Unit 3: Properties of Substances and Mixtures

In this unit, students will explore how atoms come together to create solids, liquids, and gases, and how forces between particles govern the properties of everything around you.

Intermolecular Forces (IMFs)

Intermolecular forces are the attractive forces between separate molecules or particles, and they determine many physical properties such as boiling point, melting point, vapor pressure, and solubility. These forces are weaker than the covalent or ionic bonds that hold atoms together within a molecule, but they are crucial for understanding phase changes and material properties.

London Dispersion Forces (LDFs) occur due to temporary fluctuations in electron distribution, creating instantaneous dipoles that induce dipoles in neighboring molecules. They are present in all molecules and atoms, but are the only IMF in nonpolar substances. The strength of LDFs increases with greater molar mass and surface area because larger electron clouds are more polarizable.

Dipole-Dipole Interactions occur between molecules that have permanent dipoles due to differences in electronegativity between bonded atoms. The partially positive end of one molecule is attracted to the partially negative end of another, leading to stronger attractions than LDFs for molecules of similar size. This explains why polar molecules generally have higher boiling points than nonpolar molecules of similar molar mass.

Hydrogen Bonding is a particularly strong form of dipole-dipole interaction that occurs when hydrogen is covalently bonded to highly electronegative atoms like nitrogen, oxygen, or fluorine. This creates an unusually large partial positive charge on hydrogen, allowing strong attraction to lone pairs on nearby electronegative atoms. Hydrogen bonding explains unique properties such as water’s high boiling point and ice’s lower density compared to liquid water.

Ion-Dipole Interactions occur between ionic compounds and polar molecules, most commonly between salts and water in solution. These interactions are stronger than typical dipole-dipole forces because they involve full charges interacting with partial charges. They play a key role in dissolving ionic compounds and in biochemical processes like protein folding in aqueous environments.
Relative IMF strength follows the general trend: Ion-dipole > Hydrogen bonding > Dipole-dipole > London dispersion forces (for similar-sized molecules). However, molecular size can make LDFs dominate in very large nonpolar molecules, which can even exceed the strength of hydrogen bonding in smaller molecules.
IMFs are directly related to phase change temperatures because stronger intermolecular attractions require more energy to overcome. For example, ethanol (which can hydrogen bond) has a higher boiling point than dimethyl ether (which cannot), even though they have the same molar mass and formula (C₂H₆O).
Understanding IMFs also helps explain solubility: polar solvents dissolve polar solutes and ionic compounds well (“like dissolves like”), while nonpolar solvents dissolve nonpolar substances better. This principle is essential in predicting reaction environments, purifications, and biochemical interactions.

Structure and Properties of Solids

Solids are characterized by fixed shapes and volumes due to the close packing of particles and strong intermolecular or intramolecular forces holding them in place. The arrangement of particles and the types of forces present determine a solid’s melting point, hardness, electrical conductivity, and other physical properties.
Covalent Network Solids consist of atoms connected by an extensive network of covalent bonds throughout the structure. Examples include diamond (carbon atoms in a tetrahedral network) and quartz (SiO₂ with a 3D network of Si–O bonds). These solids are extremely hard, have very high melting points, and are generally poor electrical conductors because they lack free-moving charged particles.
Ionic Solids are made of cations and anions held together by strong electrostatic attractions in a regular repeating lattice. Properties include high melting points, brittleness, and the ability to conduct electricity only when molten or dissolved in water, since ions must be mobile for conduction to occur. Lattice energy, determined by Coulomb’s law, is a key factor in their stability.
Metallic Solids consist of metal cations surrounded by a “sea” of delocalized valence electrons that can move freely throughout the lattice. This electron mobility explains metals’ electrical and thermal conductivity, malleability, and ductility. The strength of metallic bonding varies with the number of delocalized electrons and the size of the metal cations.
Molecular Solids are composed of individual molecules held together by intermolecular forces such as London dispersion forces, dipole-dipole interactions, or hydrogen bonds. They generally have lower melting points than network or ionic solids because IMFs are weaker than covalent or ionic bonds, and they are often soft and poor conductors of electricity.
The arrangement of particles in solids can be either crystalline or amorphous. Crystalline solids have a well-ordered, repeating pattern, leading to sharp melting points. Amorphous solids, such as glass and many polymers, lack a long-range order, resulting in gradual softening over a temperature range rather than a sharp melting point.
The density and packing efficiency of solids are influenced by crystal structure. Close-packed structures like face-centered cubic (FCC) and hexagonal close-packed (HCP) arrangements maximize space utilization, while body-centered cubic (BCC) arrangements are less dense but can be more structurally rigid for certain metals.
Understanding solid structure helps explain materials science applications. For example, diamond’s rigid covalent network makes it ideal for cutting tools, while metallic solids like copper are used for electrical wiring due to their conductivity and ductility.

