States of Matter: Solid and Liquid

• All objects are called matter which occupies space and has mass.
• Matter exists in three common states: gas, liquid, and solid.
• Water exists in three physical states: steam (gas), water (liquid), and ice (solid).
• Heating transforms solids to liquids and liquids to gases.
• Liquids and gases are fluids because they flow freely.
• Solids and liquids are condensed states due to their higher densities than gases.
• Changes in the state of matter occur by varying temperature and/or pressure.

Solid State

• Solids are rigid with a fixed volume and shape, and they cannot be compressed.
• Solids have higher densities than liquids because their particles are closely packed and highly organized.
• Solid particles vibrate at a fixed point and stay in a fixed position.
• Metals, chalk, sand, soil, and stone are examples of solids.

Liquid State

• Liquids have a definite volume but no fixed shape, assuming the shape of their container.
• They are slightly compressible and can flow.
• Liquid particles vibrate and constantly shift positions.
• Water, ethanol, mercury, bromine, and oil are examples of liquids.

Gas State

• Gases have neither definite volume nor shape. They conform to their container's shape and volume.
• Gases are highly compressible due to large empty spaces between particles.
• Air, hydrogen, oxygen, carbon dioxide, and nitrogen are gases.

Plasma State

• Plasma is the fourth state of matter, existing at extremely high temperatures (millions of degrees Celsius).
• At these temperatures, atoms are stripped of their electrons, forming an ionized gas.
• Plasmas have no fixed shape or volume and are less dense than solids or liquids.

Kinetic Theory of Matter

• The kinetic theory explains the three states of matter and their physical differences.
• It posits that every substance consists of numerous small particles (ions, atoms, and molecules) in continuous, random motion.
• This motion increases with temperature.

Assumptions of Kinetic Theory

• All matter is composed of constantly moving particles.
• All particles possess kinetic and potential energy, with solids having the least and gases the most energy.
• Differences in energy content and particle motion differentiate the three states of matter. Phase changes occur with energy changes.
• Temperature measures the average kinetic energy of particles.
• Spaces exist between particles, increasing from solid to liquid to gas phases.
• Attractive forces (intermolecular forces) between particles strengthen as they move closer.

Properties of Gases

• Gases have no fixed shape or volume and assume the shape and volume of their containers.
• Gases are easily compressed due to large interparticle spaces.
• Gases have low densities because their particles are far apart.
• Gases exert pressure in all directions due to collisions between molecules and container walls.
• Gases easily flow and diffuse.

Properties of Liquids

• Liquids have a definite volume but no definite shape, taking the shape of their container.
• Liquids have higher densities than gases due to closer particle arrangement.
• Liquids are slightly compressible due to minimal free space between particles.
• Liquids are fluids that flow, usually downwards due to gravity, but slower than gases.

Properties of Solids

• Solids have a definite shape and volume due to strong interparticle attraction.
• Solids generally have higher densities than gases and liquids because their particles are very close together.
• Solids are extremely difficult to compress due to strong interparticle forces.
• Solids are not fluids. Solids do not flow because their particles are rigidly held in position.

Kinetic Molecular Theory of Gases

• Gases consist of particles (molecules) in constant, straight-line motion.
• Gas particles do not attract or repel each other. They collide and bounce back.
• Gas particles are small compared to the space between them, making their volume negligible.
• The average kinetic energy of gas particles is proportional to the gas temperature in kelvins, expressed as KE = 1/2 mv².

General Properties of Gases

• Gases are highly compressible.
• Gases exert pressure equally in all directions.
• Gases have lower densities than solids and liquids.
• The volume and shape of gases are not fixed.
• Gases mix evenly and completely.

Gas Laws

• Gas laws relate measurable properties of gases, including pressure, volume, temperature, and mass.
• Pressure is the force per unit area from gas particle collisions (Pressure = Force/Area).
• Volume is the space occupied by a substance, measured in cubic meters (m3) or liters (L).
• Temperature indicates hotness or coldness, measured in Celsius (°C), Fahrenheit (°F), or Kelvin (K).
• Conversions: K = °C + 273 and °C = (°F – 32) 5/9.
• Standard Temperature and Pressure (STP) conditions are 1 atmosphere (atm) and 0 °C (273 K).
• The molar volume of any gas at STP is 22.4 liters.

Boyle's Law

• Boyle's Law (1662): At constant temperature, the pressure of a fixed amount of gas varies inversely with its volume.
• PV = k, where k is a constant at a specific temperature.
• This is written as: P1V1 = P2V2, where P1 and V1 are the initial pressure and volume, and P2 and V2 are the final pressure and volume.

Charles' Law

• Charles' Law (1787): At constant pressure, the volume of a gas increases linearly with its temperature.
• Absolute zero is 0 K (-273 °C), the theoretically lowest attainable temperature which can't be attained.
• Charles’ law states that the volume of a fixed mass of gas at constant pressure varies directly with the Kelvin temperature.
• Expressed as: V/T = k
• Applied as: V1/T1 = V2/T2

Gay-Lussac's Law

• Gay-Lussac's Law: At constant volume, the pressure of a fixed amount of gas varies directly with the temperature.
• Expressed mathematically as: P/T = k.
• For two conditions: P1/T1 = P2/T2.

Combined Gas Law

• Combines Boyle’s, Charles’, and Gay-Lussac’s laws into a single equation.
• Expressed as: (P1V1)/T1 = (P2V2)/T2.

Avogadro's Law

• Avogadro's Law (1811): At the same temperature and pressure, equal volumes of gases contain equal numbers of moles.
• Volume is directly proportional to the number of moles of gas: V = kn
• Expressed as: V1/n1 = V2/n2

Ideal Gas Equation

• Ideal gases obey gas laws closely at high temperatures and low pressures.
• The Ideal Gas Law combines Boyle’s, Charles’ and Avogadro’s laws: PV = nRT
• R is the ideal gas constant, with a value of 0.082 L atm/(mol K).

Graham's Law of Diffusion

• Diffusion is the spreading of gas molecules throughout a container.
• Graham’s Law: At constant temperature and pressure, the rate of diffusion of a gas is inversely proportional to the square root of its density or molar mass.
• r ∝ 1/√d or r ∝ 1/√M
• Expressed as: r1/r2 = √(d2/d1) = √(M2/M1)
• The rate at which a gas diffuses is also inversely proportional to the time it takes to move a certain distance.
• Expressed as: r ∝ 1/t.
• For two gases: t1/t2 = √(M1/M2)

Explore the states of matter, focusing on solids and liquids. Learn how heating transforms solids to liquids, and understand the unique properties of each state. Discover examples of each state and their characteristics.