Unit 22 · Chemistry: Atoms, Bonding and Reactions

In 1909 Hans Geiger and Ernest Marsden, working under Ernest Rutherford in Manchester, fired alpha particles at a sheet of gold foil only a few hundred atoms thick. Alpha particles are fast, positively charged pieces of helium atoms thrown out by radioactive elements. The accepted picture of the atom at that time was a soft ball of positive charge with electrons dotted through it. If that picture were right, the heavy alpha particles should sail straight through, nudged only slightly.

Most of them did. But roughly one in several thousand bounced back at a large angle, some almost straight toward the source. A spread-out positive charge could not deflect anything that hard. Only a tiny, dense, positively charged core could push a passing alpha particle away so violently. The rare bounces were the important result, not the thousands of straight paths.

In 1911 Rutherford published the nuclear model. Nearly all of an atom’s mass and all of its positive charge sit in a nucleus roughly ten thousand times smaller across than the atom itself. The electrons occupy the rest of the space, which is almost entirely empty. If an atom were the size of a football stadium, the nucleus would be about the size of a marble at the center. Everything you touch is mostly empty space held together by electric forces.

Heat hydrogen gas in a glass tube with an electric current and it glows pink. Send that glow through a prism and you do not get a rainbow. You get four sharp visible lines: red, blue-green, violet and deep violet. This set of lines is hydrogen’s emission spectrum. Every element has its own set, which is why a flame test can identify an element by color. The lines never shift. Something inside the atom must be fixed.

In 1913 Niels Bohr offered an explanation. Suppose an electron can sit only at certain energy levels, like steps on a staircase rather than a ramp. An electron that absorbs energy jumps up to a higher step. When it falls back, it gives out the extra energy as a single packet of light, a photon. The photon’s energy equals the difference between the two steps, and that energy sets the color. Bigger drop, more energy, shorter wavelength.

Because the steps are fixed, only certain drops are possible, so only certain colors appear. Physicists say the energy is quantized. Bohr’s formula for hydrogen predicted the four visible lines and more lines in the ultraviolet and infrared. It also explained absorption. A cool gas in front of a bright source steals exactly the photons that lift its electrons up a step, leaving dark lines in the same places. That is what Kirchhoff saw in sunlight.

Bohr’s staircase worked beautifully for hydrogen and failed for nearly everything else. In 1926 Erwin Schrödinger wrote an equation that treated the electron as a wave spread through space. Its solutions are orbitals: three-dimensional regions where an electron is likely to be found. An orbital is not a path. It is more like a fuzzy cloud whose density shows probability. This quantum model is the one chemists use today, and it explains the periodic table.

Orbitals come in shapes and sizes. The s orbitals are spheres. The p orbitals look like two lobes on opposite sides of the nucleus, and there are three of them pointing along three directions. The d orbitals are more complicated, and there are five. Orbitals are grouped into shells numbered 1, 2, 3 and so on, with higher shells farther from the nucleus and higher in energy. Each orbital can hold at most two electrons, a rule called the Pauli exclusion principle.

An atom’s electron configuration lists which orbitals its electrons fill, starting from the lowest energy. Sodium, with 11 electrons, is 1s2 2s2 2p6 3s1: two in the first shell, eight in the second, and a single electron alone in the third. The electrons in the outermost shell are the valence electrons. They sit farthest from the nucleus and take part in bonding. Sodium’s single valence electron is why it reacts so eagerly, and why it gives the flame its yellow line.

In 1869 Dmitri Mendeleev, a chemistry professor in St. Petersburg, wrote each known element on a card with its atomic mass and its main properties. He laid the cards out in order of mass and noticed something. Properties repeated. After a soft, reactive metal like lithium came a run of elements that changed step by step, and then another soft, reactive metal, sodium. This repetition is the periodic law: arrange the elements in order and their properties recur at regular intervals.

Mendeleev did two bold things. Where the pattern demanded an element that no one had found, he left a gap. And he predicted the missing elements’ properties in detail: their masses, densities and the compounds they would form. In 1875 a French chemist discovered gallium, which matched Mendeleev’s ’eka-aluminum’. In 1886 germanium filled the gap he had called ’eka-silicon’. Predictions that come true are the strongest evidence that a pattern is real.