Structure and Properties of Liquids

Liquids have a definite volume but take the shape of their container, indicating that their particles are close together but not held in fixed positions. Intermolecular forces are strong enough to keep the particles in close contact but weak enough to allow them to slide past one another, giving liquids fluidity.
Viscosity is a measure of a liquid’s resistance to flow. Stronger intermolecular forces, larger molecular sizes, and more complex molecular shapes increase viscosity because particles have more difficulty moving around each other. Temperature also affects viscosity, with higher temperatures generally lowering viscosity due to increased kinetic energy.
Surface tension results from cohesive forces between liquid molecules at the surface, causing the surface to behave like a stretched elastic sheet. Liquids with stronger intermolecular forces, such as water with hydrogen bonding, have higher surface tensions, which explains phenomena like water droplets forming spheres.
Cohesion and Adhesion describe the forces within and between substances. Cohesion is the attraction between molecules of the same substance, while adhesion is the attraction between a liquid’s molecules and another surface. The balance between these forces determines phenomena like meniscus shape in a graduated cylinder.
Capillary action occurs when adhesion between a liquid and a solid surface is stronger than the cohesion between the liquid’s molecules, causing the liquid to rise or fall in a narrow tube. This principle is important in biological systems like water transport in plants through xylem tissue.
The evaporation rate of a liquid depends on intermolecular force strength, surface area, and temperature. Stronger IMFs slow evaporation, while higher temperatures increase it by providing more molecules with enough kinetic energy to escape into the gas phase.
Boiling point and vapor pressure are inversely related: liquids with stronger intermolecular forces have higher boiling points and lower vapor pressures. Understanding this relationship is critical in distillation, pressure cooking, and predicting solvent behavior in reactions.
Liquids are incompressible compared to gases because their particles are already closely packed. This property is exploited in hydraulic systems, where a small applied force can transmit pressure effectively throughout the liquid.

Structure and Properties of Gases

Gases have neither a definite shape nor a definite volume, expanding to fill any container they occupy. This behavior is due to the large distances between gas particles compared to their sizes, which makes intermolecular forces negligible under most conditions.
The Kinetic Molecular Theory explains gas behavior by assuming particles move in constant, random motion, collide elastically, and have negligible volume compared to the space between them. It also assumes that the average kinetic energy of gas particles is directly proportional to the gas’s temperature in kelvins.
Pressure in gases results from collisions of particles with container walls. The frequency and force of these collisions depend on temperature, volume, and the number of particles present, as described by the ideal gas law \( PV = nRT \).
Real gases deviate from ideal gas behavior at high pressures and low temperatures. Under these conditions, particle volume is no longer negligible and intermolecular attractions affect motion, leading to lower pressures than predicted by the ideal gas law. The van der Waals equation accounts for these deviations.
Diffusion describes the mixing of gas particles due to random motion, while effusion is the movement of gas particles through a tiny hole without collisions. Graham’s law shows that lighter gases diffuse and effuse faster than heavier gases because velocity is inversely proportional to the square root of molar mass.
Gas density can be determined from the ideal gas law by rearranging it to \( \rho = \frac{PM}{RT} \), where \( M \) is molar mass. This relationship is useful for identifying unknown gases or verifying purity.
Temperature changes impact particle motion significantly. Increasing temperature increases kinetic energy, which can overcome intermolecular forces in condensed phases, while decreasing temperature slows motion and can cause condensation into liquids or solids.
Understanding gas properties connects to solution chemistry, reaction kinetics, and thermodynamics. For example, gas solubility in liquids often follows Henry’s law, which links gas pressure above a liquid to its concentration in the liquid phase.

