PCC 12803 – General Chemistry for the Life Sciences
3. Compounds and chemical bonding: bringing atoms together
3.11 Ionic versus covalent bonding
No bonding: no transfer of electrons, inert atom
Partial transfer: Covalent bonding --> electron transfer occurs less readily
Total transfer: Ionic bonding --> electrons are transferred from one atom to another atom very readily
Element’s electronegativity value (Χ) indicates how strongly an atom of that element can attract an electron:
high/low value indicates that the atom strongly/weakly attracts an electron.
If the electronegativity values of two atoms are similar (≤1.7) they are most likely to undergo covalent bonding.
If there is a large difference in the electronegativity values of two atoms (≥1.7) they are most likely to undergo
ionic bonding.
The principal biological elements have similar electronegativity values and undergo covalent bonding.
3.12 Blurring the boundaries: polarized bonds
A polar bond is a covalent bond in which the electrons are not evenly shared between the two joined atoms.
Distribution of electrons in a polar bond is governed by the electronegativity and is skewed towards the most
electronegative of the two atoms.
If the difference in electronegativity of two atoms is large/small, the bond joining them will be highly/slightly
polarized.
If the electrons in a covalent bond are shared equally, the bond is non-polar --> dipole moment of zero.
4. Molecular interactions: holding it all together
4.1 Chemical bonding versus non-covalent forces
Non-covalent interactions operate between parts of one molecule or between separate molecules (over short
distances). They are weaker than covalent bonds and can therefore be disrupted more readily. Individual noncovalent interactions are weak, but many interactions may operate between two molecules, with a large
overall effect.
Intramolecular interactions operate between separate parts of the same molecule.
Intermolecular interactions operate between different molecules.
Molecular interactions can act to stabilize a molecule, or the association of neighbouring molecules.
4.2 Electrostatic forces: the foundations of molecular interactions
Non-covalent molecular interactions are primarily electrostatic in nature: they are based on the notion of
opposite charges attracting one another.
Polarization is the process by which electrons are unevenly distributed within a molecule.
The uneven distribution of electrons throughout a molecule gives rise to partial positive and partial negative
charges. Dipole moment is the difference between these partial positive and negative charges.
A dipole is a molecule or part of a molecule that possesses a region of partial negative charge and a region of
partial positive charge. A permanent dipole is one in which the uneven distribution of electrons is permanent.
If electrons in a covalent bond are shared equally, the bond is non-polar. Molecules containing polar bonds
may be non-polar overall if it comprises identical atoms that are symmetrically (i.e. CO2).
4.3 The van der Waals interaction
Attractive forces:
- Dispersion forces are weak forces of attraction that operate over short distances between all covalent
molecules. They exist because electrons are always slightly unevenly distributed within molecules, generating
areas of partial negative and partial positive charge, which form a (temporary) induced pole.
The prevalence of dispersion forces between two molecules is influenced by their shape (closely together:
stronger) and size (larger: greater).
Dispersion forces are short-lived forces arising from temporary induced dipoles.
- Permanent dipolar interactions (dipole-dipole) are the forces of attraction that exist between opposite partial
changes on polar molecules. These are long-lived forces arising from permanent dipoles.
Repulsive force:
- Steric repulsion=two areas of like charge experience repulsion, which acts to offset attractive interactions to
some extent; operates over only very short distances.
4.4 Beyond van der Waals: other biologically essential interactions
Hydrogen bond is a special type of dipolar interaction in which a hydrogen atom on one molecule interacts
strongly with one of three electronegative atoms either on another molecule.
Hydrogen bond can only form between a hydrogen atom and atom of O, F or N.
The hydrogen atom must itself be covalently bonded to O, F or N.
The three nuclei that participate in a hydrogen bond must lie in a straight line and the distance between the
two electronegative atoms must fall within a narrow range of values.
Hydrogen bonds facilitate the formation of the double-helical structure of DNA, and underpin the specific base
pairing upon which the semi-conservative replication of DNA depends.
Hydrogen bonds operate between different parts of the peptide backbone of polypeptides to generate
conserved three-dimensional motifs.
Hydrophilic forces operate between polar molecules and water molecules to make polar molecules soluble in
aqueous media; they are mediated by hydrogen bonds between the polar molecule and water.
Ionic forces operate between ionic species carrying full positive and negative charges. They may operate
between parts of a covalent molecule that possess full positive and negative charges.
