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Final Exam Terms 2
| Question | Answer |
|---|---|
| Wave Nature of Matter (de Broglie, 1925) | Electrons have wave properties. Quantized energies explained by: only wavelengths that fit perfectly around the orbit are allowed. His orbit idea was wrong, but electron wave behavior is real. |
| Davisson and L.H. Germer Dual Slit Experiment (1927) | Shot electrons at a nickel crystal. Electrons produced a diffraction/interference pattern. Proves electrons behave like waves (wave–particle duality). |
| Heisenberg Uncertainty Principle (1927) | Cannot know position and velocity of a particle exactly at the same time. Applies significantly only to small particles like electrons. Bohr model violates it because it assumes a precise orbit. |
| Heisenberg Equation | Δx · mΔv ≥ 5.27 × 10⁻³⁵ kg·m²/s. Δx = uncertainty in the position of the particle. m = mass of particle. Δv = uncertainty in the velocity of the particle. |
| Quantum Mechanical Model (Schrödinger, 1926) | Treats electrons as 3D waves. Schrödinger equation gives: Allowed energies & probability of where the electron is likely found. Atomic orbitals = regions of high probability, not paths. |
| Schrödinger Equation | Written as: HΨ = EΨ. Ψ (psi) = wavefunction → describes electron behavior. H = mathematical operator. E = electron energy |
| Principal Quantum Number (n) | Positive integer. Defines the shell; indicates orbital size and energy level. Larger 𝑛 = larger orbital, higher energy, electron farther from nucleus. |
| Angular Momentum Quantum Number (ℓ) | Defines the subshell; determines orbital shape. Values range from 0 to 𝑛 − 1 |
| Magnetic Quantum Number (mℓ) | Defines orbital orientation in space. Values range from −ℓ to + ℓ. Number of values = number of orbitals in subshell. |
| Spin Quantum Number (ms) | Describes electron spin direction. Values are +½ (spin-up) or –½ (spin-down). |
| s orbital (ℓ = 0) | Spherical shape; one per shell. |
| p orbital (ℓ = 1) | Dumbbell shape; three per shell (px, py, pz). |
| d orbital (ℓ = 2) | Clover shape (with one donut-shaped variant); five per shell. |
| f orbital (ℓ = 3) | Complex shapes; seven per shell. |
| Number of Orbital in Subshell | Number of mℓ values. Ex; -1, 0, 1 = 3. |
| Total Number of Orbitals in Shells | = n^2 |
| Pauli Exclusion Principle | No two electrons can have the same set of four quantum numbers; each orbital holds max two electrons with opposite spins. |
| Orbital Naming | Orbitals are named by shell number 𝑛 and subshell letter (ℓ): e.g., 1s, 2p, 3d, 4f. |
| Energy vs. Shell Number | Energy increases with increasing shell number. |
| Hydrogen & One | Electron Ions → Orbitals in the same shell (same n) have equal energy. |
| Multi‑Electron Atoms | Subshells differ in energy due to electron–electron repulsion (shielding). |
| Shielding Effect | s electrons penetrate closer to nucleus, shielding p and d electrons, raising their energy. |
| Aufbau Principle | Electrons fill lowest‑energy orbitals first. |
| Orbital Capacity | Each orbital holds max two electrons with opposite spins. |
| Orbital Diagrams | Box notation shows electron placement in orbitals. Half up and down arrows. Always do all the up arrows first then down arrows. |
| Condensed Configurations | Use noble gas core, then add remaining electrons (e.g., Mg: [Ne] 3s²). |
