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Unit 1
7–9% of examEverything in chemistry starts with atoms. In this unit you learn how chemists count atoms using the mole, how to work out what a sample is made of from its mass, and how the way electrons are arranged around the nucleus explains the patterns in the periodic table.
Longer videos that cover the whole unit. Good for a first pass or a final review.
Atoms are far too small to count one by one, so chemists count them in moles: one mole is 6.022 × 10²³ particles (Avogadro's number). The molar mass of a substance in grams per mole has the same number as the average mass of one particle in atomic mass units (amu), so you can convert between grams, moles and particles.
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A mass spectrum of an element shows one peak for each isotope: where the peak sits gives the isotope's mass, and how tall it is shows how common that isotope is. The element's average atomic mass is a weighted average of its isotope masses, so it sits nearest the masses of the most common isotopes.
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A pure compound always contains the same elements in the same proportions by mass (the law of definite proportions). From the percent of each element by mass you can work out the empirical formula, the simplest whole-number ratio of atoms in the compound.
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Unlike a pure substance, a mixture can contain its parts in any proportion. By measuring how much of one element is in a sample (elemental analysis), you can work out how much of each substance is present and how pure the sample is.
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An atom is a tiny, positive nucleus of protons and neutrons surrounded by electrons in shells and subshells. You write ground-state electron configurations, like 1s² 2s² 2p⁶ for neon, by filling the lowest-energy subshells first (the Aufbau principle), and Coulomb's law explains why electrons closer to a more highly charged nucleus are harder to remove.
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Photoelectron spectroscopy (PES) uses high-energy light to knock electrons out of atoms and measures how much energy it took to remove them. Each peak stands for a subshell: its position shows the binding energy, and its relative height shows how many electrons are in that subshell, so a PES spectrum is direct evidence for electron configurations.
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Atomic radius, ionization energy, electron affinity and electronegativity all change in predictable ways across the periodic table. Across a period, the effective nuclear charge rises, so atoms get smaller and hold their electrons more tightly; down a group, outer electrons sit in shells farther from the nucleus and are more shielded, so atoms get bigger and lose outer electrons more easily.
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Valence electrons decide how an element bonds, so elements in the same group form similar compounds. You can predict an ion's usual charge from its group: group 1 metals form 1+ ions and group 17 halogens form 1− ions, which is why NaCl and KBr have the same 1:1 formula.
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