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 non-covalent 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 interactions / Shape of substance / Degree of move-ment of molecules / Relative energySolid / 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/Tin 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-orderreaction 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:
- Increasing the concentration of the reactants ( increase the overall number of molecules)
- 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.