General Chemistry 1 Final Exam Master Deck

Complete set of 131 vocabulary flashcards covering General Chemistry 1 concepts including atomic structure, quantum theory, periodic trends, bonding, stoichiometry, aqueous solutions, oxidation-reduction, and thermochemistry.

Accuracy and Precision

Accuracy is the closeness of a measurement to the accepted value; precision is how close repeated measurements are to one another.

Significant Figures (Zero Rules)

Leading zeros do not count; zeros between nonzero digits count; trailing zeros count when a decimal point makes them significant.

Significant Figures in Calculations

For multiplication and division, round the result to the fewest significant figures. For addition and subtraction, round to the fewest decimal places. Round only the final result.

Temperature Conversions

Key formulas: $K = ^\text{o}\text{C} + 273.15$ and $^\text{o}\text{F} = \frac{9}{5}(^\text{o}\text{C}) + 32$.

Density

Mass per unit volume ($d = \frac{m}{V}$); commonly expressed in units of $\text{g/mL}$ or $\text{g/cm}^3$.

Classification of Matter

Classified by state as solid, liquid, or gas; classified by composition as pure substances (elements and compounds) versus mixtures (homogeneous and heterogeneous).

Physical vs Chemical Changes

A physical change leaves chemical identity unchanged (such as melting); a chemical change results in the formation of new substances (such as combustion and rusting).

Intensive and Extensive Properties

An intensive property does not depend on sample amount (such as density and temperature); an extensive property depends on sample amount (such as mass and volume).

Conservation of Mass and Definite Proportions

Reaction mass is conserved; a given compound always contains elements in a fixed mass ratio.

Atomic Number ($Z$) and Mass Number ($A$)

Atomic number ($Z$) is the number of protons; mass number ($A$) is protons plus neutrons. For neutral atoms, electrons equal $Z$.

Isotopes and Ions

Isotopes have the same number of protons but different numbers of neutrons; ions have a different number of electrons from the neutral atom.

Electrons in an Ion

$\text{Electrons} = Z - \text{positive charge}$, or $Z + \text{magnitude of negative charge}$.

Average Atomic Mass

Calculated as $\sum(\text{isotopic mass} \times \text{fractional natural abundance})$, converting percent abundance to a decimal first.

Mole and Avogadro's Number

$1\,\text{mol} = 6.022 \times 10^{23}$ specified particles.

Grams, Moles, and Particles Conversions

$\text{mol} = \frac{\text{grams}}{\text{molar mass}}$; $\text{particles} = \text{mol} \times 6.022 \times 10^{23}$; reverse these operations for opposite conversions.

Wavelength, Frequency, and Photon Energy

Related by $c = \lambda\nu$ and $E = h\nu = \frac{hc}{\lambda}$. Shorter wavelength corresponds to higher frequency and higher energy.

Speed of Light ($c$) and Planck's Constant ($h$)

$c \approx 3.00 \times 10^8\,\text{m/s}$ and $h \approx 6.626 \times 10^{-34}\,\text{J}\cdot\text{s}$.

Photoelectric Effect

Demonstrates that light comes in photons; electrons are ejected only when photon frequency exceeds a threshold ($E = h\nu - \text{work function}$).

Atomic Line Spectrum

Produced when electrons absorb energy to jump up levels and emit photons when they drop down; only specific energy differences are allowed.

Bohr Energy Formula for Hydrogen

$E = -\frac{2.18 \times 10^{-18}\,\text{J}}{n^2}$. Ground state is $n = 1$; $n \rightarrow \infty$ gives $0$ (unbound electron).

Hydrogen Transition Photon Energy

$\Delta E_{\text{atom}} = E_{\text{final}} - E_{\text{initial}}$; photon energy equals $|\Delta E_{\text{atom}}| = \frac{hc}{\lambda}$. In emission, the atom loses energy.

de Broglie Wavelength

$\lambda = \frac{h}{p} = \frac{h}{mv}$; demonstrates that moving matter exhibits wave properties.

Heisenberg Uncertainty Principle

States that position and momentum cannot both be precisely known: $\Delta x \Delta p \ge \frac{h}{4\pi}$.

Orbital

A probability distribution described by a wave function, as opposed to a fixed circular path in a Bohr orbit.

Quantum Numbers

Permissible values: principal $n = 1, 2, \dots$; angular momentum $\ell = 0, \dots, n - 1$; magnetic $m_\ell = -\ell, \dots, +\ell$; spin $m_s = +\frac{1}{2}$ or $-\frac{1}{2}$.

