Grade 11 Chemistry Vocabulary Flashcards

Unit 1: Atomic Structure and Periodic Properties of the Elements

• Historical Development of Atomic Theory:

  • Democritus (460–370 BC): Suggested matter consists of tiny, indestructible particles called atomos ("indivisible"). His ideas were philosophical, not experimental.
  • John Dalton (1808): Developed atomic theory based on the law of conservation of mass and the law of definite proportions.

• Postulates of Dalton’s Atomic Theory:

  1. Elements are made of extremely small particles called atoms.
  2. All atoms of a given element are identical; atoms of different elements are different.
  3. Atoms cannot be subdivided, created, or destroyed.
  4. Atoms of different elements combine in simple whole-number ratios to form chemical compounds.
  5. In chemical reactions, atoms are combined, separated, or rearranged.

• Fundamental Laws of Chemistry:

  • Law of Conservation of Mass: Mass is neither created nor destroyed in a chemical reaction.
  • Law of Definite Proportions: A chemical compound contains the same elements in exactly the same proportions by mass regardless of the size or source of the sample.
  • Law of Multiple Proportions: If two elements form more than one compound, the ratios of the masses of the second element that combine with a fixed mass of the first element are small whole numbers.

• Discovery of Subatomic Particles:

  • Electrons (J.J. Thomson, 1897): Used cathode ray tubes. Observed that cathode rays bend toward a positive plate and away from a negative plate. Thomson calculated the mass-to-charge ratio ($m_e/e$) as $-5.686 \times 10^{-12} \, kg \, C^{-1}$.
  • Charge of Electron (Robert Millikan, 1909): Used the oil-drop experiment to determine the charge of an electron as $e = -1.602 \times 10^{-19} \, C$. Calculated electron mass as $9.109 \times 10^{-31} \, kg$.
  • Nucleus (Ernest Rutherford, 1911): Gold foil experiment. Observed alpha ($\alpha$) particle deflection. Concluded atoms have a tiny, dense, positively charged center called the nucleus.
  • Protons (Rutherford, 1919): Discovered hydrogen nuclei (protons) form when alpha particles strike light elements like nitrogen. Mass: $m_p = 1.67262 \times 10^{-27} \, kg$.
  • Neutrons (James Chadwick, 1932): Bombarded beryllium with alpha rays. Discovered neutral particles with mass nearly equal to protons ($m_n = 1.67493 \times 10^{-27} \, kg$).

• Radioactivity: Spontaneous emission of particles/radiation from unstable nuclei.

  • Alpha ($\alpha$) rays: Positively charged, identical to helium nuclei.
  • Beta ($\beta$) rays: High-speed electrons from the nucleus.
  • Gamma ($\gamma$) rays: High-energy radiation with no charge.

• Atomic Constants and Definitions:

  • Atomic Number ($Z$): Number of protons in the nucleus.
  • Mass Number ($A$): Total number of protons and neutrons.
  • Isotopes: Atoms of the same element with different numbers of neutrons (e.g., $^{12}C, ^{13}C, ^{14}C$).
  • Relative Atomic Mass ($A$): Average mass calculated by weighting isotopes by fractional abundance: $A = A_1 f_1 + A_2 f_2 + \dots + A_n f_n$.

• Electromagnetic Radiation (EMR):

  • Wave properties: Wavelength ($\lambda$), frequency ($\nu$), and speed ($c$).
  • Calculation: c = \nu\times\text{\lambda}, where $c = 3.0 \times 10^8 \, m/s$.
  • Quantum Theory (Max Planck, 1900): Energy is discontinuous and emitted in packets called quanta. $E = h\nu$, where $h = 6.63 \times 10^{-34} \, J \cdot s$.
  • Photoelectric Effect (Einstein, 1905): Light consists of photons. Electrons are ejected if frequency exceeds threshold ($\nu_o$). Kinetic Energy of electron ($KE_e$): $KE_e = \frac{1}{2} mv^2 = h\nu - h\nu_o$.

