Physical chemistry

This is a good book, I really liked it, just as I really liked the other book in the series which I read by the same author, the one about the laws of thermodynamics (blog coverage here). I know much, much more about physics than I do about chemistry and even though some of it was review I learned a lot from this one. Recommended, certainly if you find the quotes below interesting. As usual, I’ve added some observations from the book and some links to topics/people/etc. covered/mentioned in the book below.

Some quotes:

“Physical chemists pay a great deal of attention to the electrons that surround the nucleus of an atom: it is here that the chemical action takes place and the element expresses its chemical personality. […] Quantum mechanics plays a central role in accounting for the arrangement of electrons around the nucleus. The early ‘Bohr model’ of the atom, […] with electrons in orbits encircling the nucleus like miniature planets and widely used in popular depictions of atoms, is wrong in just about every respect—but it is hard to dislodge from the popular imagination. The quantum mechanical description of atoms acknowledges that an electron cannot be ascribed to a particular path around the nucleus, that the planetary ‘orbits’ of Bohr’s theory simply don’t exist, and that some electrons do not circulate around the nucleus at all. […] Physical chemists base their understanding of the electronic structures of atoms on Schrödinger’s model of the hydrogen atom, which was formulated in 1926. […] An atom is often said to be mostly empty space. That is a remnant of Bohr’s model in which a point-like electron circulates around the nucleus; in the Schrödinger model, there is no empty space, just a varying probability of finding the electron at a particular location.”

“No more than two electrons may occupy any one orbital, and if two do occupy that orbital, they must spin in opposite directions. […] this form of the principle [the Pauli exclusion principleUS] […] is adequate for many applications in physical chemistry. At its very simplest, the principle rules out all the electrons of an atom (other than atoms of one-electron hydrogen and two-electron helium) having all their electrons in the 1s-orbital. Lithium, for instance, has three electrons: two occupy the 1s orbital, but the third cannot join them, and must occupy the next higher-energy orbital, the 2s-orbital. With that point in mind, something rather wonderful becomes apparent: the structure of the Periodic Table of the elements unfolds, the principal icon of chemistry. […] The first electron can enter the 1s-orbital, and helium’s (He) second electron can join it. At that point, the orbital is full, and lithium’s (Li) third electron must enter the next higher orbital, the 2s-orbital. The next electron, for beryllium (Be), can join it, but then it too is full. From that point on the next six electrons can enter in succession the three 2p-orbitals. After those six are present (at neon, Ne), all the 2p-orbitals are full and the eleventh electron, for sodium (Na), has to enter the 3s-orbital. […] Similar reasoning accounts for the entire structure of the Table, with elements in the same group all having analogous electron arrangements and each successive row (‘period’) corresponding to the next outermost shell of orbitals.”

“[O]n crossing the [Periodic] Table from left to right, atoms become smaller: even though they have progressively more electrons, the nuclear charge increases too, and draws the clouds in to itself. On descending a group, atoms become larger because in successive periods new outermost shells are started (as in going from lithium to sodium) and each new coating of cloud makes the atom bigger […] the ionization energy [is] the energy needed to remove one or more electrons from the atom. […] The ionization energy more or less follows the trend in atomic radii but in an opposite sense because the closer an electron lies to the positively charged nucleus, the harder it is to remove. Thus, ionization energy increases from left to right across the Table as the atoms become smaller. It decreases down a group because the outermost electron (the one that is most easily removed) is progressively further from the nucleus. […] the electron affinity [is] the energy released when an electron attaches to an atom. […] Electron affinities are highest on the right of the Table […] An ion is an electrically charged atom. That charge comes about either because the neutral atom has lost one or more of its electrons, in which case it is a positively charged cation […] or because it has captured one or more electrons and has become a negatively charged anion. […] Elements on the left of the Periodic Table, with their low ionization energies, are likely to lose electrons and form cations; those on the right, with their high electron affinities, are likely to acquire electrons and form anions. […] ionic bonds […] form primarily between atoms on the left and right of the Periodic Table.”

