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Gen Chem (7)
| Question | Answer |
|---|---|
| What is the difference between intramolecular and intermolecular forces? | Intramolecular forces exist within a molecule and hold atoms together (e.g., covalent, ionic, metallic bonds). Intermolecular forces exist between neighboring molecules. Intramolecular forces are always stronger than intermolecular forces |
| What are the three types of intermolecular forces, ranked weakest to strongest? | London Dispersion Forces (weakest) → Dipole-Dipole Forces → Hydrogen Bonds (strongest) |
| What are London Dispersion Forces (LDFs)? | Transient dipole moments caused by non-uniform distribution of charges (temporary electron fluctuations). They exist between ALL molecules and are the weakest IMF. |
| What are dipole-dipole forces? | Attractive forces between two polar molecules with permanent dipoles. |
| What are hydrogen bonds? | A strong form of dipole-dipole interaction between a hydrogen atom covalently bonded to N, O, or F and another electronegative molecule. |
| What are the four types of intramolecular forces? | Polar covalent, nonpolar covalent, ionic, and metallic bonds. |
| What is a polar covalent bond? | Forms between two non-metals of slightly different electronegativities. Electrons are shared unequally — the more electronegative atom holds the shared electrons more strongly. |
| What is a nonpolar covalent bond? | Forms between the same non-metals or non-metals with very similar electronegativities. Electrons are shared equally. |
| What is an ionic bond? | Forms between a metal and a non-metal. There is a complete transfer of valence electrons from the metal (becomes cation) to the non-metal (becomes anion). The oppositely charged ions attract each other through electrostatic forces. |
| What is a metallic bond? | Forms between metal cations and freely moving valence electrons throughout a lattice. The loosely held outermost electrons form a "sea" of delocalized electrons, which is why metals conduct electricity. |
| Are all intramolecular forces stronger than all intermolecular forces? | Yes. Even the weakest intramolecular force (polar covalent) is stronger than the strongest intermolecular force (hydrogen bonding). |
| What is fusion? | The phase change from solid → liquid. It is endothermic (absorbs heat). Organization decreases as the solid structure breaks apart. |
| What is crystallization (freezing)? | The phase change from liquid → solid. It is exothermic (releases heat). Freely moving molecules become organized into a fixed pattern. |
| What is vaporization? | The phase change from liquid → gas. It is endothermic (absorbs heat). Particles become more disorderly as they escape the liquid. |
| What is condensation? | The phase change from gas → liquid. It is exothermic (releases heat). Gas particles slow down and become more orderly. |
| What is sublimation? | The phase change from solid → gas, skipping the liquid phase entirely. It is endothermic. A common example is dry ice (CO₂). |
| What is deposition? | The phase change from gas → solid, skipping the liquid phase entirely. It is exothermic. Gas particles acquire a very orderly configuration as heat is released. |
| What is the difference between a physical change and a chemical change during a phase transition? | Phase changes are physical changes — intermolecular bonds are broken, but intramolecular bonds are NOT broken or formed. |
| Rank the enthalpy changes: fusion, vaporization, sublimation. | ΔH°fusion < ΔH°vaporization < ΔH°sublimation. Sublimation requires the most energy; fusion requires the least. |
| Why is heat of vaporization greater than heat of fusion? | Vaporization must break almost ALL intermolecular forces between liquid molecules. Fusion only needs enough energy for molecules to escape their crystal lattice sites — other attractions remain intact in the liquid. |
| Why is heat of sublimation greater than heat of vaporization? | Solids have stronger intermolecular forces than liquids. It therefore takes more energy to bring a solid all the way to a gas than to bring a liquid to a gas. |
| What is entropy (S)? | A measure of disorder/randomness in a system. Positive ΔS = increased disorder. Negative ΔS = decreased disorder (increased order). |
| What is enthalpy (H) and ΔH? | Enthalpy measures heat/energy in a thermodynamic system. ΔH is heat gained or lost at constant pressure. ΔH > 0 = endothermic; ΔH < 0 = exothermic. It is a state function (path-independent). |
| What is the Standard Enthalpy of Formation (ΔHf)? | The enthalpy change when exactly 1 mole of a substance is formed from its pure elements at STP. Reactants must be pure elements at STP; only 1 mole of a single product can form. |
| What are the axes of a phase diagram? | X-axis = temperature (°C); Y-axis = pressure (kPa). |
