Atomic Structure and the Periodic Table
Every chemistry course starts with atoms. Mastering atomic structure, electron configuration, and periodic trends builds the foundation for bonding, reactions, and everything that follows.
Core Atomic Components
Atoms contain protons (positive charge, nucleus), neutrons (neutral, nucleus), and electrons (negative charge, orbitals). The atomic number tells you the number of protons. The mass number equals protons plus neutrons. Isotopes are atoms of the same element with different neutron counts.
Electron Organization Rules
Electron configuration follows three key principles. The Aufbau principle means you fill lowest energy orbitals first. The Pauli exclusion principle limits each orbital to two electrons with opposite spins. Hund's rule says you singly occupy orbitals before pairing electrons. The order is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p and so on.
Periodic Law and Trends
The periodic law states that element properties repeat periodically with atomic number. Periods are rows (same energy level). Groups are columns (same valence electron count).
Key trends moving left to right across a period and top to bottom down a group:
- Atomic radius decreases left-to-right (more protons pull electrons closer); increases down groups (new energy levels)
- Ionization energy (energy to remove an electron) increases left-to-right; decreases down
- Electron affinity (energy change when gaining an electron) generally becomes more negative moving left-to-right
- Electronegativity (tendency to attract shared electrons) increases left-to-right; decreases down
Critical Ion Groups to Memorize
Lock in these common ions early. They appear in nearly every reaction.
- Alkali metals (Group 1) form +1 ions: Li+, Na+, K+
- Alkaline earth metals (Group 2) form +2 ions: Mg2+, Ca2+, Ba2+
- Halogens form -1 ions: F-, Cl-, Br-, I-
- Transition metals vary: Fe2+/Fe3+, Cu+/Cu2+, Zn2+, Ag+
- Polyatomic ions: NH4+, NO3-, SO42-, PO43-, CO32-, OH-
Naming Ionic and Molecular Compounds
Ionic nomenclature uses the cation name plus the anion name. Monatomic anions use the -ide ending (Cl- becomes chloride). Transition metals require Roman numerals (Fe3+ is iron(III)).
Molecular nomenclature uses prefixes plus element names. Prefixes include mono, di, tri, tetra, penta, hexa. You omit the mono prefix on the first element. Example: CO is carbon monoxide, N2O4 is dinitrogen tetroxide.
Acid nomenclature follows two patterns. Binary acids (H plus one nonmetal) use hydro-[root]-ic (HCl is hydrochloric acid). Oxyacids with -ate anions use -ic (HNO3 is nitric acid). Oxyacids with -ite anions use -ous (HNO2 is nitrous acid).
| Term | Meaning |
|---|---|
| Atomic Structure | Protons (+, nucleus, ~1 amu), neutrons (0, nucleus, ~1 amu), electrons (−, orbitals, ~1/1836 amu). Atomic number (Z) = # protons; mass number (A) = protons + neutrons; isotopes differ in neutrons. |
| Isotopes | Atoms of same element with different numbers of neutrons. Same chemical behavior; different mass. Average atomic mass weighted by natural abundance. Some isotopes are radioactive (carbon-14, uranium-235). |
| Electron Configuration | Aufbau principle (fill lowest energy first), Pauli exclusion (2 e⁻ per orbital, opposite spins), Hund's rule (singly fill degenerate orbitals first). Order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p... |
| Quantum Numbers | Principal (n, energy level), angular momentum (ℓ, subshell: s=0, p=1, d=2, f=3), magnetic (m_ℓ, orbital orientation), spin (m_s = +½ or −½). No two electrons in an atom share all four. |
| Periodic Law | Properties of elements recur periodically with increasing atomic number. Organized into periods (rows, same energy level) and groups (columns, same valence electron count). Mendeleev first major organizer. |