• Bohr Model of the Hydrogen Atom:

  • Electrons travel in circular orbits. Energy is proportional to distance from the nucleus.
  • Energy Levels: $E_n = -R_H (\frac{1}{n^2})$, where $R_H = 2.18 \times 10^{-18} \, J$.
  • Bohr Radius ($a_o$): 0.53 \, \text{\mathring{A}}. Radius for orbit $n$: $r = n^2 a_o$.
  • Emission Spectrum Series: Lyman ($n_f=1$, UV), Balmer ($n_f=2$, Visible/UV), Paschen ($n_f=3$, IR), Brackett ($n_f=4$, IR).

• Quantum Mechanical Model:

  • Wave-Particle Duality (de Broglie): $\lambda = \frac{h}{mv}$.
  • Heisenberg Uncertainty Principle: Impossible to know position ($\Delta x$) and momentum ($\Delta p$) simultaneously: (\Delta x)(\Delta p) \bge \frac{h}{4\pi}.
  • Quantum Numbers:
    1. Principal ($n$): Energy level/shell ($1, 2, 3 \dots$). Total orbitals = $n^2$. Capacity = $2n^2$.
    2. Angular Momentum ($l$): Shape of orbital. Range $0$ to $n-1$. ($l=0$: s; $l=1$: p; $l=2$: d; $l=3$: f).
    3. Magnetic ($m_l$): Orientation. Range $-l$ to $+l$.
    4. Spin ($m_s$): Direction of electron spin. Value: $+\frac{1}{2}$ or $-\frac{1}{2}$.

• Electron Configuration Principles:

  • Aufbau Principle: Electrons fill lowest energy orbitals first.
  • Pauli Exclusion Principle: No two electrons in an atom can have the same four quantum numbers.
  • Hund’s Rule: Degenerate orbitals are each occupied by one electron before pairing.
  • Exceptions: Chromium ($Cr, Z=24$) is $[Ar] 4s^1 3d^5$; Copper ($Cu, Z=29$) is $[Ar] 4s^1 3d^{10}$.

• Periodic Trends:

  • Atomic Radius: Increases down a group (increasing $n$), decreases across a period (increasing effective nuclear charge $Z_{eff}$).
  • Ionization Energy (IE): Energy to remove an electron. Increases across a period, decreases down a group.
  • Electron Affinity (EA): Energy change when adding an electron. Becomes more negative across a period.
  • Electronegativity: Ability to attract bond pairs. Increases across a period, decreases down a group. Fluorine is highest ($4.0$).

Unit 2: Chemical Bonding

• Types of Chemical Bonds:

  • Ionic Bond: Electrostatic attraction between cations (usually metals) and anions (non-metals). Formed by electron transfer.
  • Covalent Bond: Formed by sharing electron pairs between non-metals.
  • Metallic Bond: Attraction between metal cations and a delocalized "sea" of valence electrons.

• Ionic Bonding Details:

  • Lattice Energy ($U$): Enthalpy change to separate 1 mol of ionic solid into gaseous ions. U \bpropto \frac{q_1 \times q_2}{r}.
  • Born-Haber Cycle: Thermodynamics steps to calculate lattice energy using Hess’s Law.

• Covalent Bonding Details:

  • Lewis Structures: Uses dots for valence electrons and dashes for bonds.
  • Octet Rule Exceptions:
    1. Incomplete octet (e.g., $BeCl_2, BF_3$).
    2. Expanded octet (e.g., $PCl_5, SF_6$).
    3. Odd-electron molecules (e.g., $NO, NO_2$ - free radicals).
  • Coordinate Covalent Bond: One atom donates both electrons to the bond (e.g., $NH_4^+, H_3O^+$).
  • Resonance: Occurs when more than one valid Lewis structure can be drawn; the true structure is a hybrid.

• VSEPR Theory (Molecular Geometry):

  • Predicts shape based on minimizing repulsion between electron pairs.
  • Repulsion Strength: Lone pair vs lone pair > lone pair vs bonding pair > bonding pair vs bonding pair.
  • Shapes:
    • $AB_2$: Linear (180^\bcirc).
    • $AB_3$: Trigonal planar (120^\bcirc).
    • $AB_2E$: Bent.
    • $AB_4$: Tetrahedral (109.5^\bcirc).
    • $AB_3E$: Trigonal pyramidal (NH_3, 107.3^\bcirc).
    • $AB_2E_2$: Bent (H_2O, 104.5^\bcirc).
    • $AB_5$: Trigonal bipyramidal (90^\bcirc, 120^\bcirc).
    • $AB_6$: Octahedral (90^\bcirc).

