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Describe the four types of intermolecular forces.
London Dispersion Forces: Weak, temporary attractions caused by fluctuations in electron density; present in all molecules but stronger in larger ones. 2. Dipole-Dipole Forces: Stronger attractions between polar molecules due to their permanent dipoles. 3. Hydrogen Bonds: A specific, strong type of dipole-dipole force occurring between hydrogen and highly electronegative atoms (F, O, N). 4. Ion-Dipole Forces: The strongest intermolecular forces that occur between ionic compounds and polar molecules.
Predict the relative strength of dispersion forces based on molecular shape and size.
Dispersion forces increase with the size of the molecule and are stronger in molecules with larger electron clouds. Compact shapes experience weaker dispersion forces compared to elongated shapes due to their reduced surface area.
Define the following physical properties and relate them to intermolecular force strengths: surface tension, viscosity and capillary action.
Surface Tension: The energy required to increase the surface area of a liquid; stronger intermolecular forces result in higher surface tension. Viscosity: A measure of a fluid's resistance to flow, influenced by the strength of intermolecular forces; stronger forces lead to higher viscosity. Capillary Action: The ability of a liquid to flow in narrow spaces; stronger adhesive forces enhance capillary rise.
Define phase transitions and relate phase transition temperature to intermolecular forces.
Phase transitions refer to the changes between solid, liquid, and gas phases. The temperatures at which these transitions occur depend on the strength of intermolecular forces; stronger forces typically result in higher melting and boiling points.
Define vapor pressure and the rate of vaporization and relate each to temperature, surface area and intermolecular forces.
Vapor Pressure: The pressure exerted by a vapor in equilibrium with its liquid or solid phase; it increases with temperature and decreases with stronger intermolecular forces. Rate of Vaporization: The speed at which molecules escape from the liquid phase to the vapor phase; influenced by temperature and surface area, with higher temperatures and larger surface areas leading to faster vaporization.
Use the Clausius-Clapeyron equation and the vapor pressure of a chemical at a specific temperature to solve for the vapor pressure at a different temperature.
The Clausius-Clapeyron equation relates vapor pressure to temperature. By knowing the vapor pressure at one temperature and the heat of vaporization, you can calculate the vapor pressure at another temperature.
ln(P1P2)=R−ΔHvap(T21−T11)
Identify phases, phase changes, triple point and critical point on a phase diagram and describe a supercritical fluid.
A phase diagram illustrates the states of a substance at various temperatures and pressures. The triple point is where all three phases coexist, while the critical point marks the end of distinct liquid and gas phases. A supercritical fluid exists beyond the critical point, exhibiting properties of both gases and liquids.
Use a phase diagram to describe the effect of changes in temperature or pressure on the phase of a substance.
In a phase diagram, increasing temperature generally pushes substances towards the gas phase, while increasing pressure tends to push them towards the liquid or solid phases.
Boiling Point Elevation: ΔTb=iKbm
Freezing Point Depression:
ΔTf=iKfm
Define the terms solute, solvent, solution and aqueous solution.
Solute: The substance being dissolved in a solution. Solvent: The dissolving medium, typically in greater quantity. Solution: A homogeneous mixture of solute and solvent. Aqueous Solution: A solution where water is the solvent.
Define and differentiate between nonelectrolytes, weak electrolytes and strong electrolytes.
Nonelectrolytes: Substances that do not dissociate into ions in solution. Weak Electrolytes: Partially dissociate into ions, resulting in a relatively low level of conductivity. Strong Electrolytes: Completely dissociate into ions, leading to high conductivity.
Define solubility and differentiate unsaturated, saturated and supersaturated solutions.
Solubility: The maximum amount of solute that can dissolve in a given quantity of solvent at a specific temperature. Unsaturated: Contains less solute than can be dissolved. Saturated: Contains the maximum amount of solute that can dissolve at that temperature. Supersaturated: Contains more solute than can be dissolved, temporarily stable under certain conditions.
Describe the effect of temperature and pressure on solubility.
Generally, solubility of solids in liquids increases with temperature, while solubility of gases decreases as temperature increases. Pressure significantly affects gas solubility; increased pressure enhances solubility.
Define, solve for, and convert between molarity, molality and mole fraction.
Molarity: Concentration expressed as moles of solute per liter of solution (mol/L). Molality: Concentration expressed as moles of solute per kilogram of solvent (mol/kg). Mole Fraction: The ratio of the moles of one component to the total moles in the solution; calculations can convert between these measures.
Define and provide examples of colligative properties.
Colligative properties are properties that depend on the number of solute particles in a solution rather than their identity. Examples include vapor pressure lowering, boiling point elevation, and freezing point depression.