| What is the first curve on a phase diagram (leftmost/lowest region boundary)? | The sublimation curve — the solid-gas boundary. It indicates conditions where a substance transitions directly between solid and gas phases. |
| What is the second curve on a phase diagram? | The fusion curve (melting curve) — the solid-liquid boundary. It indicates conditions where a substance transitions between solid and liquid. For water, this curve has a negative slope; for most other substances, it has a positive slope. |
| What is the third curve on a phase diagram (extending to the upper right)? | The vaporization curve — the liquid-gas boundary. It indicates conditions where a substance transitions between liquid and gas. It ends at the critical point. |
| What is the triple point? | The specific temperature and pressure where all three curves meet and all three phases (solid, liquid, gas) coexist simultaneously. |
| What is the critical point and supercritical fluid? | The critical point is where the liquid and gas phases become indistinguishable. Beyond it, the substance is a supercritical fluid. To convert it back to gas, decrease pressure. |
| What are normal melting and boiling points? | The melting point at 1 atm (101 kPa) and the temperature at which vapor pressure equals 1 atm. Some substances lack these if they can't melt/boil at 1 atm. |
| Why does water's fusion curve have a negative slope (unlike most substances)? | Ice is less dense than liquid water due to its spaced-out crystal lattice. Increased pressure at constant temperature collapses the lattice, turning ice into liquid. Most substances have denser solids, so increased pressure gives them a positive slope. |
| What is viscosity and how does temperature affect it? | Viscosity is a fluid's resistance to flow. It decreases with increasing temperature because faster-moving particles have weaker intermolecular interactions and less friction. |
| What is surface tension and how does temperature affect it? | Surface tension is a liquid surface's resistance to external forces, caused by strong IMFs at the surface. It decreases with increasing temperature because IMFs weaken |
| What is vapor pressure? | The equilibrium pressure exerted by vapor above its liquid in a closed system. Increases with temperature. When vapor pressure equals atmospheric pressure, the liquid boils. Weaker IMFs = higher vapor pressure = lower boiling point. |
| What is miscibility? | The degree to which two liquids mix. Miscible liquids blend completely (e.g., alcohol + water). Immiscible liquids form separate layers (e.g., oil + water). |
| What are ionic solids? | Hard, brittle, non-conductive in solid form, and have very high melting points due to strong ionic intramolecular forces. Example: table salt (NaCl). |
| What are metallic solids? | Malleable, ductile, highly conductive, and have high luster. They consist of positively charged metal ions surrounded by a "sea" of delocalized electrons. Melting points and hardness vary. |
| What are covalent network solids? | Hard, non-conductive, and have very high melting points. Atoms are held together by covalent bonds throughout the entire structure. Examples: diamond and graphite. |
| What are molecular solids? | Soft, non-conductive, and have low melting points. Molecules are held together by intermolecular forces (not intramolecular forces), making them weak. Example: ice. |
| What is the difference between amorphous and crystalline solids? | Amorphous solids (e.g., glass) have no long-range molecular order. Crystalline solids have a repeating long-range pattern defined by unit cells (the smallest repeating unit of a crystal lattice). |
| What is a simple cubic unit cell? | Has 1 atom per unit cell. One atom is partially located at each of the 8 corners of the cube, with each corner contributing 1/8 of an atom. Calculation: 1/8 × 8 = 1 atom. |
| What is a body-centered cubic (BCC) unit cell? | Has 2 atoms per unit cell. There is 1 full atom in the center of the cube, plus 8 corners each contributing 1/8 of an atom. Calculation: 1 + (1/8 × 8) = 2 atoms. |
| What is a face-centered cubic (FCC) unit cell? | Has 4 atoms per unit cell. Six faces each contribute 1/2 of an atom (totaling 3), and 8 corners each contribute 1/8 of an atom (totaling 1). Calculation: (1/2 × 6) + (1/8 × 8) = 4 atoms. |
| What fraction of an atom does a corner position contribute to a unit cell? | 1/8 — because each corner is shared among 8 adjacent unit cells. |
| What fraction of an atom does a face position contribute to a unit cell? | 1/2 — because each face is shared between 2 adjacent unit cells. |
| What fraction of an atom does a center position contribute to a unit cell? | 1 (a full atom) — because the center atom is not shared with any other unit cell. |
| Liquids generally have lower density than solids — what is the major exception? | Water. Ice is less dense than liquid water due to its spaced-out crystal lattice, which is why ice floats. |
Created by:
smurtab