Unit 19 · Physical Science: Matter and Reactions

Take a pinch of salt and imagine cutting one grain in half, then in half again, over and over. Eventually you reach a piece that cannot be split without turning into something else. That piece is an atom. Atoms are astonishingly small. A single drop of water holds more of them than there are stars in the Milky Way. Yet each one has parts, and those parts follow rules.

In the center sits the nucleus, a dense knot made of two kinds of particles. Protons carry a positive electric charge. Neutrons carry no charge at all. Nearly all of an atom’s mass is packed into this nucleus, but it takes up almost none of the atom’s space. If the atom were a football stadium, the nucleus would be a marble on the fifty-yard line.

Around the nucleus move the electrons. Each has a negative charge exactly equal and opposite to a proton’s, and a mass about 1,800 times smaller. In a neutral atom the number of electrons matches the number of protons, so the charges cancel. Electrons are what touch when two atoms meet, so they control almost everything about how an element behaves.

The one number that names an element is its proton count, the atomic number. Every atom with 6 protons is carbon. Every atom with 8 is oxygen. Change the proton count and you change the element itself. Change the number of neutrons or electrons and you still have the same element, just a different version of it.

Every carbon atom has 6 protons. But if you could weigh carbon atoms one at a time, you would find that they do not all weigh the same. Most have 6 neutrons. About one in a hundred has 7. A tiny fraction, roughly one in a trillion, has 8. These versions are isotopes of carbon. They are the same element, with the same chemistry, but a different mass.

Scientists name an isotope by its mass number, the protons and neutrons added together. Carbon with 6 protons and 6 neutrons is carbon-12. With 8 neutrons it is carbon-14. The atomic mass printed on a periodic table is an average of all the isotopes, weighted by how common each one is. That is why carbon’s listed mass is 12.011 rather than a whole number.

Some isotopes are stable and will sit unchanged forever. Others have a nucleus that is unbalanced, with too many or too few neutrons to hold together. These radioactive isotopes eventually change, shooting out a particle or a burst of energy. Carbon-14 is one of them. Its slow, steady decay is the clock used to date ancient wood and bone, which you will meet in the next chapter.

Isotopes matter in daily life more than most people realize. Doctors use a short-lived isotope of technetium to make images of the heart and bones. Smoke detectors in most homes contain a speck of americium-241. Water made with the heavy hydrogen isotope, deuterium, is used in some nuclear reactors. Same element, different neutron count, very different job.

Electrons do not orbit the nucleus like planets. They occupy fuzzy regions called energy levels or shells, and each shell holds only so many. The first shell, closest to the nucleus, holds 2 electrons. The second holds 8. The third holds 8 before the next shell starts to fill. Electrons take the lowest open spot first, the way water fills a glass from the bottom.

Sodium has 11 electrons. Two fill the first shell, eight fill the second, and one lone electron sits in the third. Chlorine has 17: two, then eight, then seven. Neon has 10: two and eight, with the outer shell exactly full. These outermost electrons are the valence electrons, and they are the ones that meet other atoms.

Here is the key pattern. Atoms with a full outer shell, like neon and argon, are stable and almost never react. Atoms with one or two electrons more than a full shell tend to give them away. Atoms one or two short of a full shell tend to take or share. Sodium wants to lose one. Chlorine wants to gain one. That is why they combine so eagerly to form salt.

Chemists draw this with a simple sketch: the element’s symbol surrounded by dots, one for each valence electron. Sodium gets one dot. Chlorine gets seven. Oxygen gets six. The sketch is a model, not a photograph, but it predicts an enormous amount of chemistry from a handful of dots.

Mendeleev ordered his elements by atomic weight, and it mostly worked. In 1913 a young English physicist, Henry Moseley, used X-rays to show that the true order was by atomic number, the proton count. A few elements that had seemed out of place snapped into position. The modern periodic table runs from hydrogen, with 1 proton, to elements with more than 100.

Each horizontal row is a period. Moving across a period, one proton and one electron are added at a time, and a new shell begins with each new row. Each vertical column is a group or family. Every element in a group has the same number of valence electrons. That is why they behave alike. Lithium, sodium and potassium each have one valence electron, and each reacts hard with water.