| Atomic Radius | Decreases left-to-right across period (more protons pulling same energy level). Increases down group (new energy levels). Cations smaller than parent atom; anions larger. |
| Ionization Energy | Energy to remove an electron. Increases left-to-right, decreases down. Successive ionizations increase; large jump after noble gas configuration reached. Used to predict most common ion charges. |
| Electron Affinity | Energy change when atom gains an electron. Generally more negative across period, less negative down group. Halogens have most negative (favor gaining electron); noble gases positive (unfavorable). |
| Electronegativity | Tendency to attract shared electrons in a bond. Pauling scale: F = 4.0 (most), Cs = 0.7 (least). Increases left-to-right, decreases down. Determines bond polarity. |
| Groups of the Periodic Table | Alkali metals (Group 1, +1 charge), alkaline earth metals (Group 2, +2), halogens (Group 17, −1), noble gases (Group 18, inert). Transition metals (Groups 3-12) have variable oxidation states. |
| Metals vs. Nonmetals vs. Metalloids | Metals: left of staircase, lose electrons, conduct heat/electricity, malleable, ductile. Nonmetals: right, gain electrons, poor conductors, brittle solids or gases. Metalloids: on staircase (B, Si, Ge, As, Sb, Te). |
| Common Ions (Memorize) | Alkali: Li+, Na+, K+. Alkaline earth: Mg²⁺, Ca²⁺, Ba²⁺. Halides: F−, Cl−, Br−, I−. Transition: Fe²⁺/Fe³⁺, Cu+/Cu²⁺, Zn²⁺, Ag+. Polyatomic: NH₄⁺, NO₃⁻, SO₄²⁻, PO₄³⁻, CO₃²⁻, OH−. |
| Ionic Nomenclature | Cation + anion. Monatomic cations use element name; anions use -ide ending (Cl− = chloride). Transition metals: Roman numerals (Fe³⁺ = iron(III)) or stock/-ous/-ic system. Polyatomic ions keep their names. |
| Molecular Nomenclature | Prefix + element + prefix + element-ide. Prefixes: mono, di, tri, tetra, penta, hexa. Mono omitted on first element. Examples: CO = carbon monoxide, N₂O₄ = dinitrogen tetroxide. |
| Acid Nomenclature | Binary acids (H + one other): hydro-[root]-ic acid (HCl = hydrochloric acid). Oxyacids with -ate anion: -ic acid (HNO₃ = nitric). With -ite anion: -ous acid (HNO₂ = nitrous). |
| Average Atomic Mass | Weighted average of isotope masses by natural abundance. Formula: Σ(isotope mass × fractional abundance). Reported on periodic table. Not a whole number because most elements have multiple isotopes. |
Bonding, Molecules, and Stoichiometry
Bonding determines molecular geometry, polarity, and reactivity. Stoichiometry converts between moles, mass, and particles, providing the quantitative foundation for every reaction calculation.
Three Types of Chemical Bonds
Ionic bonds form when metals transfer electrons to nonmetals. This creates electrostatic attraction between oppositely charged ions. Ionic compounds form crystalline solids with high melting points and conduct electricity when molten or dissolved.
Covalent bonds form when nonmetals share electrons. Nonpolar covalent bonds have roughly equal sharing (electronegativity difference less than 0.4). Polar covalent bonds have unequal sharing (difference between 0.4 and 1.7). These typically form molecules or network solids like diamond.
Metallic bonds involve a sea of delocalized valence electrons shared among cations. This structure explains why metals conduct electricity and heat, and why they're malleable and ductile.
Predicting Molecular Geometry
Use VSEPR theory (Valence Shell Electron Pair Repulsion) to predict shapes. Electron pairs repel and arrange to maximize distance.
- Linear geometry: two bonding pairs (AX2)
- Trigonal planar: three bonding pairs (AX3)
- Tetrahedral: four bonding pairs (AX4)
- Trigonal bipyramidal: five electron pairs (AX5)
- Octahedral: six electron pairs (AX6)
Lone pairs distort the ideal geometry. A bent molecule (like water) has two bonding pairs and two lone pairs around the central atom.