“Although the Schrödinger equation is too difficult to solve for molecules, powerful computational procedures have been developed by theoretical chemists to arrive at numerical solutions of great accuracy. All the procedures start out by building molecular orbitals from the available atomic orbitals and then setting about finding the best formulations. […] Depictions of electron distributions in molecules are now commonplace and very helpful for understanding the properties of molecules. It is particularly relevant to the development of new pharmacologically active drugs, where electron distributions play a central role […] Drug discovery, the identification of pharmacologically active species by computation rather than in vivo experiment, is an important target of modern computational chemistry.”

Work […] involves moving against an opposing force; heat […] is the transfer of energy that makes use of a temperature difference. […] the internal energy of a system that is isolated from external influences does not change. That is the First Law of thermodynamics. […] A system possesses energy, it does not possess work or heat (even if it is hot). Work and heat are two different modes for the transfer of energy into or out of a system. […] if you know the internal energy of a system, then you can calculate its enthalpy simply by adding to U the product of pressure and volume of the system (H = U + pV). The significance of the enthalpy […] is that a change in its value is equal to the output of energy as heat that can be obtained from the system provided it is kept at constant pressure. For instance, if the enthalpy of a system falls by 100 joules when it undergoes a certain change (such as a chemical reaction), then we know that 100 joules of energy can be extracted as heat from the system, provided the pressure is constant.”

“In the old days of physical chemistry (well into the 20th century), the enthalpy changes were commonly estimated by noting which bonds are broken in the reactants and which are formed to make the products, so A → B might be the bond-breaking step and B → C the new bond-formation step, each with enthalpy changes calculated from knowledge of the strengths of the old and new bonds. That procedure, while often a useful rule of thumb, often gave wildly inaccurate results because bonds are sensitive entities with strengths that depend on the identities and locations of the other atoms present in molecules. Computation now plays a central role: it is now routine to be able to calculate the difference in energy between the products and reactants, especially if the molecules are isolated as a gas, and that difference easily converted to a change of enthalpy. […] Enthalpy changes are very important for a rational discussion of changes in physical state (vaporization and freezing, for instance) […] If we know the enthalpy change taking place during a reaction, then provided the process takes place at constant pressure we know how much energy is released as heat into the surroundings. If we divide that heat transfer by the temperature, then we get the associated entropy change in the surroundings. […] provided the pressure and temperature are constant, a spontaneous change corresponds to a decrease in Gibbs energy. […] the chemical potential can be thought of as the Gibbs energy possessed by a standard-size block of sample. (More precisely, for a pure substance the chemical potential is the molar Gibbs energy, the Gibbs energy per mole of atoms or molecules.)”

“There are two kinds of work. One kind is the work of expansion that occurs when a reaction generates a gas and pushes back the atmosphere (perhaps by pressing out a piston). That type of work is called ‘expansion work’. However, a chemical reaction might do work other than by pushing out a piston or pushing back the atmosphere. For instance, it might do work by driving electrons through an electric circuit connected to a motor. This type of work is called ‘non-expansion work’. […] a change in the Gibbs energy of a system at constant temperature and pressure is equal to the maximum non-expansion work that can be done by the reaction. […] the link of thermodynamics with biology is that one chemical reaction might do the non-expansion work of building a protein from amino acids. Thus, a knowledge of the Gibbs energies changes accompanying metabolic processes is very important in bioenergetics, and much more important than knowing the enthalpy changes alone (which merely indicate a reaction’s ability to keep us warm).”

“[T]he probability that a molecule will be found in a state of particular energy falls off rapidly with increasing energy, so most molecules will be found in states of low energy and very few will be found in states of high energy. […] If the temperature is low, then the distribution declines so rapidly that only the very lowest levels are significantly populated. If the temperature is high, then the distribution falls off very slowly with increasing energy, and many high-energy states are populated. If the temperature is zero, the distribution has all the molecules in the ground state. If the temperature is infinite, all available states are equally populated. […] temperature […] is the single, universal parameter that determines the most probable distribution of molecules over the available states.”

