pH, pOH, pKa and pKb (acids and bases)
pH, pOH, pKa, pKb
(“p” stands for power)
• pH is a scale measuring the acidity or alkalinity of a solution, defined as the
negative logarithm of the hydrogen ion concentration.
• The pH scale ranges from 0 to 14, where 7 is neutral, values below 7 are
acidic, and values above 7 are basic.
• pOH is the negative logarithm of the hydroxide ion (OH⁻ concentration) in a
solution, denoted by the formula pOH = −log[OH⁻]. It serves as a scale to measure
the basicity or acidity of a solution.
• With a neutral solution having a pOH of 7, basic solutions having a pOH less
than 7, and acidic solutions having a pOH greater than 7.
• The pOH is directly related to pH by the equation pH + pOH = 14 at 25°C.
• pKa is a measure of an acid's strength, defined as the negative base-10 logarithm
of the acid dissociation constant (Ka (tendency to lose protons)).
• A lower pKa value indicates a stronger acid, meaning it donates a proton (H+)
more readily.
• The pKa value helps determine how easily a molecule releases its hydrogen
ion and is crucial for calculations involving acid-base titrations and buffer
solutions.
• pKb is a measure of a base’s strength, defined as the negative base-10 logarithm
of the basic dissociation constant (Kb (tendency to attract protons)).
• A lower pKb value indicates a stronger base, meaning it attracts a proton (H+)
more easily.
• The pKb value helps determine how easily a molecule attracts a hydrogen ion
and is crucial for calculations involving acid-base titrations and buffer
solutions.
• In a buffer solution, the pKa is the pH of the buffer
ACIDS AND BASES
Theories:
• Arrhenius: an acid is a substance that dissolved in water releases H+ ions; while
a base is a substance that dissolved in water releases OH- ions.
• There are no free H+ ions because they always bond instantly with a water
molecules, with which forms the H3O+ ion.
• Brønsted and Lowry: an acid is a donor (a species that gives protons) and a base
is an acceptor (a species that receives protons)
• Chantelier: a substance changes the equilibrium when acid or bases are added so
that the equilibrium.
• In distilled water there are little free ions that derive from the ionization of its
molecules, called auto ionization (H2O <=> H+ + OH-).
• Since water can be considered a pure liquid, it can be omitted from the
expression of Kc (the constant of equilibrium), giving a new constant Kw,
ionic product of water: Kw = [H+] x [OH-] = 10^-14 at room temperature (25°).
• Instead of saying this number, scientists prefer keeping things short, so they
calculate the pKw= -logKw = -log10^-14=14
,Definitions:
• A strong acid is a substance that in aqueous solution it dissociates completely.
A weak acid is a substance that in aqueous solution it dissociates partly.
• A strong base is a substance that in aqueous solution gives place to a total
ionization.
A weak base is a substance that in aqueous solution gives place to a partial
ionization.
Low H+ => high pH => basic High H+ => low pH => acid
• Many bases are metal hydroxides, e.g., KOH, Ca(OH)2 ans Mg(OH)2
→ Dissolve in water and ions separate: NaOH (s) –H2O → Na+ (aq) + OH- (aq)
• Acids form conjugate bases after donating the proton, and bases form conjugate
acids after accepting the proton
• Acids are classified as monoprotic, diprotic or triprotic depending on the number
of protons they give
• Amphiprotic = a substance that can be either an acid or a base (e.g., water)
Useful formulas
• pH =-log([H+]) • pOH = -log ([OH-]) • pKw= pH + pOH = 14
• pKa = -log(Ka) • pKb = -log(Kb) • Ka = [H+]*[B] / [A] • Kb = [OH-]*[A]/ [B]
Buffers
• A buffer solution is a mixture of a weak acid and its conjugate base, or a weak
base and its conjugate acid. Buffers resist drastic pH changes when small
amounts of strong acids or bases are added, until a certain point.
• Buffer range = the range of pH values over which a buffer is most effective
(pKa ± 1). If too much acid/base is added, the buffer components are “used up.”
• Once one component (acid or base) is nearly consumed, the buffer loses
effectiveness.
• Buffer capacity = the amount of acid/base a buffer can neutralize before its pH
changes significantly (≈ 1 pH unit).
• Higher concentrations of buffer components = higher buffer capacity.
• Our body has buffers bc they need to keep the equilibrium so that the
concentration of protons and stuff are kind the same
• Equation: henderson-hasselbach -> pH= pKa + log (B/A)
Good buffers:
Rule 1:
• A good buffer has roughly equal concentrations of weak acid and conjugate base.
• A buffer is no longer effective if one falls below ≈10% of the other.
Rule 2:
• Weak acid + salt = best for pH < 7.