The table has neighborhoods. The far-left column, group 1, holds the alkali metals, soft and violently reactive. The column just before the right edge, group 17, holds the halogens, which grab electrons. The far-right column holds the noble gases, which barely react at all. The wide block in the middle holds the transition metals, including iron, copper and gold.

Metals sit on the left and center and make up most of the table. Nonmetals cluster in the upper right. A staircase of elements between them, such as silicon and germanium, are metalloids with a mix of properties. Germanium, the element Mendeleev predicted, sits on that staircase, which is exactly why it later became useful in electronics.

Drop a small piece of lithium in water and it fizzes. Sodium fizzes harder and can burst into flame. Potassium ignites almost at once. Rubidium and cesium react explosively. Going down group 1, each element has its one valence electron in a shell farther from the nucleus. That electron is held more loosely, so it leaves more easily. Looser electron, faster reaction.

The halogens in group 17 run the other way. Fluorine, at the top, is the most reactive nonmetal known; it attacks glass and most metals. Chlorine is a little tamer, bromine tamer still, iodine the mildest. Each wants one more electron to fill its shell. The smaller the atom, the closer that empty spot is to the pulling nucleus, so the top of the group grabs hardest.

This is reactivity: how readily an element takes part in a chemical change. It is not a random property. It comes from how tightly the nucleus holds its outer electrons, and that depends on the atom’s size and the number of protons pulling. Chemists use the term electronegativity for an atom’s pull on shared electrons. Fluorine has the highest value; the metals at the lower left have the lowest.

The noble gases in group 18 anchor the far right. Helium, neon and argon have full outer shells and almost no tendency to gain or lose anything. That is why helium fills balloons safely while hydrogen, which looks similar, once filled airships that burned. Knowing where an element sits tells you a great deal about how carefully to handle it.

You might expect an atom with more protons and electrons to be bigger. Across a period, the opposite happens. Lithium is larger than beryllium, which is larger than boron, and so on to neon. Each added proton pulls harder on the same shell, tightening the atom. Down a group, atoms do grow, because each new period adds a whole new shell farther out. Cesium is one of the largest atoms; helium is among the smallest.

Ionization energy is the energy it takes to pull one electron completely off an atom. Think of it as the price of stealing an electron. Sodium’s price is low, about 496 kilojoules per mole, because its lone outer electron is far out and shielded by inner shells. Neon’s price is high, over 2,000 kilojoules per mole, because its shell is full and close in.

The trend is a mirror of atomic size. Small atoms with many protons guard their electrons fiercely; large atoms with few outer electrons let them go cheaply. So ionization energy rises across a period and falls down a group. Fluorine and neon in the upper right are the hardest to ionize. Cesium in the lower left is the easiest.

These trends are why the table can predict behavior. A chemist who has never handled francium can still say with confidence that it is a very large atom, gives up its electron easily, and would react ferociously. Mendeleev predicted the densities and compounds of missing elements by exactly this kind of reasoning.

Put sodium metal and chlorine gas together and they react with a flash, producing a white solid: ordinary table salt. Sodium’s lone valence electron jumps to chlorine, which had room for exactly one more. Now sodium has a full outer shell but one fewer electron than protons, so it carries a positive charge. Chlorine has gained one and carries a negative charge. Charged atoms are called ions.

Opposite charges attract. The positive sodium ions and negative chloride ions pull toward each other from every direction and settle into a repeating three-dimensional pattern, a crystal lattice. Each sodium ion is surrounded by six chloride ions, and each chloride by six sodium ions. This electrical attraction holding the lattice together is an ionic bond.

Ionic compounds share a set of traits that follow from the lattice. They are hard and brittle: hit a salt crystal and it shatters along flat planes as like charges are shoved next to each other. They have high melting points, because every ion is held by many neighbors. Solid, they do not conduct electricity, since the ions cannot move. Melted or dissolved in water, they conduct well, because the ions are free.

Ionic bonds form between elements from opposite sides of the table: a metal with low ionization energy and a nonmetal with high electronegativity. Magnesium gives two electrons to oxygen to make magnesium oxide. Calcium gives two, one each, to two chlorine atoms to make calcium chloride, one of the salts spread on Chicago sidewalks before a snowstorm.