A few elements still sat out of order by mass. In 1913 Henry Moseley measured the X-rays given off by different metals and found that each element has a whole-number position, its atomic number, which equals the number of protons in the nucleus. Ordering by atomic number fixed every problem. Today’s table has vertical groups of elements with similar valence electrons and horizontal periods, each period filling one more shell.

Atoms have no hard edge, but chemists can measure a useful atomic radius from the distance between bonded nuclei. Two trends stand out. Going down a group, atoms get bigger, because each row adds a whole new shell farther from the nucleus. Going across a period from left to right, atoms get smaller, even though electrons are being added. That second trend surprises people. The reason is the nucleus.

Across a period, each new element adds one proton and one electron, but the electron goes into the same shell. Electrons in the same shell do a poor job of shielding each other from the nucleus. So the effective nuclear charge, the pull the outer electrons actually feel, rises across the period. The stronger pull draws the whole cloud inward. Fluorine is much smaller than lithium, even with more electrons.

The same idea explains ionization energy, the energy needed to remove an electron from a gas-phase atom. Small atoms with high effective nuclear charge hold their electrons tightly, so ionization energy rises across a period. Large atoms hold their outer electron loosely, so ionization energy drops down a group. Cesium, big and heavily shielded, gives up an electron more easily than any other stable element. Helium, tiny and unshielded, holds on hardest of all.

When two atoms share electrons, they rarely share fairly. Electronegativity measures how strongly an atom pulls shared electrons toward itself. In 1932 Linus Pauling built a scale from bond energies. Fluorine, small and with a high effective nuclear charge, sits at the top near 4.0. Cesium and francium, large and heavily shielded, sit at the bottom near 0.7. The trend follows ionization energy: up and to the right means a stronger pull.

The trend explains the personalities of whole groups. The alkali metals in group 1 have one loosely held valence electron and low electronegativity, so they give that electron away. Drop a piece of sodium in water and it fizzes and may burst into flame. The halogens in group 17 are one electron short of a full shell and pull hard, so they grab electrons from almost anything. Fluorine reacts with nearly every element. The noble gases in group 18 already have full shells and mostly ignore other atoms.

Comparing two electronegativities tells you what kind of bond to expect. A large difference, as between sodium and chlorine, means one atom takes the electron outright and an ionic bond forms. A small difference, as between two carbon atoms, means a fair share and a nonpolar covalent bond. In between, as with oxygen and hydrogen, the electrons are shared but lopsided. That lopsided sharing, as the next sections show, is behind most of water’s strange behavior.

A sodium atom that loses its valence electron becomes a sodium ion with a charge of 1+. A chlorine atom that gains it becomes a chloride ion with a charge of 1−. Opposite charges attract, and that attraction is an ionic bond. But the bond is not a pair. Each sodium ion pulls on every chloride ion near it, and each chloride pulls on every sodium. The ions stack into a repeating three-dimensional pattern called a crystal lattice, with each ion surrounded by six of the opposite charge.

The lattice explains the properties. Pulling millions of strong attractions apart takes a lot of energy, so ionic compounds have high melting points. Sodium chloride melts at about 801 °C. The solid does not conduct electricity, because the ions are locked in place. Melt it, or dissolve it in water, and the ions can move and carry current. A solution that conducts this way is called an electrolyte. That is why the sports drink and the ocean both conduct, and why wet skin makes electric shocks more dangerous.

Ionic solids are hard but brittle. Tap a salt crystal with a blade and it splits along flat planes instead of bending. Bending would slide one layer of ions over the next, bringing like charges face to face, and the layers push apart. Every winter, Chicago spreads salt on its roads because dissolved ions lower the freezing point of water. The same dissolved ions then corrode steel and concrete, a trade-off the city lives with each year.

Two chlorine atoms each need one more electron. Neither can take it from the other, since they pull equally. So they share. Each contributes one electron to a pair that sits between the two nuclei, and both atoms count the pair as their own. This shared pair is a covalent bond. In 1916 the American chemist Gilbert Lewis proposed that atoms bond by sharing electrons until each has eight in its outer shell, a guideline called the octet rule. Hydrogen is the exception; it is full with two.