• Intermolecular Forces:

  • Dipole-Dipole: Between polar molecules.
  • Hydrogen Bonding: Strong dipole-dipole force involving H bonded to N, O, or F.
  • London Dispersion Forces: Weakest, between all molecules due to temporary dipoles.

• Bonding Theories:

  • Valence Bond (VB) Theory: Bonds form via orbital overlap (sigma $\sigma$ or pi $\pi$).
  • Hybridization: Mixing atomic orbitals ($sp, sp^2, sp^3, sp^3d, sp^3d^2$).
  • Molecular Orbital (MO) Theory: Formation of bonding and antibonding (\sigma^*, \bpi^*) orbitals.
  • Bond Order: $\frac{1}{2} (\text{bonding electrons} - \text{antibonding electrons})$.

• Crystalline Solids:

  • Ionic: High MP, brittle, conduct only when molten/aqueous (e.g., $NaCl$).
  • Molecular: Low MP, non-conductors (e.g., $CO_2$, Ice).
  • Covalent Network: Extremely high MP, hard (e.g., Diamond, Graphite, $SiO_2$).
  • Metallic: Malleable, ductile, high conductivity (e.g., $Cu, Fe$).

Unit 3: Physical States of Matter

• Kinetic Molecular Theory of Matter:

  1. Matter is made of constantly moving particles.
  2. Particles possess Kinetic Energy ($KE = \frac{1}{2}mv^2$) and Potential Energy.
  3. Temperature measures average KE.
  4. Spaces exist between particles.
  5. Attractive forces (intermolecular) vary by state.

• Properties by State:

  • Solids: Fixed shape/volume, high density, incompressible.
  • Liquids: Fixed volume, no fixed shape, slightly compressible, high density.
  • Gases: No fixed shape/volume, low density, highly compressible.
  • Plasma: Ionized gas at extreme temperatures; mixture of positive ions and electrons.

• Gas Laws:

  • Pressure: $P = \frac{\text{Force}}{\text{Area}}$. $1 \, atm = 760 \, mmHg = 101,325 \, Pa$.
  • Boyle’s Law: Constant $T, n$: $P_1V_1 = P_2V_2$.
  • Charles’ Law: Constant $P, n$: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$ ($T$ must be in Kelvin: K = \bcirc C + 273).
  • Gay-Lussac’s Law: Constant $V, n$: $\frac{P_1}{T_1} = \frac{P_2}{T_2}$.
  • Combined Gas Law: $\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}$.
  • Avogadro’s Law: $\frac{V_1}{n_1} = \frac{V_2}{n_2}$. Molar volume of ideal gas at STP (0^\bcirc C, 1 \, atm) = $22.4 \, dm^3/mol$.
  • Ideal Gas Equation: $PV = nRT$. Gas constant $R = 0.082 \, atm \cdot L \cdot mol^{-1} \cdot K^{-1}$ or $8.314 \, J \cdot mol^{-1} \cdot K^{-1}$.
  • Graham’s Law of Diffusion: Rate of diffusion $r$ is inversely proportional to $\sqrt{M}$ or $\sqrt{d}$: \frac{r_1}{r_2} = \bsqrt{\frac{M_2}{M_1}}.

• Liquid State Phenomena:

  • Evaporation: Surface molecules escaping liquid phase.
  • Vapor Pressure: Partial pressure of vapor in dynamic equilibrium with liquid.
  • Boiling Point: Temperature where vapor pressure equals external atmospheric pressure.
  • Energetics: \Delta H_{vap} = -\bDelta H_{cond}. Endothermic: vaporization; Exothermic: condensation.

• Phase Changes:

  • Sublimation: Solid to gas.
  • Deposition: Gas to solid.
  • Fusion (Melting): Solid to liquid. $\Delta H_{fus}$ is energy required to melt 1 mol.
  • Heating Curve: A plot of temperature vs. heat added. Plateaus indicate phase changes where temperature remains constant.