Two hydrogen atoms each have one electron and each want two. Neither can take from the other, because they pull equally. So they share. The two electrons spend their time between both nuclei, attracted to both, and hold the pair of atoms together. This shared pair is a covalent bond. The unit it makes, H2, is a molecule: a definite group of atoms bonded together.

Oxygen needs two more electrons, so it shares two pairs with another oxygen, a double bond. Nitrogen needs three and forms a triple bond, one of the strongest in chemistry. Water is one oxygen sharing a pair with each of two hydrogens. Methane, the main part of natural gas, is one carbon sharing a pair with each of four hydrogens. Chemists draw each shared pair as a short line between symbols.

Covalent bonds join nonmetals, which sit on the right side of the table. Because molecules are separate units rather than an endless lattice, covalent substances often have low melting points, and many are gases or liquids at room temperature. Water, carbon dioxide, oxygen and sugar are all covalent. They generally do not conduct electricity, since there are no free ions.

Sharing is not always equal. In water, oxygen pulls the shared electrons harder than hydrogen does, so the oxygen end of the molecule is slightly negative and the hydrogen ends slightly positive. This lopsided sharing is called polar. It is why water molecules cling to each other, why water has such a high boiling point for its size, and why it dissolves salt so well.

Copper wire carries electricity to every outlet in a building. Aluminum foil bends around a sandwich without cracking. A polished spoon reflects your face. Ionic and covalent bonds cannot explain any of this. Metals bond a third way. Their outer electrons are held so loosely that they leave their atoms altogether and drift through the entire piece of metal.

Picture a grid of positive metal ions sitting in a shared pool of moving electrons. The ions attract the electron pool, and the pool attracts the ions, so the whole solid holds together. This is a metallic bond. No single electron belongs to any single atom. Chemists call the pool a sea of electrons, and the picture explains three properties at once.

First, conductivity. Because the electrons are already free to move, a small push from a battery sends them flowing as an electric current. They also carry heat quickly, which is why a metal spoon in hot soup warms fast. Second, malleability. Hammer a metal and the ion layers slide past each other while the electron sea flows around them, so the metal bends instead of shattering like salt. Third, luster: the free electrons absorb and re-emit light, giving metals their shine.

Alloys are metals mixed with other elements to tune these properties. Steel is iron with a little carbon, whose atoms jam the sliding layers and make the metal stronger. Brass is copper and zinc. Bronze, the metal of ancient tools and statues, is copper mixed with tin. Same electron sea, but foreign atoms in the grid change how easily the layers slide.

A chemical formula is a recipe written in symbols. H2O means two hydrogen atoms bonded to one oxygen. CO2 means one carbon and two oxygens. NaCl means sodium and chlorine in a one-to-one ratio. The small numbers after a symbol count that atom; no number means one. The formula does not show how the atoms are arranged, but it fixes exactly what is inside.

Ionic compounds are named metal first, nonmetal second, with the nonmetal’s ending changed to -ide. Sodium and chlorine make sodium chloride. Magnesium and oxygen make magnesium oxide. The formula follows from the charges: the total positive charge must cancel the total negative. Magnesium ions are 2+, chloride ions are 1-, so it takes two chlorides for each magnesium: MgCl2.

Covalent compounds between two nonmetals use prefixes to count atoms, because the same pair can bond in more than one ratio. Carbon and oxygen make both carbon monoxide, CO, a poisonous gas, and carbon dioxide, CO2, the gas you breathe out. Mono- means one, di- means two, tri- means three, tetra- means four. Dinitrogen tetroxide is N2O4. The prefix mono- is dropped on the first element.

Some groups of atoms travel together as a charged unit, a polyatomic ion. Sulfate is SO4 with a 2- charge; nitrate is NO3 with a 1- charge; ammonium is NH4 with a 1+ charge. They are named as a unit: sodium sulfate, calcium nitrate, ammonium chloride. Fertilizer bags on Illinois farms list exactly these names, because the names tell a farmer precisely which ions the soil will receive.