A Lewis structure is a map of valence electrons. You count the valence electrons of every atom, connect the atoms with lines, each line standing for one shared pair, and then place the leftover electrons as dots so each atom reaches eight. The dots that are not shared are lone pairs. Water has two bonds and two lone pairs on oxygen. Ammonia has three bonds and one lone pair on nitrogen. Those lone pairs will matter for shape.

Sometimes one pair is not enough. Oxygen atoms share two pairs, a double bond. Nitrogen atoms share three, a triple bond. The number of shared pairs is the bond order. Higher bond order means a shorter, stronger bond. The nitrogen triple bond takes about 945 kJ to break per mole, one of the strongest bonds known. That strength is why the nitrogen in the air, nearly four fifths of every breath, is so hard for living things and chemists to use.

A copper wire, a steel beam and a gold ring share a way of holding together that is neither ionic nor covalent. Metal atoms hold their few valence electrons loosely. In a chunk of metal, those electrons leave their atoms and roam through the whole piece. What remains is a lattice of positive metal ions sitting in a sea of delocalized electrons, ones not tied to any one atom. The attraction between the ions and the electron sea is the metallic bond.

Free electrons carry charge, so metals are excellent electrical conductors, and they carry heat quickly for the same reason. Free electrons also absorb and re-emit light of nearly every color, which gives metals their shine. Most important for engineering, the bond is not directional. Push one layer of ions past the next and the electron sea simply flows around the new arrangement. The metal bends instead of shattering. A material that can be hammered into shape is malleable.

Pure metals are often too soft, so metalworkers mix them. An alloy is a metal blended with other elements. Steel is iron with a little carbon; the carbon atoms wedge between iron atoms and make layers harder to slide. Brass is copper and zinc. Chicago’s skyline stands on steel, and for most of the twentieth century steel mills lined the Calumet River on the city’s Southeast Side. Every skyscraper frame is a bet on the strength of the metallic bond.

A Lewis structure is flat, but molecules are not. In 1957 Ronald Gillespie and Ronald Nyholm turned a simple idea into a tool. Electron pairs repel each other, so the pairs around a central atom, both bonding pairs and lone pairs, spread out as far as they can. This is VSEPR, short for valence shell electron pair repulsion. Count the electron groups around the center, and you can predict the shape.

Two groups point in opposite directions, giving a straight line with a bond angle of 180°, as in carbon dioxide. Three groups form a flat triangle with 120° angles, as in boron trifluoride. Four groups reach into three dimensions and form a tetrahedral shape, a pyramid with a triangular base, with angles of 109.5°. Methane, CH4, is a perfect tetrahedron. Carbon’s tetrahedral bonding is the reason carbon can build the long, branching chains of living things.

Lone pairs count as groups but take up a bit more room, since they belong to one atom and spread wider. Ammonia has four groups, three bonds and one lone pair, so it is a squat pyramid with angles near 107°. Water has two bonds and two lone pairs, giving a bent shape with an angle near 104.5°. These angles are not guesses. Spectroscopy and X-ray diffraction measure them directly, and they match the predictions closely. A model that predicts numbers is a model worth keeping.

When oxygen shares electrons with hydrogen, it pulls them closer. The oxygen end of the bond becomes slightly negative and the hydrogen end slightly positive. A bond with two charged ends like this is a polar bond, and the separated charge is called a dipole. Any bond between atoms of different electronegativity is polar to some degree. A bond between identical atoms, like the two in oxygen gas, is nonpolar.

Whether the whole molecule is polar depends on shape. Carbon dioxide has two polar bonds, but the molecule is straight, so the two pulls point in opposite directions and cancel. Carbon dioxide is nonpolar. Water also has two polar bonds, but its bent shape means the pulls do not cancel. The oxygen side of the molecule carries a lasting negative end and the hydrogens a positive end. Water is strongly polar, and almost everything unusual about it follows from that.

