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Definition: chemistry is the science of matter and of the transformations it can undergo. Nothing material lies outside it: living or dead, mineral or vegetable, on Earth or in stars. The three levels 1. Macroscopic: what can be seen and measured, such as colour, mass, a gas bubbling, a precipitate forming. 2. Microscopic (particulate): the atoms, ions and molecules that explain what is seen. 3. Symbolic: formulas, equations and numbers, for example 2 H2 + O2 -> 2 H2O. A good chemical explanation moves between all three. Example: a metal dissolving in acid (macroscopic) is electron loss by metal atoms (microscopic), written as Zn -> Zn2+ + 2 e- (symbolic). Fields of chemistry - Organic: carbon compounds. - Inorganic: all other elements and their compounds. - Physical: the laws and principles (thermodynamics, kinetics). - Biological chemistry (biochemistry): chemistry of living organisms. - Analytical: identifying substances and measuring amounts. - Theoretical and computational: models and calculations of molecular structure and properties. - Pharmaceutical: chemistry applied to making drugs. The course roadmap Matter is made of atoms -> atomic structure and quantum mechanics -> the periodic table -> chemical bonds -> states of matter (gas, liquid, solid) -> solutions, reactions, equilibria, acids and bases, thermodynamics and kinetics. This subject covers the first part of that map. Clinical tip: laboratory medicine is applied analytical chemistry. A blood gas, a glucose reading or a mass spectrum is a macroscopic measurement that you interpret through microscopic ideas. Exam trap: "chemical" does not mean "artificial". Water, oxygen and sodium chloride are all chemicals.

The International System (SI, adopted in 1960 from the older MKS system of metre, kilogram, second) has seven base quantities. - length: metre (m) - mass: kilogram (kg) - time: second (s) - electric current: ampere (A) - temperature: kelvin (K) - amount of substance: mole (mol) - luminous intensity: candela (cd) Derived quantities are combinations of the base ones. - area: m2 - volume: m3 (in the lab, 1 L = 1 dm3 = 10^-3 m3) - velocity: m s^-1; acceleration: m s^-2 - density: mass / volume, kg m^-3 (in the lab g/mL) - momentum: kg m s^-1 - force: mass x acceleration, kg m s^-2 = newton (N) - pressure: force / area, N m^-2 = pascal (Pa) - energy or work: force x distance, N m = joule (J) - impulse: force x time, N s - surface tension: force / length, N m^-1 - torque: force x distance, N m - electric charge: current x time, A s = coulomb (C) Useful prefixes: kilo 10^3, centi 10^-2, milli 10^-3, micro 10^-6, nano 10^-9, pico 10^-12. Aim: convert units by multiplying by factors equal to one. Example: 250 mL = 250 x 10^-3 L = 0.250 L = 250 cm3. Exam trap: a joule is a newton-metre, so torque and energy share units but are different quantities. Clinical tip: concentration and dose errors are most often unit errors (mg versus ug, mL versus L). Write the unit at every step of a calculation.

Uncertainty - A measurement is written value +/- uncertainty, for example 12.34 +/- 0.02 mL. - The last digit of a reported number is the uncertain one. Counting significant figures - All non-zero digits are significant. - Zeros between non-zero digits are significant: 30.045 has five. - Leading zeros only place the decimal point and are not significant: 0.000030 has two. - Trailing zeros after a decimal point are significant: 2.500 has four. - Trailing zeros in a whole number without a decimal point are ambiguous. Use scientific notation: 3.0 x 10^2 has two, 3.00 x 10^2 has three. Rounding - If the first digit dropped is less than 5, keep the previous digit unchanged: 0.243 -> 0.24. - If it is greater than 5, increase the previous digit by one: 0.248 -> 0.25. - If it is exactly 5, the usual convention is to round up (some laboratories round to the nearest even digit). Calculations - Multiplication and division: the result has as many significant figures as the factor with the fewest. - Addition and subtraction: the result has as many decimal places as the term with the fewest. - Round only at the end, keeping one or two extra digits in intermediate steps. - Exact numbers (counted objects, defined conversions like 1000 mL per L) have unlimited significant figures. Example: 4.56 x 1.4 = 6.384, reported as 6.4 (two significant figures). Example: 12.11 + 18.0 + 1.013 = 31.123, reported as 31.1 (one decimal place). Accuracy and precision - Accuracy: closeness to the true value. - Precision: closeness of repeated measurements to each other. - A scale that always reads 0.5 g too high is precise but not accurate. Exam trap: 100 mL measured with a graduated cylinder is not "one significant figure" if the cylinder reads to the millilitre. State it as 100. mL or 1.00 x 10^2 mL.

Matter: anything that has mass and occupies space. Physical properties are shown without changing the substance or making it react. Examples: colour, melting point, boiling point, density, electrical conductivity. Melting ice, H2O(s) -> H2O(l), is a physical change: it is still water. Chemical properties describe the ability to interact with other substances and change identity. Examples: flammability, corrosivity, reactivity with acids. Burning hydrogen, 2 H2 + O2 -> 2 H2O, is a chemical change: new substances form. Intensive versus extensive properties - Intensive: do not depend on the amount of sample. Examples: density, temperature, colour, physical state, melting point. - Extensive: depend on the amount. Examples: mass, volume, energy content, heat capacity of the whole sample. - The ratio of two extensive properties is intensive: density = mass / volume. Tips to decide - Halve the sample. If the property does not change, it is intensive. - Signs of a chemical change: colour change, gas evolved, precipitate, heat or light released, new smell. None is proof by itself (boiling produces bubbles yet is physical). Exam trap: dissolving sugar in water is physical (the sugar molecules stay intact and can be recovered), but dissolving a metal in acid is chemical. Clinical tip: density and boiling point are intensive properties used to identify substances. Body temperature and blood pH are intensive, while total body water and blood volume are extensive.

Solid - Rigid, with its own shape and volume. - Particles sit at fixed positions and only vibrate. - Incompressible: volume barely changes with pressure or temperature. Liquid - Own volume, but takes the shape of the container. - Particles are close together but can slide past each other. - Practically incompressible. Gas - No fixed shape or volume; fills the whole container. - Particles are far apart and move fast and randomly. - Compressible, low density, mixes completely with other gases, diffuses quickly. Energy and the state - In a gas the kinetic energy of the particles exceeds the attractive forces between them. - Removing energy (lowering temperature) lets the attractions win: gas -> liquid -> solid. - Adding heat reverses this: solid -> liquid (melting) -> gas (boiling or evaporation). - Direct changes: sublimation (solid -> gas) and deposition (gas -> solid). Three parameters define a gas: pressure P, volume V, temperature T (with the amount n in moles). Clinical tip: a gas that is compressible is dangerous in a cylinder. Oxygen and nitrous oxide cylinders are kept upright and secured, and stored liquid gases (liquid nitrogen for cryotherapy) expand hundreds of times when they vaporise. Exam trap: liquids are not perfectly incompressible, but for exam purposes treat liquid and solid volume as constant with pressure.

Energy is the capacity to do work. SI unit: joule (J). Older unit: calorie (1 cal = 4.186 J). Kinetic energy: energy of motion. E_kin = 1/2 m v^2 Potential energy: energy of position or of composition. - gravitational: E = m g h - electrostatic between two charges: E = q1 q2 / (4 pi eps0 r) - chemical: energy stored in the arrangement of atoms in bonds. Key ideas - Kinetic and potential energy can be converted into each other. - In an isolated system the total energy is constant (conservation of energy). - The electrostatic formula is the basis of chemical bonding: like charges (E > 0) repel, opposite charges (E < 0) attract, and stability means lower energy. - Reactions either release energy to the surroundings (exothermic) or absorb it (endothermic). Example: burning methane releases heat because the products (CO2, H2O) have stronger, lower-energy bonds than the reactants. Temperature is a measure of the average kinetic energy of the particles. Heat is energy transferred because of a temperature difference. Clinical tip: the body's ATP hydrolysis is an example of stored chemical (potential) energy converted into mechanical work in muscle and into heat that maintains body temperature. Exam trap: temperature and heat are not the same thing. A cup of boiling water and a bath of warm water can hold the same heat with very different temperatures.

Element (pure substance) - Made of one kind of atom only. - Cannot be decomposed into simpler substances by chemical methods. - Examples: Na, Cu, O2, N2, H2. Note that O2 and N2 are molecules of an element. Compound (pure substance) - Made of two or more different elements in a fixed, definite ratio. - Can be decomposed by chemical methods into its elements. - Has properties different from those of the elements that form it. - Examples: H2O, C6H12O6 (glucose), NaCl, FeS2 (pyrite: 46.55 % Fe and 53.45 % S by mass). Sucrose example: a white crystalline solid, whereas its elements are black carbon powder and the gases hydrogen and oxygen. Compound versus mixture (see the mixtures note) - In a compound the proportion is fixed (Proust) and the elements cannot be separated physically. - In a mixture the proportion is variable and components keep their properties. - Iron + sulfur powder mixed in any ratio is a mixture (a magnet removes the iron). Heated, they react to form iron sulfide, a compound with new properties. Exam trap: water is a compound; hydrogen and oxygen gases are elements. Air is not a compound but a mixture, even though it is uniform. Clinical tip: many medicines are pure compounds, and a purity specification (for example 99.5 % pure) is a claim about the fraction of that single compound in the sample.

Homogeneous mixture (solution) - One phase; the same composition and properties at every point. - Examples: air, sugar dissolved in water, saline solution (NaCl in water). - Components are dispersed at the molecular or ionic scale. - Composition can vary from sample to sample (saline 0.9 % versus 3 %). Heterogeneous mixture - Two or more phases with different properties. - Examples: milk, oil and water, sand and salt, granite. Air, as a homogeneous mixture, is about 78 % N2, 21 % O2, about 0.93 % argon and 0.04 % CO2, plus traces of neon, helium, methane, krypton, hydrogen and xenon. Solubility and saturation - Solubility is the maximum amount of solute that dissolves in a given amount of solvent at a given temperature. - A solution at that limit is saturated. Excess solute forms a separate phase (bottom body). Physical separation methods (the slides show them only as diagrams) - Filtration: separates an insoluble solid from a liquid. - Decantation: pours off a liquid from settled solid, or separates two immiscible liquids using a separating funnel. - Extraction: a solvent that dissolves one component preferentially; shake in a separating funnel, let layers separate, drain the denser one. - Distillation: separates liquids by different boiling points. The mixture is heated, vapour passes through a condenser cooled by water and is collected. - Evaporation and crystallisation: recover a dissolved solid. - Chromatography: separates components by their different affinity for a stationary and a mobile phase. - Centrifugation: separates by density under acceleration. Clinical tip: centrifugation of blood separates plasma, the buffy coat (white cells and platelets) and red cells, a physical separation of a heterogeneous mixture. Exam trap: a homogeneous mixture is still a mixture. Uniform appearance does not make it a compound.

1. Law of conservation of mass (Lavoisier) - The total mass of the substances is unchanged in a chemical reaction. - Atoms are rearranged, not created or destroyed. - Example: 2 H2 + O2 -> 2 H2O. 4.0 g of H2 and 32.0 g of O2 give 36.0 g of water. 2. Law of definite (constant) composition (Proust) - A given compound always contains the same elements in the same proportion by mass. - Example: water is always 11.2 % H and 88.8 % O by mass, whatever its source. - Pyrite, FeS2, is always 46.55 % Fe and 53.45 % S. 3. Law of multiple proportions (Dalton) - When two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in small whole-number ratios. - Example: carbon monoxide CO has 12 g C per 16 g O; carbon dioxide CO2 has 12 g C per 32 g O. The oxygen masses per 12 g of carbon are 16 and 32, a ratio 1 : 2. - Example: N2O, NO, NO2 have 28 g N per 16, 32 and 64 g O, ratio 1 : 2 : 4. Why these matter: the only simple explanation is that elements consist of atoms of fixed mass, which combine in whole numbers (Dalton's atomic theory). Worked example (Proust): if 5.00 g of a compound of X and Y contains 3.00 g of X, then a 20.0 g sample contains 12.0 g of X, because the ratio is fixed. Exam trap: the law of multiple proportions applies to different compounds of the same two elements. Do not apply it to one compound. Note: Dalton assumed atoms are indivisible. The next chapter shows that this part of his theory had to be revised.