Determining Polarity
Molecular polarity requires both polar bonds AND asymmetric geometry. CO2 is nonpolar despite polar bonds because it's linear and symmetric. H2O is polar because it's bent. CCl4 is nonpolar even though each bond is polar because the tetrahedral geometry is symmetric.
Intermolecular forces determine boiling point, viscosity, and surface tension. Stronger forces mean higher boiling points. The three types are London dispersion forces (all molecules, strongest for large or polarizable molecules), dipole-dipole forces (polar molecules), and hydrogen bonding (H bonded to N, O, or F).
The Foundation: Moles and Molar Mass
One mole equals 6.022 x 10^23 particles (Avogadro's number). Molar mass (grams per mole) equals the sum of atomic masses. This converts between mass, particles, and moles.
Percent composition tells you the percentage by mass of each element in a compound. Divide the mass of an element by the compound's molar mass and multiply by 100.
Empirical formulas show the simplest whole-number ratio of atoms (like CH2O for glucose). Molecular formulas show the actual number of atoms (C6H12O6). The molecular formula is always a whole-number multiple of the empirical formula.
Balancing Equations and Stoichiometry
Balance equations by ensuring equal atoms of each element on both sides. Adjust coefficients only, never subscripts. Balance the hardest element first, saving oxygen and hydrogen for last.
Stoichiometry uses mole ratios from balanced equations to convert quantities. The process is: convert given mass to moles, use the mole ratio from the equation, convert moles of the target substance back to the desired unit.
Limiting reactants determine the maximum product. Calculate how much product each reactant can produce. The reactant that produces the least product is limiting. The other reactant(s) are in excess.
Percent yield compares actual to theoretical product: (actual yield / theoretical yield) × 100%. Values below 100% reflect side reactions, incomplete reactions, or purification losses.
| Term | Meaning |
|---|---|
| Ionic Bond | Electron transfer between metal (cation) and nonmetal (anion). Results from electrostatic attraction of oppositely charged ions. Forms crystalline solids with high melting points; conducts electricity when molten or dissolved. |
| Covalent Bond | Electron sharing between nonmetals. Nonpolar covalent: equal sharing (ΔEN < 0.4). Polar covalent: unequal sharing (ΔEN 0.4-1.7). Typically forms molecules or network solids (diamond, SiO₂). |
| Metallic Bond | Sea of delocalized valence electrons shared among lattice of cations. Explains metallic properties: electrical/thermal conductivity, malleability, ductility, luster. |
| Lewis Structures | Dot diagrams showing valence electrons and bonds. Steps: sum valence electrons, place central atom, connect with single bonds, distribute lone pairs to satisfy octets, form double/triple bonds if needed. |
| Resonance Structures | Multiple valid Lewis structures when electrons are delocalized (e.g., O₃, CO₃²⁻, NO₃⁻). Actual structure is a hybrid; average bond lengths and strengths. Indicated by double-headed arrows. |
| Formal Charge | Formal charge = valence e⁻ − nonbonding e⁻ − ½(bonding e⁻). Best Lewis structure minimizes formal charges and places negative charges on most electronegative atoms. |
| VSEPR Theory | Electron pairs repel and arrange to maximize distance. Shapes: linear (AX2), trigonal planar (AX3), tetrahedral (AX4), trigonal bipyramidal (AX5), octahedral (AX6). Lone pairs distort geometry. |
| Molecular Polarity | Polar molecules have net dipole moment. Requires polar bonds AND asymmetric geometry (lone pairs or unequal substituents). CO₂ nonpolar (linear, symmetric); H₂O polar (bent); CCl₄ nonpolar (tetrahedral, symmetric). |
| Intermolecular Forces | London dispersion (all molecules, strongest for large/polarizable), dipole-dipole (polar molecules), hydrogen bonding (H bonded to N, O, or F). Stronger IMFs → higher boiling point, viscosity, surface tension. |
| The Mole | SI unit of amount. 1 mole = 6.022 × 10²³ particles (Avogadro's number). Molar mass (g/mol) = sum of atomic masses. Converts between mass, particles, and moles. |