Stir salt into water and it vanishes. Stir sand into water and it sinks. Both are hard, white-ish crystals. The difference is what water can do to their bonds. Water molecules are polar, with a slightly negative oxygen end and slightly positive hydrogen ends. When water meets a salt crystal, the negative ends crowd around the positive sodium ions, and the positive ends crowd around the chloride ions. The ions are tugged loose one by one. Salt dissolves.

Sand is silicon dioxide, a covalent network in which every silicon is bonded to four oxygens and every oxygen to two silicons, on and on through the whole grain. There are no separate ions for water to pull away, and the covalent bonds are far too strong for water to break. So sand stays sand. Oil will not dissolve salt either, because oil molecules are not polar and cannot surround an ion.

Carbon shows most clearly that structure decides properties. In graphite, the soft gray material in a pencil, each carbon bonds to three neighbors in flat sheets. The sheets slide over each other easily, which is why pencils leave marks. In diamond, each carbon bonds to four neighbors in a rigid three-dimensional network. To scratch a diamond you must break covalent bonds, so diamond is the hardest natural material. Same element, different bonding, opposite behavior.

This is the payoff of the whole chapter. Metals conduct because their electrons roam. Salts shatter and dissolve because they are lattices of ions. Molecular substances melt easily because the molecules are only weakly attracted to each other. Network solids like diamond and quartz are hard because bonds run through the whole crystal. If you know the bonding, you can predict the behavior, and if you know the behavior, you can reason back to the bonding.

Burn a log in a fireplace and by morning only a handful of ash remains. It looks as though most of the wood simply disappeared. In the 1770s the French chemist Antoine Lavoisier began weighing everything, including the gases, in sealed glass vessels before and after a reaction. He found that the total mass never changed. The wood had not vanished; it had become carbon dioxide and water vapor that drifted up the chimney.

This is the law of conservation of mass: in a chemical reaction, the total mass of the starting materials equals the total mass of the products. The reason is simple once you know about atoms. A reaction only rearranges them. Bonds between atoms break, and new bonds form, but every carbon, hydrogen and oxygen atom that went in comes out somewhere. Atoms are like letters; a reaction spells new words with the same letters.

Chemists write this rearrangement as a chemical equation. The reactants, the starting materials, go on the left. An arrow means reacts to form. The products go on the right. Burning methane, the main part of natural gas, is written CH4 + 2 O2 → CO2 + 2 H2O. One carbon, four hydrogens and four oxygens on the left; one carbon, four hydrogens and four oxygens on the right. Nothing lost.

Conservation of mass lets you do real accounting. If 16 grams of methane burn completely, they use 64 grams of oxygen and make exactly 80 grams of carbon dioxide and water combined. A car’s exhaust weighs more than the gasoline it burned because it includes the oxygen pulled from the air. The mass does not appear from nowhere; it was in the air all along.

Hydrogen gas burns in oxygen to make water. A first attempt at the equation, H2 + O2 → H2O, has a problem. Two oxygen atoms go in but only one comes out. That would violate conservation of mass. You cannot fix it by changing H2O to H2O2, because that is a different substance, hydrogen peroxide. The formulas are fixed by nature. Only the amounts can change.

The fix is a coefficient, a whole number in front of a formula that counts molecules. Write 2 H2O and now there are two oxygens on the right. But there are also four hydrogens on the right and only two on the left, so put a 2 in front of H2 as well. The balanced equation is 2 H2 + O2 → 2 H2O. Count it: four hydrogens and two oxygens on each side.

A reliable method: list each element and count it on both sides. Balance the element that appears in the fewest formulas first. Save hydrogen and oxygen for last, since they often show up in several places. Never touch a subscript. Adjust one coefficient, recount, and repeat until the columns match. Rust forming on an iron nail, 4 Fe + 3 O2 → 2 Fe2O3, takes two or three rounds.

The coefficients are more than bookkeeping. They give the ratio in which the substances react. Two molecules of hydrogen for every one of oxygen, or two truckloads for every one, the ratio holds. A rocket engineer sizing fuel tanks and a baker adjusting a recipe are doing the same arithmetic on different scales. The balanced equation is the recipe card.

