Kinetic Molecular Theory (KMT)

• Kinetic Molecular Theory (KMT) explains the behavior of atoms in matter by stating particles are in constant motion, which correlates to energy.
• KMT primarily focuses on the properties of gases but can be used to describe all states of matter.

Key Assumptions of KMT

• Gases are composed of many tiny particles spaced far apart, making most of the volume empty space.
• Gas particles are in constant, random motion at high speeds, which contributes to their kinetic energy.
• Collisions between gas particles and container walls are elastic, where there is no loss of kinetic energy during collisions.

Ideal vs. Real Gases

• An ideal gas conforms to all the assumptions of KMT, and serves as a comparison standard.
• Real gases deviate from ideal behavior under high pressure and low temperature, because intermolecular forces become significant.

Characteristics of Gases

• Gases lack a definite shape or volume, expanding to fill any container because of the lack of attractive forces between particles.
• The low density of gases is a result of particles being far apart.

Gas Pressure

• Gas pressure is the force exerted by gas particles per unit area, influenced by collisions with container walls.
• A vacuum has no particles, resulting in no pressure due to the absence of collisions.
• Gases are compressible due to the large spaces between particles.
• Diffusion is the spontaneous mixing of gas particles, and Effusion refers to the passage of gas through a small opening.

Properties of Liquids

• Liquids have particles closer together than gases, resulting in lower kinetic energy and more order, but less than solids.
• Liquids are relatively incompressible because of the proximity of particles.

Intermolecular Forces in Liquids

• The behavior of liquids is influenced by intermolecular forces (IMFs), like London dispersion forces, dipole-dipole interactions, and hydrogen bonding.
• The strength of IMFs affects properties such as boiling point, viscosity, and surface tension.
• Liquids can flow and mold to containers, demonstrating fluidity, which is shared with gases.
• Diffusion in liquids occurs more slowly than in gases but increases with temperature as kinetic energy rises.

Freezing and Vaporization

• Freezing is the process where a liquid turns into a solid as kinetic energy decreases, allowing attractive forces to hold particles in a fixed arrangement.
• Vaporization includes both evaporation (surface phenomenon) and boiling (throughout the liquid), where vapor pressure equals external pressure at the boiling point.

Evaporation vs. Boiling

• Evaporation is a cooling process as higher energy particles escape first, lowering the average kinetic energy of the remaining liquid.
• The boiling point of a liquid varies with external pressure; for example, water boils at 100°C at 1 atm but at lower temperatures at higher altitudes.

Capillary Action and Surface Tension

• Capillary action results from adhesive forces between liquid and solid surfaces, allowing liquids to rise in narrow spaces.
• Surface tension is the cohesive force that pulls adjacent parts of a liquid's surface together.

Boiling Point and External Pressure

• The boiling point of a liquid is defined as the temperature at which its vapor pressure equals the external pressure.
• Normal boiling point for water is 100 °C at 1.00 atm pressure.
• In high-altitude locations like Denver (1600 m above sea level), the atmospheric pressure is lower (0.84 atm), resulting in a lower boiling point of water (95 °C).
• Pressure cookers increase the pressure above the liquid, allowing water to boil at temperatures above 100 °C.

Evaporation vs. Boiling

• Evaporation occurs at any temperature and involves molecules gaining kinetic energy to escape into the vapor phase, while boiling occurs throughout the liquid at a specific temperature.
• In a closed container, vapor pressure is created as some liquid particles evaporate and collide with the container walls, leading to a dynamic equilibrium where the rate of evaporation equals the rate of condensation.
• The equilibrium vapor pressure increases with temperature because the higher the kinetic energy of the liquid particles, the easier it is to escape into the vapor phase.

LeChatelier's Principle and Vapor Pressure

• LeChatelier's Principle states that if a change is imposed on a system at equilibrium, the system will adjust to counteract that change.
• An increase in temperature in a closed system results in an increase in vapor pressure as more particles gain sufficient energy to escape the liquid phase.

Characteristics of Solids

• Solids have a definite shape and volume due to strong intermolecular forces, resulting in low kinetic energy compared to liquids and gases.
• Crystalline solids have a well-ordered structure (crystal lattice), while amorphous solids have a random arrangement of particles.
• Solids exhibit a definite melting point, high density, and incompressibility.

Allotropes of Elements

• Allotropes are different structural forms of the same element in the same physical state, differing in the arrangement of atoms.
• Carbon allotropes: diamond, graphite, and buckminsterfullerene.
• The properties of allotropes can vary significantly due to differences in atomic arrangement, affecting their physical and chemical behavior.

Types of Phase Changes

• Phase changes include melting (solid to liquid), freezing (liquid to solid), vaporization (liquid to gas), condensation (gas to liquid), sublimation (solid to gas), and deposition (gas to solid).
• Each phase change involves energy transfer and is influenced by temperature and pressure conditions.
• The phase diagram illustrates the states of a substance at various temperatures and pressures, indicating regions of solid, liquid, and gas.

Heating Curves and Temperature Effects

• A heating curve plots temperature against time, showing how a substance changes state as it absorbs heat.
• During phase changes, temperature remains constant while the substance absorbs or releases energy, indicating a change in state rather than temperature increase.
• The rate of evaporation and vapor pressure equilibrium are influenced by surface area, temperature, and intermolecular forces (IMFs).

Intermolecular Forces (IMFs)

• IMFs are the forces that hold molecules together; they include dipole-dipole interactions, hydrogen bonding, and London dispersion forces (LDFs).
• Stronger IMFs result in lower vapor pressure and higher boiling points, as more energy is required to overcome these forces during phase changes.
• The relationship between temperature and vapor pressure is direct. As temperature increases, vapor pressure increases due to higher kinetic energy of particles.

Impact of IMFs on Phase Changes

• Weaker IMFs lead to higher rates of evaporation and higher equilibrium vapor pressure, while stronger IMFs result in lower rates of evaporation and lower vapor pressure.
• The boiling point is affected by external pressure; lower external pressure results in a lower boiling point because the vapor pressure required for boiling is reduced.
• Condensation is influenced by temperature and IMFs; lower temperatures increase the likelihood of gas particles condensing into a liquid.

This lesson explains the Kinetic Molecular Theory (KMT), which describes the behavior of atoms in matter. It focuses on the assumptions of KMT, including particle motion and elastic collisions. The lesson also discusses the differences between ideal and real gases and their characteristics.