Polarity controls solubility, what dissolves in what. Polar water molecules surround ions and other polar molecules, tugging them into solution, which is why salt and sugar dissolve. Oil and grease are made of nonpolar hydrocarbon chains that water cannot grip, so they float apart. The rule of thumb is ’like dissolves like’. Soap works because each of its molecules has a polar head that faces water and a nonpolar tail that buries itself in grease, dragging the grease into the wash water.

Covalent bonds hold a molecule together. A different, weaker set of attractions holds molecules to each other, and these intermolecular forces decide whether a substance is a gas, liquid or solid at room temperature. Boiling a liquid does not break any covalent bonds. It only pulls the molecules apart from one another. So the boiling point is a direct measure of how strongly the molecules cling.

The weakest attraction is the dispersion force. Electrons in any molecule slosh around, creating brief, flickering dipoles that tug on neighbors. Bigger molecules with more electrons have stronger dispersion forces, which is why methane is a gas and candle wax, built from the same kinds of atoms in longer chains, is a solid. Polar molecules add dipole-dipole attraction, with the positive end of one molecule facing the negative end of the next. A gecko’s foot clings to glass using dispersion forces alone, spread over millions of tiny hairs.

The strongest of these forces is the hydrogen bond. It forms when a hydrogen atom bonded to nitrogen, oxygen or fluorine is attracted to a lone pair on a nearby molecule. Water is the champion. Each molecule can form up to four hydrogen bonds, which is why water boils at 100 °C when its heavier chemical cousins boil far below zero. Hydrogen bonds also hold ice in an open lattice that floats, lock the two strands of DNA together, and fold proteins into their working shapes. Life runs on the weakest bond in this chapter.

Strike a match and wood becomes ash, smoke and heat. Leave a bicycle in the rain and shiny steel turns to orange rust. In every chemical reaction, bonds break and new bonds form, but the atoms themselves survive. Antoine Lavoisier established this in the 1780s by weighing everything that went into a sealed reaction and everything that came out. The totals matched. This is the law of conservation of mass, and it is why a chemical equation must be balanced, with the same count of each atom on both sides.

Chemists group reactions by their shape. In a synthesis reaction, two or more substances join to make one, as when hydrogen and oxygen form water. In decomposition, one substance splits into simpler ones, as when hydrogen peroxide breaks down into water and oxygen. In a single replacement, one element takes the place of another in a compound. In a double replacement, two compounds swap partners, often dropping an insoluble solid, called a precipitate, out of solution.

Combustion is a reaction with oxygen that releases heat and light. Burning methane gives carbon dioxide and water: CH4 + 2 O2 → CO2 + 2 H2O. Count the atoms. One carbon, four hydrogens and four oxygens on each side. The coefficient 2 in front of O2 is not decoration. It tells you that one methane molecule needs two oxygen molecules. Balanced equations let chemists predict how much product a given amount of fuel, ore or medicine will yield before anything is mixed.

Many reactions are really about electrons changing hands. When iron rusts, each iron atom gives up electrons and becomes an iron ion, while oxygen takes those electrons. Losing electrons is oxidation. Gaining electrons is reduction. The two always happen together, because an electron lost must go somewhere. Chemists call the pair a redox reaction. A handy memory aid is OIL RIG: oxidation is loss, reduction is gain.

To track electrons, chemists assign each atom an oxidation number, a bookkeeping charge. A free element is 0. In compounds, oxygen is usually −2 and hydrogen +1, and the numbers in a neutral compound add to zero. When iron metal (0) becomes iron in rust (+3), its oxidation number rose, so it was oxidized. Oxygen went from 0 to −2, so it was reduced. The substance that takes electrons, here oxygen, is the oxidizing agent. The one that gives them, iron, is the reducing agent.

Redox runs the world. Burning any fuel oxidizes carbon and hydrogen. Your cells oxidize glucose slowly and capture the released energy. Photosynthesis runs the same chemistry in reverse, reducing carbon dioxide into sugar using sunlight. A battery is a redox reaction with the electrons forced to travel through a wire on their way from one substance to the other. Bleach whitens by oxidizing colored molecules. Understanding who gives and who takes electrons is the first step in every one of these.

