
CH₃COOH and CH₃COO⁻.
Ionization constants or acid dissociation constants.
Weak acids partially ionize to release a hydrogen ion, thus lowering the pH of the aqueous solution.
[H+] = Ka * ([HA] / [A-])
By hyperventilating and ingesting sodium bicarbonate.
It absorbs either H+ or OH- through the reversibility of the dissociation of acetic acid.
It becomes the corresponding proton acceptor, in this case, the acetate anion (CH3COO−).
Weak bases accept a hydrogen ion, increasing the pH.
Equal numbers of hydrogen ions (hydronium ions, H3O+) and hydroxide ions.
It is corrected for buffer concentration and measured at physiological temperature, making it a closer approximation to the relevant value in warm-blooded animals.
It contains a reserve of bound H+, which can be released to neutralize an addition of OH- to the system, forming H2O.
At a pH close to its pKa of 6.86.
Acidosis
CH₃COOH, H₂PO₄⁻, and NH₄⁺.
Substitute pH for -log[H+] and pKa for -logKa.
They calculated that it would take a gallon and a half of dilute HCl to get the desired effect, which would dissolve their teeth and burn their throats.
Buffers are aqueous systems that tend to resist changes in pH when small amounts of acid or base are added.
The ionic state of ionizable groups is determined by the pH of the surrounding medium, and when sequestered in the middle of a protein, their apparent pKa can differ significantly from their pKa in water.
Buffer systems.
The pH of the mixture is plotted against the amount of NaOH added, expressed as a fraction of the total NaOH required to convert all the acetic acid to its deprotonated form, acetate (CH3COO-).
A proton donor and a proton acceptor.
HA ⇌ H⁺ + A⁻
1.0 mL of 10.0 M NaOH.
The titration curves of three weak acids with very different ionization constants: acetic acid, dihydrogen phosphate, and ammonium ion.
A glass electrode that is selectively sensitive to H+ concentration but insensitive to Na+, K+, and other cations.
CH3COOH ⇌ CH3COO− + H+
7.2
6.86.
It is balanced exactly by an increase in the other component.
H+ ions are direct participants in many reactions, but even in reactions where there is no apparent role for H+ ions, pH changes can still affect the process.
The concentration of the proton donor (acetic acid) exactly equals that of the proton acceptor (acetate), and the buffering power is maximal.
Many amino acids with functional groups that are weak acids or weak bases.
Acids are proton donors, and bases are proton acceptors.
The stronger the tendency to dissociate a proton, the stronger the acid and the lower its pKₐ.
The stronger the base, the larger the pKa of its conjugate acid.
Because pH affects the structure and activity of biological macromolecules, and a small change in pH can cause a large change in the structure and function of a protein.
Because they can cause significant changes in the structure and function of biological macromolecules.
25°C
Specific conditions of concentration (components at 1 M) and temperature (25°C).
The resulting shortness of breath mimicked that in diabetic acidosis or end-stage kidney disease.
The Henderson-Hasselbalch equation.
pH against the amount of NaOH added.
The imidazole group.
Compounds that can give up two protons, such as carbonic acid and glycine.
K2 = [H2CO3] / ([CO2(aq)][H2O])
0.010 mol.
pOH = 2.0.
The signal from the glass electrode placed in a test solution is amplified and compared with the signal generated by a solution of accurately known pH.
Alkalosis
Ka = [H+][A-] / [HA]
They experimented on themselves.
No, only their ratio changes.
From the volume and concentration of NaOH added.
The pH can be calculated using the Henderson-Hasselbalch equation.
Around pH 4.76
Carbon dioxide is a gas under normal conditions.
H₂PO₄⁻ as proton donor and HPO₄²⁻ as proton acceptor
pH = pKa + log([A-]/[HA])
CO₂ in the gas phase affects the concentration of H₂CO₃ in the aqueous phase, influencing the pH.
A proton donor and its corresponding proton acceptor, related by a reversible reaction.
6.86
Keq = [H+][OH-] / [H2O]
The result is a small change in the ratio of the relative concentrations of the weak acid and its anion, leading to a small change in pH.
Almost every biological process is pH-dependent; a small change in pH produces a large change in the rate of the process.
The relatively flat zone extending about 1 pH unit on either side of the midpoint pH of 4.76 in the titration curve of acetic acid.
Constancy of pH is achieved primarily by biological buffers, which are mixtures of weak acids and their conjugate bases.
Their characteristic equilibrium constants.
Nucleotides such as ATP and many metabolites of low molecular weight
The equilibria described by Equations 2-5 and 2-6.
That a weak acid and its anion—a conjugate acid-base pair—can act as a buffer.
Because -log(10^-14) = -log[H+] + -log[OH-], which simplifies to 14 = pH + pOH.
Because at that point, [HA] = [A-], making the log term zero.
Because they are not completely ionized when dissolved in water and play important roles in metabolism and its regulation.
pKₐ is analogous to pH and is defined by the equation pKₐ = -log Kₐ.
9.25.
pH = pKa - log([HA] / [A-])
A buffer system consists of a weak acid (the proton donor) and its conjugate base (the proton acceptor).
Ionic interactions are among the forces that stabilize a protein molecule and allow an enzyme to recognize and bind to its substrate.
Histidine residues in proteins help buffer effectively near neutral pH.
CH3COOH and CH3COO-
CO2(aq) + H2O ⇌ H2CO3
The phosphate and bicarbonate systems
pOH is the negative logarithm of the concentration of OH-.
The concentration of HCO₃⁻ in the aqueous phase and by pCO₂ in the gas phase.
The concentration of the proton donor [HAc] equals that of the proton acceptor [Ac-].