“Mixing adds disorder and increases the entropy of the system and therefore lowers the Gibbs energy […] In the absence of mixing, a reaction goes to completion; when mixing of reactants and products is taken into account, equilibrium is reached when both are present […] Statistical thermodynamics, through the Boltzmann distribution and its dependence on temperature, allows physical chemists to understand why in some cases the equilibrium shifts towards reactants (which is usually unwanted) or towards products (which is normally wanted) as the temperature is raised. A rule of thumb […] is provided by a principle formulated by Henri Le Chatelier […] that a system at equilibrium responds to a disturbance by tending to oppose its effect. Thus, if a reaction releases energy as heat (is ‘exothermic’), then raising the temperature will oppose the formation of more products; if the reaction absorbs energy as heat (is ‘endothermic’), then raising the temperature will encourage the formation of more product.”

“Model building pervades physical chemistry […] some hold that the whole of science is based on building models of physical reality; much of physical chemistry certainly is.”

“For reasonably light molecules (such as the major constituents of air, N2 and O2) at room temperature, the molecules are whizzing around at an average speed of about 500 m/s (about 1000 mph). That speed is consistent with what we know about the propagation of sound, the speed of which is about 340 m/s through air: for sound to propagate, molecules must adjust their position to give a wave of undulating pressure, so the rate at which they do so must be comparable to their average speeds. […] a typical N2 or O2 molecule in air makes a collision every nanosecond and travels about 1000 molecular diameters between collisions. To put this scale into perspective: if a molecule is thought of as being the size of a tennis ball, then it travels about the length of a tennis court between collisions. Each molecule makes about a billion collisions a second.”

“X-ray diffraction makes use of the fact that electromagnetic radiation (which includes X-rays) consists of waves that can interfere with one another and give rise to regions of enhanced and diminished intensity. This so-called ‘diffraction pattern’ is characteristic of the object in the path of the rays, and mathematical procedures can be used to interpret the pattern in terms of the object’s structure. Diffraction occurs when the wavelength of the radiation is comparable to the dimensions of the object. X-rays have wavelengths comparable to the separation of atoms in solids, so are ideal for investigating their arrangement.”

“For most liquids the sample contracts when it freezes, so […] the temperature does not need to be lowered so much for freezing to occur. That is, the application of pressure raises the freezing point. Water, as in most things, is anomalous, and ice is less dense than liquid water, so water expands when it freezes […] when two gases are allowed to occupy the same container they invariably mix and each spreads uniformly through it. […] the quantity of gas that dissolves in any liquid is proportional to the pressure of the gas. […] When the temperature of [a] liquid is raised, it is easier for a dissolved molecule to gather sufficient energy to escape back up into the gas; the rate of impacts from the gas is largely unchanged. The outcome is a lowering of the concentration of dissolved gas at equilibrium. Thus, gases appear to be less soluble in hot water than in cold. […] the presence of dissolved substances affects the properties of solutions. For instance, the everyday experience of spreading salt on roads to hinder the formation of ice makes use of the lowering of freezing point of water when a salt is present. […] the boiling point is raised by the presence of a dissolved substance [whereas] the freezing point […] is lowered by the presence of a solute.”

“When a liquid and its vapour are present in a closed container the vapour exerts a characteristic pressure (when the escape of molecules from the liquid matches the rate at which they splash back down into it […][)] This characteristic pressure depends on the temperature and is called the ‘vapour pressure’ of the liquid. When a solute is present, the vapour pressure at a given temperature is lower than that of the pure liquid […] The extent of lowering is summarized by yet another limiting law of physical chemistry, ‘Raoult’s law’ [which] states that the vapour pressure of a solvent or of a component of a liquid mixture is proportional to the proportion of solvent or liquid molecules present. […] Osmosis [is] the tendency of solvent molecules to flow from the pure solvent to a solution separated from it by a [semi-]permeable membrane […] The entropy when a solute is present in a solvent is higher than when the solute is absent, so an increase in entropy, and therefore a spontaneous process, is achieved when solvent flows through the membrane from the pure liquid into the solution. The tendency for this flow to occur can be overcome by applying pressure to the solution, and the minimum pressure needed to overcome the tendency to flow is called the ‘osmotic pressure’. If one solution is put into contact with another through a semipermeable membrane, then there will be no net flow if they exert the same osmotic pressures and are ‘isotonic’.”