A hand warmer grows hot; a cold pack grows icy. Both are chemical reactions, and the difference lies in bonds. Breaking a bond always costs energy, because the atoms are being pulled apart against their attraction. Forming a bond always releases energy. If the new bonds are stronger than the old ones, the reaction gives off heat overall and is exothermic. If the new bonds are weaker, the reaction must take heat from its surroundings and is endothermic.

The energy of a chemical system is called its enthalpy, and the change in enthalpy during a reaction is written as ΔH. A negative ΔH means energy left the system as heat: exothermic. A positive ΔH means energy entered: endothermic. Burning methane has a ΔH of about −890 kJ per mole, a large release, which is why natural gas heats homes across Illinois. Dissolving ammonium nitrate in water is endothermic, and that is what makes a cold pack cold.

You can estimate ΔH from bond energies, the average energy needed to break a particular bond. Add up the energy to break every bond in the reactants, subtract the energy released by forming every bond in the products, and the result is roughly ΔH. Hydrogen and chlorine, for example, break an H-H bond and a Cl-Cl bond and form two much stronger H-Cl bonds, so the reaction is strongly exothermic. This is the accounting behind every fuel: the products, usually carbon dioxide and water, have exceptionally strong bonds.

Iron rusts over years; a match head burns in a second. Both are reactions with oxygen. The reaction rate, how fast reactants turn into products, varies enormously. Collision theory explains why. For molecules to react they must first collide. But most collisions are useless. The molecules must hit hard enough to start breaking bonds and must be lined up so the right atoms touch. Only these effective collisions count.

Anything that raises the number of effective collisions raises the rate. Higher concentration packs more molecules into the same space, so collisions happen more often. Higher temperature makes molecules move faster, so they collide more often and, more importantly, hit harder. Near room temperature, a rise of about 10 °C often doubles the rate of a reaction, which is why food spoils quickly on a summer counter and slowly in the refrigerator.

For solids, only the surface can react. Grinding a solid into powder multiplies its surface area, and the rate climbs. A log burns slowly; sawdust of the same wood can explode. Grain elevators across Illinois have been destroyed by dust explosions for exactly this reason: fine grain dust suspended in air presents so much surface that a spark sets off the whole cloud at once. Mills fight it with ventilation and by keeping dust from building up.

Gasoline sits in a tank for months without burning, even though burning would release a huge amount of energy. Why? Before the strong new bonds can form, the old bonds must start to break, and that costs energy up front. The minimum energy colliding molecules need to react is the activation energy. Think of a hill between the reactants’ valley and the products’ valley. A spark gives a few molecules enough energy to cross, and the heat they release pushes the rest over.

In 1889 Svante Arrhenius showed that reaction rate depends on the fraction of molecules with energy above this hill, and that this fraction rises steeply with temperature. That is the real reason heating speeds reactions so much. It is not just that molecules collide more often; a far larger share of them collide hard enough. A reaction with a high activation energy can be almost frozen at room temperature and fast at a few hundred degrees.

A catalyst is a substance that speeds a reaction without being used up. It works by offering a different route with a lower hill. The iron in the Haber process, the platinum and rhodium in a car’s catalytic converter that turn exhaust gases into less harmful ones, and the enzymes in your cells are all catalysts. An enzyme can make a reaction go millions of times faster at body temperature. Without catalysts, digestion would take years and the chemical industry would grind to a halt.

A balanced equation shows the start and the end, not the route. Most reactions do not happen in one collision. They proceed through a reaction mechanism, a sequence of elementary steps, each of them a single collision or a single bond breaking. Between steps, short-lived intermediates form and are used up. Chemists piece together mechanisms by measuring how the rate changes when they change each concentration.

In a multistep mechanism, one step is usually much slower than the others. That rate-determining step controls the overall rate, just as the slowest cashier controls how fast a line moves no matter how quick the others are. A rate law is an equation that links the rate to the concentrations of the substances involved in that slow step. If doubling the concentration of a reactant doubles the rate, the rate is first order in that reactant. If it quadruples, second order.

