Liquids and Solids | General Chemistry 2

Intramolecular vs. intermolecular forces:

• Intramolecular forces: Forces within a molecule, such as covalent, ionic, and metallic bonds.
• Intermolecular forces: Forces between molecules, responsible for the physical properties of substances.

Types of intermolecular forces:

• Ion-dipole forces (40-600 kJ.mol^-1)
• Hydrogen bonding (10-40 kJ.mol^-1)
• Dipole-dipole interactions (5-25 kJ.mol^-1)
• Ion-induced dipole forces (3-15 kJ.mol^-1)
• Dipole-induced dipole forces (2-10 kJ.mol^-1)
• London dispersion forces (0.05-40 kJ.mol^-1)

Hydrogen bonding:

Hydrogen bonding is a strong electrostatic interaction that occurs between a hydrogen atom bonded to highly electronegative atoms (like N, O, F) and an electronegative atom that has a lone pair of electrons. Hydrogen bonding is crucial for the properties of water and biological molecules like DNA.
 

Dipoles and induced dipoles:

• Dipoles: Molecules with a permanent molecular dipole due to an uneven distribution of electrons.
• Induced dipoles: Nonpolar molecules that develop a temporary dipole moment when in the presence of a charged or polar molecule.
   

An isolated O₂ molecule is nonpolar.
However, when a polar water molecule approaches, the electron distribution in O₂ is rearranged, creating an induced dipole.

London dispersion forces:

London dispersion forces are electrostatic interactions that occur between two nonpolar molecules. 
Electrons within a molecule are constantly moving. At any given moment, they may be asymmetrically distributed, creating an instantaneous dipole moment. This temporary dipole can induce a dipole in a neighboring nonpolar molecule, leading to weak electrostatic interactions known as London dispersion forces.

Importance of intermolecular forces:

• Boiling and melting points: Substances with stronger intermolecular forces have higher boiling and melting points.
• Solubility: Polar solvents dissolve polar solutes, and nonpolar solvents dissolve nonpolar solutes, based on the principle "like dissolves like."
• Viscosity and surface tension: Stronger intermolecular forces result in higher viscosity and surface tension in liquids.

Viscosity (η):

Viscosity is a measure of a liquid's resistance to flow. It is influenced by the internal friction between layers of molecules as they move past one another.

Factors affecting viscosity are:

• Intermolecular forces: Stronger intermolecular forces lead to higher viscosity because molecules are more strongly attracted to each other, making it harder for them to move past one another.
• Temperature: As temperature increases, viscosity decreases because the increased kinetic energy allows molecules to overcome intermolecular forces more easily.
• Molecular structure: Long, flexible molecules tend to have higher viscosities than shorter-chain or spherical molecules due to increased entanglement and intermolecular interactions.

Surface Tension (in J.m²):

Surface tension is the energy required to increase the surface area of a liquid. It results from the imbalance of intermolecular forces at the surface of a liquid, where molecules experience a net inward force.

Factors affecting surface tension are:

• Intermolecular forces: Liquids with stronger intermolecular forces have higher surface tension. For example, water has a high surface tension due to hydrogen bonding.
• Surfactants: Surfactants are substances that reduce surface tension by decreasing the intermolecular forces between adjacent molecules at the surface, allowing the liquid to spread more easily.

Capillary Action:

Capillary action is the ability of a liquid to flow in narrow spaces without the assistance of external forces, such as gravity. This phenomenon is a result of the adhesive forces between the liquid and the surface of the capillary, as well as the cohesive forces within the liquid.

Density and compressibility:

• Density: Liquids have a higher density than gases because their molecules are more closely packed. The density of a liquid generally decreases slightly with increasing temperature as the liquid expands.
• Compressibility: Liquids are nearly incompressible, meaning that their volume does not change significantly under pressure due to the close packing of molecules.

Vapor pressure:

Vapor pressure is the pressure exerted by the vapor in equilibrium with its liquid phase at a given temperature in a closed system. This pressure arises as molecules escape from the liquid surface into the vapor phase.

The pressure exerted by the vapor on the liquid increases until equilibrium is reached, where the rate of evaporation equals the rate of condensation. At this point, the vapor pressure remains constant and is referred to as the equilibrium vapor pressure.