1. Name the three levels at which chemistry describes a phenomenon and give one example of each for rusting iron. 2. Write the seven SI base units. Express the joule and the pascal in base units. 3. How many significant figures are in 0.004050, 30,045 and 1.20 x 10^3? 4. Calculate 12.11 + 18.0 + 1.013 and 4.56 x 1.4, reporting the correct number of significant figures. 5. Classify each as physical or chemical property: melting point, flammability, density, reactivity with acid. 6. Classify as intensive or extensive: mass, temperature, volume, density. 7. Compare solid, liquid and gas in terms of shape, volume and compressibility. 8. What is the difference between an element, a compound and a mixture? Classify air, water, oxygen gas, saline, and sucrose. 9. Give two examples each of homogeneous and heterogeneous mixtures. Which separation method would you use for (a) water and oil, (b) salt dissolved in water, (c) ethanol and water? 10. Which law is illustrated by each: (a) 10 g of reactants give 10 g of products; (b) water is always 88.8 % oxygen by mass; (c) CO and CO2 contain oxygen in a 1 : 2 ratio for the same mass of carbon? 11. Convert 72 km/h to m/s and 250 mL to litres. 12. Why does a compound have properties different from its elements, while a mixture keeps the properties of its components?

1. Macroscopic: a red-brown flaky coating forms on the iron. Microscopic: iron atoms lose electrons to oxygen, forming Fe3+ ions. Symbolic: 4 Fe + 3 O2 -> 2 Fe2O3. 2. m, kg, s, A, K, mol, cd. Joule = kg m^2 s^-2. Pascal = kg m^-1 s^-2. 3. 0.004050 has four (leading zeros do not count, the trailing zero after the decimal point does). 30,045 has five. 1.20 x 10^3 has three. 4. 12.11 + 18.0 + 1.013 = 31.123, reported 31.1 (one decimal place). 4.56 x 1.4 = 6.384, reported 6.4 (two significant figures). 5. Physical: melting point, density. Chemical: flammability, reactivity with acid. 6. Extensive: mass, volume. Intensive: temperature, density. 7. Solid: own shape and volume, incompressible. Liquid: own volume, takes the container's shape, nearly incompressible. Gas: neither shape nor volume, compressible. 8. Element: one kind of atom. Compound: two or more elements in fixed ratio, separable only chemically. Mixture: variable proportions, separable physically. Air: homogeneous mixture. Water: compound. Oxygen gas: element. Saline: homogeneous mixture. Sucrose: compound. 9. Homogeneous: air, sugar solution. Heterogeneous: milk, oil and water. (a) decantation with a separating funnel; (b) evaporation or distillation; (c) fractional distillation. 10. (a) Lavoisier, conservation of mass. (b) Proust, definite composition. (c) Dalton, multiple proportions. 11. 72 km/h = 20 m/s (divide by 3.6). 250 mL = 0.250 L. 12. In a compound the atoms are chemically bonded, forming new arrangements with new properties. In a mixture no bonds form between components, so each keeps its own properties.

Dalton's atom (1808): a tiny, indivisible sphere. Several experiments overturned the indivisible part. Thomson (1897): cathode rays - A high voltage across two electrodes in an evacuated glass tube produces a beam (cathode rays) from the negative to the positive electrode. - The beam is deflected by electric and magnetic fields towards the positive plate, so it carries negative charge. - The beam is the same for any cathode metal, so the particles (electrons) are universal parts of all matter. - He measured charge/mass and found the electron to be about 1/1000 the mass of a hydrogen atom (really 1/1837). Millikan (1906): the oil-drop experiment - Measured the charge of the electron: e = 1.602 x 10^-19 C. - Showed that every charge is a whole multiple of this elementary charge (charge is quantised). Marie Curie (1898) - Isolated polonium and radium, which emit radiation spontaneously. - Radioactive atoms break down as they emit, so atoms cannot be indivisible and elements can change identity in nuclear (not chemical) processes. Thomson's model (1903): the plum-pudding atom. A uniform sphere of positive charge, the size of the atom, with electrons embedded like raisins, giving a neutral atom. The three fundamental particles - proton: charge +1.602 x 10^-19 C (+1 e), mass 1.673 x 10^-27 kg - neutron: charge 0, mass 1.675 x 10^-27 kg - electron: charge -1.602 x 10^-19 C (-1 e), mass 9.109 x 10^-31 kg Key comparison: a proton is about 1837 times heavier than an electron. Essentially all of an atom's mass is in protons and neutrons. Exam trap: the elementary charge e is the same size for protons and electrons, opposite in sign. The neutron is not "half a proton": it has similar mass and no charge. Clinical tip: cathode ray tubes are the ancestor of the X-ray tube. Electrons accelerated across a high voltage and stopped in a metal target produce X-rays.

Rutherford's experiment (1911, carried out by Geiger and Marsden) - A beam of alpha particles (He2+) hits a very thin gold foil, and a fluorescent screen records where they land. - Most particles pass straight through. - A few are deflected at large angles, and a very few bounce almost straight back. Conclusion (the nuclear atom) - The positive charge and almost all of the mass are concentrated in a very small central nucleus. - The electrons occupy the rest of the volume, far from the nucleus. - The plum-pudding model could not produce large deflections, so it was discarded. Sizes - atom radius: about 10^-10 m (1 angstrom) - nucleus radius: about 10^-14 to 10^-15 m - volume ratio: (10^-14 / 10^-10)^3 = 10^-12, so the nucleus is one trillionth of the atom's volume yet holds over 99.9 % of its mass. What holds the nucleus together - Protons repel each other electrostatically. - The strong nuclear force acts only over very short distances (about 1 fm) and attracts protons and neutrons (nucleons) to each other, overcoming the repulsion. - Neutrons add attraction without adding repulsion. Light stable nuclei need about as many neutrons as protons; heavier ones need more. Problem with the Rutherford model: classical physics says an orbiting electron should radiate energy and spiral into the nucleus. Atoms do not collapse. Quantum mechanics (next chapter) resolves this. Clinical tip: alpha, beta and gamma radiation used in nuclear medicine and radiotherapy come from unstable nuclei. Understanding nuclear structure is the first step of radiobiology. Exam trap: "the atom is mostly empty" refers to the space around the nucleus. It is not truly empty: it holds the electron cloud.

Definitions - Atomic number Z: number of protons in the nucleus. It is fixed for an element and determines its chemical identity. - Mass number A: total number of nucleons, A = Z + N, so N = A - Z. - Element: species defined by a specific Z. - Nuclide: species defined by a specific Z and A. - Notation: (A over Z) X, for example carbon-12 is 12 C with Z = 6. - In a neutral atom the number of electrons equals Z. Example: oxygen-18 has Z = 8, A = 18, so 8 protons, 10 neutrons, 8 electrons. Example: uranium-235 has 92 protons and 143 neutrons; uranium-238 has 92 protons and 146 neutrons. Isotopes: atoms with the same Z and different A. - Same chemical properties (same electron structure). - Different physical properties (mass, nuclear stability, rates of diffusion). - Some elements have a single natural isotope (Na-23, F-19, Al-27). Others have several. Natural abundances - Hydrogen: protium 99.985 %, deuterium (D, heavy hydrogen) 0.015 %, tritium (T, radioactive) trace. - Carbon: C-12 98.89 %, C-13 1.10 %, C-14 trace (radioactive). - Oxygen: O-16 99.762 %, O-17 0.038 %, O-18 0.200 %. - Chlorine: Cl-35 about 75 %, Cl-37 about 25 %. - Bromine: Br-79 about 50 %, Br-81 about 50 %. Ions - Atoms can lose or gain electrons to reach a stable configuration; the number of protons never changes. - Cation: positive ion, formed by losing electrons (Na+, Ca2+, H+). - Anion: negative ion, formed by gaining electrons (Cl-, S2-). - Example: Fe3+ has 26 protons and 23 electrons. Modern atomic theory (summary) - Atoms are divisible: a small nucleus of protons and neutrons, surrounded by electrons that fill most of the volume. - Chemical reactions do not change one element into another (a nuclear reaction is needed). - Atoms of one element share the same Z; the electrons determine chemical behaviour. - Compounds form when elements combine in fixed ratios. Clinical tip: radioisotopes of the same element behave identically in the body. Iodine-123 and iodine-131 are taken up by the thyroid exactly like stable iodine-127, which is why they are used for imaging and therapy. Carbon-14 and fluorine-18 (in FDG for PET) are also used as tracers, and carbon-13 for breath tests. Exam trap: for ions, the protons stay the same and only the electrons change. Neutrons are irrelevant to charge.

The atomic mass unit (amu, u, or dalton) - 1 u is defined as exactly 1/12 of the mass of one carbon-12 atom. - 1 u = 1.6605 x 10^-24 g = 1.6605 x 10^-27 kg. - Carbon-12 is assigned a mass of exactly 12 u. Relative atomic masses are pure numbers that tell how many times heavier an atom is than 1/12 of a carbon-12 atom. - Example: the atomic mass of carbon-13 is 13.003354. Average atomic mass A_r = sum of (fractional abundance x isotopic mass), over all natural isotopes. Worked example: boron - B-10: 10.0129 u, 19.91 % (0.1991) - B-11: 11.0093 u, 80.09 % (0.8009) - A_r = 10.0129 x 0.1991 + 11.0093 x 0.8009 = 1.993 + 8.817 = 10.81 u. Worked example: chlorine with 75 % Cl-35 (34.97 u) and 25 % Cl-37 (36.97 u) gives about 35.45 u. Mass defect (the slide's note, in outline) - The real mass of an atom is slightly less than the sum of the masses of its protons, neutrons and electrons. - The lost mass turns into binding energy according to E = mc2. - Nuclear binding energies are about a million times larger per mole than chemical bond energies, which is why nuclear reactions release far more energy than combustion. Why the atomic mass is not a whole number: (1) mass defect, and (2) averaging over isotopes. Exam trap: the mass number A is a whole number for one nuclide. The atomic mass in the table is an average, a decimal. Clinical tip: the molar masses used to prepare a drug solution are averages of natural isotopes, so they are correct for any real sample.

How it works (in outline) 1. The sample is vaporised. 2. It is ionised, typically by a stream of electrons that removes an electron, giving positive ions. 3. The ions are accelerated by a high-voltage field into a beam. 4. The beam passes through a magnetic field, which bends ions along a curved path. Lighter ions bend more than heavier ones of the same charge. 5. A detector records the ions that arrive, giving a spectrum of intensity versus mass-to-charge ratio (m/z). What the spectrum tells you - The position of each peak is the isotopic mass. - The height of each peak is the relative abundance. - From these, the average atomic mass is computed as above. - For molecules, the peak of the intact ion gives the molecular mass, and fragment peaks give structural clues. Example: chlorine gives two atomic peaks at 35 and 37 with heights in a ratio of about 3 : 1, showing the 75 % / 25 % abundances. Dichlorine gives three peaks at 70, 72, 74. Clinical tip: mass spectrometry is widely used in the clinical laboratory: newborn screening for metabolic disorders, therapeutic drug monitoring, toxicology, hormone assays and identification of bacteria (MALDI-TOF) all rely on it. Exam trap: mass spectrometry separates by mass-to-charge, not mass alone. A doubly charged ion of mass 70 appears at m/z 35.

Molecular (formula) mass - Sum of the atomic masses of all atoms in the formula. - H2O: 2 x 1.0079 + 15.9994 = 18.0152 u. - Na2CO3: 2 x 22.99 + 12.01 + 3 x 16.00 = 105.99 u. The mole - One mole is the amount of substance containing as many particles as there are atoms in 12 g of carbon-12. - That number is Avogadro's number, N_A = 6.022 x 10^23 particles per mole. - The particles can be atoms, molecules, ions, electrons or photons; always specify which. Molar mass M - The mass in grams of one mole, in g/mol. - Numerically equal to the atomic or molecular mass in u. - Examples: C 12.011 g/mol; H2 2.016 g/mol; H2O 18.016 g/mol; CO2 44.01 g/mol; NaCl 58.44 g/mol; Pb 207.2 g/mol. - Older terms: gram-atom (mole of atoms) and gram-molecule (mole of molecules). Be exact about the particle - 1.008 g of hydrogen = 1 mol of H atoms. - 2.016 g of hydrogen = 1 mol of H2 molecules = 2 mol of H atoms. Conversions - n (mol) = m (g) / M (g/mol) - number of particles = n x N_A - Reaction coefficients count moles or molecules, not grams: 2 H2 + O2 -> 2 H2O. Worked example: how many molecules are in 9.008 g of water? n = 9.008 / 18.016 = 0.500 mol, so 0.500 x 6.022 x 10^23 = 3.01 x 10^23 molecules. Worked example: how many grams in 0.25 mol of NaCl? 0.25 x 58.44 = 14.6 g. Clinical tip: normal saline is 0.9 % NaCl, 9.0 g in a litre, or 9.0 / 58.44 = 0.154 mol/L, about 154 mmol/L each of Na+ and Cl-. Being able to switch between grams and moles is what lets you read laboratory reference values. Exam trap: 1 mol of H2 is 2 g, not 1 g. Always ask whether the slide or question means atoms or molecules.

