The attractions between particles
London dispersion acts in all particles and generally grows with polarizability. Dipole–dipole attraction occurs between polar molecules. Hydrogen bonding is a particularly important case involving H bonded to N, O, or F interacting with a lone pair. These are attractions between particles; breaking them in a phase change is not the same as breaking every covalent bond inside a molecule.
Dispersion attraction arises because electrons fluctuate, briefly creating uneven charge distributions that induce similar shifts nearby. It is present even in nonpolar particles and can be substantial in large, easily distorted electron clouds. A permanent dipole gives another way for molecules to align and attract. Hydrogen bonding is not an extra covalent bond between whole water molecules; it is a particularly strong, directional intermolecular interaction in a familiar set of cases.
Boiling points often give clues about these forces, but molecular mass alone does not decide the order. Shape changes surface contact, branching can reduce it, and hydrogen bonding can produce a striking deviation. Compare similar-size molecules when isolating one effect. The observed temperature is the result of many particles' collective behavior under a stated pressure, not a direct reading of one isolated pairwise force.
Gases and solutions
An ideal gas model treats particles as tiny, widely spaced objects with negligible attraction. Real gases deviate most at high pressure or low temperature. Dissolving requires disrupting old attractions and forming new ones; “like dissolves like” is a useful starting point, not a universal law. Concentration reports how much solute occupies a specified solution volume.
The gas-law model becomes more plausible when particles are far apart and moving fast enough that their attractions are relatively unimportant. At high pressure, finite particle volume matters; at low temperature, attractions matter more and condensation may approach. Different gases deviate by different amounts. A calculated ideal pressure is therefore a model prediction to compare with conditions, not a promise that all gases behave identically.
In solutions, the solute and solvent must interact favorably enough to compensate for disrupting their prior arrangements. Dissolution can be endothermic and still occur because entropy also matters. A saturated solution is in dynamic balance with undissolved material under fixed conditions. Changing temperature may change solubility, but the direction depends on the particular substance; the blanket claim that heating always dissolves more is false.
Separating and seeing
Chromatography separates substances because they partition differently between mobile and stationary phases. Spectrophotometry relates absorbed light to concentration under appropriate conditions. A calibration curve is evidence connecting instrument response to known concentrations, while a clear solution is not necessarily pure.
In chromatography, one component may spend more time interacting with the stationary material while another travels mostly with the moving solvent. Separation depends on this difference, not simply on one component being “heavier.” The same sample can behave differently if the solvent or stationary phase changes. A spot's position is evidence only under the specified setup, and two substances can sometimes overlap.
For spectrophotometry, select light that the substance absorbs and compare its response with standards. A blank accounts for the solvent and container's contribution. If absorbance is too high, diluting into the instrument's reliable range can improve the measurement; the original concentration is then reconstructed using the dilution factor. Interfering substances, scattering, or chemical changes can make a neat-looking calibration misleading.
Changes of state
Vapor pressure reflects particles escaping into the gas above a liquid and returning. Boiling begins when vapor pressure matches external pressure, so a liquid can boil at a lower temperature at high altitude. A phase diagram maps which phase is stable over ranges of temperature and pressure. Temperature may pause during a phase transition as energy changes particle arrangement instead of average kinetic energy.
At equilibrium above a liquid, evaporation and condensation continue at equal rates. Vapor pressure rises with temperature because a greater fraction of molecules can escape the liquid's attractions. Boiling occurs throughout the liquid when its vapor pressure equals the surrounding pressure; it is not defined as reaching 100 °C. This explains why water boils at different temperatures on a mountain and in a pressure cooker.
A phase diagram is a map of stable phases under combinations of pressure and temperature. Lines show coexistence, not thick regions where a substance is “halfway” between phases. The triple point is a condition where three phases coexist; the critical point marks the end of the ordinary liquid–gas boundary. Real heating curves also depend on whether pressure stays fixed and how quickly heat is supplied.
Signals and concentrations
The Beer–Lambert relationship can connect absorbance with concentration when light path length and chemical behavior are controlled. A spectrophotometer does not directly “see moles”; it measures a signal that must be interpreted through a model or calibration. Dilution changes concentration while conserving the amount of solute transferred, provided nothing reacts or is lost.
Absorbance is logarithmically related to transmitted light: a sample that absorbs more sends less of the selected wavelength through to the detector. Within an appropriate range, absorbance can be proportional to concentration. This is the basis of the Beer–Lambert model, but its calibration can fail if molecules associate, change chemical form, or scatter light. Selecting a suitable wavelength and preparing standards matter as much as writing the equation.
To dilute a solution, a known amount of solute is transferred into a larger final volume. The relationship between initial and final molarity follows conservation of solute moles. It is important to use the final total solution volume, not merely the amount of water added. Mixing two reacting solutions is a different situation: the solute may be consumed, so a simple dilution relationship no longer describes the final composition.