Factors affecting vapor pressure:

• Temperature: Vapor pressure increases with temperature. As temperature rises, more molecules gain sufficient kinetic energy to overcome intermolecular forces and escape into the vapor phase, thus increasing the vapor pressure. This phenomenon can be quantitatively described by the Clausius-Clapeyron equation:
   

• Intermolecular forces: Substances with stronger intermolecular forces have lower vapor pressures at a given temperature because more energy is required for molecules to escape into the vapor phase.

• Vapor pressure of water: 4.58 mmHg at 273 K versus 17.5 mmHg at 293 K
• Vapor pressure of Et₂O (weak dipole-dipole interactions and dispersion forces): 440 mmHg at 273K
  Vapor pressure of water (strong hydrogen bonds): 4.58 mmHg at 273K

Boiling point and vapor pressure:

• The boiling point of a liquid is the temperature at which its vapor pressure equals the external atmospheric pressure. At this point, bubbles of vapor form within the liquid, leading to boiling.
• Normal boiling point: The temperature at which the vapor pressure of a liquid equals 1 atmosphere (760 mmHg).

Volatility:

A substance with a high vapor pressure at normal temperatures is considered volatile. These substances evaporate quickly and are often associated with strong odors or quick evaporation, like ethanol or acetone.

Types of phase changes:

Endothermic phase transitions (processes that absorb heat from the surroundings):

• Melting (fusion): Transition from solid to liquid (ΔH_fus)
• Vaporisation: Transition from liquid to gas (ΔH_vap)
• Sublimation: Transition from solid to gas without passing through the liquid phase (ΔH_sub = ΔH_fus + ΔH_vap according to Hess’s law).

Exothermic phase transitions (processes that release heat to the surroundings):

• Freezing: Transition from liquid to solid (-ΔH_fus)
• Condensation: Transition from gas to liquid (-ΔH_vap)
• Deposition: Transition from gas to solid without passing through the liquid phase (-ΔH_sub)

Heating curves:

A heating curve illustrates how the temperature of a substance changes as heat is added at a constant rate. Different segments of the curve represent different phases and phase transitions.
 

• Segment A (solid phase): The substance is in the solid phase. As heat is added, the temperature increases according to:
   

• Segment B (melting): The substance undergoes a phase transition from solid to liquid. During this process, the temperature remains constant as the substance absorbs heat:
   

Phase diagrams:

A phase diagram displays the regions of all phases (solid, liquid, gas) of a pure substance across different temperatures and pressures. It maps out the conditions under which each phase is stable and where transitions between phases occur.
 

Components of a phase diagram:

• Axes: pressure (y-axis) and temperature (x-axis).
• Regions: solid region, liquid region, and gas region.
• Phase boundaries: The lines that separate different regions on the diagram, representing the conditions (temperature and pressure) at which two phases coexist in equilibrium (solid-liquid, liquid-gas, and solid-gas).

Key points on a phase diagram:

• Triple point: The point on the phase diagram where all three phases coexist in equilibrium. Each substance has a unique triple point defined by a specific temperature and pressure.
• Critical point: The point on the phase diagram where the liquid and gas phases become indistinguishable from one another. Beyond this point, the substance exists as a supercritical fluid with properties of both liquids and gases.

Crystals:

Crystals are solids with an ordered, repeating arrangement of atoms, ions, or molecules. They can be classified based on the type of forces between their constituent particles, which determine their physical properties. The main types of crystals include ionic, molecular, network, metallic, and amorphous crystals.

Ionic crystals:

Ionic crystals are composed of cations and anions held together by coulombic (charge-charge) attractions, also known as ionic bonds.

Properties:

• Hard and brittle due to the strong electrostatic forces between ions.
• High melting points because a large amount of energy is required to break the ionic bonds.
• Poor electrical conductors in solid form but conduct electricity when molten or dissolved in water because the ions are free to move.
   

Sodium chloride (NaCl), Magnesium oxide (MgO).

Covalent (network) crystals:

Covalent crystals are composed of atoms connected by an extended network of covalent bonds, forming a continuous, three-dimensional structure.

Properties:

• Very hard and have high melting points because covalent bonds are strong and require significant energy to break.
• Typically non-conductive due to the lack of free electrons.
   

Diamond (carbon, C), Silicon dioxide (quartz, SiO₂).