- Empirical formula: the elements present and the smallest whole-number ratio between their atoms. Glucose: CH2O. - Molecular formula: the actual number of each atom in a molecule. Glucose: C6H12O6. It is a whole-number multiple of the empirical formula. - Structural formula: shows how the atoms are connected (Lewis structure), for example H-O-H. - Steric (spatial) formula: shows how atoms are arranged in space, using VSEPR theory, for example water is bent, methane tetrahedral. Examples - Benzene: empirical CH, molecular C6H6. - Hydrogen peroxide: empirical HO, molecular H2O2. - Ethene C2H4 and cyclopropane C3H6 share the empirical formula CH2. Ionic compounds have only an empirical (formula-unit) formula: NaCl does not exist as separate molecules. Isomers share a molecular formula but differ in structure: ethanol and dimethyl ether are both C2H6O. Exam trap: the empirical formula cannot tell isomers apart. A structural formula is required. Clinical tip: drug names and structures matter. Molecular formulas identify the compound, but only the structural and steric formula explain activity, because receptors recognise three-dimensional shape.

Method for the empirical formula 1. Assume 100 g of compound if percentages are given, so that percentages become grams. 2. Convert each mass to moles: n = m / M. 3. Divide every mole value by the smallest one. 4. If the ratios are not whole numbers, multiply all by a small integer: 0.5 -> x2; 0.33 or 0.67 -> x3; 0.25 or 0.75 -> x4. 5. Write the formula from the whole-number ratios. Method for the molecular formula 1. Compute the empirical formula mass. 2. Divide the given molar (formula) mass by it: k = M / M_empirical. 3. Multiply every subscript in the empirical formula by k. Worked example 1: 63.52 % Fe and 36.48 % S Fe: 63.52 / 55.85 = 1.137 mol. S: 36.48 / 32.07 = 1.138 mol. Ratio 1 : 1. Empirical formula: FeS. Worked example 2: 26.56 % K, 35.41 % Cr, rest O Oxygen = 100 - 26.56 - 35.41 = 38.03 %. K: 26.56 / 39.10 = 0.679. Cr: 35.41 / 52.00 = 0.681. O: 38.03 / 16.00 = 2.377. Divide by 0.679: K 1, Cr 1, O 3.5. Multiply by 2: K2Cr2O7. This is potassium dichromate. Worked example 3: tetraethyllead, 60.00 g sample with 38.43 g Pb, 17.83 g C and 3.74 g H Pb: 38.43 / 207.2 = 0.1855. C: 17.83 / 12.01 = 1.485. H: 3.74 / 1.008 = 3.71. Divide by 0.1855: Pb 1, C 8.0, H 20.0. Empirical formula: PbC8H20, that is Pb(C2H5)4. Worked example 3b: 170.00 g sample with 29.84 g Na, 67.49 g Cr and 72.67 g O Na: 29.84 / 22.99 = 1.298 mol. Cr: 67.49 / 52.00 = 1.298 mol. O: 72.67 / 16.00 = 4.542 mol. Divide by 1.298: 1 : 1 : 3.5. Multiply by 2: Na2Cr2O7 (sodium dichromate). Worked example 4: molecular formula from mass Empirical formula CH (mass 13.02 u), formula mass 78.11 u. k = 78.11 / 13.02 = 6. Molecular formula: C6H6 (benzene). Empirical formula NH2 (16.02 u), formula mass 32.06 u: k = 2, molecular formula N2H4 (hydrazine). Worked example 5: 0.44 g H and 6.92 g O, formula mass 34.00 u H: 0.44 / 1.008 = 0.437 mol. O: 6.92 / 16.00 = 0.4325 mol. Ratio 1 : 1, empirical HO (17.01 u). k = 34.00 / 17.01 = 2, so H2O2. Exam trap: use the right atomic masses and do not round the ratios too early. A ratio like 1.98 is 2, and 1.49 or 1.51 is 1.5. But 1.2 is not 1: try multiplying by 5. Clinical tip: the same calculation, run in reverse, gives percentage composition and is used to check drug purity by elemental analysis.

1. What did Thomson's cathode-ray experiment show, and what did Millikan's add? 2. Describe Rutherford's experiment and what each of its observations means. 3. Give the charge and relative mass of a proton, a neutron and an electron. What is the ratio of proton mass to electron mass? 4. Estimate the fraction of an atom's volume occupied by its nucleus. 5. How many protons, neutrons and electrons are in Fe-56, in Fe3+, and in Cl-37 (as Cl-)? (Z: Fe 26, Cl 17) 6. Define isotope. Why do isotopes of one element have the same chemistry? 7. Naturally occurring copper is 69.2 % Cu-63 (62.93 u) and 30.8 % Cu-65 (64.93 u). Calculate the average atomic mass. 8. What does a mass spectrometer measure? 9. How many moles and molecules are in 36.0 g of water? 10. What is the mass of 0.500 mol of CO2? How many atoms of oxygen does it contain? 11. Explain the difference between empirical and molecular formulas, with the example of glucose. 12. A compound is 40.0 % C, 6.7 % H, 53.3 % O and has molar mass 180 g/mol. Find the empirical and molecular formulas. 13. Why is 1 mole of H2 gas 2.016 g and not 1.008 g?

1. Cathode rays are negatively charged particles common to all matter (electrons), about 1/1837 of a hydrogen atom's mass. Millikan measured the electron charge, 1.602 x 10^-19 C, and showed that charge is quantised. 2. Alpha particles on gold foil. Most passed through: the atom is mostly empty space. A few were deflected strongly or reflected: the positive charge and mass are concentrated in a tiny dense nucleus. 3. Proton: +1 e, 1.673 x 10^-27 kg. Neutron: 0, 1.675 x 10^-27 kg. Electron: -1 e, 9.109 x 10^-31 kg. Proton/electron mass is about 1837. 4. (10^-14 / 10^-10)^3 = 10^-12. 5. Fe-56: 26 protons, 30 neutrons, 26 electrons. Fe3+: 26 protons, 30 neutrons, 23 electrons. Cl-37 anion: 17 protons, 20 neutrons, 18 electrons. 6. Isotopes: same Z, different A. The chemistry depends on the electrons, which are the same. 7. 0.692 x 62.93 + 0.308 x 64.93 = 43.55 + 20.00 = 63.55 u. 8. The mass-to-charge ratio of ions, and hence isotopic masses and relative abundances. 9. n = 36.0 / 18.016 = 2.00 mol; molecules = 2.00 x 6.022 x 10^23 = 1.20 x 10^24. 10. M = 44.01 g/mol, so 22.0 g. Molecules = 0.500 x 6.022 x 10^23 = 3.01 x 10^23, each with 2 oxygen atoms: 6.02 x 10^23 oxygen atoms. 11. Empirical: simplest whole-number ratio, CH2O. Molecular: actual atom count, C6H12O6. 12. Moles per 100 g: C 3.33, H 6.65, O 3.33. Ratio 1 : 2 : 1, empirical CH2O (30.03). k = 180 / 30.03 = 6, so C6H12O6. 13. The mole of molecules of H2 contains 2 mol of H atoms, 2 x 1.008 = 2.016 g. 1.008 g is a mole of H atoms.

Maxwell (1865): electromagnetic radiation is a set of electric and magnetic fields oscillating perpendicular to each other and to the direction of travel. In vacuum all of it travels at the speed of light: c = lambda x nu = 2.998 x 10^8 m/s Properties of a wave - Wavelength (lambda): distance between two successive crests, in metres. - Frequency (nu): number of cycles per second, in hertz (s^-1). - Amplitude: height of the wave; its square gives the intensity (brightness). - Nodes: points of zero amplitude. - Speed: c = lambda x nu, so wavelength and frequency are inversely related. The electromagnetic spectrum (increasing frequency and energy, decreasing wavelength) radio waves -> microwaves -> infrared -> visible (about 400 to 700 nm, violet to red) -> ultraviolet -> X-rays -> gamma rays. Energy carried: E = h x nu = h x c / lambda, with Planck's constant h = 6.626 x 10^-34 J s. High frequency means high energy. Ultraviolet carries more energy than visible light, and visible more than infrared, because its wavelength is shorter. Worked example: yellow light at 589 nm. nu = 2.998 x 10^8 / 589 x 10^-9 = 5.09 x 10^14 Hz. E = 6.626 x 10^-34 x 5.09 x 10^14 = 3.37 x 10^-19 J per photon. Wave behaviour in classical physics - Refraction: a wave changes direction when passing from one medium to another (air to water). - Diffraction: a wave passing through a slit spreads out. A beam of classical particles would not. - These were the tests that separated waves from particles, until 1900. Clinical tip: the danger of radiation follows the spectrum. Radio waves and microwaves heat tissue, infrared warms, ultraviolet damages skin DNA, and X-rays and gamma rays are ionising. MRI uses radiofrequency, imaging uses X-rays and gamma photons, and phototherapy for neonatal jaundice uses blue light of about 450 nm. Exam trap: light does not slow in vacuum with colour; all colours share c in vacuum. What differs is wavelength and frequency.

Bohr's three postulates (1913) 1. The electron can occupy only certain allowed circular orbits (stationary states), each with a fixed energy. 2. While in a stationary state the electron does not radiate energy. 3. The electron changes state only by absorbing or emitting a photon whose energy equals the energy difference: delta E = E2 - E1 = h x nu. Energy levels of hydrogen E_n = -2.18 x 10^-18 J / n^2, with n = 1, 2, 3, ... (about -13.6 eV / n^2). - n = 1 is the ground state, the lowest energy. Energies are negative because the electron is bound; E = 0 corresponds to a free electron (n = infinity). - Levels get closer together as n increases. - Orbit radius grows as n^2 (r_1 = 0.53 angstrom, the Bohr radius). Emission and absorption - Absorption: an incoming photon of exactly the right energy lifts the electron to a higher level (excited state). - Emission: the electron falls to a lower level and releases a photon. Worked example: the 656.3 nm red line (n = 3 -> 2). delta E = 2.18 x 10^-18 (1/4 - 1/9) = 2.18 x 10^-18 x 0.1389 = 3.03 x 10^-19 J. lambda = h c / delta E = 6.626 x 10^-34 x 3.00 x 10^8 / 3.03 x 10^-19 = 6.56 x 10^-7 m = 656 nm. It matches. Ionisation energy of hydrogen: from n = 1 to infinity, 2.18 x 10^-18 J per atom, about 1312 kJ/mol. Limits of the Bohr model - It works only for one-electron systems (H, He+, Li2+). - It treats the electron as a particle with a precise orbit, ignoring its wave nature. - It cannot explain the spectra of multi-electron atoms. Clinical tip: fluorescence and lasers rest on the same emission process, and fluorescent labels (used in flow cytometry and immunofluorescence) absorb at one wavelength and emit at a longer one because part of the energy is lost as heat in between. Exam trap: the energies are negative. A larger n means a less negative, higher energy.

Schroedinger (1926) wrote a wave equation for the electron in an atom: (h^2 / 8 pi^2 m) x (d2 psi/dx2 + d2 psi/dy2 + d2 psi/dz2) + (E - V) psi = 0 where psi is the wave function, m the electron mass, E the total energy, V the potential energy (electrostatic attraction to the nucleus). Key points - The equation has many solutions. Only certain wave functions are physically acceptable (like standing waves on a guitar string). Each has a well-defined energy, so energy is quantised as a natural result of the mathematics, not as an added assumption. - Each solution is an atomic orbital: a wave function psi(x, y, z) with a definite energy, shape and orientation. - psi itself has no direct physical meaning. psi squared is the probability density: the probability of finding the electron at a given point. - The electron cloud is a picture of psi squared. Usually the orbital is drawn as the region enclosing about 90 % of the probability. - The lowest-energy solution is the ground state; the others are excited states (infinite in number). Orbit versus orbital - Bohr's orbit: a defined trajectory of a particle around the nucleus. Wrong. - Orbital: a mathematical function; the region of space where the electron is likely to be found. No path is implied. Radial probability distribution: the probability of finding the electron at a certain distance r from the nucleus, summed over all directions. For 1s it peaks at the Bohr radius (0.53 angstrom). Clinical tip: the orbital picture explains why molecules have shapes and why drugs bind to receptors: bonds form by overlap of orbitals. Exam trap: an orbital is not a place the electron travels along. It is a probability distribution, and up to two electrons can share it.