Mechanisms explained one of the great environmental problems of the late twentieth century. In 1974 Mario Molina and Sherwood Rowland showed that chlorine atoms released from CFC refrigerants high in the atmosphere destroy ozone in a two-step cycle. A chlorine atom takes an oxygen from ozone, then gives it up again and is regenerated. Because the chlorine is not consumed, one atom can destroy thousands of ozone molecules. That catalytic mechanism led to the 1987 Montreal Protocol, a treaty phasing out CFCs, and the ozone layer has slowly begun to recover.

Seal nitrogen dioxide, a brown gas, in a flask. The color fades partway as molecules pair up into colorless dinitrogen tetroxide, and then it stops fading. It never becomes fully colorless. Both gases are present, and their amounts hold steady. This is equilibrium. But the molecules have not stopped. The forward reaction continues, and so does the reverse. They just run at the same rate, so nothing changes overall. The state is dynamic, like a crowd where people enter and leave a room at exactly equal rates.

Every reversible reaction in a closed container heads toward equilibrium. Where it settles depends on the reaction. Some sit far toward products, some far toward reactants. Chemists describe the position with the equilibrium constant, K, the ratio of product concentrations to reactant concentrations at equilibrium, each raised to the power of its coefficient. A large K means mostly products. A tiny K means the reaction barely goes. K depends on temperature and nothing else.

Equilibrium is everywhere you look. Carbon dioxide dissolved in a soda is in equilibrium with the gas above it, until you open the cap and let the gas escape. Oxygen binds to hemoglobin in your lungs, where it is plentiful, and unbinds in your muscles, where it is scarce. Stalactites grow as dissolved calcium carbonate slowly comes out of equilibrium with cave water. Understanding equilibrium means understanding why reactions stop where they do, and what it would take to move them.

In 1884 the French chemist Henry Le Châtelier stated a rule that lets chemists steer reactions. If a system at equilibrium is disturbed, it shifts in the direction that partly counteracts the disturbance. This is Le Châtelier's principle. Add more reactant, and the system uses some of it up by making more product. Remove a product as it forms, and the reaction keeps making more to replace it. Industry uses that second trick constantly.

Pressure matters for gases. Squeeze a gas mixture and the equilibrium shifts toward whichever side has fewer molecules, because fewer molecules take up less room and relieve some of the pressure. In the ammonia reaction, four molecules of gas, one N2 and three H2, become two molecules of NH3. High pressure therefore pushes hard toward ammonia. That is why Haber and Bosch needed vessels rated for hundreds of atmospheres.

Temperature is the subtle one. Heat behaves like a product in an exothermic reaction. Raising the temperature is like adding product, so the equilibrium shifts back toward reactants. Lowering it shifts toward products. Ammonia formation is exothermic, so cooling favors ammonia. But temperature is also the only disturbance that changes K itself, and cold reactions are slow. A chemist who wants both a good yield and a fast rate is caught between kinetics and equilibrium. The next lesson shows how Haber cut that deal.

The ammonia reaction, N2 + 3 H2 ⇌ 2 NH3, is exothermic and shrinks four gas molecules to two. Le Châtelier’s principle says the best yield comes at low temperature and high pressure. But at low temperature the nitrogen triple bond almost never breaks; the reaction is far too slow to be useful. At high temperature the reaction is fast but the equilibrium lies toward the reactants. Neither extreme works. Haber’s insight was that the answer had to be a compromise.

Modern plants run at roughly 400 to 500 °C and at pressures of about 150 to 250 atmospheres. The temperature is high enough for a reasonable rate and low enough for a usable yield. An iron catalyst lowers the activation energy so the rate at that temperature is acceptable. The catalyst does not change the equilibrium position, only how quickly it is reached. Even so, only a fraction of the gas turns to ammonia on each pass.

The rest of the trick is engineering. The mixture leaving the reactor is cooled until ammonia liquefies and is drained off, which removes product and, by Le Châtelier’s principle, keeps pulling the reaction forward. The unreacted nitrogen and hydrogen are pumped back through in a loop. Nothing is wasted. The cost is energy. Making the hydrogen from natural gas and compressing the gases takes a large amount of fuel, and the process is estimated to use about 1 to 2 percent of the world’s energy supply. Feeding the world has never been free.