| Percent Composition | % by mass of each element in compound. Formula: (mass of element in 1 mole / molar mass of compound) × 100%. Used to find empirical formulas from experimental data. |
| Empirical vs. Molecular Formulas | Empirical: simplest whole-number ratio of atoms (CH₂O). Molecular: actual formula (C₆H₁₂O₆ for glucose). Molecular = n × empirical, where n = (molar mass)/(empirical formula mass). |
| Balancing Equations | Conservation of mass: atoms of each element equal on both sides. Balance by inspection: hardest element first, O and H last (usually). Never change subscripts, only coefficients. |
| Stoichiometry | Quantitative relationships in reactions. Convert given to moles → use mole ratio from balanced equation → convert moles of target to desired unit. Foundation for limiting reactant, yield calculations. |
| Limiting Reactant | Reactant fully consumed first, determining maximum product. Method: calculate moles of product each reactant can produce; smaller = limiting. Excess reactant has leftover after reaction. |
| Percent Yield | (Actual yield / theoretical yield) × 100%. Usually < 100% due to side reactions, incomplete reactions, or purification losses. > 100% indicates impurity or measurement error. |
Reactions, Thermodynamics, and Equilibrium
Understanding reactions, energy changes, and equilibrium prepares you for the most heavily tested chemistry topics, including acid-base and redox chemistry.
Classifying Reactions
The six main reaction types are:
- Synthesis (A + B → AB): combining two substances
- Decomposition (AB → A + B): breaking one substance into two
- Single replacement (A + BC → AC + B): one element displaces another
- Double replacement (AB + CD → AD + CB): ions swap partners
- Combustion (CxHy + O2 → CO2 + H2O): burning in oxygen
- Redox and acid-base: electron or proton transfer
Predicting Products with Solubility Rules
Memorize these solubility rules to predict which compounds form precipitates.
Always soluble compounds include Group 1 salts, NH4+ compounds, NO3- compounds, and ClO4- compounds. Most halides are soluble except with Ag+, Pb2+, and Hg2+. Most carbonates, phosphates, and sulfides are insoluble except with Group 1 metals or NH4+.
Net ionic equations remove spectator ions (ions that don't participate) and show only reacting species. Strong electrolytes like HCl dissociate completely. Weak electrolytes and precipitates remain intact.
Oxidation-Reduction Basics
Oxidation numbers track electron transfer. Rules: elements in pure form have 0, monatomic ions equal the charge, H is usually +1 (except -1 with metals), O is usually -2 (except -1 in peroxides), and the sum equals the overall charge.
Oxidation is an increase in oxidation number (electron loss). Reduction is a decrease in oxidation number (electron gain).
Energy and Enthalpy
Exothermic reactions release heat (ΔH < 0). Endothermic reactions absorb heat (ΔH > 0). Enthalpy (ΔH) is heat released or absorbed at constant pressure.
Hess's law states that the enthalpy change of a reaction equals the sum of enthalpy changes for individual steps. This lets you calculate ΔH without running the experiment. Use: ΔH_rxn = sum of ΔH_f(products) minus sum of ΔH_f(reactants).
Entropy and Spontaneity
Entropy (ΔS) measures disorder. Entropy increases when solids become liquids, liquids become gases, gas moles increase, or substances dissolve and mix. Use Gibbs free energy (ΔG = ΔH - TΔS) to predict spontaneity.
Negative ΔG means the reaction is spontaneous forward. Positive ΔG means it's nonspontaneous. Zero ΔG means equilibrium.
Chemical Equilibrium
At chemical equilibrium, forward and reverse reaction rates are equal. The equilibrium constant (K) is [products]/[reactants] raised to stoichiometric coefficients. Pure solids and pure liquids don't appear in the K expression.
Le Chatelier's principle predicts how systems respond to stress. Adding a reactant or removing a product shifts the system right. Increasing pressure favors the side with fewer gas molecules. Raising temperature favors the endothermic direction.
