First-year courses · Sciences

General chemistry I: atoms, bonding and gases

Atomic structure, electron configurations, periodic trends, Lewis structures, VSEPR, stoichiometry, gas laws and solutions in first-year chemistry flashcards.

45 cards6 topicsAbout 20 minutes a roundTypical first-year general chemistry, first term (structure, bonding, stoichiometry, gases, solutions)

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About this deck

This deck matches the first term of a typical first-year general chemistry course. It runs from atomic structure, isotopes and the mole, through light and electron configurations, periodic trends, ionic and covalent bonding, Lewis structures, formal charge and resonance, VSEPR shapes, hybridization and intermolecular forces, to stoichiometry (limiting reactant, percent yield, empirical formulas), the gas laws and solution concentration.

Grade 12 chemistry is the usual starting point. You should be able to balance a simple equation, convert between grams and moles, and use scientific notation and significant figures. Keep a periodic table nearby; several cards ask you to reason from position on the table rather than memorize values.

A good routine is to answer out loud, then explain why in one sentence (for example, why ionization energy rises across a period). Explaining the reason, not just the trend, is what exam questions reward. For calculation cards, cover the example, do the numbers with units, and compare. Once the Bonding and structure tag is solid, draw the Lewis structure and predict the shape for every molecule you meet in other cards.

Students commonly mix up electron-domain geometry with molecular shape (NH3 has four electron domains but a trigonal pyramidal shape), use Celsius instead of kelvin in gas law calculations, and pick the limiting reactant by comparing grams instead of moles adjusted for the coefficients. Another frequent slip is assuming that a molecule with polar bonds must be polar overall; symmetry can cancel the bond dipoles, as in CO2.

Before you start: Grade 12 chemistry (or Chemistry 30 / Chemistry 12) and comfort with algebra and scientific notation.

All 45 cards

Cover the definitions with a sheet of paper and test yourself, or use the study modes above.

Atomic structure

Atomic number and mass number: Z counts the protons in the nucleus and identifies the element; A counts protons plus neutrons.Example: Carbon-14 has Z = 6 and A = 14, so it has 6 protons and 8 neutrons.
Isotopes: Atoms of the same element (same proton count) with different numbers of neutrons, so different masses but nearly identical chemistry.Example: Chlorine-35 and chlorine-37.
Average atomic mass: The abundance-weighted mean of the masses of an element's naturally occurring isotopes, as listed on the periodic table.Example: Cl: 0.758(34.97) + 0.242(36.97) = 35.45 u.
Mole and Avogadro's number: The amount of substance containing 6.022 x 10^23 particles (atoms, molecules or ions).Example: 0.50 mol of water contains 3.011 x 10^23 molecules.
Molar mass: Mass of one mole of a substance in g/mol, found by adding the atomic masses in the formula.Example: H2O: 2(1.008) + 16.00 = 18.02 g/mol, so 36.04 g is 2.000 mol.
Photon energy: Light comes in packets with energy E = h nu = hc/lambda, where h = 6.626 x 10^(-34) J s and c = 3.00 x 10^8 m/s.Example: A 500 nm photon carries E = 3.97 x 10^(-19) J.
Quantum numbers: Four labels for an electron: n (shell, size), l (subshell shape, 0 to n - 1), ml (orbital orientation, -l to +l) and ms (spin, +1/2 or -1/2).Example: A 3p electron has n = 3 and l = 1, and ml can be -1, 0 or +1.
Pauli exclusion principle: No two electrons in an atom can have the same set of four quantum numbers, so an orbital holds at most two electrons, with opposite spins.Example: A p subshell (3 orbitals) holds at most 6 electrons.
Hund's rule: Electrons fill degenerate orbitals singly with parallel spins before any orbital gets a second electron.Example: Nitrogen's three 2p electrons occupy three separate p orbitals.
Aufbau principle: Electrons occupy the lowest-energy available orbitals first, roughly in the order 1s 2s 2p 3s 3p 4s 3d 4p ...Example: Iron (26 electrons): [Ar] 4s^2 3d^6.
Valence electrons: The electrons in the outermost shell (for main-group elements, the highest n s and p electrons); they take part in bonding.Example: Oxygen, 1s^2 2s^2 2p^4, has 6.