• Weak base + salt = best for pH > 7.
• Biological example – Blood:
• Blood pH ≈ 7.35 - 7.45 (slightly basic).
• Buffer system: Carbonic acid (H₂CO₃) / bicarbonate ion (HCO₃⁻).
• Keeps blood pH stable; deviations of ±0.1 are normal, ±0.4 can be fatal.
Buffering in Blood
• Normal blood pH ≈ 7.4 maintained by the carbonate buffer system:
CO2 + H2O ⇌ H2CO3- + H3O+
,• Concentrations: [H₂CO₃] ≈ 0.0012 M, [HCO₃⁻] ≈ 0.024 M.
• Ratio ensures extra buffering against metabolic acids.
• Lactic acid from exercise is neutralized by HCO₃⁻ → H₂CO₃ → broken down to CO₂
and H₂O (removed by breathing).
• Breathing rate also regulates blood pH:
• Faster breathing = removes CO₂, lowers acidity.
• Slower breathing = retains CO₂, increases acidity.
Example: bicarbonate buffer system
2 types of buffers in our body:
Extracellular: (40%)
Bicarbonate buffer system
• Works in the blood to maintain pH
• Blood composed mainly of water, which breaks apart into H+ and OH- ions
→ Hydrogen ions determine pH
• When an acidic substance enters the bloodstream, the bicarbonate ions
neutralize the hydronium ions forming carbonic acid and water
• Equation: CO2 (g) + H2O (l) <-> H2CO3 (aq) <-> H+ (aq) + HCO3- (aq)
Plasma proteins
• Plasma proteins act as an important buffer in the extracellular fluid by binding
to or releasing hydrogen ions as needed.
• They contribute to buffering because they can ionize or combine with H+ to
resist changes in pH.
Phosphate buffer saline (PBS)
• Its major role is as a buffer inside cells and in the urine, but inorganic phosphate
contributes to extracellular buffering as well.
• Water-based salt solution containing disodium hydrogen phosphate, sodium
chloride, potassium chloride and potassium dihydrogen phosphate
• Osmolarity and ion concentration similar to those of the human body
Intracellular: (60%)
Proteins:
• The high concentration of proteins inside cells gives them significant buffering
capacity.
• Their amino acid side chains have acidic and basic groups that can accept or
donate protons (H+) to resist changes in pH.
• For example: hemoglobin- tends to connect to protons without additional e-
• When there is a lot of CO2 it tends to allow it to go inside the cells and connect
to the water molecules
Phosphate buffer system:
• This system is critical in the intracellular fluid.
• It relies on the equilibrium between dihydrogen phosphate (H2PO4-) and
hydrogen phosphate (HPO4 2-) ions to neutralize added acids or bases.
• H2PO4- (weak acid): H2PO4- <-> H+ + HPO4 2-
Organic phosphates:
• Molecules like ATP and ADP are also important intracellular organic phosphate
buffers.
, Osmosis, molarity, osmorality, osmolality
• Osmosis is the net movement of solvent molecules (typically water) across a
semipermeable membrane from an area of lower solute concentration to an
area of higher solute concentration, or from an area of high water concentration
to low water concentration, to equalize the solute concentrations on both sides
of the membrane.
• This passive process is crucial for many biological functions, such as water
uptake by plant roots and the proper functioning of human kidneys and cells.
Key Components of Osmosis
• Semipermeable Membrane: This membrane has pores that allow solvent
molecules, like water, to pass through but block
larger solute molecules (such as salts or sugar).
• Solvent: The substance that dissolves, usually water in biological contexts.
• Solute: The substance that is dissolved in the solvent, like salt or sugar.
• Concentration Gradient: The difference in solute concentration between the two
sides of the membrane. Water moves to dilute the side
with higher solute concentration.
How Osmosis Works
Semi-permeable
membrane: membrane
that lets somethings
pass through but not
everything.
(The same probability of
something going to the
left to the right is for
something that goes
from the right to the left
in any given period of
time)
• The water acts as a solvent (the what there is more of (liquid usually)) (with
some solutes in it (particles dissolved in a liquid)).
• The solutes are not able to diffuse since they are too big to pass through the gaps
in the membrane.
• But the water molecules still move back and forth, but the water molecules in
the side with a higher concentration of the solutes will not move to the other
side as much as the molecules on the other side will.
• This bc the solutes also bounce around, and they can block maybe the free
passage of the membrane (so there is a lower probability for the water in their
side to move).
• Maybe the solute doesn’t necessarily block the passage, maybe it has a positive
or negative charge and since water has partial negative charge on the O and
partial positive charges on the Hs, these solute molecules might stick more to
the water → water molecules can’t move through the membrane
• There is a net migration of water from where there is less concentration of the
solute to where there is more (=osmosis)
pH, pOH, pKa, pKb
(“p” stands for power)
• pH is a scale measuring the acidity or alkalinity of a solution, defined as the
negative logarithm of the hydrogen ion concentration.