Unit 4: Chemical Kinetics

• Reaction Rate: The change in concentration of a reactant or product per unit time.

  • Formula: For $aA + bB \rightarrow cC+dD$, \text{Rate} = -\frac{1}{a} \frac{\bDelta[A]}{\bDelta t} = \frac{1}{c} \frac{\bDelta[C]}{\bDelta t}.
  • Units: $mol \, dm^{-3}s^{-1}$.

• Rate Measurement:

  • Average Rate: Change over a time interval.
  • Instantaneous Rate: Rate at a specific time, found from the slope of the tangent line to the concentration-time curve.

• Collision Theory:

  • Reaction occurs only if:
    1. Particles collide.
    2. Particles have correct orientation.
    3. Collision energy \bge Activation Energy ($E_a$).

• Factors Affecting Rate:

  1. Nature of Reactants: Ionic reactions in aqueous solution are very fast.
  2. Surface Area: Higher area (powdered solids) increases collision frequency.
  3. Concentration: Higher concentration increases collision frequency. Pressure increases rate for gaseous reactions.
  4. Temperature: Increases both collision frequency and the fraction of molecules with energy \bge E_a. Rate often doubles for every 10^\bcirc C increase.
  5. Catalyst: Increases rate by providing an alternative path with lower $E_a$. Not consumed.

Unit 5: Chemical Equilibrium

• Nature of Equilibrium:

  • Reversible Reactions: Can proceed in both forward and reverse directions ($\rightleftharpoons$).
  • Dynamic Equilibrium: Rates of forward and reverse reactions are equal. Concentrations of reactants and products remain constant but not necessarily equal.

• Equilibrium Constants:

  • Law of Mass Action: For $aA + bB \rightleftharpoons cC+dD$:
  • $K_C = \frac{[C]^c[D]^d}{[A]^a[B]^b}$.
  • $K_P = \frac{(P_C)^c(P_D)^d}{(P_A)^a(P_B)^b}$.
  • Relationship: K_P = K_C(RT)^{\bDelta n}, where $\Delta n = (c+d)-(a+b)$ for gaseous species.

• Applications of K:

  • $K > 10^{10}$: Reaction goes to completion.
  • $K < 10^{-10}$: Reaction does not proceed.
  • Reaction Quotient ($Q$): Calculated using non-equilibrium concentrations.
    • $Q < K$: Shift right (forward).
    • $Q > K$: Shift left (reverse).
    • $Q = K$: At equilibrium.

• Le Châtelier’s Principle: If a system at equilibrium is disturbed, it shifts to minimize the disturbance.

  • Concentration: Add reactant/remove product $\rightarrow$ shifts Right.
  • Pressure (Gas): Increase pressure $\rightarrow$ shifts to side with fewer gas moles.
  • Temperature: Increase temp $\rightarrow$ shifts toward endothermic direction. Decrease temp $\rightarrow$ shifts toward exothermic direction.
  • Catalyst: Affects rate only; does not shift position or change $K$.

• Industrial Processes:

  • Haber Process (Ammonia): $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g) \, \Delta H = -92 \, kJ/mol$. Optimized at 200\text{--}400 \, atm, 500^\bcirc C, with iron catalyst.
  • Contact Process (Sulphuric Acid): $2SO_2(g) + O_2(g) \rightleftharpoons 2SO_3(g) \, \Delta H = -196 \, kJ/mol$. Uses $V_2O_5$ catalyst at 450^\bcirc C.

Unit 6: Oxygen-Containing Organic Compounds

• Alcohols ($R-OH$):

  • Classification: Monohydric (1 -OH), dihydric (2 -OH), trihydric (3 -OH).
  • Primary (1^\bcirc), Secondary (2^\bcirc), Tertiary (3^\bcirc) based on C attachment.
  • Physical Properties: High BP due to hydrogen bonding. Solubility decreases with increasing C chain.
  • Preparation: Alkenes hydration, alkyl halide hydrolysis, ester hydrolysis.
  • Reactions:
    • Oxidation (1^\bcirc): 1^\bcirc \text{ Alcohol} \rightarrow \text{Aldehyde} \rightarrow \text{Carboxylic Acid}.
    • Oxidation (2^\bcirc): 2^\bcirc \text{ Alcohol} \rightarrow \text{Ketone}.
    • Dehydration: $Alcohol \rightarrow \text{Alkene} + H_2O$ (High temp/Acid).
    • Active Metals: $2ROH + 2Na \rightarrow 2RONa + H_2$.