Chemists sort reactions into a few families by the shape of the change. In synthesis, two or more substances join to make one. Hydrogen and oxygen make water. Iron and sulfur heated together make iron sulfide. In decomposition, one substance breaks into two or more. An electric current splits water into hydrogen and oxygen. Heat breaks calcium carbonate, which is limestone, into lime and carbon dioxide.

In a single replacement reaction, one element trades places with another in a compound. Drop a zinc strip into blue copper sulfate solution and the zinc dissolves while reddish copper metal appears: Zn + CuSO4 → ZnSO4 + Cu. A more reactive metal pushes a less reactive one out. This is how the reactivity trends from the periodic table show up in a beaker.

In a double replacement reaction, two compounds swap partners. Mix clear solutions of silver nitrate and sodium chloride and a white solid, silver chloride, appears instantly and sinks. The solid that forms from two solutions is called a precipitate. Water treatment plants, including those along Lake Michigan, use this kind of reaction to pull unwanted ions out of drinking water.

Combustion is a reaction with oxygen that releases heat and light. A fuel made of carbon and hydrogen burns to carbon dioxide and water. Fires, engines and furnaces all run on combustion. It is also the reaction that has changed the atmosphere most since the 1800s. Knowing the family of a reaction lets you predict its products before you run it, which is a large part of what chemistry is for.

A steel ship rusts over decades. Steel wool held in a flame burns in seconds. Both are iron reacting with oxygen. Why does one take a lifetime and the other a moment? The reaction rate depends on how often particles hit each other with enough energy to break bonds. Chemists call this collision theory. No collision, no reaction. Weak collision, no reaction either.

Every reaction has an energy hill to climb before it can go downhill. The height of the hill is the activation energy. Wood does not burst into flame on the shelf because room-temperature collisions rarely have enough energy to get over the hill. Touch a match to it and the local temperature rises, collisions get harder, and enough of them make it over that the reaction can heat the next bit of wood itself.

Four factors change the rate. Temperature: hotter particles move faster and collide harder and more often. Concentration: more particles in the same space means more collisions; a fire in pure oxygen is far fiercer than in air. Surface area: only the outside of a solid can react, so powder burns faster than a lump. Flour dust in a grain elevator can explode, which is why Illinois elevators control dust so carefully. Sawdust is dangerous where a plank is not.

The fourth factor is a catalyst, a substance that lowers the activation energy without being used up. It gives the particles an easier path over the hill. A car’s catalytic converter uses platinum to speed the conversion of exhaust gases. Every living cell runs on catalysts called enzymes, which let reactions that would take years happen in a fraction of a second at body temperature.

Crack open a chemical hand warmer and it grows hot in your pocket for hours. Squeeze an instant cold pack and it turns icy against a sprained ankle. In one, a reaction gives off energy; in the other, a process soaks it up. Both come down to the same accounting: it costs energy to break a bond, and forming a bond pays energy back. The difference between the cost and the payback is the energy you feel.

A reaction that releases more energy than it absorbs is exothermic. Burning methane breaks the C–H and O=O bonds, which costs energy, then forms C=O and O–H bonds, which returns more. The surplus leaves as heat and light. The products hold less stored energy than the reactants did. Rusting, combustion, and the reaction of sodium with water are all exothermic. Hand warmers use iron rusting fast.

A reaction that absorbs more energy than it releases is endothermic. Baking soda mixed with vinegar makes the cup feel cool. Photosynthesis is the largest endothermic reaction on Earth: a leaf uses sunlight to push carbon dioxide and water uphill into sugar and oxygen. The energy does not vanish. It is stored in the new bonds and comes back out when you burn the wood or digest the sugar.

Chemists draw this as an energy diagram. Reactants sit at one height, products at another, with the activation energy hill between them. An exothermic reaction ends lower than it started; an endothermic one ends higher. Energy, like mass, is conserved. What a reaction gives off, it took from its bonds; what it absorbs, it stores. The total in the universe is unchanged.