| Valence Electrons | Highest occupied shell (main group) or highest s/d orbitals (transition metals). Mainly the s and p. But sometimes s and d if no p. |
| Core Electrons | All other electrons not involved in bonding. |
| Group Trends | Elements in same group share similar valence configurations → similar chemical properties. |
| Excited Electrons | Absorb energy and move to higher orbitals (e.g., Li: 1s²2s¹ → 1s²2p¹). |
| Transition Metal Exceptions | Some elements deviate from expected filling (e.g., Cu actual: [Ar] 4s¹3d¹⁰ vs. expected [Ar] 4s²3d⁹). |
| Stability Rule | Half‑filled or fully filled d subshells are often more stable. |
| Trend | Exceptions become more common in larger period numbers. |
| Valence Electrons (Main Group) | Electrons in the highest occupied shell. |
| Valence Electrons (Transition Metals) | Electrons in highest s and d orbitals after noble gas core. |
| Core Electrons | All other electrons not involved in bonding. |
| Bonding Role | Valence electrons participate in chemical reactions. |
| Group Pattern | Valence configurations follow same pattern → similar chemical properties. |
| Excitation | Highest‑energy electron absorbs energy and moves to higher orbital. |
| Main Group Ions | Gain/lose electrons to achieve filled s + p subshell (noble gas configuration). |
| Group Charges | IA = 1⁺, IIA = 2⁺, IIIA = 3⁺, VA = 3⁻, VIA = 2⁻, VIIA = 1⁻. |
| Ions | Ex: Na Neutral → 1s² 2s² 2p⁶ 3s¹. Na⁺ → 1s² 2s² 2p⁶ ([Ne]). F Neutral → 1s² 2s² 2p⁵. F⁻ → 1s² 2s² 2p⁶ ([Ne]). Fe Neutral → [Ar] 4s² 3d⁶. Fe²⁺ → [Ar] 3d⁶. Fe³⁺ → [Ar] 3d⁵. |
| Atomic Size Factors | Nuclear attraction (protons pull electrons), electron repulsion, and shell number (n). |
| Effective Nuclear Charge (Zeff) | Z – shielding factor; higher Zeff = stronger pull, smaller atom. |
| Effective Nuclear Charge (Zeff) (Across a Row) | Size decreases (more protons increase pull). |
| Effective Nuclear Charge (Zeff) (Down a Group) | Size increases (electrons in higher shells farther from nucleus). |
| Periodic Trends Definition | Trends in physical and chemical properties across rows or down columns of the periodic table. |
| Periodic Trends Factors | Determined by nuclear attraction (protons pulling electrons), electron–electron repulsion, and shell number (n). |
| Atomic Size Definition | Radius of atom = half the distance between nuclei of two bonded atoms. |
| Atomic Size (Across a Row) | Size decreases; more protons increase nuclear pull, outweighing added electron repulsion. |
| Atomic Size (Down a Group) | Size increases; valence electrons occupy higher shells, farther from nucleus. |
| Ionic Radii Cations | Lose electrons; fewer electrons, same nuclear charge → smaller radius than neutral atom. |
| Cation Charge Effect | Higher positive charge = smaller radius (e.g., Fe²⁺ vs. Fe³⁺). |
| Ionic Radii Anions | Gain electrons; more repulsion, same nuclear charge → larger radius than neutral atom. |
| Ionic Radii (Trend Down Group) | Both cations and anions increase in size; valence electrons in larger shells. |
| Isoelectronic Species Definition | Atoms/ions with same electron configuration. |
| Isoelectronic Species Rule | Greater nuclear charge = smaller radius (e.g., Na⁺ vs. Ne). |
| Isoelectronic Species Practice | Na⁺ smaller than F⁻; K⁺ smaller than Cl⁻; Cu²⁺ smaller than Cu⁺, Na smaller than K. |
