Lecture 1 & 2

The Nature of Physical Chemistry | Matter | System, State, and Equilibrium | Energy | State of Matter: Gas

Atkins (2006)

Develops the principles of chemistry to explain physical and chemical properties of matter.

Bulthelezi (2007)

• Studies macroscopic, atomic, subatomic, and particulate phenomena using laws and concepts of physics.

• (e.g., motion, energy, force, thermodynamics, quantum chemistry, statistical mechanics, dynamics, and equilibrium).

The Lincoln Library of Essential Information (2011)

• Applies physics methods to chemical problems.

• Studies transformation of substances through qualitative and quantitative approaches (experimental & theoretical).

Microscopic Scale

• Only observable using magnification (e.g., microscope).

• Includes crystal structures, which impact bulk material behavior.

Macroscopic Scale

• Scale describe how relatively large quantities of substances beave.

including:

• Melting & boiling points

• Latent heats of fusion & vaporization

• Thermal conductivity & specific heat capacity

• Coefficient of linear thermal expansion

Atomic Scale

• Deals specifically on atomic properties.

• Example: Boron (B), atomic number 5, atomic mass 10.81 au (for a free neutral atom in ground state).

Subatomic Scale

Scale that generally involve using theories, measurements, and techniques.

Purpose of Physical Chemistry

• Collects data on gases, liquids, solids, solutions, and colloidal dispersions.

• Systemizes this data into scientific laws.

• Develops theoretical foundations for explaining chemical behavior.

differs from Organic & Inorganic Chemistry by focusing less on substance descriptions and more on theoretical principles and quantitative problems.

Systematic Approach

Begins with the very basic constituents of matter—fundamental particle and builds up to larger systems.

Phenomenological Approach

• Investigation of macroscopic material.

• Measures observable properties like Pressure (P) and Volume (V).

Thermodynamics

• deals with heat, work, and temperature.

Atkins (2006): Studies energy interconversion in materials and their molecular properties.

Kinetics

Studying rates of reactions.

Quantum Mechanics

• Study the behavior of particles.

• The phenomena at molecular level.

Statistical Mechanics

Macroscopic behavior of physical systems in terms of dynamic laws.

Spectroscopy

Study of the absorption and emissions of light.

Photochemistry

Interaction of light and matter.

Matter

• Composed of electrons, neutrons, and protons.

• Further divisible into subatomic particles.

• Made of atoms, where each substance has a unique number of protons, neutrons, and electrons.

Phases of Matter

• A property of matter identified by distinct physical characteristics.

• Three common phases: Solid, Liquid, Gas.

• Two less common but important phases: Plasma, Bose-Einstein Condensate (BEC).

Solid

• Molecules are closely bound by strong molecular forces.

• Holds a fixed shape and volume.

Liquid

• Molecular forces weaker than in solids.

• Takes the shape of its container with a free surface in gravitational fields.

• Has a fixed volume, even in microgravity.

Gas

• Molecular forces are very weak.

• Fills both the shape and volume of its container.

Plasma

• Gas that can carry electrical charge.

• Atoms exist in an excited state, giving off light.

• Contains free ions and electrons, allowing it to conduct electricity.

Bose-Einstein Condensate (BEC)

• Forms near absolute zero (0K or -273°C).

• Atoms overlap and behave like a single wave, creating a super atom.

Substance

A pure form of matter.

Amount of Substance

Measured in moles.

Extensive Property

• Depends on the amount of matter

• (e.g., mass, volume).

Intensive Property

• Independent of the amount

• (e.g., density, pressure, temperature).

System

• Defined by physical chemists as the object of study in an investigation.

• Can be solid, liquid, gas, or a combination of these phases.

Macroscopic system

Consists of many components.

Microscopic system

Focuses on individual atoms and molecules.

Surroundings

Everything outside the system.

Universe

System + Surroundings

Intensive Property

• Independent of the amount of matter.

• Examples: Pressure, Temperature, Refractive Index.

Extensive Property

• Depends on the amount of matter.

• Examples: Volume, Mass.

State of a System

• A thermodynamic system is in a certain state when all measured properties are fixed.

• Changes in properties alter the state of the system.

• These properties are called State Variables (or State Functions / Thermodynamic Parameters).

Equation of State

Relates empirical data that are summarized terms of experimentally defines variables.

Independent Variables

Amount of substance, Temperature, Pressure.

Dependent Variable

Volume (determined by the equation of state).

System is in equilibrium

Its state variables remain unchanged over time.

Thermal Equilibrium

No state change occurs upon contact.

Zeroth Law of Thermodynamics

A (Iron Block) = B (Copper Block) = C (Flask of Water)

• If A is in thermal equilibrium with B, and B is in thermal equilibrium with C, then A is in thermal equilibrium with C

Thermodynamics

• Studies energy flow (heat, work) into or out of a system.

• Deals with the relationships between heat, work, temperature, and energy.

• Describe energy changes and whether a system can perform work.

Thermal Equilibrium

Uniform temperature throughout the system, equal to surroundings.

Mechanical Equilibrium

No mechanical work is done by one part of the system on another.

Chemical Equilibrium

No net chemical change occurs within the system.

Heterogeneous System

Each phase’s state variables remain constant.

Important system properties in thermodynamics

Pressure, Volume, Concentration, and Temperature.

Energy

Capacity to do work and to heat.

Work

• Defined as force (F) causing mechanical displacement (dx) on a body.