1. Why would an orbiting electron collapse into the nucleus in classical physics? What resolves this? 2. Calculate the frequency and energy of a photon of wavelength 500 nm. (c = 3.00 x 10^8 m/s, h = 6.626 x 10^-34 J s) 3. Rank X-rays, visible light, microwaves and ultraviolet by increasing energy. 4. What is the photoelectric effect, and why can't a classical wave explain it? 5. What did Planck assume to explain black-body radiation? 6. Why do atomic spectra consist of lines? 7. State Bohr's three postulates and give two limits of his model. 8. Calculate the de Broglie wavelength of an electron moving at 6 x 10^6 m/s (m = 9.11 x 10^-31 kg). 9. State the uncertainty principle and explain why an electron has no definite orbit. 10. What is the difference between an orbit and an orbital? What does psi squared mean? 11. List the allowed values of l and m_l for n = 3. How many orbitals are in the third shell? 12. Which of these are not allowed, and why: 2d, 3p, 1p, 4f? 13. Write the electron configuration of N, Fe, Fe3+, Cl- and Cu. How many unpaired electrons does nitrogen have? 14. State the Pauli principle, Hund's rule and the Aufbau principle. 15. Why is the 4s orbital filled before 3d, and which electrons does Fe lose to become Fe2+?

1. An accelerating charge radiates energy, so the electron would lose energy and spiral inwards. Quantum mechanics: the electron is a standing wave with quantised energies, and the lowest state cannot go lower; the uncertainty principle also prevents confinement to the nucleus. 2. nu = c / lambda = 3.00 x 10^8 / 500 x 10^-9 = 6.00 x 10^14 Hz. E = h nu = 6.626 x 10^-34 x 6.00 x 10^14 = 3.98 x 10^-19 J. 3. Microwaves < visible < ultraviolet < X-rays. 4. Metals emit electrons only when the light frequency exceeds a threshold, whatever the intensity. A classical wave's energy depends on amplitude, so intensity should matter; Einstein's photons of energy h nu explain the threshold. 5. That energy is exchanged in discrete quanta E = h nu. 6. Atomic energy levels are quantised, so only photons whose energy equals the difference between two levels are emitted or absorbed. 7. (1) Only certain orbits with fixed energy are allowed; (2) no radiation in a stationary state; (3) a jump between states involves a photon of energy h nu = E2 - E1. Limits: only for one-electron species; treats the electron as a particle with a defined orbit, ignoring wave nature. 8. lambda = 6.626 x 10^-34 / (9.11 x 10^-31 x 6 x 10^6) = 1.2 x 10^-10 m. 9. delta x times delta p is at least h / (4 pi). Position and momentum cannot both be known exactly, so a definite trajectory does not exist; only probabilities do. 10. Orbit: a defined path (Bohr). Orbital: a wave function; psi squared is the probability density of finding the electron at a point. 11. l = 0, 1, 2. For l = 0: m = 0; l = 1: -1, 0, +1; l = 2: -2, -1, 0, +1, +2. Total orbitals = 3^2 = 9. 12. 2d (l must be less than n) and 1p are not allowed. 3p and 4f are allowed. 13. N: 1s2 2s2 2p3 (three unpaired). Fe: [Ar] 3d6 4s2. Fe3+: [Ar] 3d5. Cl-: [Ne] 3s2 3p6 = [Ar]. Cu: [Ar] 3d10 4s1. 14. Pauli: at most two electrons per orbital, with opposite spins. Hund: degenerate orbitals are filled singly with parallel spins before pairing. Aufbau: fill from lowest energy upwards. 15. s orbitals penetrate closer to the nucleus, so 4s is lower in energy than 3d in neutral atoms. Fe becomes Fe2+ by losing the two 4s electrons, giving [Ar] 3d6.

Before the table - Early attempts grouped elements by affinity or by similar properties (for example the eighteenth-century tables of affinity). Mendeleev (1869) - Noticed that properties repeat periodically when elements are placed in order of increasing atomic mass. - Arranged similar elements in the same column. - Left gaps for elements not yet discovered and predicted their properties ("eka-aluminium" was later found as gallium, "eka-silicon" as germanium). - Some pairs were out of order by mass (for example tellurium and iodine). Moseley (1913) - Using X-ray spectra, showed that the fundamental quantity is the atomic number Z (proton count). - Ordered the table by Z, fixing the inconsistencies. - Predicted further missing elements. - Modern statement: the properties of the elements are periodic functions of atomic number. The reason for periodicity: it is the periodic repetition of valence electron configurations as the shells fill. Modern layout - 7 horizontal rows, the periods. Elements of one period have the same highest n and progressively more electrons in that shell. - 18 vertical columns, the groups. Elements of one group have similar outer configurations and therefore similar chemistry. Old notation: 1A to 8A for main groups and 1B to 8B for transition metals; modern numbering 1 to 18. - Groups with names: 1 alkali metals (Li, Na, K); 2 alkaline-earth metals (Mg, Ca); 17 halogens (F, Cl, Br, I); 18 noble gases (He, Ne, Ar). The blocks - s block: groups 1 and 2 (filling s orbitals), plus He. - p block: groups 13 to 18 (filling p). - d block: groups 3 to 12, the transition metals. - f block: lanthanides and actinides (filling 4f and 5f), shown at the bottom. - Period length: 2, 8, 8, 18, 18, 32, 32 elements, matching the capacity of the orbitals that fill in that period. Number of valence electrons of a main-group element equals the old group number (1A to 8A). Example: sulfur is in group 16 (6A), with configuration 3s2 3p4, six valence electrons. Clinical tip: many biologically important elements sit in a few columns. Na and K (group 1), Mg and Ca (group 2), and the halogens F, Cl, I. Elements in a group with similar chemistry can substitute for one another, which is why lithium (group 1) affects sodium-related processes, and why strontium and radium accumulate in bone like calcium. Exam trap: hydrogen sits in group 1 by configuration (1s1) but is not an alkali metal.

Metals (left and centre of the table, about 80 % of elements) - Shiny solids (mercury is a liquid). - High melting points. - Good conductors of heat and electricity. - Ductile (drawn into wires) and malleable (beaten into sheets). - Tend to lose electrons and form positive ions (cations). - Their oxides are basic. Non-metals (upper right) - Solid, liquid or gas (C, Br2, O2). - Low melting points (except network solids like diamond). - Not shiny, poor conductors of heat and electricity (insulators; graphite is an exception). - Brittle, neither ductile nor malleable. - Tend to gain electrons and form negative ions (anions), or share electrons. - Their oxides are acidic. Metalloids (along the diagonal: B, Si, Ge, As, Sb, Te) - Properties intermediate between metals and non-metals. - Semiconductors (silicon, germanium), important for electronics and medical devices. Trend: metallic character decreases from left to right across a period and increases from top to bottom in a group. The most metallic element is at the bottom left (caesium, francium), and the most non-metallic at the top right (fluorine, excluding noble gases). Acid-base behaviour of oxides - Metal oxide + water gives a base: Na2O + H2O -> 2 NaOH. - Non-metal oxide + water gives an acid: SO3 + H2O -> H2SO4; CO2 + H2O -> H2CO3. - Oxides of elements in between (Al2O3) can be both (amphoteric). Clinical tip: the diagonal semiconductors are the basis of the detectors in digital X-ray and CT, and gadolinium, platinum (in cisplatin) and gold are metals whose chemistry is exploited in imaging and therapy. Exam trap: "metal" is a chemical property (tendency to lose electrons), not a description of hardness. Sodium is a metal that is soft enough to cut with a knife.

Definition: half the distance between the nuclei of two identical bonded atoms (covalent radius) or of neighbouring atoms in a metal (metallic radius). - Smallest: helium (0.32 angstrom in the slide's data). Largest: caesium (2.65 angstrom). Across a period (left to right): decreases. - The number of protons increases while electrons are added to the same shell. - Electrons in the same shell shield each other poorly, so the effective nuclear charge Z_eff increases and pulls the shell inward. - Example (angstrom): Li 1.52, Be 1.13, B 0.88, C 0.77, N 0.70, O 0.66, F 0.64. Down a group: increases. - Each period adds a new shell, farther from the nucleus. - Inner electrons shield the outer ones, so Z_eff hardly changes. - Example: Li 1.52, Na 1.86, K 2.27, Rb 2.47, Cs 2.65 angstrom. Ionic radius - Cations are smaller than their parent atoms. Losing electrons (often an entire shell) reduces electron-electron repulsion, and the remaining electrons feel more Z_eff. Example: Na 1.86 angstrom (186 pm), Na+ 1.02 angstrom (102 pm). - Anions are larger than their parent atoms. Extra electrons increase repulsion and the cloud swells. Example: Cl 99 pm, Cl- 181 pm. - Among isoelectronic ions (same number of electrons), the one with the higher nuclear charge is smaller: O2- > F- > Na+ > Mg2+ > Al3+. - For a cation with several charges, the higher charge gives the smaller radius: Fe2+ is larger than Fe3+. Clinical tip: the similar sizes of K+ and Rb+, or of Ca2+ and Sr2+, explain why body pumps and bone can confuse them, while the difference in size between Na+ (102 pm) and K+ (138 pm) is what lets ion channels distinguish them. Exam trap: a positive ion is smaller and a negative ion is larger than its atom. Do not reverse this.

Definition: the energy required to remove one electron from the outermost shell of an isolated, neutral gaseous atom in its ground state, taking it to infinite distance. X(g) + E_i -> X+(g) + e-, with E_i > 0 (always endothermic). Units: kJ/mol or eV. Successive ionisation energies - E_i1 is the first, E_i2 the second, and so on: E_i1 < E_i2 < E_i3, because removing an electron from a positive ion is harder. - Beryllium: E_i1 = 900, E_i2 = 1760, E_i3 = 14850 kJ/mol. The huge jump at the third shows that the third electron comes from the inner 1s shell. This proves that Be has two valence electrons and forms Be2+. Trends - Across a period: E_i increases (Z_eff rises, radius falls, electrons held more tightly). Alkali metals have the lowest, noble gases the highest. - Down a group: E_i decreases (larger radius, more shielding, outer electrons easier to remove). - Small irregularities: E_i drops from Be to B (the 2p electron is higher and better shielded than 2s), and from N to O (the fourth 2p electron is paired, and repulsion makes it easier to remove). - Example, first E_i (kJ/mol): Li 520, Na 496, K 419; He 2372, Ne 2081, Ar 1521. Meaning: low E_i means strong metallic character (the atom easily forms a cation). Clinical tip: the very low ionisation energy of sodium and potassium is why they exist in the body only as ions, Na+ and K+, and why the pure metals react violently with water. Exam trap: the ionisation energy is required, not released. It is always positive.

Definition: the energy change when a gaseous neutral atom in its ground state acquires an electron. X(g) + e- -> X-(g) + energy released, delta E = A_e. By the sign convention used in the slides, A_e < 0 means energy is released (stable anion) and A_e > 0 means energy must be supplied (unstable anion). Examples (kJ/mol) - Cl -349, F -328, Br -325, I -295, O -141 (first electron), H -73, Li -60, Na -53. - Noble gases and group 2 elements (Be, Mg): positive, they do not form stable anions, because the added electron goes into a new, higher subshell (or a full shell). - Nitrogen is close to zero, since the added electron would have to pair in a half-filled 2p subshell. - Fluorine is slightly less negative than chlorine, because the small 2p cloud makes the added electron repel the existing ones strongly. Trends - Across a period: electron affinity becomes more negative (higher Z_eff and smaller radius). The halogens gain electrons most readily. - Down a group: generally less negative, because the added electron is farther from the nucleus and shielded. Why it matters - Only elements with a large negative electron affinity and high ionisation energy (the non-metals) tend to form anions. - Together, E_i and A_e explain the type of bond an atom will form: low E_i with a high A_e partner gives an ionic bond. Second electron affinity: adding a second electron to an anion (O- to O2-) always requires energy, since the anion repels the electron. O2- is stable in solids only because the lattice energy pays for it. Exam trap: do not confuse the sign conventions. Some textbooks quote the electron affinity as a positive number for released energy. Check the convention. Clinical tip: halogens with high electron affinity are strong oxidants. Hypochlorite (bleach) and iodine act as disinfectants because they take electrons from bacterial components.