Subshells ($\ell = 0, 1, 2, 3$)

Designated as $s, p, d, f$ respectively; they contain $1, 3, 5, 7$ orbitals.

Subshell Electron Capacities

Maximum capacities are $2, 6, 10, 14$ electrons for $s, p, d, f$ subshells, respectively.

Aufbau, Pauli, and Hund Rules

Aufbau: fill lower-energy orbitals first; Pauli: at most two opposite-spin electrons per orbital; Hund: occupy equal-energy orbitals singly with parallel spins before pairing.

Orbital Filling Order

Filling sequence through typical general chemistry elements: $1s\,2s\,2p\,3s\,3p\,4s\,3d\,4p\,5s\,4d\,5p\,6s\,4f\,5d\,6p$.

Chromium and Copper Exceptions

Ground-state anomalies: $\text{Cr}: [\text{Ar}]\,4s^1 3d^5$ and $\text{Cu}: [\text{Ar}]\,4s^1 3d^{10}$, rather than simple Aufbau predictions.

Transition-Metal Cations

Configured by removing $ns$ electrons before $(n - 1)d$ electrons (for example, $\text{Fe}^{2+} = [\text{Ar}]\,3d^6$).

Paramagnetic and Diamagnetic

Paramagnetic species contain $\ge 1$ unpaired electron; diamagnetic species have all electrons paired.

Shielding and Effective Nuclear Charge

Inner electrons reduce the attraction of the nucleus for outer electrons; effective nuclear charge is the net positive attraction felt by an electron.

Atomic Radius Trend

Generally decreases left to right across a period (due to stronger effective nuclear charge) and increases down a group (due to additional shells).

Cation and Anion Sizes

Cations are smaller than their parent atoms; anions are larger than their parent atoms.

Isoelectronic Series Size Ranking

For species with the same electron count, more protons produce a smaller radius (for example, $\text{O}^{2-} > \text{F}^- > \text{Na}^+ > \text{Mg}^{2+}$).

First Ionization Energy

The energy required to remove an electron from a gas-phase atom; generally increases from left to right across a period and decreases down a group.

First Ionization Energy Exceptions

Exceptions occur where $\text{Be} > \text{B}$ and $\text{N} > \text{O}$ (similarly $\text{Mg} > \text{Al}$ and $\text{P} > \text{S}$), caused by subshell energy differences and electron pairing.

Successive Ionization Energies

A large jump indicates the removal of a core electron after valence electrons have been removed.

Electron Affinity

The energy change when a gas-phase atom gains an electron; generally becomes more favorable across a period, with important exceptions.

Metallic Character Trend

Increases down a group and toward the left of the periodic table; metals tend to lose electrons.

Ionic and Covalent Bonds

Ionic: electrostatic attraction between ions, usually metal/nonmetal. Covalent: electron sharing, usually between nonmetals.

Lattice Energy

The energy associated with separating or forming an ionic crystal from gaseous ions; becomes stronger with larger ionic charges and smaller ion sizes.

Ionic Formula Writing

Balance cation and anion charges so that the total charge is zero; express using the lowest whole-number ratio.

Variable-Charge Metal Nomenclature

Named using a Roman numeral to indicate the charge on the metal cation (for example, $\text{FeCl}_3$ is iron(III) chloride).

Binary Molecular Compound Nomenclature

Named using Greek prefixes for atom counts, ending the second element in -ide, and usually omitting mono- on the first element.

Greek Numerical Prefixes (1–10)

mono ($1$), di ($2$), tri ($3$), tetra ($4$), penta ($5$), hexa ($6$), hepta ($7$), octa ($8$), nona ($9$), deca ($10$).

Common Polyatomic Ions

$\text{NH}_4^+$ ammonium; $\text{OH}^-$ hydroxide; $\text{NO}_3^-$ nitrate; $\text{NO}_2^-$ nitrite; $\text{SO}_4^{2-}$ sulfate; $\text{SO}_3^{2-}$ sulfite; $\text{CO}_3^{2-}$ carbonate; $\text{PO}_4^{3-}$ phosphate; $\text{C}_2\text{H}_3\text{O}_2^-$ acetate.

Oxyanion Naming (-ate/-ite, per-/hypo-)

-ate has more oxygen atoms than -ite; per- has one more oxygen than -ate; hypo- has one fewer oxygen than -ite.