Acids and Bases
Strong acids completely dissociate: HCl, HBr, HI, HNO3, H2SO4, HClO4. Weak acids partially dissociate: HF, acetic acid, phosphoric acid. Strong acids have weak conjugate bases.
The pH scale is pH = -log[H+]. Each unit represents a tenfold change. At 25°C, pH + pOH = 14, and Kw = [H+][OH-] = 1.0 x 10^-14. Neutral solutions have pH 7, acidic solutions have pH < 7, and basic solutions have pH > 7.
Gas Laws and Colligative Properties
The ideal gas law is PV = nRT (R = 0.0821 L·atm/(mol·K)). The law combines Boyle's, Charles's, and Avogadro's laws.
Colligative properties depend only on solute particle count, not identity. These include boiling point elevation (ΔTb = iKbm), freezing point depression (ΔTf = iKfm), vapor pressure lowering, and osmotic pressure.
| Term | Meaning |
|---|---|
| Types of Reactions | Synthesis (A + B → AB), decomposition (AB → A + B), single replacement (A + BC → AC + B), double replacement (AB + CD → AD + CB), combustion (CxHy + O₂ → CO₂ + H₂O), acid-base, redox. |
| Net Ionic Equations | Remove spectator ions (don't change state) from complete ionic equation. Show only species that actually react. Strong electrolytes dissociate; weak electrolytes and precipitates do not. |
| Solubility Rules | Always soluble: group 1 salts, NH₄⁺, NO₃⁻, ClO₄⁻, most C₂H₃O₂⁻. Halides soluble except Ag+, Pb²⁺, Hg₂²⁺. Carbonates, phosphates, sulfides mostly insoluble except with group 1 or NH₄⁺. |
| Oxidation Numbers | Rules: element in pure form = 0, monatomic ion = charge, H = +1 (−1 with metals), O = −2 (−1 in peroxides), sum = charge of species. Oxidation: increase in oxidation number; reduction: decrease. |
| Thermochemistry Basics | System: what's being studied. Surroundings: everything else. Exothermic: releases heat (ΔH < 0). Endothermic: absorbs heat (ΔH > 0). q = mcΔT for temperature changes. |
| Enthalpy (ΔH) | Heat change at constant pressure. State function: depends only on initial and final states. Hess's Law: ΔH for reaction = sum of ΔH for steps. ΔH_rxn = ΣΔH_f(products) − ΣΔH_f(reactants). |
| Entropy (ΔS) | Measure of disorder. ΔS > 0 for: solid → liquid → gas, increase in moles of gas, dissolving, mixing. Third law: ΔS of pure perfect crystal at 0 K = 0. Units: J/(mol·K). |
| Gibbs Free Energy | ΔG = ΔH − TΔS. Negative ΔG = spontaneous forward. Positive ΔG = nonspontaneous forward. Zero ΔG = equilibrium. ΔG° = −RT ln K relates standard free energy to equilibrium constant. |
| Chemical Kinetics | Study of reaction rates and mechanisms. Rate depends on concentration (rate law), temperature (Arrhenius), and catalysts (lower Ea). Collision theory: reactions require proper orientation and sufficient energy. |
| Chemical Equilibrium | Dynamic state where forward and reverse rates are equal. K (equilibrium constant) = [products]/[reactants] raised to stoichiometric coefficients. Pure solids and liquids excluded from K expression. |
| Le Chatelier's Principle | System at equilibrium shifts to oppose applied stress. Adding reactant or removing product shifts right. Increasing pressure favors side with fewer moles of gas. Temperature: raise T → shift endothermic direction. |
| Acids and Bases | Arrhenius: H+ donors / OH− producers. Brønsted-Lowry: proton donor / acceptor. Lewis: electron pair acceptor / donor. Conjugate acid-base pairs differ by one H+. Amphoteric species can be either. |
| Strong vs. Weak Acids | Strong (fully dissociate): HCl, HBr, HI, HNO₃, H₂SO₄, HClO₄, HClO₃. Weak (partially dissociate): HF, CH₃COOH (acetic), H₃PO₄, HNO₂. Strong acids have weak conjugate bases. |
| pH Scale | pH = −log[H+]. pOH = −log[OH−]. pH + pOH = 14 at 25°C. Kw = [H+][OH−] = 1.0 × 10⁻¹⁴. Neutral pH = 7; acidic < 7; basic > 7. Each unit is a 10-fold change. |