• The pH scale ranges from 0 to 14, where 7 is neutral, values below 7 are
acidic, and values above 7 are basic.
• pOH is the negative logarithm of the hydroxide ion (OH⁻ concentration) in a
solution, denoted by the formula pOH = −log[OH⁻]. It serves as a scale to measure
the basicity or acidity of a solution.
• With a neutral solution having a pOH of 7, basic solutions having a pOH less
than 7, and acidic solutions having a pOH greater than 7.
• The pOH is directly related to pH by the equation pH + pOH = 14 at 25°C.
• pKa is a measure of an acid's strength, defined as the negative base-10 logarithm
of the acid dissociation constant (Ka (tendency to lose protons)).
• A lower pKa value indicates a stronger acid, meaning it donates a proton (H+)
more readily.
• The pKa value helps determine how easily a molecule releases its hydrogen
ion and is crucial for calculations involving acid-base titrations and buffer
solutions.
• pKb is a measure of a base’s strength, defined as the negative base-10 logarithm
of the basic dissociation constant (Kb (tendency to attract protons)).
• A lower pKb value indicates a stronger base, meaning it attracts a proton (H+)
more easily.
• The pKb value helps determine how easily a molecule attracts a hydrogen ion
and is crucial for calculations involving acid-base titrations and buffer
solutions.
• In a buffer solution, the pKa is the pH of the buffer
ACIDS AND BASES
Theories:
• Arrhenius: an acid is a substance that dissolved in water releases H+ ions; while
a base is a substance that dissolved in water releases OH- ions.
• There are no free H+ ions because they always bond instantly with a water
molecules, with which forms the H3O+ ion.
• Brønsted and Lowry: an acid is a donor (a species that gives protons) and a base
is an acceptor (a species that receives protons)
• Chantelier: a substance changes the equilibrium when acid or bases are added so
that the equilibrium.
• In distilled water there are little free ions that derive from the ionization of its
molecules, called auto ionization (H2O <=> H+ + OH-).
• Since water can be considered a pure liquid, it can be omitted from the
expression of Kc (the constant of equilibrium), giving a new constant Kw,
ionic product of water: Kw = [H+] x [OH-] = 10^-14 at room temperature (25°).
• Instead of saying this number, scientists prefer keeping things short, so they
calculate the pKw= -logKw = -log10^-14=14
,Definitions:
• A strong acid is a substance that in aqueous solution it dissociates completely.
A weak acid is a substance that in aqueous solution it dissociates partly.
• A strong base is a substance that in aqueous solution gives place to a total
ionization.
A weak base is a substance that in aqueous solution gives place to a partial
ionization.
Low H+ => high pH => basic High H+ => low pH => acid
• Many bases are metal hydroxides, e.g., KOH, Ca(OH)2 ans Mg(OH)2
→ Dissolve in water and ions separate: NaOH (s) –H2O → Na+ (aq) + OH- (aq)
• Acids form conjugate bases after donating the proton, and bases form conjugate
acids after accepting the proton
• Acids are classified as monoprotic, diprotic or triprotic depending on the number
of protons they give
• Amphiprotic = a substance that can be either an acid or a base (e.g., water)
Useful formulas
• pH =-log([H+]) • pOH = -log ([OH-]) • pKw= pH + pOH = 14
• pKa = -log(Ka) • pKb = -log(Kb) • Ka = [H+]*[B] / [A] • Kb = [OH-]*[A]/ [B]
Buffers
• A buffer solution is a mixture of a weak acid and its conjugate base, or a weak
base and its conjugate acid. Buffers resist drastic pH changes when small
amounts of strong acids or bases are added, until a certain point.
• Buffer range = the range of pH values over which a buffer is most effective
(pKa ± 1). If too much acid/base is added, the buffer components are “used up.”
• Once one component (acid or base) is nearly consumed, the buffer loses
effectiveness.
• Buffer capacity = the amount of acid/base a buffer can neutralize before its pH
changes significantly (≈ 1 pH unit).
• Higher concentrations of buffer components = higher buffer capacity.
• Our body has buffers bc they need to keep the equilibrium so that the
concentration of protons and stuff are kind the same
• Equation: henderson-hasselbach -> pH= pKa + log (B/A)
Good buffers:
Rule 1:
• A good buffer has roughly equal concentrations of weak acid and conjugate base.
• A buffer is no longer effective if one falls below ≈10% of the other.
Rule 2:
• Weak acid + salt = best for pH < 7.
• Weak base + salt = best for pH > 7.