Binary Acid Nomenclature

Named as hydro- + nonmetal root + -ic acid (for example, $\text{HCl}(aq)$ is hydrochloric acid).

Oxyacid Nomenclature

Named by converting an -ate polyatomic ion to -ic acid, and an -ite polyatomic ion to -ous acid, without using a hydro- prefix.

Hydrate

An ionic compound containing a fixed number of water molecules per formula unit; named using a Greek prefix + hydrate.

Empirical vs Molecular Formula

Empirical formula gives the simplest whole-number atom ratio; molecular formula gives actual atom counts and is a whole-number multiple of the empirical formula.

Lewis Structure Workflow

Count valence electrons (adjust for charge), connect skeleton, fill terminal octets, place remaining electrons on the central atom, form multiple bonds if needed, and check formal charges.

Octet Rule

Main-group atoms often gain, lose, or share electrons to reach eight valence electrons; hydrogen follows the duet rule.

Octet Rule Exceptions

Hydrogen requires $2$ electrons; boron and beryllium may have incomplete octets; period $3+$ central atoms can have expanded octets; odd-electron species are exceptions.

Formal Charge

$\text{FC} = \text{valence } e^- - \text{nonbonding } e^- - \frac{1}{2}(\text{bonding } e^-)$. The sum of formal charges equals total ion/molecule charge.

Resonance Contributor Preference

Favor complete octets (where applicable), small formal charges, and negative formal charge placed on more electronegative atoms.

Resonance

The existence of multiple valid Lewis drawings differing only in electron placement, not atom positions; the actual structure is a resonance hybrid.

Bond Order (Lewis Structure)

Single = $1$, double = $2$, triple = $3$; higher bond order generally corresponds to a shorter, stronger bond.

VSEPR Theory

Electron domains repel and arrange to minimize repulsion; lone pairs influence molecular shape and affect bond angles.

Electron Domain

A region of electron density where a single, double, or triple bond counts as one domain, and each lone pair counts as one domain.

Electron-Domain Geometries (2–6 Domains)

2 domains: linear; 3 domains: trigonal planar; 4 domains: tetrahedral; 5 domains: trigonal bipyramidal; 6 domains: octahedral.

Molecular Shapes for 3 Domains

$\text{AX}_3$: trigonal planar ($\approx 120^\text{o}$); $\text{AX}_2\text{E}$: bent ($< 120^\text{o}$).

Molecular Shapes for 4 Domains

$\text{AX}_4$: tetrahedral ($109.5^\text{o}$); $\text{AX}_3\text{E}$: trigonal pyramidal ($\approx 107^\text{o}$); $\text{AX}_2\text{E}_2$: bent ($\approx 104.5^\text{o}$).

Molecular Shapes for 5 Domains

Includes $\text{AX}_5$ trigonal bipyramidal, $\text{AX}_4\text{E}$ seesaw, $\text{AX}_3\text{E}_2$ T-shaped, and $\text{AX}_2\text{E}_3$ linear.

Molecular Shapes for 6 Domains

Includes $\text{AX}_6$ octahedral, $\text{AX}_5\text{E}$ square pyramidal, and $\text{AX}_4\text{E}_2$ square planar.

Electron Geometry vs Molecular Geometry

Electron geometry accounts for all bonds and lone pairs; molecular geometry describes the spatial positions of atoms only.

Bond Polarity vs Molecular Polarity

Polar bonds result from electronegativity differences; a molecule is polar if individual bond dipoles do not cancel given its three-dimensional shape.

Orbital Hybridization (2–6 Domains)

Introductory model corresponds to $sp, sp^2, sp^3, sp^3d, sp^3d^2$ for 2, 3, 4, 5, and 6 electron domains, respectively.

Sigma ($\sigma$) and Pi ($\pi$) Bonds

$\sigma$: formed by end-to-end orbital overlap; $\pi$: formed by side-by-side overlap of parallel unhybridized $p$ orbitals.

Sigma and Pi Bond Counts

Single bond: $1\sigma$; double bond: $1\sigma + 1\pi$; triple bond: $1\sigma + 2\pi$.

Molecular-Orbital Bond Order

$\text{Bond order} = \frac{\text{bonding electrons} - \text{antibonding electrons}}{2}$.

MO Theory Magnetism

Unpaired electrons in molecular orbitals cause paramagnetism; all electrons paired causes diamagnetism (for example, $\text{O}_2$ is paramagnetic).