Two factors control every trend 1. Increased nuclear charge increases the attraction between the nucleus and electrons. 2. Repulsion between electrons (shielding by core electrons) reduces the effective attraction. Along a period (left to right) - Z_eff increases; valence electrons are in the same shell and screen each other almost not at all. - Atomic radius decreases. - Ionisation energy increases. - Electron affinity becomes more negative. - Metallic character decreases. Down a group (top to bottom) - New shells are added, and core electrons shield strongly. - Atomic radius increases. - Ionisation energy decreases. - Electron affinity generally becomes less negative. - Metallic character increases. Quick reference - Most reactive metal: bottom left (Cs, Fr). - Most reactive non-metal: top right (F). - Largest atom: bottom left. Smallest atoms: top right, apart from noble gases. Formation of monoatomic ions: tendency to reach the configuration of the nearest noble gas. - s-block metals lose their 1 or 2 s electrons: Na+, Ca2+. - p-block metals lose the p electrons and then the s electrons of the outer shell, never the d electrons: Al3+, Ga3+. (Heavy p-block metals often keep the s pair: Pb2+, Sn2+.) - d-block metals lose the outer s and then some d electrons: Fe2+, Fe3+, Cu+, Cu2+, Zn2+. - Non-metals gain electrons to reach the next noble gas: O2-, S2-, Cl-, Br-, I-, N3-. - Simple charge rule: groups 1, 2, 13 form +1, +2, +3; groups 15, 16, 17 form -3, -2, -1. Worked example: predict the charges of Mg, Al, S and Br. Mg 2+ ([Ne]), Al 3+ ([Ne]), S 2- ([Ar]), Br - ([Kr]). Clinical tip: the ions that run physiology (Na+, K+, Ca2+, Mg2+, Cl-, HPO4 2-) are all products of these rules, and knowing that Ca2+ is bigger than Mg2+ helps explain why they bind different proteins. Exam trap: transition metals do not follow the simple group rule. Iron is Fe2+ and Fe3+, both stable.

1. What is the ordering principle of the modern periodic table and who established it? 2. How does the electron configuration determine the group and the period? Give the period, group and block of sulfur and of iron. 3. Name three properties of metals and three of non-metals. 4. Rank in order of increasing atomic radius: Na, Mg, K, Cl. 5. Explain why a cation is smaller and an anion larger than the parent atom. 6. Order by size: O2-, F-, Na+, Mg2+. 7. Define first ionisation energy. Why does it increase across a period and decrease down a group? 8. The successive ionisation energies of an element (kJ/mol) are 738, 1451, 7733, 10543. To which group does it belong? 9. Why is the ionisation energy of oxygen lower than that of nitrogen? 10. Which element in each pair has the more negative electron affinity: Cl or Na; F or Ne; Cl or Br? 11. Predict the most common ions of K, Ca, Al, O, Cl and N. 12. Why do metal oxides form basic solutions and non-metal oxides acidic solutions? Give one example each.

1. Atomic number Z (Moseley, correcting Mendeleev's mass-based table). 2. Period = highest occupied shell n; main-group number = valence electrons. S [Ne] 3s2 3p4: period 3, group 16 (6A), p block. Fe [Ar] 3d6 4s2: period 4, group 8 (8B), d block. 3. Metals: shiny, conduct electricity, malleable and ductile, form cations. Non-metals: brittle, insulating, low melting, form anions. 4. Cl < Mg < Na < K. 5. Cation: fewer electrons, less repulsion, and same nuclear charge pulls the cloud in. Anion: more electrons, more repulsion, so the cloud expands. 6. Mg2+ < Na+ < F- < O2- (isoelectronic; larger nuclear charge means smaller ion). 7. Energy needed to remove the outermost electron from a gaseous atom. Across a period Z_eff increases and radius falls; down a group the radius and shielding increase. 8. The jump between the second and third values shows that two electrons are valence electrons: group 2 (magnesium). 9. Oxygen's fourth 2p electron is in an orbital with another electron, and the repulsion makes it easier to remove than nitrogen's half-filled, unpaired 2p3 electron. 10. Cl; F (Ne is positive); Cl (the smaller atom in the third period has slightly more negative affinity than bromine). 11. K+, Ca2+, Al3+, O2-, Cl-, N3-. 12. Metal oxides give hydroxide ions: Na2O + H2O -> 2 NaOH. Non-metal oxides give acids: CO2 + H2O -> H2CO3.

Core versus valence electrons - Core electrons: in the inner, filled shells; they do not take part in bonding. - Valence electrons: in the outermost shell; they form the bonds. When does a bond form? - As two atoms approach, the potential energy falls (attraction between one nucleus and the other atom's electrons) until it reaches a minimum at the equilibrium distance. If they come closer, repulsion (between nuclei and between electron clouds) makes it rise again. - A bond exists if the aggregate has lower energy than the separated atoms. The energy released is the bond energy. Energy must be supplied to break it. The noble-gas rule - Noble gases are stable and unreactive because their outer shell is complete (s2 p6, eight electrons; helium has two). - Other elements gain, lose or share electrons to reach the configuration of the nearest noble gas (octet rule; Lewis, 1916). Three limiting bond types 1. Ionic: metal + non-metal. Complete transfer of one or more electrons from one atom to the other; the resulting ions are held by electrostatic attraction. Example NaCl. 2. Covalent: non-metal + non-metal. Sharing of electron pairs between atoms forming discrete molecules. Example H2O, Cl2. 3. Metallic: metal + metal. Valence electrons delocalise over the whole metal (a "sea" of electrons around positive ions). Example Cu, Na. Coulomb's law - The energy between two point charges is E = k q1 q2 / r. If the charges have the same sign, E is positive (repulsion, energy absorbed); if opposite, E is negative (attraction, energy released). The bigger the charges and the smaller the distance, the larger the energy. - All bonds are ultimately electrostatic. Predicting bonds from the periodic table: a metal (low ionisation energy) plus a non-metal (high electron affinity) gives an ionic bond; two non-metals give a covalent bond; two metals give a metallic bond. Exam trap: real bonds are rarely pure. Most lie between the ionic and covalent limits, which is described by electronegativity difference (see later). Clinical tip: bond types set behaviour in the body: ionic salts dissolve into ions that carry current in nerves, covalent bonds hold organic molecules together, and metal ions such as Fe and Zn coordinate to proteins.

Electronegativity (chi): the ability of an atom in a compound to attract the bonding electrons towards itself. - Pauling scale (1932): fluorine is the highest at 4.0, then O 3.5, N 3.0, Cl 3.0, Br 2.8, C 2.5, H 2.1, Na 0.9, Cs about 0.7. - Mulliken: chi is proportional to the average of the ionisation energy and the electron affinity. - Trends: increases across a period and decreases down a group, like ionisation energy. The most electronegative element is F, at the top right. Bond polarity - Non-polar (homopolar) covalent: identical atoms or delta chi near 0 (H-H, Cl-Cl, C-H approx.). Electrons shared equally. - Polar (heteropolar) covalent: unequal sharing. The more electronegative atom carries a partial negative charge (delta minus) and the other delta plus. H-Cl: Cl is delta minus. - Ionic: transfer of electrons; net charges. - The polarity is measured by the dipole moment mu = charge x distance, in debye (D). Guide by delta chi - Less than about 0.4: essentially non-polar covalent. - About 0.4 to 1.7: polar covalent. - Greater than about 1.7 to 2.0: predominantly ionic. - All bonds are a continuum. Ionic and covalent are limiting forms. Halogen hydrides (data from the slides) - HF: dipole moment 1.98 D, delta chi 1.9, about 45 % ionic character. - HCl: 1.03 D, delta chi 0.9, 17 % ionic. - HBr: 0.79 D, delta chi 0.7, 12 % ionic. - HI: 0.38 D, delta chi 0.4, 5 % ionic. - Ionic character falls as the halogen becomes less electronegative. Order of ionic character among salts: KI < KBr < NaCl < LiCl (approx.), with all mainly ionic. Some compounds of metals, such as AlCl3 or LiI, show significant covalent character. Clinical tip: polarity governs solubility and membrane passage: polar and ionic drugs dissolve in blood but cross membranes poorly, while non-polar drugs are lipid soluble. The O-H, N-H and C=O bonds of biomolecules are all polar and can form hydrogen bonds. Exam trap: polar bond and polar molecule are different things. See the dipole moment note.

Simple picture - Metal atoms have few valence electrons and low ionisation energy. - The valence electrons are shared by all atoms: positive ions in a "sea" of mobile electrons. Properties explained - Electrical and thermal conductivity: the electrons flow freely. - Malleability and ductility: metal ions slide past each other and the electron sea keeps holding them, so the crystal deforms without breaking. - Lustre: free electrons absorb and re-emit light. - Compact (close-packed) structures, high density. Band theory (molecular orbital view) - Combining two atomic orbitals gives two molecular orbitals: one bonding, one antibonding. N atomic orbitals give N molecular orbitals. - For a very large N (of order 10^23) the levels are so close that they form a virtually continuous band. In lithium, N atoms give N/2 bonding-type levels, filled, and N/2 antibonding-type, empty. - Bands arise from s orbitals (Li), p orbitals (Al) and d orbitals (Cu), and can overlap. - Valence band: the highest band with electrons. Conduction band: the lowest empty band. - Conduction requires electrons to move into empty levels. Conductors, semiconductors, insulators - Metal (conductor): the valence and conduction bands overlap or the band is partly filled. Electrons move with negligible energy. - Insulator: a large energy gap (band gap) between full valence band and empty conduction band. Electrons cannot jump; no conductivity. - Semiconductor: a small gap (around 1 eV, silicon 1.1 eV), so some electrons are excited thermally or by light or doping; conductivity rises with temperature (opposite to metals). - Graphene has a band structure with zero gap; MoS2 monolayer has a gap of about 1.8 eV. Clinical tip: semiconductor detectors and LEDs power modern medical devices: pulse oximeters (LEDs plus photodiode), digital radiography detectors and lasers. Chromium, nickel and titanium alloys used in implants and stents are metals chosen for their strength and inertness. Exam trap: as temperature rises, the conductivity of a metal falls (more scattering), but that of a semiconductor rises.

1. Why do atoms form bonds, and how does potential energy vary with distance between two atoms? 2. Predict the bond type (ionic, covalent, metallic) in NaCl, Cl2, Cu and H2O. 3. Use the Born-Haber values for NaCl (sublimation +107, ionisation +494, half Cl2 bond +122, electron affinity -349, lattice -787 kJ/mol) to compute the enthalpy of formation. 4. Rank by lattice energy: NaCl, MgO, KCl, LiF. Explain. 5. Why is an ionic crystal brittle but a metal malleable? 6. Draw Lewis structures for H2O, NH3, CO2, HCN and SO4 2-. How many lone pairs on the central atom? 7. Order by bond length and bond energy: C-C, C=C, C triple bond C. 8. Give an example of an incomplete octet, an expanded octet and an odd-electron molecule. 9. Which is more polar, H-Cl or H-F, and why? Is CO2 polar? Is H2O? 10. Use VSEPR to predict the shape and approximate angles of BF3, NH3, H2O, SF4 and XeF4. 11. Why is diamond hard with a very high melting point whereas CO2 is a gas? 12. In band theory, what distinguishes a metal, a semiconductor and an insulator? 13. Name the strongest intermolecular force in: CH4, HCl, CH3OH, NaCl in water. 14. Why does ice float, and why does H2O boil so much higher than H2S? 15. Explain why CH3OCH3 has a lower boiling point than CH3CH2OH.

1. Bonding lowers the energy of the system. Energy falls as atoms approach (attraction), reaches a minimum at the bond length and then rises steeply (repulsion). 2. NaCl ionic; Cl2 covalent (non-polar); Cu metallic; H2O polar covalent. 3. 107 + 494 + 122 - 349 - 787 = -413 kJ/mol. 4. MgO > LiF > NaCl > KCl. MgO has 2+/2- charges; LiF has small ions; KCl has larger ions. 5. In an ionic crystal, shifting a layer brings like charges together and it cleaves. In a metal, the electron sea keeps the ions bound as they slide. 6. H2O: 2 lone pairs. NH3: 1. CO2: 0 on C (two lone pairs on each O). HCN: 0. SO4 2-: 0 (four S-O bonds, or expanded octet with double bonds). 7. Length: C-C > C=C > C triple bond C. Energy: C triple bond C > C=C > C-C. 8. BF3 (incomplete), PF5 or SF6 (expanded), NO (odd electron). 9. H-F is more polar (larger delta chi). CO2 is non-polar (linear, dipoles cancel). H2O is polar (bent). 10. BF3 trigonal planar, 120 degrees. NH3 trigonal pyramidal, about 107. H2O bent, 104.5. SF4 see-saw. XeF4 square planar, 90. 11. Diamond is a covalent network (each C bonded to four others), so melting means breaking covalent bonds. CO2 is a molecular substance held by weak London forces. 12. Metal: partly filled or overlapping bands. Semiconductor: small gap. Insulator: large gap between filled and empty bands. 13. CH4 London; HCl dipole-dipole; CH3OH hydrogen bond; NaCl in water ion-dipole. 14. Ice has an open hydrogen-bonded lattice that is less dense than the liquid. H2O forms hydrogen bonds whereas H2S does not, and these must be broken to boil. 15. Ethanol has an O-H that forms hydrogen bonds; dimethyl ether cannot donate one, so it has only weaker dipole-dipole and London forces.