Salt bridge=ionic force that operates between oppositely charged amino acid side chains in proteins.
Hydrophobic forces arise when hydrophobic entities cluster together to shield themselves from water.
The clustering of hydrophobic entities during hydrophobic interactions is stabilized by dispersion forces.
Non-polar molecules are hydrophobic and are immiscible with water.
Polar molecules are hydrophilic and can interact readily with water.
Hydrophobic interactions can be explained by thermodynamic principles.
4.5 Breaking molecular interactions: the three states of matter
The extent of the non-covalent forces that exist between molecules dictates a substance’s physical state.
Nr of molecular
Shape of substance Degree of moveRelative energy
interactions
ment of molecules
Solid
Many
Fixed
Virtually none
Low
Liquid
Few
Variable
Moderate
Medium
Gas
Virtually none
Unrestricted
High
High
As the energy of a molecule increases, the number of non-covalent interactions decreases (neg. relationship).
Melting point=temperature at which a compound makes the transition from a solid to a liquid.
Boiling point=temperature at which a compound makes the transition from a liquid to a gas.
Melting=transition from solid to liquid, associated with a decrease in intermolecular forces.
Vaporization=transition from liquid to gas, associated with a further decrease in intermolecular forces.
Condensation=transition from gas to liquid, associated with an increase in intermolecular forces.
Freezing=transition from liquid to solid, associated with a further increase in intermolecular forces.
Non-polar molecules experience few non-covalent interactions and have low melting and boiling points.
Polar molecules experience more non-covalent interactions and have higher melting and boiling points.
Hydrogen bonds contribute more to the elevation of a compound’s melting and boiling point than dispersion
forces or dipolar interactions.
12. Chemical analysis 2: how do we know how much is there?
12.1 The mole
One mole is just a number, equal to 6 x 10^23 (Avogadro constant).
We can have one mole of any substance (atom, molecule, ion etc.).
The mass of one mole of a substance is its molar mass in g/mol (which has the same value as its atomic mass).
Number of moles present in a sample=mass of sample/molar mass
12.2 Concentrations
Concentration of a solution tells us how much of a substance is present in a particular volume of the solution.
Molarity is the concentration of a chemical substance = number of moles present in 1 L of solution in mol/dm3.
Number of moles (mol)=concentration (mol/L) x volume (L)
12.3 Changing the concentration: solutions and dilutions
Solvent is the medium in which a substance (the solute) is dissolved in aqueous solutions, the solvent is water.
When we dilute a solution, the number of moles of the solute remains the same, but the total volume
increases, so the concentration decreases.
Serial dilution reduces the concentration of a solution in a series of steps, such that the concentration
gradually falls from step to step. (In a graph the curve has an exponential decay)
12.4 Measuring concentrations (spectroscopic approaches)
UV-visible spectrophotometry can measure the concentration of a compound in a solution and measures the
absorbance of a solution at a given, specific wavelength.
Spectrophotometer compares the intensity of light shone onto the sample with the amount that has passed
through it to calculate the amount of light absorbed by the sample=absorbance, A. The absorbance is
proportional to the concentration of the compound.
The relationship between absorbance and concentration is described by the Beer-Lambert law: A=ε * c * l
(absorbance=molar absorptivity (in L/(mol*cm)) * concentration in (mol/L) * path length/width cuvette (in cm))
Atomic spectroscopy measures the emission of energy from a sample following its irradiation with
electromagnetic radiation.
The intensity of the light emitted is proportional to the concentration of atoms present in the sample.
Atomic emission spectroscopy is used to determine the concentration of a specific element in a sample. By
using a calibration curve (intensity vs. conc) the concentration of an unknown sample can be determined.
Fluorescence is the emission of light by a substance, following irradiation with electromagnetic radiation,
whose wavelength is shorter than the wavelength used to excite the substance in the first place.
Fluorescence spectroscopy exploits the way that certain groups of atoms (fluorophores), absorb light of one
wavelength and emit light of another longer wavelength.
The intensity of light emitted is proportional to the concentration of the compound present.
It is very sensitive and can be used to determine very low concentrations of biological molecules.
12.5 Using chemical reactions to measure concentration (chemical approaches)
Titrations exploit the change in colour that occurs during a certain chemical reaction to help us determine the
concentration of a compound of interest.
Indicators are often used to help track the progress of a titration.
Electrochemical sensors exploit the transfer of electrons during redox reactions to give a measure of the
concentration of a particular compound: the greater the flow of electrons, the higher the concentration of the
compound of interest.