| Gas Laws | Boyle's (P₁V₁ = P₂V₂ at constant T, n). Charles's (V₁/T₁ = V₂/T₂ at constant P, n). Avogadro's (V proportional to n). Combined: PV = nRT (ideal gas law, R = 0.0821 L·atm/(mol·K)). |
| Colligative Properties | Depend only on number of solute particles, not identity. Include boiling point elevation (ΔTb = iKbm), freezing point depression (ΔTf = iKfm), vapor pressure lowering, osmotic pressure. |
How to Study chemistry Effectively
Mastering chemistry requires the right study approach, not just more hours. Research in cognitive science shows three techniques produce the best outcomes: active recall (testing yourself rather than re-reading), spaced repetition (reviewing at scientifically-optimized intervals), and interleaving (mixing related topics instead of studying one in isolation). FluentFlash is built on all three.
Why Passive Review Fails
The most common mistake is relying on passive methods. Re-reading notes, highlighting textbook passages, or watching lectures feels productive but yields only 10-20% of the retention that active recall achieves. Flashcards force your brain to retrieve information, strengthening memory far more than recognition alone.
Pair flashcards with spaced repetition scheduling, and you learn in 20 minutes what would take hours of passive review.
A Practical 4-Week Study Plan
Follow this proven approach:
- Create 15-25 flashcards covering the highest-priority concepts from your current unit
- Review them daily for the first week using the FSRS algorithm
- As cards become easier, intervals automatically expand from minutes to days to weeks
- Stay at the edge of your knowledge instead of wasting time on material you've already mastered
- After 2-3 weeks of consistent practice, concepts become automatic rather than effortful
The key is consistency. Daily 15-20 minute sessions beat weekend marathons.
Pairing Flashcards with Problem Practice
Flashcards alone aren't enough. Chemistry requires both memorization and problem-solving fluency. Spend 30-45 minutes weekly on problem sets for each topic. This forces you to apply concepts, not just recall definitions.
Teach concepts aloud or write them down. If you can explain why a reaction proceeds or how Le Chatelier's principle works, you understand it deeply.
- 1
Generate flashcards using FluentFlash AI or create them manually from your notes
- 2
Study 15-20 new cards per day, plus scheduled reviews
- 3
Use multiple study modes (flip, multiple choice, written) to strengthen recall
- 4
Track your progress and identify weak topics for focused review
- 5
Review consistently, daily practice beats marathon sessions
Why Flashcards Work Better Than Other Study Methods for chemistry
Flashcards aren't just for vocabulary. They're one of the most research-backed study tools for any subject, including chemistry. The reason lies in how memory works.
The Testing Effect
When you read a textbook passage, your brain stores that information in short-term memory. Without retrieval practice, it fades within hours. Flashcards force retrieval, which transfers information from short-term to long-term memory.
The testing effect, documented in hundreds of peer-reviewed studies, shows that flashcard students consistently outperform re-readers by 30-60% on delayed tests. Every time you successfully recall a chemistry concept, you're strengthening the neural pathway that stores it.
FSRS Scheduling Amplifies Retention
FluentFlash uses the FSRS algorithm, a modern spaced repetition system that schedules reviews at mathematically optimal intervals based on your performance. Cards you find easy get pushed further into the future. Cards you struggle with come back sooner.
Students using FSRS-based systems typically retain 85-95% of material after 30 days. Compare that to roughly 20% retention from passive review alone.