Open a bottle of soda and it fizzes. Inside the sealed bottle, carbon dioxide gas was dissolving into the liquid at exactly the same rate that dissolved gas was escaping back out. Nothing looked like it was happening, but two opposite processes were running at full speed and canceling. This balance is equilibrium. Open the cap, the pressure drops, and the balance tips toward escape.

Many chemical reactions are reversible. The products can react to re-form the reactants. When the forward and reverse rates become equal, the amounts stop changing, though individual molecules keep switching back and forth. It is like a crowded room with one door: people go in and out constantly, but the number inside holds steady.

Push on an equilibrium and it pushes back. Add more of a reactant and the forward reaction speeds up until a new balance forms with more product. Remove a product as it forms and the reaction keeps making more to replace it. Raise the temperature and the balance shifts toward the direction that absorbs heat. Chemists use this to steer reactions toward the product they want.

The most important example feeds the world. Nitrogen and hydrogen combine to make ammonia, the base of most fertilizer, but the reaction is reversible and slow. In the early 1900s Fritz Haber and Carl Bosch found that high pressure, moderate heat and an iron catalyst push the balance toward ammonia and make it fast enough to be practical. Much of the nitrogen in the corn and soybeans grown across Illinois arrives through this one equilibrium.

In 1896 the French physicist Henri Becquerel left a uranium salt on a wrapped photographic plate in a drawer. When he developed the plate, it was fogged as if by light. The uranium was giving off something invisible on its own, with no sunlight and no reaction. Marie and Pierre Curie tracked the effect, found two new elements that did it far more strongly, and named the phenomenon radioactivity.

Chemical reactions never touch the nucleus. Radioactive decay is a change in the nucleus itself. An unstable nucleus throws out a piece and becomes a different, more stable nucleus. Alpha decay ejects a bundle of two protons and two neutrons; the atom drops two places on the periodic table. Beta decay turns a neutron into a proton and shoots out an electron; the atom moves up one place. Gamma rays are pure energy, released as the nucleus settles.

You cannot predict when one particular nucleus will decay. But a large group decays at a perfectly steady rate, and that rate is described by the half-life: the time for half the atoms in a sample to decay. After one half-life, half remain. After two, a quarter. After three, an eighth. The half-life is a fixed property of the isotope. Heat, pressure and chemical bonding do not change it at all.

Half-lives range from fractions of a second to billions of years. Iodine-131, used in treating thyroid disease, has a half-life of about 8 days, so it is gone from a patient within weeks. Uranium-238 has a half-life of about 4.5 billion years, close to the age of Earth, so about half of the original uranium in the planet is still here. Long half-lives make radioactive waste a long-term problem; short ones make some isotopes safe to use in medicine.

High in the atmosphere, cosmic rays strike nitrogen atoms and turn a few of them into carbon-14. That radioactive carbon mixes into carbon dioxide, and plants breathe it in during photosynthesis. Animals eat the plants. So every living thing carries a small, steady trace of carbon-14, about one atom in a trillion, replaced as fast as it decays. The moment the organism dies, the intake stops and the clock starts.

Carbon-14 decays by beta emission back into nitrogen with a half-life of about 5,730 years. A piece of wood that has half the carbon-14 of a living tree died about 5,730 years ago. A quarter means about 11,460 years. By counting the decays in a sample, or by counting the atoms directly, a lab can date bone, wood, cloth, charcoal or seeds. This is radiocarbon dating.

The chemist Willard Libby worked out the method at the University of Chicago in the late 1940s. It earned him the Nobel Prize in 1960. He tested it on wood from Egyptian tombs whose ages were already known from written records. The dates matched. That check against independent evidence is what turned a clever idea into a trusted tool. Archaeologists use it to date the burned corn and charcoal of the ancient mound-builders at Cahokia, near Collinsville, Illinois.

The method has limits, and honest scientists state them. After about ten half-lives, roughly 50,000 years, too little carbon-14 remains to measure, so dinosaur bones cannot be carbon dated. The amount of carbon-14 in the air has varied a little over time, so raw dates are corrected using tree rings of known age. Other isotopes with longer half-lives, such as uranium and potassium, date rocks that are millions or billions of years old.