13. Energy: what makes reactions go?
13.1 What is energy?
Energy is the capacity to do work. The total energy of a system is its’ internal energy, U.
Total internal energy U = work w + heat q
Energy cannot be created or destroyed, but only converted from one form to another as energy is conserved.
Kinetic energy is the energy an object has due to its motion; it depends on its mass and velocity.
Potential chemical energy is the energy a chemical compound has stored in its chemical bonds.
The amount of energy stored in a bond=bond energy.
When a bond is broken, an amount of energy equal to the bond energy is consumed.
When a bond is formed, an amount of energy equal to the bond energy is liberated.
13.2 Energy transfer
Energy transfer occurs between a system (=particular thing we’re interested in, contained within a boundary)
and its surroundings (=everything else in contact with the surroundings).
Open system: matter and energy can be transferred between the system and its surroundings.
Closed system: only energy can be transferred between the system and its surroundings.
Isolated system: matter and energy cannot be transferred between the system and its surroundings.
Work is any process that can lift a weight and heat is the transfer of energy from hot to cold.
The transfer of energy as heat – from hot to cold – is a spontaneous process.
A substance’s temperature gives a measure of its energy: a system with a high temperature has a large amount
of energy and a system with a low temperature has a smaller amount of energy.
! Temperature: chemical (25 C 298 K) and biochemical (37 C of 310 K)
13.3 Enthalpy
The energy change associated with a chemical reaction is called the enthalpy change of reaction, ΔH.
ΔH=difference between the energy consumed to break bonds and energy liberated when bonds are formed.
H=U + pV (enthalpy=internal energy + pressure*volume)
Exothermic reaction has a negative value for ΔH: heat is transferred from the system to its surroundings (heat
is released).
Endothermic reaction has a positive value for ΔH: heat is transferred from the surroundings to a system (heat is
absorbed).
ΔH=ΣE (reactants) – ΣE (products)
We cannot measure the enthalpy of a compound directly. We can only measure the change in a compound’s
enthalpy when some kind of chemical reaction occurs. We can depict the energy change by using an energy
diagram. We can measure the enthalpy change for an exothermic reaction using a bomb calorimeter.
Enthalpy change of formation is the enthalpy change when one mole of a compound forms from its
constituent elements in their standard states.
Enthalpy change of combustion is the enthalpy change when one mole of a compound burns completely in
excess oxygen.
For a reaction whose ΔH is negative, the products are more stable than the reactants.
For a reaction whose ΔH is positive, the reactants are more stable than the products.
13.4 Entropy: the distribution of energy as the engine of change
Entropy gives us a measure of the distribution of energy in a system: the greater the spread of energy (and the
greater the entropy), the greater the disorder in the system.
Gases have a larger entropy than liquids, which have a larger entropy than solids.
More disorganized systems have a larger entropy than organized ones: unfolded proteins have a larger entropy
than folded ones.
Systems with more components have an inherently larger entropy than those with fewer components.
Entropy of a substance increases as its energy increases.
 During an exothermic reaction, the entropy of the system decreases but the entropy of the surroundings
increases as energy transfers from the system to the surroundings.
 During an endothermic reaction, the entropy of the system increases but the entropy of the surroundings
decreases as energy transfers from the surroundings to the system.
The extent of the increase in entropy of the surroundings depends on the temperature of the surroundings:
the higher the temperature, the smaller the increase in entropy.
ΔS= Δq/T (change in entropy in J/K=change in heat energy/temperature)
When the surroundings are at constant pressure: q=H thus ΔS= ΔHsur/T in J/K*mol
For a reaction to be spontaneous, the overall entropy change – ΔStotal – must be greater than zero. That is,
there must be a net increase in entropy.
13.5 Spontaneous versus non-spontaneous processes: how much energy do we need?
A spontaneous process proceeds to completion without an input of energy and only happens in one direction.
If we want to reverse a spontaneous process, we must provide a source of energy to drive the process.
13.6 Gibbs free energy: the driving force of chemical reactions
The change in Gibbs free energy for a reaction, ΔG, is the energy that is free to be harnessed to do something
useful – that is, to do work.
ΔG= ΔH – TΔS (change in Gibbs free energy in J/mol=change in enthalpy – energy required to drive the change
in entropy)
For a reaction in the system to be spontaneous, ΔG must be negative.
