Matter, Chemical Trends & Bonding: Test Review

Key facts

Where marks get lost

• ×Treating the Bohr model as physically real — electrons don't actually travel in neat circular orbits like planets; it's a useful simplification, not a photograph of reality. (Atomic Models & Bohr Theory)
• ×Forgetting the shell-3/shell-4 exception for K and Ca (writing 2, 8, 9 instead of 2, 8, 8, 1 for potassium). (Atomic Models & Bohr Theory)
• ×Mixing up atomic number (protons) with mass number (protons + neutrons) when finding electron count. (Atomic Models & Bohr Theory)
• ×Thinking emission and absorption spectra look the same — absorption removes specific colours from a continuous spectrum (dark lines), emission produces only specific bright colours on a black background. (Atomic Models & Bohr Theory)
• ×Assuming all electrons in an atom have the same energy — electrons in different shells have different, discrete energies. (Atomic Models & Bohr Theory)
• ×Rounding an isotope's mass number and calling it the exact atomic mass — mass number is a whole-number count of nucleons, not the precise isotope mass used in averaging. (Isotopes, Atomic Mass & Isotopic Abundance)
• ×Averaging isotope masses with a simple (unweighted) average instead of weighting by abundance. (Isotopes, Atomic Mass & Isotopic Abundance)
• ×Forgetting to convert percent abundance to a decimal (dividing by 100) before multiplying. (Isotopes, Atomic Mass & Isotopic Abundance)
• ×Confusing atomic number (protons, defines the element) with mass number (protons + neutrons, defines the isotope). (Isotopes, Atomic Mass & Isotopic Abundance)
• ×Assuming the two isotopes in a two-isotope problem must be equally abundant (50/50) — abundances are found from the given average, not assumed. (Isotopes, Atomic Mass & Isotopic Abundance)
• ×Assuming hydrogen is an alkali metal because it's drawn above Group 1 — it's a nonmetal. (Organization of the Periodic Table)
• ×Using the group number directly as the valence electron count for groups 13–18 without subtracting 10. (Organization of the Periodic Table)
• ×Thinking transition metals (groups 3–12) follow the simple group-number-equals-valence-electrons rule — they don't; their bonding behaviour is more complex and covered separately. (Organization of the Periodic Table)
• ×Confusing period (row, shells) with group (column, valence electrons) — they describe different things. (Organization of the Periodic Table)
• ×Assuming all metalloids behave identically — their properties vary and are genuinely intermediate, not a fixed single behaviour. (Organization of the Periodic Table)
• ×Thinking more protons always means a smaller atom regardless of position — this is only true moving across a period, not down a group. (Periodic Trends)
• ×Forgetting that ionization energy has a jump-down at the start of each new period (the new outer electron sits in a bigger shell, heavily shielded by all the full shells beneath it). (Periodic Trends)
• ×Calling noble gases an exception to the ionization-energy trend — they fit it, with the highest ionization energy in each period. They are the exception for electron affinity and electronegativity instead. (Periodic Trends)
• ×Confusing electron affinity (energy change when gaining an electron) with ionization energy (energy needed to remove an electron) — they describe opposite processes. (Periodic Trends)
• ×Assuming electronegativity trends apply to noble gases — noble gases are generally excluded since they don't typically form bonds. (Periodic Trends)
• ×Drawing Lewis symbols with core (inner-shell) electrons included — only valence electrons are shown. (Ionic Bonding & Compounds)
• ×Forgetting that when a metal becomes a cation, it typically loses its entire outer shell, becoming smaller than the neutral atom (not larger). (Ionic Bonding & Compounds)
• ×Believing ionic compounds are made of individual, separate molecules — they are giant lattices, so formulas represent a ratio, not a molecule. (Ionic Bonding & Compounds)
• ×Thinking solid ionic compounds conduct electricity — they only conduct when the ions are mobile (molten or dissolved). (Ionic Bonding & Compounds)
• ×Crossing charges without reducing to lowest whole-number ratio (e.g. writing Mg₂O₂ instead of the reduced MgO). (Ionic Bonding & Compounds)
• ×Adding prefixes (di-, tri-) to binary ionic compound names — prefixes are only for molecular (covalent) compounds. (IUPAC Nomenclature)
• ×Forgetting brackets around a polyatomic ion when more than one is needed, e.g. writing CaOH2 instead of Ca(OH)2. (IUPAC Nomenclature)
• ×Determining a multivalent metal's Roman numeral from the metal's common charge instead of working backward from the anion in that specific formula. (IUPAC Nomenclature)
• ×Mixing up similarly named polyatomic ions, especially the '-ate' (more oxygens) vs '-ite' (fewer oxygens) pairs, like sulfate SO₄²⁻ vs sulfite SO₃²⁻. (IUPAC Nomenclature)