Burning a kilogram of coal releases enough energy to boil a bathtub or two of water. Splitting a kilogram of uranium releases millions of times more. The difference is that chemical reactions only rearrange electrons, while nuclear reactions rearrange the nucleus, where the forces are vastly stronger. In a nuclear change, a tiny fraction of the mass disappears, and it reappears as energy according to Einstein’s equation E = mc².

Fission is the splitting of a heavy nucleus. When a neutron is absorbed by uranium-235, the nucleus wobbles and breaks into two smaller nuclei, such as barium and krypton, and releases two or three more neutrons plus a burst of energy. If those neutrons hit other uranium-235 nuclei, the process repeats and grows. This is the chain reaction that Fermi’s pile achieved under Stagg Field.

Fusion is the opposite: joining light nuclei into a heavier one. In the core of the sun, at about 15 million degrees Celsius, hydrogen nuclei are slammed together so hard that they fuse into helium. The helium weighs slightly less than the hydrogen that made it, and that missing mass is the sunlight that reaches Illinois eight minutes later. Fusion powers every star and releases even more energy per kilogram than fission.

Fission is used in every nuclear power plant today because it can be started and controlled at ordinary temperatures. Fusion needs star-like heat and pressure to force the positive nuclei past their repulsion. Laboratories have produced brief bursts of fusion energy, and research reactors are being built, but no fusion plant yet delivers electricity to a grid. Its fuel is plentiful and its waste is far less long-lived, which is why the effort continues.

Chicago Pile-1 produced less power than a flashlight. A modern reactor produces enough electricity for a city, but the idea is the same. Uranium fuel, enriched so that it holds a few percent uranium-235, is sealed in metal rods. Water flows around the rods. It serves two jobs at once: it slows the neutrons, as Fermi’s graphite did, and it carries away the heat. A large reactor is mostly plumbing.

Control rods made of neutron-absorbing materials such as boron or cadmium slide in and out among the fuel. Pushed in, they soak up neutrons and the chain reaction slows. Pulled out, more neutrons survive to cause fission and the power rises. Operators keep the reaction exactly balanced, so that each fission leads on average to exactly one more. Push the rods fully in and the reaction stops within seconds.

The heat boils water into steam, either directly or through a second loop of pipes. The steam spins a turbine, which spins a generator that makes electricity. From the turbine onward, a nuclear plant is identical to a coal or gas plant. The difference is the source of heat: a nuclear reactor burns nothing, so it releases no carbon dioxide or smoke while running.

Illinois has more nuclear reactors than any other state. They supply roughly half of the electricity generated here. One of the first commercial nuclear plants in the country opened in 1960 at Dresden, near Morris, southwest of Chicago. The concrete domes visible from the interstate house reactors that trace directly back to the pile under Stagg Field.

Every energy source is a trade. The evidence on nuclear power’s benefits is clear. A reactor runs day and night regardless of wind or weather, and it emits no carbon dioxide while operating. A single fuel pellet the size of a fingertip yields about as much energy as a ton of coal. Measured per unit of electricity, the deaths caused by nuclear power over its history are far fewer than those from coal, which kills through air pollution and mining every year.

The risks are also real. Spent fuel stays dangerously radioactive for thousands of years, and the United States still has no permanent underground storage site; waste sits in pools and steel casks at the plants, including those in Illinois. Accidents are rare but severe. The 1986 explosion at Chernobyl in the Soviet Union spread radioactive material across Europe. The 2011 earthquake and tsunami at Fukushima in Japan caused three reactor cores to melt and forced the evacuation of more than a hundred thousand people.

Cost is a third factor. Nuclear plants are expensive and slow to build, often taking a decade, while the fuel itself is cheap. Wind and solar have become cheaper to build but produce nothing on a still night. Natural gas is cheap and flexible but releases carbon dioxide. Engineers compare these using the same tools you would: cost, benefit, and the constraints of place, time and safety.

Reasonable people looking at the same evidence reach different conclusions. They weigh a certain, ongoing harm against a rare, catastrophic one differently. A good argument in this debate states the evidence, acknowledges what it does not settle, and says which values tip the decision. That is not a weakness of science; it is where science hands the question to citizens.