If ΔG is positive, energy must be supplied to the system for the reaction to happen.
Reactions that release Gibbs free energy are called exergonic.
Reactions that require an input of energy are called endergonic.
Catabolic reactions are typically exergonic: they involve the breakdown of large molecules into simpler
components. Anabolic reactions are typically endergonic: they involve the building up of large molecules from
simpler subunits. Anabolic and catabolic reactions can be coupled such that the energy yielded by exergonic
reactions can be used to drive endergonic reactions.
14. Kinetics: what affects the speed of a reaction?
14.1 The rate of reaction
The rate of a chemical reaction is the speed of change in concentration of reactants or products per unit time
as the reaction proceeds = change measured/change in time
We can determine the rate of a reaction by measuring the concentrations of reactants or products at different
times during the course of a reaction, and plotting on a graph the change of concentration with time.
Rate of reaction at a particular time is equal to the gradient (slope) of the graph at that time. Curve  tangent.
A  Products
Rate of reaction = k[A]x in which k is the rate constant. The rate of a reaction is dependent upon the
concentration of the reactants in a way that is determined by the order of the reaction.
x=0: For a zero-order reaction, the rate of reaction is independent of the concentration of reactants, and the
rate is constant, regardless of the concentration of reactants present. (conc. vs. time: straight line)
x=1: For a first-order reaction, the rate of reaction is directly proportional to the concentration of reactants: as
the concentration increases, the reaction rate will also increase. (conc. vs. time: slope)
The half-life tells us how long it takes for the concentration of a reactant to fall to half of its initial value.
The half-life for a zero-order reaction decreases as the concentration of reactant decreases.
The half-life for a first-order reaction stays constant, regardless of the concentration of reactant.
14.2 The collision theory of reaction rates
In order for a reaction to happen, the reacting molecules must collide.
The collision between reacting molecules must happen with sufficient energy and in the correct orientation.
Typically, only a small number of reactant molecules have enough energy to be able to react.
We can increase the number of molecules with sufficient energy to react by:
1. Increasing the concentration of the reactants ( increase the overall number of molecules)
2. Increasing the temperature of the system ( increase the average kinetic energy of the molecules in
the system & increase the rate of reaction)
14.3 The activation energy: getting reactions started
Activation energy is the minimum energy a reactant molecule must possess for the reaction to proceed.
Transition state is the highest-energy species that exists during the course of a chemical reaction; it occurs at
the point in the reaction when the activation energy is achieved.
14.4 Catalysis: lowering the activation energy
We can increase the rate of reaction by lowering its activation energy.
We can lower the activation energy of a reaction by using a catalyst=a substance that alters the rate of a
chemical reaction, but is itself chemically unchanged at the end of the reaction.
A catalyst merely influences the rate at which a spontaneous reaction occurs.
14.5 Enzymes: the biological catalysts
Enzymes are protein based catalysts which operate in biological systems; they are totally specific: a particular
enzyme catalyses only one type of reaction.
When considering enzyme-catalysed reactions, we refer to the reactant as the substrate.
The cleft is the site on the enzyme at which the substrate binds to the enzyme.
The active site is the region of the cleft that exhibits the highest degree of complementarity with the shape of
the substrate and participates in the strongest interactions with the substrate, binding it firmly in place.
1. The substrate (S) binds to the active site of the enzyme, forming an enzyme-substrate complex (ES).
2. The enzyme adjusts its conformation slightly so that the active site and substrate together take on the shape
of the transition state.
3. The transition state-enzyme complex gives rise to the product (P).
4. The product is released from the active site, and the enzyme is returned unchanged to its native
conformation, ready to bind another substrate.
Enzymes operate via the so-called induced-fit mechanism (E+S ↔ ES  E+P)
The key function of an enzyme is to facilitate a smooth, low-energy pathway from substrate to transition state.
14.6 Enzyme kinetics
The maximum rate of an enzyme-catalysed reaction is called Vmax. At Vmax the enzyme is working at full
capacity: every active site has a substrate bound to it. As Vmax is approached, any further increase in substrate
concentration does not produce any increase in reaction rate.
The number of substrate molecules processed by an enzyme per second=turnover number.
Km (Michaelis constant) tells us the concentration of substrate at which the velocity of the reaction is ½ Vmax.
A small Km value tells us that an enzyme binds strongly to its substrate and needs only a low concentration of
substrate to catalyse the formation of product effectively.