Solutions and Mixtures

A solution is a homogeneous mixture of two or more substances in which the solute is uniformly dispersed in the solvent. The particle size of the solute is at the molecular or ionic level, meaning the mixture is uniform throughout and cannot be separated by filtration.
Mixtures can be classified as homogeneous or heterogeneous. Homogeneous mixtures have uniform composition and appearance, such as saltwater or air, while heterogeneous mixtures have visibly different components or phases, like oil and water or granite.
Solubility is influenced by factors such as temperature, pressure, and the nature of the solute and solvent. For most solids, solubility increases with temperature, whereas for gases, solubility decreases with increasing temperature but increases with higher pressure (Henry’s law).
The process of dissolving involves breaking solute-solute and solvent-solvent interactions and forming solute-solvent interactions. A solution forms spontaneously if the overall energy change is favorable and entropy increases, as explained by thermodynamic principles.
Concentration measures how much solute is present in a given quantity of solvent or solution. Common units include molarity (M), molality (m), mass percent, volume percent, and mole fraction. Accurate concentration measurements are critical for stoichiometric calculations in chemical reactions.
Solutions can be unsaturated, saturated, or supersaturated. An unsaturated solution can dissolve more solute at a given temperature, a saturated solution holds the maximum amount of solute in equilibrium with undissolved solute, and a supersaturated solution contains more than the equilibrium amount, making it unstable.
Colligative properties depend on the number of solute particles rather than their identity. These include vapor pressure lowering, boiling point elevation, freezing point depression, and osmotic pressure. These effects occur because solute particles disrupt the solvent’s ability to change phases.
In ionic solutions, dissociation produces ions that can conduct electricity, creating an electrolyte solution. Strong electrolytes fully dissociate, while weak electrolytes partially dissociate, leading to lower conductivity.
Mixtures and solutions play essential roles in analytical chemistry, biological systems, and industrial processes. For example, electrolyte balance in the human body depends on ionic solutions, while chemical manufacturing relies on controlled solution concentrations for reaction efficiency.

Photoelectric Effect and Photon Properties

The photoelectric effect occurs when light of sufficient frequency strikes a metal surface and causes the ejection of electrons. This phenomenon provided crucial evidence for the particle nature of light, as it could not be explained by classical wave theory alone.
According to Albert Einstein’s explanation, light consists of discrete packets of energy called photons, each with energy given by \( E = hf \), where \( h \) is Planck’s constant and \( f \) is the frequency of light. The energy of each photon is directly proportional to its frequency and inversely proportional to its wavelength.
The ejection of an electron occurs only if the photon’s energy exceeds the material’s work function, which is the minimum energy required to remove an electron from the surface. If \( hf > \phi \), where \( \phi \) is the work function, the excess energy becomes the kinetic energy of the ejected electron: \( KE = hf - \phi \).
The photoelectric effect demonstrates that increasing light intensity increases the number of ejected electrons, but does not affect their kinetic energy unless the frequency changes. In contrast, increasing frequency above the threshold increases the kinetic energy of emitted electrons.
Photon properties are dual in nature, displaying both particle and wave characteristics. While they carry discrete quanta of energy, they also exhibit wave-like behaviors such as diffraction and interference, illustrating the concept of wave-particle duality in quantum mechanics.
The experimental results of the photoelectric effect supported quantum theory and contributed to the development of modern physics, leading to advancements in technologies like solar cells, photodetectors, and quantum communication systems.
Understanding photon energy is critical in spectroscopy, as photons can excite electrons in atoms and molecules to higher energy levels. The specific photon frequencies required for these transitions provide unique spectral lines that identify elements and compounds.

Spectroscopy Techniques

Spectroscopy is the study of the interaction between matter and electromagnetic radiation, used to determine structural and compositional information about a sample. Different techniques focus on various regions of the electromagnetic spectrum, each revealing specific details about atoms or molecules.
Ultraviolet-Visible (UV-Vis) Spectroscopy measures the absorption of UV and visible light by electrons in molecules, often promoting them from lower-energy orbitals to higher-energy orbitals. This technique is valuable for studying conjugated systems, transition metal complexes, and determining solution concentrations using Beer’s law \( A = \varepsilon bc \).
Infrared (IR) Spectroscopy detects vibrations in chemical bonds by measuring absorption of infrared light. Each bond type vibrates at characteristic frequencies, producing a spectrum that acts as a “molecular fingerprint” to identify functional groups in organic and inorganic compounds.
Nuclear Magnetic Resonance (NMR) Spectroscopy uses strong magnetic fields and radiofrequency pulses to probe the environments of atomic nuclei, typically \( ^1\text{H} \) or \( ^{13}\text{C} \). The resulting spectra reveal connectivity, molecular structure, and the number of chemically distinct atoms.
Mass Spectrometry (MS) is not a spectroscopic technique in the strictest sense but is often grouped with them due to its analytical utility. It ionizes molecules and separates ions based on mass-to-charge ratio, providing precise molecular mass and structural fragmentation patterns.
Photoelectron Spectroscopy (PES) uses high-energy photons to eject electrons from atoms or molecules and measures their kinetic energy to determine binding energies. This provides information about electron configuration, orbital energies, and chemical environments.
Each spectroscopy technique complements the others, giving a more complete picture of molecular structure and properties. For example, IR spectroscopy might identify functional groups, NMR might determine connectivity, and MS might confirm molecular mass, allowing cross-validation of results.
Spectroscopy connects directly to quantum theory, as each observed absorption or emission corresponds to a specific energy difference between quantized states. This principle makes spectroscopy one of the most important tools in both research and industry for qualitative and quantitative analysis.