• Biological example – Blood:
• Blood pH ≈ 7.35 - 7.45 (slightly basic).
• Buffer system: Carbonic acid (H₂CO₃) / bicarbonate ion (HCO₃⁻).
• Keeps blood pH stable; deviations of ±0.1 are normal, ±0.4 can be fatal.
Buffering in Blood
• Normal blood pH ≈ 7.4 maintained by the carbonate buffer system:
CO2 + H2O ⇌ H2CO3- + H3O+
,• Concentrations: [H₂CO₃] ≈ 0.0012 M, [HCO₃⁻] ≈ 0.024 M.
• Ratio ensures extra buffering against metabolic acids.
• Lactic acid from exercise is neutralized by HCO₃⁻ → H₂CO₃ → broken down to CO₂
and H₂O (removed by breathing).
• Breathing rate also regulates blood pH:
• Faster breathing = removes CO₂, lowers acidity.
• Slower breathing = retains CO₂, increases acidity.
Example: bicarbonate buffer system
2 types of buffers in our body:
Extracellular: (40%)
Bicarbonate buffer system
• Works in the blood to maintain pH
• Blood composed mainly of water, which breaks apart into H+ and OH- ions
→ Hydrogen ions determine pH
• When an acidic substance enters the bloodstream, the bicarbonate ions
neutralize the hydronium ions forming carbonic acid and water
• Equation: CO2 (g) + H2O (l) <-> H2CO3 (aq) <-> H+ (aq) + HCO3- (aq)
Plasma proteins
• Plasma proteins act as an important buffer in the extracellular fluid by binding
to or releasing hydrogen ions as needed.
• They contribute to buffering because they can ionize or combine with H+ to
resist changes in pH.
Phosphate buffer saline (PBS)
• Its major role is as a buffer inside cells and in the urine, but inorganic phosphate
contributes to extracellular buffering as well.
• Water-based salt solution containing disodium hydrogen phosphate, sodium
chloride, potassium chloride and potassium dihydrogen phosphate
• Osmolarity and ion concentration similar to those of the human body
Intracellular: (60%)
Proteins:
• The high concentration of proteins inside cells gives them significant buffering
capacity.
• Their amino acid side chains have acidic and basic groups that can accept or
donate protons (H+) to resist changes in pH.
• For example: hemoglobin- tends to connect to protons without additional e-
• When there is a lot of CO2 it tends to allow it to go inside the cells and connect
to the water molecules
Phosphate buffer system:
• This system is critical in the intracellular fluid.
• It relies on the equilibrium between dihydrogen phosphate (H2PO4-) and
hydrogen phosphate (HPO4 2-) ions to neutralize added acids or bases.
• H2PO4- (weak acid): H2PO4- <-> H+ + HPO4 2-
Organic phosphates:
• Molecules like ATP and ADP are also important intracellular organic phosphate
buffers.
, Osmosis, molarity, osmorality, osmolality
• Osmosis is the net movement of solvent molecules (typically water) across a
semipermeable membrane from an area of lower solute concentration to an
area of higher solute concentration, or from an area of high water concentration
to low water concentration, to equalize the solute concentrations on both sides
of the membrane.
• This passive process is crucial for many biological functions, such as water
uptake by plant roots and the proper functioning of human kidneys and cells.
Key Components of Osmosis
• Semipermeable Membrane: This membrane has pores that allow solvent
molecules, like water, to pass through but block
larger solute molecules (such as salts or sugar).
• Solvent: The substance that dissolves, usually water in biological contexts.
• Solute: The substance that is dissolved in the solvent, like salt or sugar.
• Concentration Gradient: The difference in solute concentration between the two
sides of the membrane. Water moves to dilute the side
with higher solute concentration.
How Osmosis Works
Semi-permeable
membrane: membrane
that lets somethings
pass through but not
everything.
(The same probability of
something going to the
left to the right is for
something that goes
from the right to the left
in any given period of
time)
• The water acts as a solvent (the what there is more of (liquid usually)) (with
some solutes in it (particles dissolved in a liquid)).
• The solutes are not able to diffuse since they are too big to pass through the gaps
in the membrane.
• But the water molecules still move back and forth, but the water molecules in
the side with a higher concentration of the solutes will not move to the other
side as much as the molecules on the other side will.
• This bc the solutes also bounce around, and they can block maybe the free
passage of the membrane (so there is a lower probability for the water in their
side to move).
• Maybe the solute doesn’t necessarily block the passage, maybe it has a positive
or negative charge and since water has partial negative charge on the O and
partial positive charges on the Hs, these solute molecules might stick more to
the water → water molecules can’t move through the membrane
• There is a net migration of water from where there is less concentration of the
solute to where there is more (=osmosis)
