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Unit 1 · Atomic Structure & Properties Unit Hub Flashcards Cheat Sheet Essentials Visual Review MC Practice SAQ Practice

AP Chemistry Unit 1 Visual Review

A topic-by-topic visual walkthrough of Atomic Structure & Properties — moles, mass spectrometry, electron configuration, PES, and periodic trends.

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TOPIC 1.1 Moles & Molar Mass The mole — chemistry's counting unit A mole is just a number, like "a dozen" — but huge: 1 mol = 6.022 × 10²³ particles That's Avogadro's number (Nₐ). It lets us convert between countable particles and weighable grams. Molar mass (M) The mass of one mole of a substance, in grams per mole. Read it straight off the periodic table (atomic mass) and add up the atoms in a formula. Example: H₂O = 2(1.01) + 16.00 = 18.02 g/mol CO₂ = 12.01 + 2(16.00) = 44.01 g/mol The key formula & a worked example n = mass ÷ molar mass n = moles · mass in g · M in g/mol Rearrange to find any one of the three. Q: How many moles are in 36.0 g of water? n = 36.0 g ÷ 18.02 g/mol = 2.00 mol H₂O That's 2.00 × 6.022 × 10²³ = 1.20 × 10²⁴ molecules. Always track units — they tell you whether to multiply or divide. The mole bridges grams ↔ particles: n = mass ÷ molar mass, and 1 mol = 6.022 × 10²³. The Review Hub · AP Chemistry Unit 1 TOPIC 1.2 Mass Spectra of Elements MASS SPECTRUM OF CHLORINE 75.8% ³⁵Cl 24.2% ³⁷Cl relative abundance mass-to-charge (m/z) → What a mass spectrum shows A mass spectrometer separates a sample's isotopes by their mass. Each PEAK is an isotope; the peak's HEIGHT is that isotope's relative abundance in nature. Average atomic mass The periodic table's atomic mass is a WEIGHTED average of all isotopes, by abundance: avg = Σ (isotope mass × fractional abundance) Worked example — chlorine avg = (35)(0.758) + (37)(0.242) = 26.5 + 8.95 = 35.5 amu — matching the periodic-table value for Cl. The average sits closer to the more-abundant isotope (³⁵Cl), which is why it's ~35.5, not 36. Peaks = isotopes, height = abundance; the average atomic mass is their weighted average. The Review Hub · AP Chemistry Unit 1 TOPIC 1.3 Elemental Composition of Pure Substances Percent composition The mass % of each element in a compound: % = (mass of element ÷ molar mass) × 100 Every pure compound has a FIXED, definite composition. Empirical vs. molecular formula Empirical — the simplest whole-number ratio of atoms. Molecular — the actual number of atoms in a molecule. Molecular is always a whole-number multiple of empirical. Example: glucose is C₆H₁₂O₆ (molecular), CH₂O (empirical). Finding an empirical formula from percent mass A compound is 40.0% C, 6.7% H, 53.3% O by mass. Find its empirical formula. 1. Assume 100 g → grams = percents: 40.0 g C, 6.7 g H, 53.3 g O 2. Convert to moles: C = 40.0/12.01 = 3.33 · H = 6.7/1.01 = 6.63 · O = 53.3/16.00 = 3.33 3. Divide by the smallest (3.33): C = 1 · H = 2 · O = 1 4. Empirical formula = CH₂O If you also know the molar mass (180 g/mol), divide by empirical mass (30) → ×6 → molecular formula C₆H₁₂O₆. Percent mass → moles → divide by smallest = empirical formula (the simplest atom ratio). The Review Hub · AP Chemistry Unit 1 TOPIC 1.4 Composition of Mixtures Pure substance vs. mixture Pure substance — fixed composition (an element or compound). Every sample is identical. Mixture — two or more substances physically combined in VARIABLE proportions; each keeps its own properties. Mixtures can be separated by physical means (no reaction). Homogeneous vs. heterogeneous Homogeneous — uniform throughout; a solution (e.g., salt water, air, brass). Heterogeneous — non-uniform; you can see distinct parts (e.g., sand in water, oil and vinegar). Composition of a mixture is not fixed — it can vary. Calculating mixture composition Because a mixture's parts don't react, its total mass is just the SUM of the components. To find the mass percent of a component: mass % = (mass of component ÷ total mass) × 100. Example: A 50.0 g solution contains 8.0 g of NaCl dissolved in water. mass % NaCl = (8.0 g ÷ 50.0 g) × 100 = 16% NaCl For moles of a component, use its own molar mass — never the