Periodic trends

Effective nuclear charge: The net positive pull felt by an electron after shielding by inner electrons; roughly Z minus the number of core electrons for valence electrons.Example: Rough estimate for Na's 3s electron: 11 - 10 = +1; for Cl's valence electrons: 17 - 10 = +7.
Atomic radius trend: Atoms shrink from left to right across a period (higher effective nuclear charge) and grow down a group (more shells).Example: Na is larger than Cl; K is larger than Na.
Ionization energy trend: The energy needed to remove an electron from a gaseous atom generally rises across a period and falls down a group.Example: Cl (about 1251 kJ/mol) is much higher than Na (about 496 kJ/mol).
Electronegativity: A measure of how strongly an atom in a bond pulls shared electrons toward itself; it increases up and to the right, with fluorine highest (about 4.0 on the Pauling scale).Example: In H-Cl, the shared pair sits closer to Cl.
Ionic radii: Cations are smaller than their parent atoms (an electron shell lost or less repulsion); anions are larger (added electron repulsion).Example: Na+ is smaller than Na; Cl- is larger than Cl.

Bonding and structure

Ionic bond: Electrostatic attraction between oppositely charged ions, usually formed when a metal transfers electrons to a nonmetal.Example: NaCl forms a lattice of Na+ and Cl- ions.
Polar covalent bond: A shared electron pair that is unequally shared because the atoms differ in electronegativity, giving partial charges.Example: In H-Cl, Cl carries a partial negative charge.
Lewis structure: A diagram showing valence electrons as dots or lines: bonding pairs between atoms and lone pairs on atoms, usually giving each atom an octet.Example: CO2: O=C=O with two lone pairs on each O.
Exceptions to the octet rule: Some atoms have fewer than eight valence electrons (B, Be), odd electron counts (NO), or more than eight when in period 3 or below (S in SF6).Example: BF3 has only 6 electrons around boron.
Formal charge: Valence electrons of the free atom minus lone-pair electrons minus half the bonding electrons; the best structure keeps these near zero.Example: In O=C=O every atom has 0; in CO, C is -1 and O is +1.
Resonance structures (chemistry): Two or more valid Lewis structures differing only in electron placement; the real molecule is an average (hybrid) of them.Example: Nitrate, NO3-, has three equivalent structures, so all three N-O bonds are the same length.
VSEPR theory: Electron domains (bonds and lone pairs) around a central atom arrange themselves as far apart as possible, which predicts molecular shape.Example: Four domains point to the corners of a tetrahedron, 109.5 degrees apart.
Electron-domain geometry versus molecular shape: Domain geometry counts lone pairs; shape describes only atom positions. Lone pairs also squeeze bond angles slightly.Example: NH3: tetrahedral domains, trigonal pyramidal shape, about 107 degrees.
Common VSEPR shapes: 2 domains linear (180), 3 trigonal planar (120), 4 tetrahedral (109.5), 5 trigonal bipyramidal (90 and 120), 6 octahedral (90 degrees).Example: SF6 is octahedral; BF3 is trigonal planar.
Bent shape of water: Four electron domains with two lone pairs on oxygen give a bent molecule with an H-O-H angle of about 104.5 degrees.Example: This bend is why water is polar.
Molecular polarity: A molecule is polar if its bond dipoles do not cancel; this depends on both bond polarity and shape.Example: CO2 is linear, so its dipoles cancel (nonpolar); H2O is bent and polar.
Hybridization: Mixing of atomic orbitals to explain bonding geometry: sp (linear), sp2 (trigonal planar), sp3 (tetrahedral).Example: Carbon in CH4 is sp3; in C2H4 it is sp2; in C2H2 it is sp.
Intermolecular forces: Attractions between molecules: London dispersion (all molecules), dipole-dipole (polar molecules) and hydrogen bonding (H on N, O or F); stronger forces raise boiling points.Example: Water boils at 100 degrees C while H2S boils near -60 degrees C because of hydrogen bonding.