• Ethers (R-O-R^\bprime):

  • Symmetrical (R groups same) or Unsymmetrical.
  • Nomenclature: Alkoxyalkane (IUPAC).
  • BP lower than alcohols (no H-bonding with self). Good solvents.
  • Preparation: Williamson Synthesis ($RX + RO^- \rightarrow ROR + X^-$).

• Aldehydes ($R-CHO$) and Ketones (R-CO-R^\bprime):

  • Contain the Carbonyl group ($C=O$).
  • Aldehyde suffix: "-al". Ketone suffix: "-one".
  • Aldehydes easily oxidize to carboxylic acids; ketones do not.

• Carboxylic Acids ($R-COOH$):

  • Nomenclature: Alkanoic acid.
  • High BP (dimerize via 2 H-bonds).
  • Acidity: Weak acids; react with metals, bases, carbonates to form salts (e.g., sodium ethanoate).
  • Preparation: Oxidation of 1^\bcirc alcohols; oxidation of alkylbenzenes.

• Esters (R-COOR^\bprime):

  • Formed via Esterification: $Acid + Alcohol \rightleftharpoons Ester + H_2O$ (acid catalyst).
  • Saponification: Base hydrolysis of esters to form soap and alcohol.
  • Pleasant odors (fruits/flowers).

• Fats and Oils:

  • Triesters of glycerol (Triglycerides).
  • Fats: Saturated, solid at RT, animal origin.
  • Oils: Unsaturated, liquid at RT, plant origin.
  • Hardening of Oils: Hydrogenation of vegetable oils to form solid fats (e.g., margarine).
  • Rancidity: Spoilage due to hydrolysis or oxidation at double bonds.

• Nature of Equilibrium:
  • Reversible Reactions: Can proceed in both forward and reverse directions ($\rightleftharpoons$).
  • Dynamic Equilibrium: Rates of forward and reverse reactions are equal. Concentrations of reactants and products remain constant but not necessarily equal.
• Equilibrium Constants:
  • Law of Mass Action: For $aA + bB \rightleftharpoons cC+dD$:
  • $K_C = \frac{[C]^c[D]^d}{[A]^a[B]^b}$.
  • $K<em>P = \frac{(P</em>C)^c(P<em>D)^d}{(P</em>A)^a(P_B)^b}$.
  • Relationship: $K<em>P = K</em>C(RT)^{\Delta n}$, where $\Delta n = (c+d)-(a+b)$ for gaseous species.
  • Applications of K:
  • K > 10^{10}: Reaction goes to completion.
  • Reaction Quotient ($Q$): Calculated using non-equilibrium concentrations.
    • Q < K: Shift right (forward).
    • Q > K: Shift left (reverse).
    • $Q = K$: At equilibrium.
• Le Châtelier’s Principle: If a system at equilibrium is disturbed, it shifts to minimize the disturbance.
  • Concentration: Add reactant/remove product $\rightarrow$ shifts Right.
  • Pressure (Gas): Increase pressure $\rightarrow$ shifts to side with fewer gas moles.
  • Temperature: Increase temp $\rightarrow$ shifts toward endothermic direction. Decrease temp $\rightarrow$ shifts toward exothermic direction.
  • Catalyst: Affects rate only; does not shift position or change $K$.
• Industrial Processes:
  • Haber Process (Ammonia): $N<em>2(g) + 3H</em>2(g) \rightleftharpoons 2NH_3(g) \, \Delta H = -92 \, kJ/mol$. Optimized at $200\text{--}400 \, atm, 500^\circ C$, with iron catalyst.
  • Contact Process (Sulphuric Acid): $2SO<em>2(g) + O</em>2(g) \rightleftharpoons 2SO<em>3(g) \, \Delta H = -196 \, kJ/mol$. Uses $V</em>2O_5$ catalyst at $450^\circ C$.