Fundamental principle (Pauling) - A covalent bond forms when an atomic orbital of one atom overlaps an orbital of another, and the overlap region between the nuclei holds a pair of electrons. - The two wave functions are in phase and add, increasing the electron density between the nuclei. - The pair must have opposite spins (Pauli principle). A set of overlapping orbitals holds at most two electrons. - The greater the overlap, the stronger and more stable the bond. - VB theory considers the valence electrons, following Lewis's idea that bonding and lone pairs are localised on atoms or between two atoms. Sigma (sigma) bond - Head-on (end-to-end) overlap along the line joining the nuclei; maximum electron density on that axis. - Formed by s+s (H2), s+p (HF), p+p along the axis (F2), or hybrid orbitals. - Free rotation around a sigma single bond. Pi (pi) bond - Sideways overlap of parallel p orbitals; electron density above and below the internuclear axis, with a nodal plane through the nuclei. - Weaker than sigma, and it prevents rotation (locks the geometry; the cause of cis/trans isomers). - Present only in double and triple bonds. Counting - Single bond: 1 sigma. - Double bond: 1 sigma + 1 pi. - Triple bond: 1 sigma + 2 pi. Examples: C2H4 (ethene) has 5 sigma (4 C-H and 1 C-C) and 1 pi. C2H2 (ethyne) has 3 sigma and 2 pi. N2 has 1 sigma and 2 pi. CO2 has 2 sigma and 2 pi. Worked example: acetaldehyde CH3-CHO has 6 sigma bonds and 1 pi bond (in C=O). Exam trap: pi bonds are not formed on their own; a sigma bond is always present first. Clinical tip: the planar, rigid peptide bond (partial double-bond character) and cis/trans double bonds in fatty acids come from pi bonding, and they determine protein folding and membrane fluidity.

1. State the fundamental principle of valence bond theory. 2. Distinguish a sigma from a pi bond. How many of each are in ethene, ethyne and N2? 3. Why do we need hybridisation to explain methane? 4. Give the hybridisation and geometry: BeCl2, BF3, CH4, NH3, H2O, PCl5, SF6. 5. What is the hybridisation of each carbon in CH3-CH=CH-C triple bond CH? 6. Why is the peptide nitrogen planar? 7. Why does rotation occur about C-C single bonds but not about C=C? 8. Explain the formation of bonding and antibonding MOs from two 1s orbitals. 9. Calculate the bond order of H2, He2, He2+, N2 and O2. Which are paramagnetic? 10. Why does He2 not exist but H2 does? 11. Why does liquid oxygen stick to a magnet, and why can't Lewis or VB theory explain it? 12. Predict the bond order and magnetic behaviour of O2-, O2 2- and O2+. 13. Compare VB and MO descriptions of benzene.

1. A covalent bond forms when two atomic orbitals overlap and the overlap region holds an electron pair with opposite spins. More overlap means a stronger bond. 2. Sigma: head-on overlap along the axis. Pi: side-on overlap above and below the axis. Ethene: 5 sigma, 1 pi. Ethyne: 3 sigma, 2 pi. N2: 1 sigma, 2 pi. 3. Carbon has two unpaired 2p electrons in the ground state, and the s and p orbitals have different shapes, but methane has four identical bonds at 109.5 degrees. Mixing gives four equivalent sp3 orbitals. 4. BeCl2 sp linear; BF3 sp2 trigonal planar; CH4 sp3 tetrahedral; NH3 sp3 pyramidal; H2O sp3 bent; PCl5 sp3d trigonal bipyramidal; SF6 sp3d2 octahedral. 5. CH3 sp3; CH= sp2; =CH- sp2; C triple bond sp; final CH sp. (Carbons: sp3, sp2, sp2, sp, sp.) 6. Resonance: the nitrogen lone pair delocalises into the C=O pi system, so N is sp2 and the amide group is planar. 7. A single sigma bond has cylindrical symmetry; rotation does not break overlap. A pi bond requires parallel p orbitals, and rotating would break it. 8. Adding the 1s wave functions in phase gives sigma1s (density between nuclei, lower energy). Subtracting gives sigma*1s (node between nuclei, higher energy). 9. H2: 1. He2: 0. He2+: 0.5. N2: 3. O2: 2. Paramagnetic: He2+, O2. 10. He2 has 2 electrons in the bonding and 2 in the antibonding MO, so BO 0 and there is no net stabilisation. 11. O2 has two unpaired electrons in degenerate pi* MOs (paramagnetic). Lewis or VB structures pair all the electrons. 12. O2-: BO 1.5, paramagnetic (1 unpaired). O2 2- (peroxide): BO 1, diamagnetic. O2+: BO 2.5, paramagnetic (1 unpaired). 13. VB: two resonance Kekule structures with localised double bonds, sp2 carbons and sigma bonds. MO: six pi electrons in three bonding MOs delocalised over the whole ring, all C-C bonds equal (140 pm).

States of matter - Solid: ordered, long-range order. Particles occupy fixed positions in a lattice and only vibrate. Fixed shape and volume, rigid, incompressible, dense. - Liquid: short-range order. Particles are close and can slide past each other. Own volume but not own shape; nearly incompressible. - Gas: disordered. Particles are far apart with high kinetic energy. No fixed shape or volume, low density, highly compressible, low viscosity, fast diffusion, gases mix completely. Variables that describe a gas: pressure P (atm), volume V (L), temperature T (K), amount n (mol). Ideal (perfect) gas: a model with these postulates (kinetic theory) 1. Particles have zero (negligible) volume compared with the container. 2. There are no interactions between particles or with the walls except elastic collisions (no energy is lost); kinetic energy depends only on temperature. 3. Particles are in constant, fast, random motion. 4. Pressure arises from collisions with the walls. When is a real gas ideal? At low pressure (few particles; particle volume negligible) and high temperature (high kinetic energy overcomes attractions). Deviations are largest for easily liquefied gases (water vapour, CO2, NH3) near their condensation point. Kinetic energy and temperature: the average kinetic energy of the particles is proportional to the absolute temperature, E = 3/2 k T per molecule (3/2 R T per mole). At the same temperature, lighter molecules move faster: H2 faster than O2 faster than CO2. Example: why can we compress a gas but not a liquid? Gas particles are mostly separated by empty space, while liquid particles are already in contact. Clinical tip: gases at body conditions (oxygen, nitrogen, nitrous oxide, carbon dioxide) behave almost ideally, so the ideal-gas equations are used routinely in respiratory physiology and in calibrating ventilators. Exam trap: at the same temperature all gases have the same average kinetic energy, but not the same speed.

Definition: P = force / area. SI unit: pascal (Pa) = N/m2. Units - 1 atm = 760 mmHg = 760 torr = 101 325 Pa = 101.325 kPa = 1.013 bar (approximately). - 1 mmHg = 133.3 Pa. - Blood pressure in mmHg; atmospheric pressure at sea level about 760 mmHg. Torricelli barometer - A tube filled with mercury is inverted in a dish of mercury. The mercury column falls until its weight balances the atmospheric pressure. At sea level, 0 degrees C, latitude 45 degrees, the height is 760 mm: P = 760 mmHg = 1 atm. - The vacuum above the column has no pressure. Manometer - U-tube used to measure the pressure of a gas in a container relative to the atmosphere. - Closed-end manometer: gas pressure equals the height difference of the mercury levels (against a vacuum). - Open-end manometer: P(gas) = P(atm) + delta h if the gas pushes the mercury up on the open side; P(gas) = P(atm) - delta h if the gas pressure is lower. Example: an open manometer with mercury 30 mm higher on the side attached to the gas, at atmospheric pressure 760 mmHg. Then the gas pressure is 760 + 30 = 790 mmHg. Conversion: 2.00 atm x 760 mmHg/atm = 1520 mmHg; 95.0 kPa / 101.3 kPa/atm = 0.938 atm. Clinical tip: blood pressure, intraocular pressure, intracranial pressure, and oxygen partial pressures are all in mmHg because the historical mercury manometer was the first measuring device. Gauge pressure (relative to atmospheric) is what an oxygen cylinder gauge shows. Exam trap: gas laws need absolute pressure, not gauge pressure. Add atmospheric pressure to a gauge reading.

Equation of state for an ideal gas: PV = nRT - P pressure, V volume, n moles, T absolute temperature (kelvin), R the gas constant. Values of R - R = 0.0821 L atm K^-1 mol^-1 (P in atm, V in L). - R = 8.314 J K^-1 mol^-1 (SI, P in Pa and V in m3). - R = 1.987 cal K^-1 mol^-1. - Derivation: R = PV / nT = (1 atm x 22.414 L) / (1 mol x 273.15 K) = 0.0821. How the laws follow - Constant T: PV = nRT = constant (Boyle). - Constant P and n: V/T constant (Charles). - Constant V and n: P/T constant (Gay-Lussac). - Constant P and T: V/n constant (Avogadro). Useful forms - Moles from mass: n = g / M, so PV = (g / M) R T and M = g R T / (P V). - Density: d = g / V = P M / (R T). - Change of state: P1 V1 / (n1 T1) = P2 V2 / (n2 T2). Worked example 1: volume of 2.00 mol of gas at 300 K and 2.00 atm. V = nRT / P = 2.00 x 0.0821 x 300 / 2.00 = 24.6 L. Worked example 2: molar mass of a gas. 1.50 g of a gas occupies 0.500 L at 1.00 atm and 27 degrees C. n = PV / RT = 1.00 x 0.500 / (0.0821 x 300) = 0.0203 mol. M = 1.50 / 0.0203 = 73.9 g/mol. Worked example 3: density of oxygen at 37 degrees C (310 K) and 1.00 atm. d = P M / R T = 1.00 x 32.0 / (0.0821 x 310) = 1.26 g/L. Problem-solving steps 1. List P, V, n, T and identify what is unknown. 2. Convert to atm, litres, mol and kelvin. 3. Choose the equation (PV = nRT, or the combined law if the state changes). 4. Check the answer: does it make sense in size and unit? Clinical tip: oxygen cylinders: a 10 L cylinder at 150 atm holds about 1500 L at 1 atm (Boyle), so at 5 L/min it lasts about 300 minutes (ignoring residual pressure). Exam trap: mixing the two R values. If pressure is in atm use 0.0821; in pascals, use 8.314.

Why real gases deviate from ideal behaviour 1. Molecules attract each other (intermolecular forces). 2. Molecules have finite size, so the free volume is smaller than the container volume. A real gas is close to ideal at high temperature and low pressure. Compressibility factor: Z = P V_m / (R T) for one mole. For an ideal gas Z = 1 at all pressures. - Z less than 1: attractions dominate (moderate pressures; the gas is more compressible than ideal). Typical of CO2, CH4 and NH3. - Z greater than 1: molecular volume dominates (very high pressures; molecules cannot be squeezed further). Hydrogen and helium show it early because their attractions are tiny. Effect of attractions: at moderately high pressure molecules are close enough to attract each other. A molecule about to hit the wall is pulled back by its neighbours, so the impact is weaker and the measured pressure is lower than the ideal value. Effect of volume: at very high pressure the molecules occupy a significant part of the container, so the free volume is V minus the volume of the molecules. van der Waals equation (P + a n^2 / V^2) (V - n b) = n R T - a n^2 / V^2 corrects the pressure for attractions. The constant a is larger for gases with strong intermolecular forces. - n b corrects the volume for the finite size of molecules. The constant b is roughly the volume of one mole of molecules. - Typical constants: helium a = 0.034 L2 atm/mol2, b = 0.024 L/mol; CO2 a = 3.6, b = 0.043. Worked example: 1.00 mol of CO2 in 1.00 L at 300 K. Ideal: P = nRT/V = 24.6 atm. van der Waals: P = nRT/(V - nb) - a n^2/V^2 = 24.6 / 0.957 - 3.6 = 25.7 - 3.6 = 22.1 atm. The attractions dominate and the real pressure is about 10 % lower. Relative density (from the ideal-gas equation): d1 / d2 = M1 / M2 under the same P and T. CO2 (44) is about 1.5 times as dense as air (mean M about 29). Clinical tip: nitrous oxide and CO2 in cylinders are stored partly as liquid, so the gauge pressure stays at the vapour pressure until the liquid is gone and does not fall in step with the contents. Oxygen cylinders hold compressed gas, so the gauge reflects what is left. Exam trap: in the van der Waals equation the attraction term is added to P (the measured P is too low), and the volume term is subtracted from V.