Chromatography

Chromatography is a separation technique that relies on differences in how substances distribute between a stationary phase and a mobile phase. Components of a mixture move at different rates depending on their interactions with each phase, allowing separation and analysis.
There are several types of chromatography, including paper chromatography, where a liquid solvent moves through paper carrying solutes at different rates, and thin-layer chromatography (TLC), which uses a coated plate as the stationary phase for faster separations and visualization with indicators.
Column chromatography uses a packed column of solid stationary phase through which a liquid mobile phase passes, separating compounds based on polarity, size, or affinity. Gas chromatography (GC) uses a gas as the mobile phase and is especially useful for volatile compounds.
The separation in chromatography is often quantified using the retention factor (Rf), which compares the distance a substance travels to the distance traveled by the solvent front. This value helps identify compounds when compared to known standards.

Distillation

Distillation is a separation technique that relies on differences in boiling points of substances. A mixture is heated until one component vaporizes, and the vapor is then condensed back into a liquid in a separate container.
Simple distillation works best when the boiling points of components differ significantly, while fractional distillation uses a fractionating column to achieve better separation when boiling points are closer together.
Distillation can also be used to purify solvents, remove impurities, or separate liquid products in chemical synthesis. For example, petroleum refining uses fractional distillation to separate hydrocarbons into useful fractions like gasoline, kerosene, and diesel.
Both chromatography and distillation are essential in analytical and industrial chemistry. Chromatography excels in separating small amounts for identification or analysis, while distillation is favored for large-scale purification and isolation of liquid mixtures.

Kinetic Molecular Theory

Kinetic Molecular Theory (KMT) explains the behavior of gases at the particle level by making a series of assumptions about gas particles and their motion. It connects microscopic motion to macroscopic properties like pressure, temperature, and volume.
The theory assumes that gas particles are in constant, random motion and that collisions between particles and with container walls are perfectly elastic, meaning no kinetic energy is lost during collisions. This explains why gases exert constant pressure on container walls.
KMT also assumes that the volume of individual gas particles is negligible compared to the total volume of the gas. This assumption holds true for ideal gases but can break down for real gases under high pressure, where particle volume becomes significant.
Another assumption is that there are no intermolecular forces between gas particles. In reality, attractive or repulsive forces exist, but they are small enough at low pressures and high temperatures that gases behave nearly ideally.
The average kinetic energy of gas particles is directly proportional to the absolute temperature of the gas. This relationship is expressed as \( KE_{\text{avg}} = \frac{3}{2}RT \) for one mole of gas particles, where \( R \) is the gas constant and \( T \) is temperature in kelvins.
Because temperature is a measure of the average kinetic energy of particles, gases at the same temperature have the same average kinetic energy regardless of their molar mass. However, lighter molecules move faster than heavier ones at the same temperature.
KMT explains gas laws such as Boyle’s law (pressure inversely proportional to volume at constant temperature) and Charles’s law (volume directly proportional to temperature at constant pressure) as consequences of particle collisions and motion.
While KMT provides a good model for ideal gases, deviations occur for real gases, particularly under low temperature or high pressure, where intermolecular forces and particle size can no longer be ignored. These deviations are accounted for by the van der Waals equation.