Kₐ
They reveal the pKₐ of weak acids.
It is the pH at the midpoint of the titration curve.
Between pH 8.25 and 10.25.
They gave ammonium chloride to babies with tetany, which cleared up the tetany in a few hours.
Sodium hydroxide (NaOH).
The protonated amino and carboxyl groups of amino acids and the phosphate groups of nucleotides function as weak acids.
They have a characteristic and nearly constant pH.
Carbonic acid (H2CO3) as proton donor and bicarbonate (HCO3-) as proton acceptor.
The pH is 7.0.
At about pH 7.0.
Given pH and the molar ratio of proton donor and acceptor.
H₂CO₃ acts as the proton donor and HCO₃⁻ acts as the proton acceptor.
Strong acids like hydrochloric, sulfuric, and nitric acids, and strong bases like NaOH and KOH, are completely ionized in dilute aqueous solutions.
Kₐ = [H⁺][A⁻] / [HA]
The stronger the acid, the smaller its pKa.
Kw = [H+][OH-] = 10^-14 M^2
He ate ammonium chloride, which breaks down in the body to release hydrochloric acid and ammonia.
Because extracellular fluids and most cytoplasmic compartments in mammals have a pH in the range of 6.9 to 7.4.
2 equivalents.
At pH 7.3.
Generally between 10% and 90% titration of the weak acid.
Because one of its components, carbonic acid (H2CO3), is formed from dissolved carbon dioxide (CO2) and water (H2O) in a reversible reaction.
Ammonia
Because the three acids have different strengths.
NaOH dissociates completely into Na+ and OH-.
Given pKa and the molar ratio of proton donor and acceptor.
They buffer cells and tissues against pH changes.
At the midpoint, exactly 0.5 equivalent of NaOH has been added per equivalent of the acid, and one-half of the original acetic acid has undergone dissociation.
-log[H+] = -logKa - log([HA] / [A-])
Because they made themselves alkaline by hyperventilating and ingesting sodium bicarbonate.
The equilibrium quickly adjusts to restore the product to 1 x 10^-14 M^2 (at 25°C).
Between about 5.9 and 7.9.
1 equivalent.
At pH 7.3.
Because histidine can exist in either the protonated or unprotonated form near neutral pH.
4.76
The stronger the acid, the greater its tendency to lose its proton.
The pH is above the pKa (6.86), indicating the acid is more than 50% titrated.
HAc dissociates further to satisfy its own equilibrium constant.
With an indicator dye or a pH meter.
HPO4^2- (the species that gains a proton).
Because it is a salt of a strong acid and strong base, and does not resist changes in pH.
pKa = 6.0.
K1 = [H+][HCO3-] / [H2CO3]
The vacuoles
The concentrations of H₂CO₃ and HCO₃⁻, the proton donor and acceptor components.
The added OH- combines with the free H+ in the solution to form H2O.
Enzymes and many molecules contain ionizable groups with characteristic pKa values, which are influenced by the pH of the surrounding medium.
The pH changes slightly, but the change is very small compared to the pH change that would result if the same amount were added to pure water or a solution of NaCl.
6.0
pH = 7.3.
Compounds that can give up three protons, such as phosphoric acid.
Ka = [CO2(aq)] / [CO2(g)]
0.032 mol of NaH2PO4 and 0.068 mol of Na2HPO4.
pH + pOH = 14.
Important quantitative relationships in the titration curve of weak acids.
The equilibria between gaseous CO₂ in the lungs and bicarbonate (HCO₃⁻) in the blood plasma.
The pH is exactly equal to the pKa of acetic acid.
The amount of a substance that will react with, or supply, one mole of hydrogen ions in an acid-base reaction.
pK1 = 1.8, pK2 (imidazole) = 6.0, and pK3 = 9.2.
The Henderson-Hasselbalch equation.
The concentrations of the proton donor (CH3COOH) and the proton acceptor (CH3COO-) are equal, and the pH is numerically equal to the pKa.
The bicarbonate system.
20 to 1
The new pH is 7.2.
At pH 4.76.
Given pH and pKa.
The pKa of acetic acid is 4.76.
H2PO4- (the species that gives up a proton).
The apparent pKa of acetic acid.
pH = pKa + log([A-]/[HA]).
The ionization of histidine, illustrating its role as a weak acid with a pKa of 6.0.
It represents the tendency of an acid to lose its proton and form its conjugate base.
About 5%
The remaining nondissociated acetic acid is gradually converted into acetate.
The final pH is 12.
Cells and organisms maintain a specific and constant cytosolic pH, usually near pH 7, to keep biomolecules in their optimal ionic state.
Either the protonated form or the unprotonated form.
The useful region of buffering power.
The fraction is 20/21, or about 95.2%
pH = pKa + log([HPO4^2-]/[H2PO4^-]).
In pure water, the pH increases from 7 to 12, while in a buffered solution, it increases from 7.0 to just 7.2.
At the midpoint of its titration, where [HA] = [A-].
pCO₂.
Compounds that can give up only one proton, such as acetic acid and ammonium ion.
It represents the pKₐ value for each reaction.
0.042 mol.
Acetic acid, with the highest Ka (lowest pKa).
0.010 mol/L or 1.0 × 10^-2 M.
The concentration of dissolved CO₂, which in turn depends on the concentration of CO₂ in the gas phase or the partial pressure of CO₂ (pCO₂).
More and more HAc ionizes, forming Ac-, as NaOH is added.
Because the H₂CO₃ of blood plasma is in equilibrium with a large reserve capacity of CO₂(g) in the air space of the lungs.
The acetic acid is already slightly ionized, to an extent that can be calculated from its ionization constant.
