Unit 1 · 7–9% of the exam
Atomic Structure and Properties
Unit 1 links the mole to real masses and builds the modern picture of the atom: isotopes, electron configurations and the Coulombic attractions that explain periodic trends. Nearly every later unit leans on mole conversions and on explaining properties with nuclear charge, distance and shielding.
Electron configuration builder
Add electrons one at a time and watch the Aufbau order, Hund's rule and the Cr/Cu exceptions play out.
Topics in this unit
1.1
Moles and Molar Mass
The mole connects the mass of a sample you can weigh to the number of particles in it: n = m/M and N = n × Nₐ.
1.2
Mass Spectroscopy of Elements
A mass spectrum shows the masses and relative abundances of an element's isotopes; their weighted average is the average atomic mass on the periodic table.
1.3
Elemental Composition of Pure Substances
A pure compound has a fixed mass ratio of its elements (law of definite proportions), so percent composition leads to its empirical formula.
1.4
Composition of Mixtures
Mixtures have variable composition; elemental analysis of a mixture lets you find how much of each component is present.
1.5
Atomic Structure and Electron Configuration
Electrons fill subshells in order of increasing energy (1s 2s 2p 3s 3p 4s 3d…), and Coulomb's law explains why electrons nearer the nucleus are held more tightly.
1.6
Photoelectron Spectroscopy
In PES, each peak is a subshell: its position gives the binding energy and its height is proportional to the number of electrons in that subshell.
1.7
Periodic Trends
Atomic radius, ionization energy and electronegativity trends follow from effective nuclear charge and the distance of valence electrons from the nucleus.
1.8
Valence Electrons and Ionic Compounds
Elements gain or lose electrons to reach a noble-gas configuration, so their position in the table predicts ion charges and the formulas of ionic compounds.
Must know
- n = m / M (moles = mass ÷ molar mass)
- N = n × Nₐ, where Nₐ = 6.022 × 1023 mol−1
- Average atomic mass = Σ (isotope mass × fractional abundance)
- Coulomb's law: force ∝ q1q2 / r2 (greater charge, smaller distance → stronger attraction)
- Aufbau order: 1s 2s 2p 3s 3p 4s 3d 4p…; transition-metal atoms lose 4s electrons before 3d
- PES: higher binding energy = closer to the nucleus; peak height ∝ number of electrons
- Across a period: radius ↓, first ionization energy ↑, electronegativity ↑ (Zeff increases, same shell)
- Down a group: radius ↑, first ionization energy ↓ (valence electrons in a higher shell, farther away)
Common mix-ups
Average atomic mass is the mass of the most common isotope.
It is a weighted average of all isotopes, so it may not match any single isotope's mass (chlorine is 35.45, but no atom has that mass).
Atoms get bigger across a period because they have more electrons.
Across a period, added electrons go into the same shell while the nuclear charge rises, so Zeff increases and the atoms get smaller.
Fe2+ is [Ar]4s23d4 because the last electrons added were 3d.
When transition metals form cations, the 4s electrons are removed first, so Fe2+ is [Ar]3d6.
Practice questions
AP-style multiple choice. Every answer comes with the reasoning, including why the tempting wrong answer is wrong.
How many moles of water are in 9.0 g of H2O (M = 18.02 g/mol)?
Free-response practice
A 2.50 g sample of a mixture of NaCl and sand is analyzed and found to contain 0.910 g of chlorine. (a) Calculate the mass of NaCl in the sample. (b) Calculate the mass percent of NaCl in the mixture. (c) The first ionization energy of Mg is greater than that of Na. Explain this using principles of atomic structure.
Flashcards
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See it in real life
Everyday reactions that use Unit 1 ideas, animated atom by atom.