Drop a pinch of salt into water and it vanishes. Drop a spoon of sugar in and it vanishes too. Yet the salt water conducts electricity and the sugar water does not, the salt can be pulled back out only by boiling the water away, and the two behave nothing alike in your cooking. Strike a match, and a stick of wood and a little powder become heat, light and smoke. Charge a phone, and something inside it moves from one side to the other and back again.
A chemist looks at all of these and asks the same two questions. What are the atoms doing with their outer electrons? And where is the energy going? General Chemistry is the course that teaches you to ask those questions in your sleep. Almost every later chemistry course, and most of biology, is these two ideas applied to a bigger puzzle.
An atom is a tiny nucleus of protons and neutrons with electrons around it. The nucleus almost never changes in a chemical reaction. That is the difference between chemistry and nuclear physics: chemistry is what happens when atoms trade, share or lend electrons and the nuclei stay put. (Radioactivity is the other business; Marie Curie won the 1903 Nobel Prize in Physics and the 1911 prize in Chemistry, the second for discovering radium and polonium, and her work sits on the border between them.)
Electrons are not little planets. Quantum mechanics (see Quantum Mechanics, 34/83, for the real thing) says each electron occupies an orbital, a cloud-shaped region where it is likely to be found, and each orbital has a fixed energy. Orbitals come in shells numbered 1, 2, 3 and so on, and in shapes called s, p, d and f. An orbital holds at most two electrons. Fill them from the lowest energy up, and you get the atom's electron configuration.
Try it. Slide to carbon (6): 1s² 2s² 2p², so four electrons in the outer shell. Oxygen (8) has six. Sodium (11) has one, alone in a third shell. Sodium's single lonely outer electron and chlorine's seven, one short of a full eight, are why the two meet so eagerly. At 24 (chromium) and 29 (copper) you will see the table's first small exceptions: a half-full or full 3d subshell is slightly more stable than the tidy filling order predicts.
Configurations by the Madelung filling order (1s 2s 2p 3s 3p 4s 3d 4p) with the chromium and copper exceptions; shell counts are the sum over subshells. Shells are drawn as rings for counting, not as orbits.
In 1869 Dmitri Mendeleev presented a table that arranged the known elements by atomic weight and, crucially, left gaps where the pattern said an element should exist, and described what the missing ones would be like. The gaps were later filled by gallium, scandium and germanium. The modern table is the same idea, ordered by atomic number, and now we know what the rows are: each row is a shell filling up, and each column holds atoms with the same number of outer electrons, which is why they behave alike. (Mendeleev was not the only person to see the pattern; Lothar Meyer published a similar one around the same time. Mendeleev's gift was trusting the pattern enough to predict.)
Three trends do most of the work, and one picture explains all three. An outer electron feels the nucleus's pull, partly cancelled by the electrons between. Move across a row and the nuclear charge grows while the new electrons go into the same shell, so the pull gets stronger and the atom shrinks. Move down a column and a whole new shell is added, so the atom is bigger and its outer electron is easier to remove.
Data: first ionisation energies in kJ/mol from the NIST Atomic Spectra Database values (via Wikipedia's data page); Pauling electronegativities; covalent radii from Cordero et al., Dalton Trans. 2008 (carbon sp3; Mn, Fe, Co low-spin). Noble gases He, Ne and Ar have no Pauling value.
Look at the ionisation energy plot: a sawtooth. Each peak is a noble gas, whose shell is full and which holds its electrons tightly; each trough is the alkali metal that follows, with one electron far out in a new shell. Sodium needs about 496 kJ per mole of electrons to lose one, neon needs 2081. That sawtooth is why sodium is violently reactive and neon is not reactive at all. Everything about metals, nonmetals and the families in between falls out of it.
Atoms bond because the shared or transferred arrangement has lower energy than separate atoms. Which kind of bond forms depends on how unevenly the atoms pull. Electronegativity is a number for an atom's pull on the electrons it shares; Linus Pauling's scale runs from 0.82 for potassium to 3.98 for fluorine. If two atoms pull almost equally (hydrogen with hydrogen), the electrons are shared: a covalent bond. If one pulls vastly harder (chlorine versus sodium), it takes the electron outright, and the two resulting ions, one plus and one minus, stick together by plain electrical attraction: an ionic bond. In between lies the common case, the polar covalent bond, where the electrons are shared but sit closer to one atom. Water's O–H bond is the standard example.
Pauling values (Z = 1 to 36, noble gases He, Ne and Ar omitted). Two metals are labelled metallic instead.
The result matters outside the lab. Salt dissolves in water because water's polar O–H bonds make its oxygen end slightly negative and its hydrogen ends slightly positive, and those poles pry ions out of the crystal. Oil and water do not mix because oil has no poles to grip. "Like dissolves like" is bond polarity in one sentence.
A molecule is not a flat diagram. Water is bent, not straight; ammonia is a little pyramid. One rule predicts most of it, the VSEPR rule (valence-shell electron-pair repulsion). Count the electron domains on the central atom, meaning each bond (single, double or triple counts as one) and each lone pair, since each domain is a cloud of negative charge. They repel, so they spread as far apart as they can on a sphere: two domains sit 180° apart, three at 120°, four at 109.5°, five and six in the arrangements below. The molecule's shape is then the arrangement of the atoms only; the lone pairs are invisible in the name but they still take up room.