Calorie (cal)

Unit of heat

Gas

• The simplest state of matter that fills any container it occupies.

• Molecules are widely spaced and move freely.

• The most studied and best understood phase of matter.

Expansibility

Gases expand limitless to fill any container.

Compressibility

Gases can be compressed by applying pressure.

Diffusibility

Gases diffuse rapidly to form homogeneous mixtures.

Pressure

Gases exert pressure in all directions on container walls.

Effect of Heat

Heating increases gas pressure (in a fixed volume) or increases volume (if the container allows expansion).

Parameters of Gases

• Volume (V) – Space occupied by gas (SI unit: m³, common: L, mL, cm³).

• Pressure (P) – Force per unit area due to molecular collisions (SI unit: Pa, common: atm, mmHg, torr).

• Temperature (T) – Measured in Kelvin (K)

• Number of moles (n) – Amount of gas

1 bar

Standard (steady) pressure.

Manometer

A U-shaped tube with liquid, where pressure difference is measured by liquid height.

Barometer

Measures atmospheric pressure using mercury (Hg).

Thermal Equilibrium

Energy flows between objects until temperatures equalize.

Temperature

Describes flow of energy.

Diathermic

Allows heat exchange (e.g., metal).

Adiabatic

No heat transfer (e.g., Styrofoam).

Perfect Gas

Robert Boyle and Jacques Charles contributed to understanding gas behavior.

Ideal (Perfect) Gas

Hypothetical gas that follows gas laws exactly and completely fills any container.

Non-Ideal (Real) Gas

Deviates from ideal behavior due to intermolecular forces and molecular volume.

Gas Laws

• Derived from experimental studies (17th-19th century).

• Establish relationships between Pressure (P), Volume (V), and Temperature (T).

Boyle’s Law (1660) – Pressure-Volume Relationship

• Robert Boyle formulated this law through experiments at room temperature.

• "At constant temperature, the volume of a fixed mass of gas is inversely proportional to its pressure."

Isotherm (P vs. V curve)

A downward curve showing inverse proportionality.

Charles’ Law (1787) – Temperature-Volume Relationship

• Jacques Charles investigated the effect of temperature on volume at constant pressure.

• "At constant pressure, the volume of a fixed mass of gas is directly proportional to its absolute temperature (Kelvin scale)."

Kinetic Theory of Gases (Brownian Motion)

Describes the motion of gas molecules.

• Gas molecules undergo random motion (Brownian Motion).

• Speed of molecules increases as temperature increases.

• Gas molecules are widely separated, with interactions only occurring during collisions.

• No intermolecular forces (assumption for an ideal gas).

Combined Gas Law

• Boyle’s Law and Charles’ Law can be combined into a single relationship known as the Combined Gas Law.

• "For a fixed mass of gas, the volume is directly proportional to Kelvin temperature and inversely proportional to pressure."

Gay-Lussac’s Law (1802) – Pressure-Temperature Relationship

Joseph Gay-Lussac established a relationship between pressure and temperature.

"At constant volume, the pressure of a fixed mass of gas is directly proportional to its absolute temperature (Kelvin scale)."

Isochor (P vs. T curve)

A straight line showing direct proportionality.

Avogadro’s Law – Volume-Mole Relationship

Amedeo Avogadro developed this law based on the relationship between gas volume and the number of moles.

"Equal volumes of gases at the same temperature and pressure contain equal numbers of molecules or moles."

Molar Volume of a Gas at STP

• 0°C (273.15 K) and 1 atm.

• 1 mole of any gas at STP = 22.4 L

• pure liquid and gas at 1M concentration.

Standard Ambient Temperature and Pressure (SATP)

• Temperature = 25°C (298.15 K).

• Pressure = 1 atm.

• Used for thermodynamic calculations.

Normal Temperature and Pressure (NTP)

• Temperature = 20°C (293.15 K).

• Pressure = 1 atm.

• Used for air and industrial calculations.

Universal Gas Law (Ideal Gas Law)

• Applies to all gases exhibiting ideal behavior.

• "The volume of a given amount of gas is directly proportional to the number of moles and temperature, and inversely proportional to the pressure."

Surface of States

• The Ideal Gas Equation represents a three-dimensional surface of possible states.

• The gas cannot exist in states outside this surface.

Dalton’s Law of Partial Pressures (1807)

• John Dalton proposed that in a mixture of gases, each gas exerts a pressure independently, as if it were alone in the container.

• Gases do not interact and contribute independent pressures.

• "The total pressure of a mixture of gases is equal to the sum of the partial pressures of all the gases present."

Amagat’s Law of Partial Volumes (1880)

• Émile Amagat formulated this law to describe gas volume behavior.

• Gases have additive volumes, with average molecular interactions.

• "The volume of a gas mixture is equal to the sum of the volumes of its individual components at the same temperature and pressure."

Diffusion

The spontaneous mixing of two gases due to random molecular motion.

Effusion

The escape of gas molecules through a small hole (without collisions) into a vacuum or low-pressure region.

Graham’s Law of Diffusion (1829)

• Thomas Graham observed that lighter gas molecules diffuse faster than heavier ones.

• Lighter gases diffuse faster than heavier gases.

• "Under the same conditions of temperature and pressure, the rate of diffusion of a gas is inversely proportional to the square root of its molecular mass."

Graham’s Law of Effusion

• Lighter gases effuse faster than heavier gases.

• Used to determine the molecular mass of an unknown gas.