“Broadly speaking, the reaction quotient [‘Q’] is the ratio of concentrations, with product concentrations divided by reactant concentrations. It takes into account how the mingling of the reactants and products affects the total Gibbs energy of the mixture. The value of Q that corresponds to the minimum in the Gibbs energy […] is called the equilibrium constant and denoted K. The equilibrium constant, which is characteristic of a given reaction and depends on the temperature, is central to many discussions in chemistry. When K is large (1000, say), we can be reasonably confident that the equilibrium mixture will be rich in products; if K is small (0.001, say), then there will be hardly any products present at equilibrium and we should perhaps look for another way of making them. If K is close to 1, then both reactants and products will be abundant at equilibrium and will need to be separated. […] Equilibrium constants vary with temperature but not […] with pressure. […] van’t Hoff’s equation implies that if the reaction is strongly exothermic (releases a lot of energy as heat when it takes place), then the equilibrium constant decreases sharply as the temperature is raised. The opposite is true if the reaction is strongly endothermic (absorbs a lot of energy as heat). […] Typically it is found that the rate of a reaction [how fast it progresses] decreases as it approaches equilibrium. […] Most reactions go faster when the temperature is raised. […] reactions with high activation energies proceed slowly at low temperatures but respond sharply to changes of temperature. […] The surface area exposed by a catalyst is important for its function, for it is normally the case that the greater that area, the more effective is the catalyst.”


John Dalton.
Atomic orbital.
Electron configuration.
S,p,d,f orbitals.
Computational chemistry.
Atomic radius.
Covalent bond.
Gilbert Lewis.
Valence bond theory.
Molecular orbital theory.
Orbital hybridisation.
Bonding and antibonding orbitals.
Schrödinger equation.
Density functional theory.
Chemical thermodynamics.
Laws of thermodynamics/Zeroth law/First law/Second law/Third Law.
Conservation of energy.
Spontaneous processes.
Rudolf Clausius.
Chemical equilibrium.
Heat capacity.
Statistical thermodynamics/statistical mechanics.
Boltzmann distribution.
State of matter/gas/liquid/solid.
Perfect gas/Ideal gas law.
Robert Boyle/Joseph Louis Gay-Lussac/Jacques Charles/Amedeo Avogadro.
Equation of state.
Kinetic theory of gases.
Van der Waals equation of state.
Maxwell–Boltzmann distribution.
Thermal conductivity.
Nuclear magnetic resonance.
Debye–Hückel equation.
Ionic solids.
Supercritical fluid.
Liquid crystal.
Benoît Paul Émile Clapeyron.
Phase (matter)/phase diagram/Gibbs’ phase rule.
Ideal solution/regular solution.
Henry’s law.
Chemical kinetics.
Rate equation/First order reactions/Second order reactions.
Rate-determining step.
Arrhenius equation.
Collision theory.
Diffusion-controlled and activation-controlled reactions.
Transition state theory.
Redox reactions.
Electrochemical cell.
Fuel cell.
Reaction dynamics.
Spectroscopy/emission spectroscopy/absorption spectroscopy/Raman spectroscopy.
Raman effect.
Magnetic resonance imaging.
Fourier-transform spectroscopy.
Electron paramagnetic resonance.
Mass spectrum.
Electron spectroscopy for chemical analysis.
Scanning tunneling microscope.


October 5, 2017 - Posted by | Biology, Books, Chemistry, Pharmacology, Physics

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