Intermolecular Forces, Thermodynamics, and Chemical Kinetics
Intermolecular Forces
F. Orientation/Permanent Dipole-Dipole
Dipoles are oriented; the positive pole of one molecule attracts the negative pole of the nearest molecule. The force (F) of attraction increases with polarity. This is present in liquids and gases. Examples of weak attractions include HCl, NH3, H2O, and ethanol.
F. Dispersion/London Forces
These forces occur between nonpolar molecules due to instantaneous dipoles caused by electron vibration. The force value increases with molecular mass and is present in gases, liquids, and solids. Examples include F2, Cl2, CF4, CCl4, CBr4, and CI4.
F. Hydrogen Bridge
Occurs when a hydrogen atom is covalently bonded to a highly electronegative atom like F, O, or N. The H+ acts as a positive pole and attracts neighboring molecules. Electrostatic in nature, present in liquids at room temperature. Examples include H2S and ice.
Covalent Solids
Atoms are united in three dimensions by covalent bonds, forming networks or atomic crystals. These solids are typically hard, tough, insoluble in water, and if polar, have high melting and boiling points. Examples include diamond (tetrahedral) and graphite (hexagonal).
Chemical Systems and Thermodynamics
Chemical Systems (Sist. Quim.)
A chemical system involves substances that react or mix with each other and the environment. Systems can be open (exchanging matter and energy), closed (exchanging only energy), or isolated (exchanging neither matter nor energy). They can also be homogeneous (uniform state) or heterogeneous (distinct phases).
Thermodynamic Variables
Variables can be extensive (depend on the amount of substance) or intensive (independent of the amount). Processes include isochoric (constant volume), isobaric (constant pressure), isothermal (constant temperature), and adiabatic (no heat exchange, Q = 0).
Internal Energy (E)
Internal energy (U) is an extensive state function representing the sum of all internal energies (kinetic, potential, etc.). ΔU = Ufinal – Uinitial.
First Law of Thermodynamics
The energy of a system can change through heat (Q) and work (W). The system can gain energy from or lose energy to the surroundings. ΔU = Q + W.
Enthalpy (H)
Enthalpy (H) is a thermodynamic quantity representing the heat exchanged by the system with its environment at constant pressure. ΔH = Hfinal – Hinitial. Hess’s Law states that the enthalpy change for a reaction is the same whether it occurs in one step or multiple steps. The overall enthalpy change is the algebraic sum of the enthalpy changes for each step.
Entropy (S)
Entropy (S) measures the degree of disorder in a system. ΔS > 0 indicates increased disorder, while ΔS < 0 indicates decreased disorder. Units are J·kg-1·K-1. ΔS° depends only on the initial and final states.
Second Law of Thermodynamics
A spontaneous process is accompanied by an increase in the total entropy of the universe.
Gibbs Free Energy (G)
Gibbs free energy (G) determines spontaneity at constant temperature and pressure. ΔG° represents the change in free energy. It is a state function depending on the initial and final states. ΔG = ΔH – TΔS. Values are often tabulated.
Chemical Kinetics
Aspects of Chemical Reactions
Reactions can be fast or slow, accelerate or decelerate, and be spontaneous or inhibited. Chemical kinetics studies the rate at which a spontaneous chemical reaction occurs and measures its velocity in the laboratory. Reactions can be rapid (e.g., acid-base neutralization) or slow (no observable changes). Reaction intermediates may be measurable (e.g., I2 + H2 → 2HI). Reactions start when reactants come into contact. Reactant concentrations decrease as they are consumed, while product concentrations increase.
Mean Velocity (Vmean): Change in concentration of product or reactant over a finite time interval. VHI = Δ[HI]/Δt.
Instantaneous Velocity (Vinst): Change in concentration of reactants or products over an infinitesimal time interval. VHI = d[HI]/dt (slope of the concentration vs. time curve). For a linear change, velocity is constant.
Rate Equation and Order of Reaction
The rate equation is a mathematical expression relating the velocity of a reaction to the concentrations of reactants: V = k[A]a[B]b.
Reaction Order: The exponent to which a reactant’s concentration is raised in the rate equation is called the partial order with respect to that reactant. The sum of the exponents is the total order of the reaction.
- Zero order: mol·L-1·s-1
- First order: s-1
- Second order: mol-1·L·s-1
- n-th order: various units
Reaction Mechanism
Complex reactions occur in successive elementary steps. The molecular equation represents the overall reaction without showing intermediates. Collision theory states that the reaction rate is proportional to the number of particle collisions with sufficient energy and proper orientation. Not all collisions are effective; only those with sufficient energy and correct orientation lead to a reaction.
Transition State Theory
This theory states that reactions proceed through an intermediate activated complex. The activated complex is formed when reactants collide with sufficient energy, weakening existing bonds and forming new ones. The energy required for reactants to reach the activated complex is the activation energy (Ea). The Arrhenius equation relates the rate constant (k) to the absolute temperature (T) and activation energy: k = Ae-Ea/RT.
Catalysts and Inhibitors
Catalysts are substances that increase the rate of a reaction without being consumed. They do not appear in the overall reaction equation. Catalysts work by providing an alternative reaction pathway with a lower activation energy, thus increasing the number of effective collisions and the reaction rate. Catalysts change the kinetic and thermodynamic variables of a reaction.
- Homogeneous catalysis: Catalyst is in the same phase as the reactants (e.g., obtaining H2SO4 by the lead chamber process).
- Heterogeneous catalysis: Catalyst is in a different phase from the reactants (e.g., liquid or solid catalyst with gaseous reactants). Reactions occur on the surface of the catalyst, and products are desorbed.
Inhibitors decrease the rate of reactions.