A large Km value tells us that the enzyme binds weakly to its substrate and needs a
high concentration of substrate to catalyse the formation of product effectively.
Michaelis-Menten equation: v = Vmax [S] / Km [S]
Lineweaver-Burk plot plots 1/v against 1/[S]:
y=ax+b  1/v = Km/Vmax * (1/[S]) + 1/Vmax
The x-axis intercept is –(1/Km) and the y-as intercept is Km/Vmax.
Enzymes operate at optimum temperatures, above which enzymes denature and lose their catalytic activity.
15. Equilibria: how far do reactions go?
15.1 Equilibrium reactions
Equilibrium reaction=reaction that can proceed in both the forward and reverse directions simultaneously.
At equilibrium, the rates of the forward and back reactions are equal; they continue to proceed, but there is no
overall change in the system=dynamic equilibrium.
At equilibrium, the concentrations of the chemical species present will usually be different from each other,
but the actual value of that concentration does not change.
The point at which a state of equilibrium is reached is the same, regardless of the direction from which we
approach it.
15.2 Forward and back reactions: where is the balance struck?
If the forward reaction predominates over the back reaction, we say the equilibrium reaction lies to the right
and the products are favoured.
If the back reaction predominates over the forward reaction, we say the equilibrium reaction lies to the left,
and the reactants are favoured.
The position of equilibrium is represented by an equilibrium constant, K (dimensionless).
For an equilibrium reaction: aA + bB ↔ cC + dD, K=[C]c [D]d / [A]a [B]b
K tells us about the ratio of products to reactants when a reaction is at a state of equilibrium and the
magnitude of K tells us the relative position of a reaction at equilibrium:
A very large value for K tells us the reaction proceeds almost completely to the right.
A very small value for K tells us the reaction lies almost completely to the left.
When K has a value between approx. 10-2 and 102, the equilibrium reaction is fairly evenly balanced.
K for a particular reaction depends on the temperature at which the reaction is performed.
15.3 The reaction quotient
Reaction quotient Q tells us the relative ratio of products to reactants at any point in a reaction.
When Q=K the reaction is at equilibrium.
If Q<K the reaction proceeds from left to right (in the direction of the forward reaction).
If Q>K the reaction proceeds from right to left (in the direction of the back reaction).
If Q=K the reaction is already at equilibrium concentrations. There is no further change.
15.4 Binding reactions in biological systems
The binding of a protein and its ligand can be represented by the general equilibrium reaction: P + L ↔ PL
If a protein and its ligand bind weakly, the binding reaction, PL  P + L (dissociation), lies to the right + Kd large.
If a protein and its ligand bind strongly, the binding reaction, P + L  PL (association), lies to the left + Kd small.
The dissociation constant Kd= [P][L]/[PL] gives a measure of the strength of binding between P and L.
15.5 Perturbing an equilibrium
When an equilibrium is perturbed (disturbed) the system acts to counteract the change so that a state of
equilibrium is re-established=Le Chatelier principle.
There are three ways to perturb an equilibrium:
1. Change the concentration of species present:
Increasing/decreasing the concentration of the reactants favours the forward/back reaction
Increasing/decreasing the concentration of the products favours the back/forward reaction
2. Change the pressure (only gases)
Increasing the pressure favours the reaction which yields the smallest number of molecules
Decreasing the pressure favours the reaction which yields the largest number of molecules
3. Change the temperature
Increasing the temperature favours endothermic reactions (they absorb energy from surroundings)
Decreasing the temperature favours exothermic reactions (they transfer energy to surroundings)
Addition of catalyst increases the rate of both the forward and back reactions, but K remains unchanged.
15.6 Free energy and chemical equilibria
When a reaction is at equilibrium the change in Gibbs free energy ΔG=0.
Van ‘t Hoff isotherm: ΔG0= -RT ln(K) (Gibbs free energy change=gas constant 8.31 (in J/K*mol) * temp (in K) *K)
If the equilibrium constant for a reaction is very large, ΔG0 is negative: the reaction is spontaneous.
If the equilibrium constant for a reaction is very small, ΔG0 is positive: the reaction is not spontaneous.
A reaction with ΔG0 > 0 lies to the left: very few products form.
A reaction with ΔG0 < 0 lies to the right: virtually all the reactants change into products.
16. Acids, bases and the aqueous environment: the medium of life
16.1 Acids and bases: making life happen
An acid, HA, is a proton donor: a proton is generated when an acid dissociates.
