States of Matter and Kinetic Molecular Theory

• Matter exists in three phases: gas, liquid, and solid, each with distinct properties.

Gas State

• Gases have indefinite shape and volume, expanding to fill their container.
• Gas particles are widely spaced, move freely, and have weak intermolecular forces.

Liquid State

• Liquids have indefinite shape but definite volume, taking the shape of their container.
• Liquid particles are loosely packed and can move past each other.
• Liquid intermolecular forces maintain a fixed volume.

Solid state

• Solids have definite shape and volume.
• Solid particles are tightly packed in a fixed arrangement and can only vibrate.

Kinetic Molecular Theory (KMT)

• Explains gas behavior based on the motion of individual molecules or atoms.
• It provides a framework for understanding gas properties.

Key Assumptions of KMT:

• All matter consists of tiny particles in constant motion.
• Gas particles move in straight lines until they collide.
• There are no attractive or repulsive forces between gas particles.
• Collisions between gas particles are elastic, conserving total kinetic energy.
• The average kinetic energy of gas particles is proportional to temperature.

Types of Forces in Matter:

Intramolecular Forces

• These forces hold atoms together within a molecule.
• Are stronger than intermolecular forces, determining chemical properties and stability.

Covalent Bonds

• Atoms share electrons to achieve a stable electron configuration.

Ionic Bonds

• Electrons are transferred from one atom to another, creating ions that attract each other.
• Example: Sodium chloride (NaCl).

Metallic Bonds

• Found in metals, where electrons are delocalized and move freely.
• Example: Copper (Cu) and iron (Fe).

Intermolecular Forces

• Weaker than intramolecular forces, but determine physical properties.
• Boiling point, melting point, viscosity, surface tension, and solubility.
• Three types: London dispersion, dipole-dipole, and hydrogen bonds.

London Dispersion Forces

• Known as the weakest type of intermolecular forces.
• Present in all molecules and arise from temporary fluctuations in electron distribution.
• Temporary dipoles induce dipoles in neighboring atoms or molecules.

Dipole-Dipole Forces

• Attractive forces between the positive end of one polar molecule and the negative end of another.
• Create higher boiling and melting points compared to nonpolar substances.
• Repulsive forces occur when like charges align, but attractions are maximized overall.

Hydrogen Bonds

• Special case of dipole-dipole interactions.
• Occur when hydrogen is covalently bonded to highly electronegative atoms (F, O, N).
• These bonds are stronger due to hydrogen's small size and high polarizability.

Molecular Structure of Water

• Water is the universal solvent.
• Colorless in small quantities, but may appear blue in large volumes due to red light absorption.
• Pure water is odorless and tasteless due to the absence of dissolved substances.

Physical Properties of Water

• Freezing Point: The temp at which water changes from liquid to solid (ice)
• Boiling Point: The temp at which water changes from liquid to gas (steam)
• Both freezing and boiling points serve as temperature scale references.
• Ice forms a hexagonal crystalline structure due to hydrogen bonding.

Properties of Solids

• Arrangement of Electrons: Distribution affects solid matter behavior.
• Strong bonds create hardness or conductivity.
• Electrons transfer in ionic solids to affect strength and solubility.

Covalent Bonds

• Atoms share electrons in solids such as diamond.

Ionic Bonds

• Positively charged ions are attracted to negative ions in salt.

Metallic Bonds

• Atoms release electrons into a "sea" to allow electrical conductivity and malleability in metals.

Intermolecular Forces

• Affect a solid's melting point, boiling point, and hardness.
• Example: Hydrogen bonds in ice make it less dense than liquid water.

Van der Waals Forces

• Allow layers of carbon atoms to slide in graphite.

Types of Solid

• Amorphous Solids: Lack a regular, repeating pattern of atoms, molecules, or ions.
• Random arrangement makes their properties isotropic (uniform in all directions).
• Unlike crystalline solids, amorphous solids soften over a range of temperatures.
• Under specific conditions such as heating, pressure, or chemical treatments, amorphous solids can be converted into crystalline.

Crystalline Solids

• Atoms, ions, or molecules are arranged in a highly ordered, repeating pattern.
• Long-range order gives distinct properties such as geometric shapes, sharp melting points, and anisotropy.
• Anisotropy: exhibit different physical properties in different directions

Molecular Solids

• Since held by relatively weak forces, easily separated and brittle.

Ionic Solids

• Hard and Brittle due to strong electrostatic forces.

Types of Atomic Solids:

Covalent Network Solids

• Atoms are bonded together by strong covalent bonds in a continuous, extended network, or as one giant molecule.

Group 18 solids

• Formed by the noble gasses

Metallic Solids

• Composed of metal atoms held together by metallic bonds.

Phase Changes

• Transition between solid, liquid, or gas states related to molecular kinetic energy.

Solid to Liquid (Melting)

• Molecules gain kinetic energy and weaken intermolecular forces.

Liquid to Gas (Evaporation/Boiling)

• Increased heat energy increases kinetic energy.
• Molecules overcome intermolecular forces and enter the gas phase.

Gas to Liquid (Condensation)

• Gas loses kinetic energy, molecules slow down and form intermolecular bonds.

Liquid to Solid (Freezing)

• Liquid loses heat, molecules lose kinetic energy.
• Movement slows, allowing intermolecular forces to form a solid structure.

Solid to Gas (Sublimation) & Gas to Solid (Deposition)

• Sublimation: solid to gas transition (ex dry ice)
• Deposition: reverse of sublimation (ex frost)

Phase Diagram

• Graphical representation of substance states relative to temperature and pressure.
• Fusion Curve: Boundary between solid and liquid phases.
• Vaporization Curve: Boundary between liquid and gas phases.
• Defines the boiling point for a given pressure, transitions to gas phase if temperature increases, or condenses into a liquid if pressure increases.
• Critical Point: Point where liquid and gas phases merge into a supercritical fluid.
• Sublimation Curve: Boundary between solid and gas phases.
• Used to define conditions for sublimation or deposition.

Triple Point

• Where solid, liquid, and gas phases coexist in equilibrium.
• An exact, specific temperature and pressure
• The lowest point at which a liquid can exist
• Example: Water exists at 0.01°C and 0.006 atm.
• Used to define the Kelvin Temperature Scale
• Used for understanding phase behavior, calibrating thermometers, and lyophilization.
• Supercritical Fluid: A state of matter above critical temperature and pressure.
• Exhibits properties of both gas and liquid.
• No distinct phase boundary, smooth transition