Stoichiometry

Balanced chemical equation: An equation with coefficients chosen so each element has the same number of atoms on both sides (conservation of mass).Example: C3H8 + 5 O2 gives 3 CO2 + 4 H2O.
Mole ratio: The ratio of coefficients in a balanced equation, used to convert moles of one substance to moles of another.Example: From the propane equation, 2 mol C3H8 produces 6 mol CO2.
Limiting reactant: The reactant that runs out first and so determines the maximum amount of product; compare moles available divided by coefficients.Example: 2 H2 + O2 gives 2 H2O with 4 mol H2 and 3 mol O2: H2 runs out first, so 4 mol H2O form.
Percent yield: Actual yield divided by theoretical yield, times 100%.Example: Theoretical 4.00 mol H2O = 72.06 g; actual 60.0 g gives 83.3%.
Empirical formula: The simplest whole-number ratio of atoms in a compound, found from mass percentages converted to moles.Example: 40.0% C, 6.7% H, 53.3% O: mole ratio 3.33 : 6.65 : 3.33, so CH2O.

Gases

Ideal gas law: PV = nRT, with T in kelvin; R = 0.08206 L atm/(mol K) or 8.314 J/(mol K).Example: 1.00 mol at 273.15 K and 1.00 atm occupies 22.4 L.
Boyle's law: At constant temperature and amount, pressure and volume are inversely proportional: P1V1 = P2V2.Example: 2.0 L at 1.0 atm compressed to 4.0 atm: 0.50 L.
Charles's law: At constant pressure and amount, volume is proportional to kelvin temperature: V1/T1 = V2/T2.Example: 1.0 L at 300 K heated to 600 K: 2.0 L.
Dalton's law of partial pressures: In a gas mixture the total pressure is the sum of the pressures each gas would exert alone; each partial pressure equals mole fraction times total pressure.Example: Air at 101.3 kPa is about 21% O2, so P(O2) is about 21.3 kPa.
Kinetic molecular theory: Model of gases as tiny particles in random motion with elastic collisions; average kinetic energy is proportional to kelvin temperature, and u_rms = sqrt(3RT/M).Example: N2 at 298 K (M = 0.028 kg/mol): u_rms = 515 m/s.
Graham's law of effusion: Effusion rate is inversely proportional to the square root of molar mass: rate1/rate2 = sqrt(M2/M1).Example: H2 effuses sqrt(32/2) = 4 times as fast as O2.
Real gas behaviour: Gases deviate from ideal behaviour at high pressure (molecules take up space) and low temperature (attractions matter); the van der Waals equation corrects for both.Example: CO2 near its condensation point has a smaller PV than the ideal gas law predicts.

Solutions

Molarity: Moles of solute per litre of solution, mol/L (written M).Example: 0.25 mol NaCl in 0.500 L of solution is 0.50 M.
Dilution equation: Adding solvent keeps moles of solute constant, so M1V1 = M2V2.Example: 10.0 mL of 6.0 M HCl diluted to 100.0 mL is 0.60 M.
Like dissolves like: Polar and ionic solutes dissolve best in polar solvents; nonpolar solutes dissolve best in nonpolar solvents.Example: Salt dissolves in water; oil dissolves in hexane but not in water.
Freezing point depression: Dissolving a solute lowers a solvent's freezing point by delta T = i Kf m, where m is molality and i the number of particles per formula unit.Example: 1.0 m NaCl in water (i = 2, Kf = 1.86 degrees C kg/mol): freezes about 3.7 degrees C lower.