| Ionization Energy Definition | Energy required to remove an electron from an atom in the gas phase. |
| First Ionization Energy (IE₁) | Removes highest‑energy electron: X + energy → X⁺ + e⁻. |
| Second Ionization Energy (IE₂) | Removes next electron: X⁺ + energy → X²⁺ + e⁻. |
| Ionization Energy (Trend Across a Row) | IE increases; atoms smaller, electrons closer to nucleus, harder to remove. |
| Ionization Energy (Trend Down a Group) | IE decreases; atoms larger, electrons farther from nucleus, easier to remove. |
| Successive Ionizations | IE₁ < IE₂ < IE₃ …; big jump occurs when removing core electrons. |
| Electron Affinity (EA) Definition | Energy change when an atom in gas phase gains an electron: X + e⁻ → X⁻. |
| Positive Electron Affinity (EA) | Energy required; atom less stable. |
| Negative Electron Affinity (EA) | Energy released; atom more stable. |
| Electron Affinity (EA) Trend | More negative EA = atom more readily gains electrons. |
| Electron Affinity (EA) Exceptions | Noble gases and alkaline earth metals resist gaining electrons (would require higher‑energy orbitals). |
| Metal Properties | Shiny, ductile, malleable, good conductors, high boiling points. |
| Metallic Behavior (Across a Row) | Metallic character decreases; reactivity of metals decreases as IE increases. |
| Metallic Behavior (Down a Group) | Metallic character increases; reactivity of metals increases as IE decreases. |
| Metallic Behavior Reactivity Rule | Metals react by losing electrons to form cations. |
| Ionic Compounds | Metal + nonmetal (or polyatomic ion). |
| Molecular (Covalent) Compounds | Nonmetal + nonmetal. |
| Polyatomic Ions | Covalently bonded clusters with overall charge (e.g., NH₄⁺, OH⁻, NO₃⁻, CO₃²⁻, SO₄²⁻, PO₄³⁻, CH₃COO⁻). |
| Ionic Compounds: Writing Formulas Charge Balance | Neutral compounds; sum of charges = 0. |
| Ionic Compounds: Writing Formulas Ratio Rule | Use smallest whole‑number ratio of ions. Examples → Li⁺ + Br⁻ → LiBr; K⁺ + N³⁻ → K₃N; Mg²⁺ + O²⁻ → MgO. Polyatomic Examples → Ca²⁺ + CO₃²⁻ → CaCO₃; Mg²⁺ + OH⁻ → Mg(OH)₂; Ca²⁺ + PO₄³⁻ → Ca₃(PO₄)₂. |
| Identifying Ions in Formulas | KI → K⁺, I⁻. CaBr₂ → Ca²⁺, Br⁻. Li₂CO₃ → Li⁺, CO₃²⁻. Al₂(SO₄)₃ → Al³⁺, SO₄²⁻. CuCl₂ → Cu²⁺, Cl⁻. K₂CO₃ → K⁺, CO₃²⁻. NaCH₃COO → Na⁺, CH₃COO⁻. |
| Ionic Compounds: Binary Ionic | Metal + nonmetal. Name = cation name + anion root + “ide.” Examples → LiBr = lithium bromide; Na₂O = sodium oxide; KF = potassium fluoride; MgCl₂ = magnesium chloride. |
| Ionic Compounds: Variable Charge Metals | Use Roman numeral for cation charge. Examples → CuCl = copper(I) chloride; CuCl₂ = copper(II) chloride; FeO = iron(II) oxide; Fe₂O₃ = iron(III) oxide. |
| Ionic Compounds: Polyatomic Compounds | Name = cation + polyatomic ion. NaNO₃ = sodium nitrate; NH₄F = ammonium fluoride; Ca₃(PO₄)₂ = calcium phosphate; BaCO₃ = barium carbonate; Li₂SO₄ = lithium sulfate; MgSO₃ = magnesium sulfite. |
| Molecular Compounds: Composition | Two nonmetals. Formula cannot be predicted from periodic table. |
| Molecular Compounds: Naming Rule | First element: prefix + name. Second element: prefix + root + “ide.” Use prefixes (mono, di, tri, tetra, etc.). Drop “mono” from first element and vowels to avoid double vowels. Examples: CO = carbon monoxide, N₂O₄ = dinitrogen tetroxide. |
Created by:
LaurenMaue
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