Mass Percent of an Element

$\%\,\text{element} = \left(\frac{\text{mass of element in } 1\,\text{mol compound}}{\text{molar mass of compound}}\right) \times 100$.

Empirical Formula Calculation

Assume a $100\,\text{g}$ sample, convert each element to moles, divide by the smallest mole value, and multiply to obtain whole-number subscripts.

Molecular Formula Calculation

Calculate the factor $n = \frac{\text{molecular molar mass}}{\text{empirical-formula mass}}$, then multiply every empirical subscript by $n$.

Combustion Analysis

Determine moles of $\text{C}$ from $\text{CO}_2$, moles of $\text{H}$ from $\text{H}_2\text{O}$ ($\times 2$), and determine $\text{O}$ by mass difference if only $\text{C}, \text{H}, \text{O}$ are present.

Stoichiometry First Step

Write and balance the chemical equation; reaction coefficients provide mole ratios, not mass ratios.

Grams-to-Grams Route

$\text{grams } A \rightarrow \text{mol } A \rightarrow \text{mol } B$ (via balanced coefficients) $\rightarrow \text{grams } B$.

Limiting Reactant

The reactant that produces the smaller calculated amount of product, thereby limiting the maximum amount of product formed.

Theoretical Yield

The maximum amount of product predicted from the limiting reactant by stoichiometry.

Percent Yield

$\%\,\text{yield} = \frac{\text{actual yield}}{\text{theoretical yield}} \times 100$.

Backward Calculation from Actual Yield

$\text{Theoretical yield needed} = \frac{\text{desired actual yield}}{\%\,\text{yield} / 100}$, followed by stoichiometric calculation.

Reaction Classifications

Synthesis combines reactants; decomposition splits a compound; combustion with $\text{O}_2$ often makes oxides; displacement replaces an element or ion.

Solute, Solvent, and Solution

Solute is the dissolved substance; solvent is the dissolving medium; solution is the resulting homogeneous mixture.

Hydration

The process by which water molecules surround and stabilize dissolved ions, enabling many ionic solids to dissociate in solution.

Electrolyte vs Nonelectrolyte

Electrolytes yield ions in solution and conduct electricity; nonelectrolytes dissolve without producing appreciable ions.

Strong vs Weak Electrolytes

Strong electrolytes exhibit essentially complete ionization/dissociation; weak electrolytes exhibit partial ionization. Electrolyte strength is distinct from concentration.

Molarity

$M = \frac{\text{moles solute}}{\text{liters of solution}}$. Always convert $\text{mL}$ to $\text{L}$.

Dilution Equation

$M_1 V_1 = M_2 V_2$, based on the condition that the amount of solute remains unchanged during dilution.

Solution Stoichiometry

Use $M \times V(\text{L})$ to obtain moles, apply the balanced chemical equation mole ratio, and convert to requested units.

Broad Solubility Rules

Group 1, $\text{NH}_4^+$, and nitrate salts are soluble; many halides are soluble except with $\text{Ag}^+$ and $\text{Pb}^{2+}$; many carbonates, phosphates, and hydroxides are insoluble with listed exceptions.

Molecular, Complete Ionic, and Net Ionic Equations

Molecular shows intact formulas; complete ionic splits strong aqueous electrolytes; net ionic removes spectator ions.

Spectator Ion

An ion present unchanged on both sides of a complete ionic equation; it is canceled out to write the net ionic equation.

Arrhenius Acids and Bases

An Arrhenius acid increases $\text{H}^+$ (or $\text{H}_3\text{O}^+$) concentration in water; an Arrhenius base increases $\text{OH}^-$ concentration in water.

Brønsted Acids and Bases

A Brønsted acid donates $\text{H}^+$; a Brønsted base accepts $\text{H}^+$. Conjugate acid/base pairs differ by one proton.

Neutralization

An acid and base react, commonly yielding water and a salt; the net ionic equation for a strong acid and strong base is $\text{H}^+ + \text{OH}^- \rightarrow \text{H}_2\text{O}$.

Seven Strong Acids

$\text{HCl}, \text{HBr}, \text{HI}, \text{HNO}_3, \text{HClO}_3, \text{HClO}_4, \text{H}_2\text{SO}_4$ (for sulfuric acid, only the first ionization is strong).

Common Strong Bases

Group 1 hydroxides and the heavy Group 2 hydroxides: $\text{Ca}(\text{OH})_2, \text{Sr}(\text{OH})_2, \text{Ba}(\text{OH})_2$.