Dalton's law (1801): in a mixture of non-reacting gases, the total pressure is the sum of the partial pressures. P_total = P_A + P_B + ... - Partial pressure P_A: the pressure that gas A would exert if it alone occupied the whole volume at the same temperature. - Molecular explanation: ideal-gas particles do not interact, so each gas hits the walls independently. Mole fraction: x_A = n_A / n_total (dimensionless; the sum of the fractions is 1). Partial pressure: P_A = x_A x P_total. In ideal gases the mole fraction equals the volume fraction and the percentage by volume. Amagat's law (partial volumes): V_total = V_A + V_B + ..., where V_A = n_A R T / P_total. Worked example: air at 750 mmHg - Composition: 78 % N2, 21 % O2, 0.93 % Ar, 0.07 % CO2 (rounded). - P(N2) = 0.78 x 750 = 585 mmHg. P(O2) = 0.21 x 750 = 157.5 mmHg. P(Ar) = 0.0093 x 750 = 6.98 mmHg. P(CO2) = 0.0007 x 750 = 0.53 mmHg. - Sum = 750 mmHg. Gas collected over water: total pressure = P(gas) + P(water vapour). Water vapour pressure at 25 degrees C is about 23.8 mmHg, at 37 degrees C about 47 mmHg. Clinical applications - Inspired oxygen: at sea level P(O2) = 0.21 x 760 = 160 mmHg. In the trachea, humidified air has a water vapour pressure of 47 mmHg, so P(O2) = 0.21 x (760 - 47) = 150 mmHg. - Altitude sickness: at high altitude the fraction of O2 is still 21 % but the total pressure is lower, so P(O2) is lower. On Everest (about 250 mmHg) it is about 52 mmHg. - Giving 100 % oxygen at 1 atm raises P(O2) to about 713 mmHg in the airway; hyperbaric oxygen at 3 atm gives over 2000 mmHg. - Gas exchange in the lungs depends on the partial-pressure gradient, and blood gas reports (PaO2, PaCO2) are partial pressures. - Decompression sickness occurs when nitrogen dissolved at high partial pressure during a dive forms bubbles on rapid ascent. Exam trap: partial pressures depend on mole fraction, not on mass fraction, and it is the partial pressure, not the concentration, that drives diffusion across the alveolar membrane.

Definition: the pressure exerted by the vapour in equilibrium with its liquid (or solid) in a closed container at a given temperature. Dynamic equilibrium - Molecules with enough kinetic energy escape from the surface (evaporation) while vapour molecules return (condensation). - At equilibrium, the rate of evaporation equals the rate of condensation. Nothing changes macroscopically, but both processes continue. Dependence - On temperature: raising T increases the fraction of molecules with enough energy to escape, so vapour pressure rises steeply (water: 23.8 mmHg at 25 degrees C, 47 at 37, 760 at 100). - On intermolecular forces: strong forces (polarity, hydrogen bonding) mean low vapour pressure and low volatility. Ether is highly volatile; water is not; mercury barely at all. - It does not depend on the amount of liquid or on the volume of the container (as long as some liquid remains). Volatility: the tendency to evaporate; high vapour pressure means high volatility. Boiling: the liquid boils when its vapour pressure equals the external pressure. So the normal boiling point (at 1 atm) of water is 100 degrees C, and at altitude (lower pressure) water boils at a lower temperature. An autoclave raises the pressure so that water boils at 121 degrees C and sterilises instruments. Clinical tip: volatile anaesthetics (sevoflurane, isoflurane) are liquids with high vapour pressure at room temperature; vaporisers use this. Rubbing alcohol cools skin as it evaporates, taking heat. In the lungs, inhaled air is saturated with water vapour at 47 mmHg at 37 degrees C. Exam trap: vapour pressure is a property of the liquid and temperature only. Do not confuse it with the total pressure of a gas mixture.

1. List the postulates of the kinetic theory of an ideal gas. Under which conditions does a real gas behave ideally? 2. Convert 1.50 atm to mmHg and kPa. Convert 37 degrees C and -196 degrees C to kelvin. 3. A gas occupies 4.0 L at 2.0 atm. What is its volume at 5.0 atm, constant temperature? 4. A gas is 5.0 L at 27 degrees C. What is its volume at 127 degrees C, constant pressure? 5. A sealed 10 L cylinder at 20 degrees C and 5.0 atm is heated to 60 degrees C. What is the new pressure? 6. How many moles are in 11.2 L of a gas at STP? What is the mass if it is CO2? 7. What volume does 0.50 mol of an ideal gas occupy at 25 degrees C and 1.0 atm? 8. 0.80 g of a gas occupies 0.25 L at 1.0 atm and 27 degrees C. Find its molar mass. 9. Calculate the density of nitrogen at STP. 10. A mixture contains 2.0 mol N2 and 0.50 mol O2 at a total pressure of 5.0 atm. Find the partial pressures. 11. Calculate the partial pressure of oxygen in dry air at 630 mmHg, and in the humidified airway at 37 degrees C and 760 mmHg total. 12. Explain why water boils at a lower temperature at high altitude and why a pressure cooker cooks faster. 13. Why is the vapour pressure of ethanol higher than that of water at the same temperature? 14. Why do real gases deviate from ideal behaviour, and at which conditions are they most ideal? What do the terms a and b in the van der Waals equation correct?

1. Particles of zero volume; no interactions except elastic collisions; constant random motion; kinetic energy depends only on T. Real gases behave ideally at low pressure and high temperature. 2. 1.50 atm = 1140 mmHg = 152 kPa. 37 degrees C = 310.15 K. -196 degrees C = 77 K. 3. P1 V1 = P2 V2: V2 = 2.0 x 4.0 / 5.0 = 1.6 L. 4. T1 = 300 K, T2 = 400 K. V2 = 5.0 x 400 / 300 = 6.7 L. 5. T1 = 293 K, T2 = 333 K. P2 = 5.0 x 333 / 293 = 5.7 atm. 6. n = 11.2 / 22.4 = 0.50 mol. CO2: 0.50 x 44.01 = 22.0 g. 7. V = nRT / P = 0.50 x 0.0821 x 298 / 1.0 = 12.2 L. 8. n = PV / RT = 1.0 x 0.25 / (0.0821 x 300) = 0.01015 mol. M = 0.80 / 0.01015 = 78.8 g/mol. 9. d = 28.02 / 22.4 = 1.25 g/L. 10. x(N2) = 2.0 / 2.5 = 0.80, so P = 4.0 atm. x(O2) = 0.20, so P = 1.0 atm. 11. Dry: 0.21 x 630 = 132 mmHg. Humidified: 0.21 x (760 - 47) = 150 mmHg. 12. Water boils when its vapour pressure equals the external pressure. At altitude, the pressure is lower, so boiling occurs at a lower temperature. A pressure cooker raises the pressure, so water boils above 100 degrees C and food cooks faster. 13. Ethanol has weaker intermolecular forces than water (less hydrogen bonding per molecule), so more molecules escape at the same temperature. 14. Because molecules attract each other and have a finite volume. Most ideal at high temperature and low pressure. a corrects the pressure for attractions. b corrects the volume for molecular size.

Definition: the formal charge an atom would have in a compound if all the bonding electrons were assigned to the more electronegative atom. - It does not represent the real charge on the atom (in covalent compounds the real charge is partial). - Positive if the atom loses electrons (oxidation, more electropositive), negative if it gains. Conservation - In a neutral compound, the sum of the oxidation numbers of all atoms is zero. Example: H2SO4: 2(+1) + (+6) + 4(-2) = 0. - In an ion, the sum equals the charge. Example: SO4 2-: (+6) + 4(-2) = -2. NO3-: (+5) + 3(-2) = -1. Empirical rules (apply in this order of priority) 1. Atoms in an element (free, uncombined) have o.n. 0: H2, O2, Mg, Na, Cl2. 2. Monoatomic ions have o.n. equal to their charge: Na+ +1, Cl- -1, Fe3+ +3. 3. Fluorine is always -1 in compounds. 4. Group 1A metals are always +1: LiCl, NaCl, KBr. 5. Group 2A metals are always +2: BeCl2, MgCl2. 6. Boron and aluminium (3A) are +3: BF3, AlCl3. 7. Hydrogen is +1 (H2O, HCl, CH4, H2SO4), except in metal hydrides where it is -1 (NaH, CaH2). 8. Oxygen is -2 (H2O, CO2), except in peroxides where it is -1 (H2O2, Na2O2), in superoxides where it is -1/2 (KO2), and in OF2 where it is +2. 9. Halogens (7A) are -1, except when bonded to oxygen or a more electronegative halogen (then positive: +1, +3, +5, +7). 10. Carbon takes any value from -4 to +4, nitrogen from -3 to +5, including fractional averages. Maximum and minimum - The maximum o.n. of a main-group element is usually the number of valence electrons (group number): S +6, N +5, Cl +7. - The minimum o.n. for a non-metal is the group number minus 8 (the number of electrons to complete the octet): S -2, N -3, Cl -1. - Metals are always positive in compounds. Worked examples - CO2: C +4, O -2. CO: C +2. CH4: C -4. C (graphite): 0. - KMnO4: K +1, O -2 x 4 = -8, so Mn = +7. - NaCl O.N. Na +1, Cl -1. NaClO (+1), NaClO2 (+3), NaClO3 (+5), NaClO4 (+7): the chlorine o.n. rises with each oxygen. - H2O2: H +1, O -1. OF2: O +2, F -1. - Na2SO3 (S +4) and Na2SO4 (S +6). K2Cr2O7: 2(+1) + 2x + 7(-2) = 0, so Cr +6. - Fe3O4: average Fe +8/3, mixed Fe2+ and Fe3+. - Manganese can show +7, +6, +4, +3, +2, 0 and -1 (see its compounds: KMnO4 +7, MnO2 +4, MnCl2 +2). Exam trap: to find an unknown o.n., first assign the fixed ones (O, H, group 1 and 2 metals), then solve for the remainder so the total equals the charge. Clinical tip: oxidation numbers track electron movement in biological redox: Fe2+ (in deoxyhaemoglobin) versus Fe3+ (in methaemoglobin, which cannot carry oxygen), and the redox pairs NAD+/NADH.

Rule for formulas of binary compounds: write the less electronegative element first, then the more electronegative. The name gives the more electronegative element first with the suffix -ide, then "of" the other (in English, the other element is named first: sodium chloride). - Naming the number of atoms: prefixes mono-, di-, tri-, tetra-, penta-, hexa- (IUPAC), or the Roman numeral of the o.n. in brackets (Stock). Hydrides: hydrogen + metal, H is -1. - General formula MHx. - NaH sodium hydride; CaH2 calcium hydride; LiH lithium hydride. - Ionic, reactive with water: NaH + H2O -> NaOH + H2. Used as a reducing agent (LiAlH4 in synthesis). Covalent hydrides of non-metals (H treated as +1 or by convention) - CH4 methane; NH3 ammonia; PH3 phosphine; SiH4 silane. These keep their traditional names. Hydracids: hydrogen + halogen or chalcogen (H is +1). In water they give acidic solutions. - HF hydrogen fluoride (hydrofluoric acid). - HCl hydrogen chloride (hydrochloric acid; gastric acid). - HBr hydrogen bromide (hydrobromic acid). - HI hydrogen iodide (hydroiodic acid). - H2S hydrogen sulfide (hydrosulfuric acid). - Naming: the gas is "hydrogen X-ide"; the aqueous acid is "hydro-X-ic acid". Example: HCl(g) is hydrogen chloride; HCl(aq) is hydrochloric acid. Clinical tip: hydrochloric acid in the stomach (pH 1 to 2) is HCl(aq). Hydrogen sulfide is a toxic gas that inhibits cytochrome oxidase. Ammonia is produced by protein and urea metabolism and is toxic in hepatic encephalopathy. Exam trap: in NaH the H is -1 (hydride ion H-), the opposite of its usual +1.