That is why water is bent: oxygen has four domains (two bonds and two lone pairs), arranged like a tetrahedron, but only two corners hold atoms. Lone pairs are slightly bulkier than bonds, so they squeeze the angle: 109.5° in methane, about 107° in ammonia, about 104.5° in water. Water's bend is why it is polar and why ice floats, which is why lakes freeze from the top down.
Directions are the ideal geometry for each number of domains; lone pairs fill the equatorial positions first in the five-domain case and opposite positions in the six-domain case, as VSEPR prescribes. Angles are textbook values; real molecules deviate by a few degrees. The tool stops at six domains (the common textbook range).
You cannot count atoms, but you can weigh them, and the mole turns one into the other. A mole is a fixed number of things, exactly 6.02214076 × 10²³ since the 2019 redefinition of the SI. The point of choosing that number is that one mole of any element weighs, in grams, its atomic weight: carbon about 12.011 g, oxygen 15.999 g (the IUPAC standard atomic weights, rounded). So grams become moles by dividing by the atomic or molar weight, and a balanced equation is a recipe in moles.
A worked number. One mole of propane (the gas in a camping stove) weighs 3 × 12.011 + 8 × 1.008 = 44.097 g. It needs five moles of oxygen, 159.99 g, and produces three moles of carbon dioxide (132.03 g) and four of water (72.06 g). In: 204.09 g. Out: 204.09 g. Nothing is lost. If there is less oxygen than five to one, the oxygen is the limiting reagent: the reaction stops when it runs out, some propane is left, and (in a real stove) you get soot and poisonous carbon monoxide, the reason you never burn a stove indoors without ventilation. Because air is about 21% oxygen, burning a volume of propane needs roughly 24 volumes of air.
Molar masses from the IUPAC abridged standard atomic weights (C 12.011, H 1.008, O 15.999). Assumes complete combustion to CO₂ and H₂O.
Why does propane burn and release heat? Because the products' bonds are stronger than the reactants'. Breaking a bond always costs energy, since you pull two atoms against their attraction. Making a bond always releases it. A reaction's energy change is the bill for breaking all the old bonds minus the payment for making all the new ones. If you get paid more than you spent, heat comes out: exothermic. Otherwise it must be supplied: endothermic.
Methane burning: break four C–H and two O=O (2648 kJ), make two C=O and four O–H (3450 kJ), net −802 kJ per mole: exothermic. The measured value is −802.5. That match is partly luck, since the numbers are averages over many molecules, as the hydrogenation of ethene (−123 estimated against −136 measured) shows.
Average bond enthalpies in kJ/mol as in common textbook tables (the LibreTexts table lists e.g. C–H 415, H–Cl 432; C=O is the value for CO₂). Measured ΔH from standard enthalpies of formation, gaseous water. N₂ + 3H₂ → 2NH₃ is the Haber reaction (measured −92.28 kJ per mole of N₂ is the quoted standard value).
An acid, at this level, is a molecule that hands a hydrogen ion (a proton) to water; a base takes one. The strength of an acid is not about how concentrated it is but about how readily it lets go. A strong acid such as hydrochloric acid gives up essentially all its protons. A weak one such as the acetic acid in vinegar gives up only a small fraction, which is why vinegar is safe to put on a salad and the same concentration of hydrochloric acid is not. The pH is minus the base-10 logarithm of the hydrogen-ion concentration, so each step down in pH is ten times more acidic.
Worked number: 0.1 M hydrochloric acid has pH 1.00. 0.1 M acetic acid, with an acid constant of 1.75 × 10⁻⁵ (pKa 4.76), has pH 2.88, and only about 1.3% of its molecules have let go. Dilute it a thousand times and the fraction rises to roughly a third, an early hint of the equilibrium ideas of Equilibrium & Kinetics (35/37). And dilute even a strong acid to 10⁻⁸ M and the pH is 6.98, not 8: water itself supplies hydrogen ions, and you cannot make an acid into a base by adding water.
Solved exactly from the charge balance [H+] = [A-] + [OH-] with Kw = 10-14 and Ka = 10-4.756 for acetic acid at 25 °C. Ideal solution (activities ignored).
Chemistry is electron bookkeeping and energy bookkeeping. An atom's outer electrons set its place on the table, how it bonds and what shape its molecules take; the mole lets you weigh your way into those counts; and the balance between breaking and making bonds says whether a reaction gives heat or takes it.
Try this: this week, look at the back of a food label or a bottle of vinegar and find one chemical name. Write its formula and count its outer electrons. (Kitchen-safe, and no mixing needed: just a pen.)
Sources. Electron configurations, VSEPR, the mole, bond enthalpies and pH are standard first-year textbook chemistry (OpenStax Chemistry 2e). Pauling electronegativities, first ionisation energies (NIST Atomic Spectra Database) and Cordero covalent radii (Dalton Trans. 2008, 2832) as tabulated on Wikipedia's data pages. Mendeleev presented his table to the Russian Chemical Society in 1869 (Wikipedia, Dmitri Mendeleev). Marie Curie: Nobel Prizes in Physics 1903 and Chemistry 1911 (nobelprize.org via Wikipedia). SI mole redefinition 2019 (BIPM). This is a textbook summary, not laboratory instruction. Go deeper.