Metallic crystals:

Metallic crystals are composed of positive metal cations arranged in a lattice structure with a "sea" of delocalized electrons that move freely throughout the structure.

Properties:

• Good electrical and thermal conductivity due to the mobility of the free electrons.
• Malleable and ductile because metal ions can slide past each other without breaking the metallic bond.
• Lustrous appearance due to the reflection of light by free electrons.
• Variable melting points, typically high but lower than ionic or covalent crystals.
   

Silver (Ag), Copper (Cu), Iron (Fe), gold (Au).

Molecular crystals:

Molecular crystals are composed of molecules held together by weak van der Waals forces, dipole-dipole interactions, or hydrogen bonds.

Properties:

• Lower melting and boiling points due to the weak intermolecular forces.
• Generally soft and have low hardness.
• Non-conductive because they lack free-moving electrons or ions.
   

Iodine (I₂), Ice (H₂O), Carbon dioxide (dry ice, CO₂).

Principle of X-ray crystallography:

X-ray crystallography relies on the diffraction of X-rays by the orderly arrangement of atoms in a crystal. As X-rays pass through the crystal, they are scattered by the electrons around the atoms. The scattered X-rays interfere with each other, producing a diffraction pattern unique to the structure of the crystal.

Bragg's law:

The relationship between the wavelength of the X-rays, the angle at which they are diffracted, and the spacing between planes in the crystal lattice is described by Bragg's Law:
 

Steps in X-ray crystallography:

1. Crystal preparation: The crystal must be pure, well-ordered, and free of defects, as imperfections can distort the diffraction pattern.
2. X-ray exposure: The crystal is exposed to a beam of X-rays. As the X-rays pass through the crystal, they are diffracted by the atomic planes within the crystal.
3. Data collection: The diffraction pattern is recorded on a detector. The pattern consists of spots that correspond to the angles and intensities of the diffracted X-rays. The position and intensity of these spots provide information about the crystal structure.
4. Data analysis: The diffraction data is analyzed using mathematical techniques, including Fourier transforms, to construct an electron density map. This map shows the distribution of electrons within the crystal, allowing scientists to determine the positions of the atoms.
5. Structure determination: The electron density map is used to build a three-dimensional model of the crystal's atomic structure. This model reveals the arrangement of atoms, the types of chemical bonds, and the geometry of the molecule or solid.

Unit cell:

The unit cell is the smallest repeating unit in a crystal lattice that, when repeated in all directions, creates the entire crystal structure. It defines the crystal’s symmetry and dimensions.

Types of unit cells:

• Primitive (Simple): Atoms are located only at the corners of the unit cell.
• Body-centered: Atoms are located at the corners and one atom at the center of the unit cell.
• Face-centered: Atoms are located at the corners and at the centers of each face of the unit cell.

Crystal lattice:

A crystal lattice is a three-dimensional arrangement of points that represent the positions of atoms, ions, or molecules in a crystal. The unit cells stack together in space to form this lattice.

The lattice parameters include the lengths of the unit cell edges (a, b, c) and the angles between them (α, β, γ), which define the geometry of the lattice.

Packing efficiency:

Packing efficiency refers to the fraction of the volume of a unit cell that is occupied by the constituent particles (atoms, ions, or molecules). It indicates how tightly the particles are packed in the crystal.
 

Cubic packing: Simple cubic (52%), body-centered cubic (68%), face-centered cubic (74%).

Coordination number:

The coordination number of a particle in a crystal structure is the number of nearest neighbors that surround it. This number affects the stability and properties of the crystal.
 

• In a face-centered cubic (FCC) structure, the coordination number is 12.
• In a body-centered cubic (BCC) structure, the coordination number is 8.

Simple cubic (SC):

• Atoms at each corner of the cube
• Coordination Number: 6
• Stacking pattern: a-a-a-a-
• Lattice parameter: length of an edge = 2 x radii of the atom

Body-centered cubic (BCC):

• Atoms at each corner and one atom at the center of the cube
• Coordination Number: 8
• Stacking pattern: a-b-a-b-
• Lattice parameter: length of the main diagonal = 4 x crystallographic radii

Face-centered cubic (FCC):

• Atoms at each corner and at the centers of all the faces of the cube.
• Coordination Number: 12
• Stacking pattern: a-b-c-a-b-c-
• Lattice parameter: length of the diagonal of a face = 4 x crystallographic radii