whole mixture's. On the exam, elemental analysis of a mixture is a common quantitative task. Mixtures have variable composition and separate physically; pure substances are fixed. The Review Hub · AP Chemistry Unit 1 TOPIC 1.5 Atomic Structure & Electron Configuration nucleus (p⁺ + n⁰) · electron shells electrons fill from the inside out The three subatomic particles Protons (+, in nucleus) set the element & atomic number. Neutrons (0, in nucleus) add mass; vary → isotopes. Electrons (−, in shells) do the chemistry. Three rules for filling orbitals Aufbau — fill the lowest-energy orbitals first (1s, 2s, 2p…). Pauli exclusion — max 2 electrons per orbital, opposite spins. Hund's rule — fill each orbital in a subshell singly before pairing. Subshells hold: s = 2, p = 6, d = 10, f = 14 electrons. Example configurations Oxygen (8 e⁻): 1s² 2s² 2p⁴ · Iron (26 e⁻): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶ · shorthand [Ar] 4s² 3d⁶ Core electrons are shielded; valence (outer) electrons determine reactivity. Electrons fill orbitals by Aufbau, Pauli, and Hund — the configuration drives an element's chemistry. The Review Hub · AP Chemistry Unit 1 TOPIC 1.6 Photoelectron Spectroscopy (PES) PES SPECTRUM OF NITROGEN (1s² 2s² 2p³) 1s² 2s² 2p³ # of electrons ← higher binding energy (closer to nucleus) How to read a PES spectrum Each peak = one subshell (1s, 2s, 2p…). Peak height = the number of electrons in that subshell. Peak position = binding energy (how tightly held). Binding energy = distance from nucleus Inner electrons (1s) are held most tightly → HIGH binding energy (plotted on the left). Outer electrons are held loosely → LOW binding energy (right). PES confirms the shell/subshell model directly. In PES: peaks = subshells, height = # electrons, position = binding energy. The Review Hub · AP Chemistry Unit 1 TOPIC 1.7 Periodic Trends PERIODIC TABLE radius ↓ · IE, EN ↑ (→ across) radius ↑ (↓ down) trends move across periods & down groups The driving force: Coulomb's law Attraction between the nucleus and electrons grows with a higher effective nuclear charge (Zeff) and shrinks with distance. Inner electrons SHIELD outer ones from the nucleus. The three trends Atomic radius: ↑ down a group, ↓ across a period (→) Ionization energy: ↓ down, ↑ across (opposite of radius) Electronegativity: ↓ down, ↑ across (highest at F, top-right) Smaller atom + higher Zeff = electrons held more tightly. Why radius decreases across a period Moving left→right, protons are added to the same shell. Higher nuclear charge pulls the electron cloud IN, so atoms get smaller — even though electron count rises. Cations are smaller than their atoms; anions are larger. Trends come from Coulombic attraction: radius shrinks →, while IE & electronegativity rise →. The Review Hub · AP Chemistry Unit 1 TOPIC 1.8 Valence Electrons & Ionic Compounds Valence electrons drive bonding Valence electrons are the outermost electrons — the ones involved in bonding. For main-group elements, the group number gives the count (e.g., O in group 16 has 6). Atoms react to reach a full outer shell (octet, 8 electrons). Forming ions Metals LOSE electrons → positive cations (e.g., Na → Na⁺). Nonmetals GAIN electrons → negative anions (Cl → Cl⁻). Opposite charges attract to form an ionic compound — electrons are TRANSFERRED, not shared. Coulombic attraction & lattice energy The strength of the attraction between two ions follows Coulomb's law: force ∝ (q₁ × q₂) ÷ r² q = ion charges · r = distance between ions Bigger charges & smaller ions → stronger bond. So: higher charge and smaller radius → higher lattice energy → higher melting point. Example: MgO (charges +2/−2) has a much higher melting point than NaCl (charges +1/−1). Lattice energy = energy released when gaseous ions form a solid ionic lattice; it measures ionic bond strength. Metals + nonmetals transfer electrons to form ions; higher charge & smaller ions bond stronger. The Review Hub · AP Chemistry Unit 1
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This is great for review the night before the exam — fast, visual, and covers everything you need to remember about Unit 1's atomic structure content.