Formula: symbol of the metal or non-metal, then oxygen, with subscripts set so that the total o.n. is zero (criss-cross the oxidation numbers and simplify). - Oxygen + metal: basic oxide. Oxygen + non-metal: acid oxide (anhydride). Naming systems - IUPAC (systematic): Greek prefixes for the number of atoms: dinitrogen pentoxide. - Stock: Roman numeral for the oxidation number of the element: iron(III) oxide. - Traditional: -ous (lower o.n.) and -ic (higher o.n.): ferrous oxide, ferric oxide. Examples - Na2O disodium oxide (sodium oxide). CaO calcium oxide. MgO magnesium oxide. Ag2O silver oxide. - Fe2O3 diiron trioxide, iron(III) oxide. FeO iron(II) oxide. - Cu2O copper(I) oxide. CuO copper(II) oxide. - CrO3 chromium trioxide, chromium(VI) oxide. B2O3 diboron trioxide. - CO carbon monoxide (C +2). CO2 carbon dioxide (C +4). - SO2 sulfur dioxide, sulfur(IV) oxide. SO3 sulfur trioxide, sulfur(VI) oxide. - P2O5 (P4O10) diphosphorus pentoxide, phosphorus(V) oxide. - Nitrogen oxides: N2O dinitrogen monoxide (nitrous oxide, N +1), NO nitrogen monoxide (nitric oxide, +2), N2O3 dinitrogen trioxide (+3), NO2 nitrogen dioxide (+4), N2O5 dinitrogen pentoxide (+5). - H2O dihydrogen oxide (water). Reactions with water - Basic oxide + H2O -> hydroxide: Na2O + H2O -> 2 NaOH; CaO + H2O -> Ca(OH)2. - Acid oxide + H2O -> oxyacid: SO3 + H2O -> H2SO4; CO2 + H2O -> H2CO3; N2O5 + H2O -> 2 HNO3. Clinical tip: N2O (laughing gas) is an anaesthetic and analgesic, NO is a vasodilator signalling molecule, CO binds haemoglobin 200 times more strongly than O2, and CO2 dissolved in blood forms carbonic acid that is buffered by bicarbonate. CaO is quicklime, Fe2O3 is rust, and titanium dioxide and zinc oxide are sunscreen ingredients. Exam trap: the Roman numeral in the Stock name is the oxidation number of the element, not the number of atoms.

Formula: metal first, then the non-metal; subscripts balance the charges (criss-cross). Name: non-metal stem + -ide, then the metal (with Roman numeral if the metal has more than one o.n.). Formation: hydroxide + hydracid -> binary salt + water. NaOH + HCl -> NaCl + H2O. Examples - NaCl sodium chloride. CaCl2 calcium chloride. BaBr2 barium bromide. - FeCl2 iron(II) chloride, ferrous chloride. FeCl3 iron(III) chloride, ferric chloride. - SnS2 tin(IV) sulfide. ZnSe zinc selenide. Zn3As2 zinc arsenide. Mg3P2 magnesium phosphide. MnI2 manganese(II) iodide. - Binary compounds of two non-metals use the Greek prefixes: PCl5 phosphorus pentachloride; BF3 boron trifluoride; ClF3 chlorine trifluoride; CCl4 carbon tetrachloride. Common anion stems: fluoride, chloride, bromide, iodide, oxide, sulfide, nitride, phosphide, hydride. Clinical tip: many drugs and reagents are binary salts: potassium chloride (K+ replacement), sodium iodide, zinc oxide, silver sulfadiazine, ferrous sulfate (a ternary salt) for anaemia. Fe2+ is ferrous, Fe3+ is ferric; only ferrous is absorbed well. Exam trap: the number of atoms in the formula comes from balancing charges. Al2O3, not AlO. Ionic compounds are usually named without Greek prefixes.

Formula: M(OH)m, where m is the oxidation number of the metal. Use brackets when more than one OH is needed: Ca(OH)2. Name: metal hydroxide (with Roman numeral if needed). Formation: basic oxide + water -> hydroxide - Na2O + H2O -> 2 NaOH (sodium hydroxide) - CaO + H2O -> Ca(OH)2 (calcium hydroxide) - Li2O + H2O -> 2 LiOH (lithium hydroxide) - Fe2O3 + 3 H2O -> 2 Fe(OH)3 (iron(III) hydroxide) Other examples: KOH potassium hydroxide, Mg(OH)2 magnesium hydroxide, Al(OH)3 aluminium hydroxide, Fe(OH)2 iron(II) hydroxide. Properties: hydroxides of groups 1 and 2 are strong bases, releasing OH- in water. Others are sparingly soluble. Neutralisation: hydroxide + acid -> salt + water. NaOH + HCl -> NaCl + H2O. Clinical tip: milk of magnesia is Mg(OH)2 (an antacid and laxative), aluminium hydroxide is an antacid and phosphate binder, and sodium hydroxide is a caustic that causes severe tissue burns. Calcium hydroxide ("slaked lime") is used in dentistry. Exam trap: OH- is a polyatomic ion, so with more than one, put it in brackets: Ca(OH)2, not CaOH2.

Formula: HxAyOz (hydrogen, non-metal, oxygen). Formation: acid oxide (anhydride) + water -> oxyacid. - SO3 + H2O -> H2SO4 (sulfuric acid, S +6) - SO2 + H2O -> H2SO3 (sulfurous acid, S +4) - N2O5 + H2O -> 2 HNO3 (nitric acid, N +5) - N2O3 + H2O -> 2 HNO2 (nitrous acid, N +3) - CO2 + H2O -> H2CO3 (carbonic acid, C +4) - B2O3 + H2O -> 2 HBO2 (boric acid, or H3BO3 with more water) - CrO3 + H2O -> H2CrO4 (chromic acid) Naming by the oxidation number of the central non-metal (traditional) - If it has two o.n.: lower -ous, higher -ic (sulfurous, sulfuric; nitrous, nitric). - If it has four (halogens): hypo-...-ous (lowest), -ous, -ic, per-...-ic (highest). Chlorine family - HClO hypochlorous acid (Cl +1) from Cl2O + H2O. - HClO2 chlorous acid (+3). - HClO3 chloric acid (+5). - HClO4 perchloric acid (+7). Other important acids: H3PO4 phosphoric acid (P +5), H2SO4, HNO3, H2CO3. Acid strength: rises with the number of oxygens not bonded to H: HClO < HClO2 < HClO3 < HClO4. H2SO4 is stronger than H2SO3. Clinical tip: carbonic acid (formed from CO2) is the central acid of the blood buffer system, phosphoric acid derivatives underlie ATP and DNA, nitrous acid forms in the stomach from nitrite, and hypochlorous acid is the active disinfectant in bleach and is also made by neutrophils to kill bacteria. Exam trap: the number of oxygens tells the acid strength and the o.n. tells the name. Perchloric > chloric > chlorous > hypochlorous.

Formation: hydroxide + oxyacid -> oxysalt + water. - NaOH + HNO3 -> NaNO3 + H2O (sodium nitrate) - NaOH + HNO2 -> NaNO2 + H2O (sodium nitrite) - 2 NaOH + H2SO4 -> Na2SO4 + 2 H2O (sodium sulfate) - Ca(OH)2 + H2CO3 -> CaCO3 + 2 H2O (calcium carbonate) - Basic oxide + oxyacid: CaO + H2CO3 -> CaCO3 + H2O. - Hydroxide + acid oxide: Ca(OH)2 + CO2 -> CaCO3 + H2O. Anion names (acid name -> anion name) - hypo-...-ous -> hypo-...-ite: HClO -> ClO- hypochlorite. - -ous -> -ite: HClO2 -> ClO2- chlorite; HNO2 -> NO2- nitrite; H2SO3 -> SO3 2- sulfite. - -ic -> -ate: HClO3 -> ClO3- chlorate; HNO3 -> NO3- nitrate; H2SO4 -> SO4 2- sulfate. - per-...-ic -> per-...-ate: HClO4 -> ClO4- perchlorate. - H2CO3 -> CO3 2- carbonate (HCO3- hydrogencarbonate, bicarbonate). - H3PO4 -> PO4 3- phosphate (HPO4 2- hydrogenphosphate; H2PO4- dihydrogenphosphate). Examples of salts - NaClO sodium hypochlorite (bleach). Ca(ClO)2 calcium hypochlorite. Ca(ClO2)2 calcium chlorite. - KClO3 potassium chlorate. Zn(ClO4)2 zinc perchlorate. - CaSO3 calcium sulfite. CaSO4 calcium sulfate. Na2SO4 sodium sulfate. K2SO3 potassium sulfite. - KMnO4 potassium permanganate (Mn +7). K2CrO4 potassium chromate (Cr +6). K2Cr2O7 potassium dichromate (Cr +6). Acid salts: if an acid loses only some of its protons, the anion keeps hydrogen: NaHCO3 sodium hydrogencarbonate, NaHSO4 sodium hydrogensulfate, Na2HPO4. Ion formation: HNO3 -> H+ + NO3-. H2SO4 -> 2 H+ + SO4 2-. HClO -> H+ + ClO-. Neutralisation summary (all give salt + water) - HCl + NaOH -> NaCl + H2O. - CuO + H2SO4 -> CuSO4 + H2O. - H2SO4 + Ca(OH)2 -> CaSO4 + 2 H2O. Clinical tip: many medical products are oxysalts: sodium bicarbonate (antacid, treatment of acidosis), calcium carbonate (antacid, calcium supplement), magnesium sulfate (eclampsia), potassium nitrate (desensitising toothpaste), sodium nitrite (cyanide poisoning), sodium hypochlorite (disinfectant), potassium permanganate (dilute solutions for wound washing), and barium sulfate (radiological contrast, which is insoluble and therefore non-toxic). Exam trap: -ite has fewer oxygens than -ate for the same element. Sulfite SO3 2-, sulfate SO4 2-; nitrite NO2-, nitrate NO3-.

1. Define oxidation number. Is it the same as real charge? 2. Determine the oxidation number of the underlined element: S in H2SO4, N in HNO3, Mn in KMnO4, Cr in K2Cr2O7, Cl in NaClO3, C in CH4, C in CO2, O in H2O2, H in CaH2. 3. Give the o.n. of chlorine in NaCl, NaClO, NaClO2, NaClO3 and NaClO4 and name each compound. 4. Name: Na2O, CaO, Fe2O3, FeO, CO, SO3, N2O5, P2O5, Cu2O. 5. Write formulas: calcium hydride, hydrogen bromide, aluminium oxide, iron(III) chloride, carbon tetrachloride, sodium hydroxide. 6. Complete and name the products: SO3 + H2O; Na2O + H2O; N2O3 + H2O; CaO + H2O. 7. Write and name the acids with chlorine in oxidation states +1, +3, +5, +7. 8. Name: NaNO2, NaNO3, K2SO3, CaSO4, KClO3, Ca(ClO)2, KMnO4, K2Cr2O7. 9. Complete and balance: HCl + NaOH; H2SO4 + Ca(OH)2; CaO + H2CO3; Ca(OH)2 + CO2. 10. What is the difference between hydrides and hydracids in terms of the hydrogen oxidation number? Give an example of each.

1. The formal charge an atom would have if the bonding electrons went to the more electronegative atom. It is a bookkeeping number and not the real charge, especially in covalent compounds. 2. S +6; N +5; Mn +7; Cr +6; Cl +5; C -4; C +4; O -1; H -1. 3. NaCl -1 sodium chloride; NaClO +1 sodium hypochlorite; NaClO2 +3 sodium chlorite; NaClO3 +5 sodium chlorate; NaClO4 +7 sodium perchlorate. 4. Sodium oxide (disodium oxide); calcium oxide; iron(III) oxide; iron(II) oxide; carbon monoxide; sulfur trioxide; dinitrogen pentoxide; diphosphorus pentoxide; copper(I) oxide. 5. CaH2; HBr; Al2O3; FeCl3; CCl4; NaOH. 6. SO3 + H2O -> H2SO4 (sulfuric acid). Na2O + H2O -> 2 NaOH (sodium hydroxide). N2O3 + H2O -> 2 HNO2 (nitrous acid). CaO + H2O -> Ca(OH)2 (calcium hydroxide). 7. HClO hypochlorous (+1); HClO2 chlorous (+3); HClO3 chloric (+5); HClO4 perchloric (+7). 8. Sodium nitrite; sodium nitrate; potassium sulfite; calcium sulfate; potassium chlorate; calcium hypochlorite; potassium permanganate; potassium dichromate. 9. HCl + NaOH -> NaCl + H2O. H2SO4 + Ca(OH)2 -> CaSO4 + 2 H2O. CaO + H2CO3 -> CaCO3 + H2O. Ca(OH)2 + CO2 -> CaCO3 + H2O. 10. Hydrides: H is -1 (bonded to a metal), for example NaH, CaH2. Hydracids: H is +1 (bonded to a non-metal), for example HCl, H2S.