In 1800 Alessandro Volta stacked disks of zinc and silver separated by cloth soaked in salt water and produced the first steady electric current. He had built a redox reaction into a machine. In any electrochemical cell, oxidation happens at one electrode, the anode, and reduction happens at the other, the cathode. Electrons released at the anode cannot cross the cell directly. They must go around through a wire, and on the way they can light a bulb or run a phone.

Different metals give up electrons with different eagerness. Zinc loses electrons far more readily than copper, so in a zinc-copper cell zinc is the anode. Chemists rank this eagerness on a table of standard reduction potentials, measured in volts against hydrogen. The farther apart two half-reactions sit on the table, the higher the voltage of a cell built from them. A salt bridge or porous barrier lets ions flow between the two halves to keep charges balanced, or the current would stop within moments.

Every battery you own is a variation on Volta’s stack. An alkaline cell uses zinc and manganese dioxide and delivers about 1.5 volts. A car’s lead-acid battery uses lead and lead dioxide, about 2 volts per cell, six cells in series. A lithium-ion cell moves lithium ions between a graphite anode and a metal oxide cathode at about 3.7 volts, and because lithium is so light it packs great energy per kilogram. Its inventors shared the 2019 Nobel Prize in Chemistry. A rechargeable battery is one whose redox reaction can be driven backward by pushing current in.

Corrosion is redox you did not ask for. On a wet steel surface, one patch of iron becomes an anode and loses electrons, which travel through the metal to a nearby patch where oxygen and water accept them. Iron ions meet hydroxide, and rust, hydrated iron oxide, flakes off, exposing fresh metal. Dry iron barely rusts, because water is needed to carry ions. Salt water is worse, because dissolved ions conduct better and complete the circuit faster. Chicago’s road salt attacks cars and bridges for exactly this reason.

One defense is to keep water and oxygen away. Paint, oil and plastic coatings all work until they are scratched. A better defense uses chemistry. Coating steel with zinc, called galvanizing, protects even when the zinc is scratched, because zinc is more easily oxidized than iron. The zinc becomes the anode and corrodes instead, while the iron becomes a cathode and stays intact. Highway guardrails and steel fence posts are galvanized for this reason.

The same idea protects ships and pipelines. A block of a very reactive metal such as magnesium or zinc is bolted to the steel hull or buried next to the pipe. It corrodes on purpose and is replaced every few years, while the steel it is wired to is spared. This is a sacrificial anode. Stainless steel takes another route: its chromium forms a thin, tough oxide film that seals the surface. Aluminum protects itself the same way, which is why aluminum cans and window frames do not crumble even though aluminum is more reactive than iron.

A battery lets a favorable redox reaction make electricity. Electrolysis runs the idea in reverse. Push current from an outside source through a liquid containing ions. You can then force electrons onto a substance that does not want them, and pull electrons off one that does not want to give them. Pass current through water with a little dissolved salt. Hydrogen bubbles rise at the cathode and oxygen bubbles rise at the anode. There is twice as much hydrogen as oxygen by volume, exactly as the formula H2O predicts.

Michael Faraday worked out the rules in 1834. The amount of substance produced at an electrode is proportional to the amount of electric charge passed. That link between charge and moles was an early clue that electricity comes in particles, and that atoms carry fixed amounts of it. Today it lets engineers calculate exactly how much metal a given current will deposit. That calculation is the basis of electroplating, the coating of a cheap metal with a thin layer of chromium, silver or gold.

Electrolysis built the modern world’s lightest structural metal. Aluminum is the most abundant metal in Earth’s crust. But it clings to oxygen so tightly that no ordinary chemical reaction frees it, and in the 1850s it cost more than silver. In 1886 Charles Hall in Ohio and Paul Héroult in France independently found the answer. Aluminum oxide dissolved in molten cryolite could be electrolyzed to give pure metal. The Hall-Héroult process still makes all the world’s aluminum. Its enormous hunger for electricity is why recycling a can saves roughly 95 percent of the energy of making a new one.