Gas Laws and Their Molecular Interpretations

Gas laws describe the relationships between pressure (\(P\)), volume (\(V\)), temperature (\(T\)), and amount of gas in moles (\(n\)), and are all consistent with the postulates of the Kinetic Molecular Theory. These laws link macroscopic measurements to microscopic particle motion.
Boyle’s Law states that \( P \propto \frac{1}{V} \) at constant \(T\) and \(n\). From a molecular perspective, decreasing volume increases the frequency of particle collisions with container walls, which increases pressure. Conversely, increasing volume reduces collision frequency, lowering pressure.
Charles’s Law states that \( V \propto T \) at constant \(P\) and \(n\). As temperature increases, the average kinetic energy of gas particles increases, causing them to move faster and collide with the walls more forcefully. To maintain constant pressure, the volume must expand.
Avogadro’s Law states that \( V \propto n \) at constant \(P\) and \(T\). Adding more moles of gas increases the number of particles, which increases total collision frequency. To keep pressure constant, the volume must expand proportionally.
Combined Gas Law integrates Boyle’s, Charles’s, and Gay-Lussac’s laws into \( \frac{PV}{T} = \text{constant} \) for a fixed amount of gas. It allows prediction of changes when more than one variable changes at a time.
Ideal Gas Law \( PV = nRT \) combines all individual gas laws into a single equation, where \( R \) is the universal gas constant. It assumes particles have negligible volume and no intermolecular forces, consistent with KMT’s assumptions.
Dalton’s Law of Partial Pressures states that the total pressure of a mixture is the sum of the partial pressures of each gas. From a molecular standpoint, each type of particle contributes independently to the total pressure through its collisions with the walls.
Understanding gas laws at the molecular level helps explain deviations from ideal behavior in real gases, which occur when assumptions like negligible particle volume or absence of intermolecular forces no longer hold, especially under high pressure or low temperature.

Beer’s Law

Beer’s Law (or Beer–Lambert Law) describes the quantitative relationship between the absorbance of light by a solution and its concentration. It is expressed as \( A = \varepsilon b c \), where \( A \) is absorbance (no units), \( \varepsilon \) is molar absorptivity (\( \text{L·mol}^{-1}\text{·cm}^{-1} \)), \( b \) is the path length of the sample cell (cm), and \( c \) is the concentration (mol/L).
Absorbance is a measure of how much light is absorbed by a solution at a specific wavelength, and it is related to transmittance (\( T \)) by the equation \( A = -\log(T) \), where \( T \) is the fraction of light that passes through the sample.
The molar absorptivity (\( \varepsilon \)) is a constant for a given substance at a specific wavelength, reflecting how strongly the substance absorbs light. Larger \( \varepsilon \) values mean stronger absorption per unit concentration and path length.
The path length (\( b \)) refers to the distance light travels through the solution in the cuvette. In most lab spectrophotometers, this is standardized at 1.00 cm, simplifying calculations.
Beer’s Law allows chemists to determine the concentration of an unknown solution by measuring its absorbance and using a calibration curve or the known values of \( \varepsilon \) and \( b \). This is widely used in analytical chemistry for colorimetric and spectrophotometric analysis.
For Beer’s Law to hold, several conditions must be met: the solution must be homogeneous, the absorbing species must not scatter light (no turbidity), and the system must follow a linear relationship between absorbance and concentration within the measured range.
Deviations from Beer’s Law can occur at high concentrations due to intermolecular interactions, chemical equilibria shifts, or instrumental limitations, leading to nonlinear calibration curves. Diluting samples often restores linearity.
Applications of Beer’s Law include determining the concentration of proteins, nucleic acids, metal ions, and other analytes, as well as monitoring reaction progress by measuring absorbance changes over time.

Common Misconceptions – Unit 3: Properties of Substances and Mixtures

Many students believe that the Kinetic Molecular Theory applies perfectly to all gases under all conditions. In reality, it only describes ideal gases, and real gases deviate significantly under high pressure or low temperature because particle size and intermolecular forces become important.
Some assume that if two gases are at the same temperature, heavier gas molecules must have greater kinetic energy. In fact, all gases at the same temperature have the same average kinetic energy, but lighter molecules move faster while heavier molecules move slower.
When using gas laws, students often forget that temperature must be converted to kelvins. Using Celsius in the formulas will produce incorrect results because Kelvin is the absolute temperature scale linked directly to particle energy.
In Beer’s Law, many confuse absorbance with transmittance. Absorbance is logarithmically related to transmittance, meaning that a small change in transmittance can cause a large change in absorbance, and vice versa.
Students sometimes think Beer’s Law applies regardless of concentration range. In reality, at high concentrations, intermolecular interactions or scattering can cause deviations from linearity, so calibration curves are only valid in a specific concentration range.
A common error in spectrophotometry is forgetting that molar absorptivity (\( \varepsilon \)) depends on wavelength. Measuring at a wavelength far from the analyte’s maximum absorbance will decrease sensitivity and accuracy.
In chromatography, students may believe that retention time depends solely on molecular size. While size can play a role, polarity and interactions with the stationary and mobile phases are often the dominant factors in separation.
Many assume that distillation perfectly separates components with different boiling points. In reality, impurities can still carry over, and fractional distillation is needed for better separation of substances with close boiling points.