A base, B is a proton acceptor: it removes a free proton from solution.
In water, an acid dissociates to generate a proton.
In water, a base accepts a proton from a water molecule.
A proton is stabilised in aqueous solution by reacting with a water molecule to form a solvated proton (H3O+).
Acids and bases must operate in pairs: an acid must have a base to which to donate a proton.
A conjugate acid-base pair comprises the acid and base versions of the same chemical species.
When an acid donates a proton, it becomes its conjugate base.
When a base accepts a proton, it becomes its conjugate acid.
An acid-base reaction involves two acid-base pairs.
Water is amphoteric: it can act as an acid and as a base.
16.2 The strength of acids and bases: to what extent does the dissociation reaction occur?
The strength of an acid tells us the extent to which the acid dissociates when dissolved in water:
 A strong acid dissociates almost completely in water.
 A weak acid only partially dissociates in water.
The strength of a base tells us the extent to which it accepts a proton from water to generate hydroxide ions.
 A strong base accepts a proton from water very readily.
 A weak base is poor at accepting a proton from water.
Strong acids have weak conjugate bases and weak acids have strong conjugate bases.
Acid dissociation constant Ka gives us a measure of the extent to which an acid dissociates in aqueous solution.
HA ↔ H+ + A- …………… Ka=[H+][A-]/[HA]
A strong acid has a large Ka value and a weak acid has a small Ka value.
Base dissociation constant Kb gives us a measure of the extent to which a base reacts with water to yield
hydroxide ions (or how readily it accepts protons from an acid).
B + H2O ↔ BH+ + OH- ……………… Kb = [BH+][OH-]/[B]
A strong base has a large Kb value and a weak base has a small Kb value.
pKa and pKb values are scaled-down representations of Ka and Kb.
pKa= -log (Ka)
pKb = - log (Kb)
16.3 Keeping things balanced: the ion product of water
The ion product of water Kw is the equilibrium constant for the dissociation of water into H + and OH- ions.
Kw=[H+][OH-]=1 x 10-14
The product of the proton concentration and hydroxide concentration for any aqueous solution is equal to the
value of Kw.
For any acid/conjugate base system: Kw = Ka * Kb
16.4 Measuring concentrations: the pH scale
The pH of a solution gives us a measure of the concentration of protons (H +-ions) in solution.
pH= - log [H+]
pH < 7  acidic solution
pH = 7  neutral solution
pH > 7  alkaline solution
The concentration of protons in a strong acid solution is equal to the concentration of the acid itself.
Henderson-Hasselbach equation: pH=pKa + log ([A-]/[HA])
Neutralization is the reaction between an acid and a hydroxide base: acid + base  salt + water
The pH of the final solution depends on the strength of the acid and base that are reacting.
The salt of a strong acid and a strong base gives a neutral solution.
The salt of a strong acid and a weak base gives an acidic solution.
The salt of a weak acid and a strong base gives an alkaline solution.
The salt of a weak acid and a weak base gives an acidic, alkaline or neutral solution depending on the pK a values
of the ions involved.
The pOH of a solution gives us a measure of the concentration of hydroxide ions in the solution.
pOH = -log [OH-]
pH + pOH = 14
16.5 The behaviour of acids and bases in biological systems
The acidity (or basicity) of a compound affects its partitioning between aqueous and hydrophobic systems.
A strong acid is most likely to partition into an aqueous (hydrophilic) phase.
A weak acid is most likely to participate into a hydrophobic/lipophilic phase.
The pH of a compound’s environment can affect its acidity (or basicity).
A decrease in pH leads to a decrease in the extent of dissociation of an acid.
An increase in pH leads to an increase in the extent of dissociation of an acid.
A decrease in pH leads to an increase in the extent of dissociation of a base.
An increase in pH leads to a decrease in the extent of dissociation of a base.
The pKa value of a molecule represents the pH at which it is 50% ionized in solution.
If the pH is lower/higher than an acid’s pKa value, the acid will be less/more than 50% dissociated.
16.6 Buffer solutions: keeping pH the same
A buffer solution is a solution that resists changes in pH; it contains large quantities of an acid-base conjugate
pair, which can suppress the addition of protons or hydroxide ions to a solution and prevent the pH from
changing.
A buffer solution must use a weak acid or base.
It consists of a weak acid and the salt of its conjugate base or a weak base and the salt of the conjugate acid.
The pH around which a solution buffers, is equal to the pKa of the acid.