For colligative property relationships, be able to conceptually describe and mathematically solve for a missing value.
a. Raoult’s Law: Used to calculate vapor pressure of a solution, taking into account the mole fraction of the solvent and vapor pressure of the pure solvent. b. Boiling Point Elevation and Freezing Point Depression: Use formulas that involve the van 't Hoff factor and number of solute particles.
Determine the expected Van’t Hoff factor given an ionic or covalent compound.
Van’t Hoff factor (i) indicates the number of particles the solute dissociates into in solution; for example, NaCl has an i of 2, while glucose has an i of 1.
Zero Order: [A]=[A]0−kt[A] = [A]_0 - kt[A]=[A]0−kt
First Order: ln[A]=ln[A]0−kt\ln[A] = \ln[A]_0 - ktln[A]=ln[A]0−kt
Second Order: 1[A]=1[A]0+kt\frac{1}{[A]} = \frac{1}{[A]_0} + kt[A]1=[A]01+kt
Use the expected or measured Van’t Hoff factor in the colligative property relationships described in Objective 7.
Apply the Van’t Hoff factor in the calculations for colligative properties to adjust for the number of particles affecting properties such as boiling point elevation and vapor pressure lowering.
Define and differentiate initial, average and instantaneous rate of reaction.
Initial Rate: The rate of reaction at the moment the reactants are mixed. Average Rate: The change in concentration over a specific time interval. Instantaneous Rate: The rate of reaction at a specific point in time, determined using calculus.
Relate rate of reaction to the rate of change for any reactant or product.
The rate of reaction can be expressed as the change in concentration of reactants consumed (negative) or products formed (positive) over time, reflecting the stoichiometry of the balanced equation.
Relate physical state, temperature, concentration and catalysis to reaction rates.
The physical state of reactants (solid, liquid, gas) affects how easily they collide; higher temperatures increase kinetic energy and collision frequency; increased concentration generally leads to more collisions and higher reaction rates; catalysts lower activation energy, increasing the rate without being consumed.
Determine a rate law from experimental evidence.
Rate laws express the relationship between reaction rate and concentrations of reactants, derived from experimental data showing how the rate changes with varying concentrations.
Use the rate law to solve for any missing variable.
With a given rate law and known quantities for rate and concentrations, apply algebraic manipulation to isolate and calculate missing variables.
Use the integrated rate laws to be able to solve for initial concentration, time, rate constant and final concentration given information on the remaining variables.
Integrated rate laws provide mathematical relationships between concentration and time; use these equations to calculate unknowns such as initial concentration or time based on rate constants and observed concentrations.
Determine reaction order from graphs of concentration versus time or from the units of the rate constant.
Review concentrations plotted over time to identify zero-order (linear), first-order (logarithmic), or second-order (inverse) relationships; units of the rate constant also indicate reaction order.
Define half-life and be able to solve for half-life, rate constant, initial concentration or reaction order given information on the remaining variables.
Half-life is the time required for half of the reactant to be consumed. It varies depending on the reaction order; apply respective equations to solve for half-life or other unknowns from given data.
Define the terms and describe the relationship between temperature, frequency factor, activation energy and the rate constant.
Temperature affects the kinetic energy of molecules; the frequency factor relates to how often molecules collide correctly; activation energy is the minimum energy required for a reaction to proceed. The Arrhenius equation links these factors to determine the rate constant.
Describe the role of orientation and collision frequency in collision theory.
Collision theory states that for a reaction to occur, reactant molecules must collide with sufficient energy and proper orientation; higher collision frequency increases the likelihood of successful reactions.
Solve for the relationship between temperature and the rate constant.
Utilize the Arrhenius equation to relate temperature to the rate constant, indicating that as temperature increases, the rate constant generally increases due to increased energy and collision frequency.
Define the terms and identify examples of: reaction mechanism, elementary step, intermediate, molecularity, rate-determining step and homogenous and heterogenous catalyst.
Reaction Mechanism: The detailed sequence of steps by which a reaction occurs. Elementary Step: A single event in the mechanism; an overall reaction can involve multiple elementary steps. Intermediate: A species that is formed and consumed in the mechanism, not present in the final equation. Molecularity: The number of reactant particles involved in an elementary step. Rate-Determining Step: The slowest step that limits the overall rate of the reaction. Homogenous Catalyst: A catalyst that exists in the same phase as the reactants, whereas Heterogeneous Catalyst exists in a different phase.
Evaluate the plausibility of a proposed reaction mechanism.
Assess the validity of a mechanism by ensuring that its elementary steps align with the overall reaction, rates of each step are feasible, and intermediates appear in a logical sequence and concentration that reflects observable measurements.