• ×Using "mono-" on the first element of a molecular compound (writing "monocarbon dioxide" instead of just "carbon dioxide"). (IUPAC Nomenclature)
• ×Forgetting to subtract used electrons from the running total at each step, leading to too many or too few electrons in the final structure. (Covalent Bonding & Lewis Structures)
• ×Placing hydrogen as a central atom — hydrogen can only ever form 1 bond and is always terminal. (Covalent Bonding & Lewis Structures)
• ×Giving hydrogen a lone pair or expecting it to have 8 electrons — hydrogen only ever needs 2 (the duet rule). (Covalent Bonding & Lewis Structures)
• ×Not going back to check the central atom's octet, missing the need for a double or triple bond. (Covalent Bonding & Lewis Structures)
• ×Applying the octet rule rigidly to every molecule, without recognizing known exceptions like BF₃ or SF₆. (Covalent Bonding & Lewis Structures)
• ×Forgetting to take the absolute value of the electronegativity difference — ΔEN is always reported as positive. (Bond Types & Electronegativity Difference)
• ×Treating the 0.4 and 1.7 cutoffs as sharp physical walls rather than useful teaching guidelines — real bonding is a continuous spectrum. (Bond Types & Electronegativity Difference)
• ×Assuming any bond between a metal and nonmetal is automatically ionic without checking ΔEN (most are, but it's the electronegativity gap that actually determines bond type). (Bond Types & Electronegativity Difference)
• ×Mixing up which atom gets the partial negative charge — it's always the more electronegative atom that pulls electron density toward itself (${\delta }^{-}$). (Bond Types & Electronegativity Difference)
• ×Confusing a coordinate covalent bond with an ionic bond — a coordinate covalent bond still shares an electron pair; it does not transfer an electron. (Bond Types & Electronegativity Difference)
• ×Assuming a molecule with polar bonds must be polar overall — symmetric molecules like CO₂ and CCl₄ have polar bonds but are nonpolar overall. (Molecular Shape (VSEPR) & Polarity)
• ×Forgetting to count lone pairs on the central atom as electron domains when determining shape. (Molecular Shape (VSEPR) & Polarity)
• ×Mixing up electron-domain geometry (counts lone pairs) with molecular shape/geometry (describes only where the atoms are) — for water, the electron-domain arrangement is tetrahedral, but the molecular shape is bent. (Molecular Shape (VSEPR) & Polarity)
• ×Assuming any tetrahedral molecule is automatically nonpolar — only true if all 4 outer atoms are identical, as in CH₄ or CCl₄, not CH₂Cl₂. (Molecular Shape (VSEPR) & Polarity)
• ×Predicting bond angles as exactly 109.5°/120°/180° in every case, ignoring that lone pairs compress the angle slightly. (Molecular Shape (VSEPR) & Polarity)
• ×Saying a substance boils because 'covalent bonds break' — boiling separates molecules from each other by overcoming IMFs; the covalent bonds within each molecule stay intact. (Intermolecular Forces)
• ×Thinking nonpolar molecules have no intermolecular forces at all — they always have London dispersion forces, just no dipole-dipole or hydrogen bonding. (Intermolecular Forces)
• ×Assuming hydrogen bonding happens whenever hydrogen is present in a molecule — it specifically requires H bonded directly to N, O, or F. (Intermolecular Forces)
• ×Forgetting that London dispersion forces can outweigh dipole-dipole forces if the nonpolar molecule is significantly larger (e.g., large nonpolar molecules can have higher boiling points than small polar ones). (Intermolecular Forces)
• ×Confusing intermolecular forces (between separate molecules) with intramolecular bonds (within one molecule, e.g. covalent bonds) — very different strengths and roles. (Intermolecular Forces)
• ×Thinking density increases on freezing for all substances — water is a notable exception because of hydrogen bonding. (The Special Properties of Water)
• ×Explaining water's high boiling point using only molar mass — its size alone would predict a much lower boiling point; hydrogen bonding is the real cause. (The Special Properties of Water)
• ×Assuming 'universal solvent' means water dissolves absolutely everything — it dissolves polar/ionic substances well, but not nonpolar ones like oils and fats. (The Special Properties of Water)
• ×Attributing surface tension to 'water being sticky' instead of to the specific unbalanced hydrogen-bond attraction at the surface. (The Special Properties of Water)
• ×Forgetting that specific heat capacity is about energy needed to change temperature, not to change state (that's latent heat, a separate concept). (The Special Properties of Water)

Practice test

22 questions from across the unit. Each one shows the lesson it comes from.

Endless drill

Fresh numbers every time, for the calculation questions in this unit.