Chapter 14 States of Matter

Molecules in Gas, Liquid, & Solid States Gas particles are further apart than liquid particles which are further apart than solid particles Attractive forces hold particles together in solid and liquid state video

Molecules in Gas, Liquid, & Solid States

Kinetic-Molecular Theory The properties of solids, liquids, and gases can be explained based on: –the speed of the molecules –the attractive forces between molecules.

Kinetic- Molecular Theory Model to help explain the behavior of substances “kinetic” refers to motion The energy that an object has because of motion is called kinetic energy Kinetic theory states that the tiny particles in all forms of matter are in constant motion

Kinetic-Molecular Theory In solids, the molecules have little freedom of motion. They are held in place by strong attractive forces. –May only vibrate

Kinetic-Molecular Theory (cont.) In liquids, the molecules more freedom of motion than solids but not enough to escape their attraction to neighboring molecules. –They can slide past one another and rotate, as well as vibrate.

Kinetic-Molecular Theory (cont.) In gases, the molecules have “complete” freedom from each other. They have enough energy to overcome “all” attractive forces.

Kinetic–Molecular Theory of Gases 1. A gas consists of small particles that: Are very far apart from one another. Have essentially no attractive (or repulsive) forces. Have very small volumes compared to the volume of the container they occupy.

Kinetic–Molecular Theory of Gases 2. Gas particles are in rapid random motion: Travel in straight paths Change direction only when collide with another object or one another –Explains why gases fill container

3. All collisions are perfectly elastic Gas particles are constantly colliding with one another and with the walls of the container –Total kinetic energy remains the same –No loss of kinetic energy Kinetic–Molecular Theory of Gases

Gas Pressure Moving particles exert a force when they collide with other particles or objects Pressure = force area Gas pressure is due to –Force of collisions –Number of collisions Vacuum –No particles are present, therefore no collisions, no pressure

Atmospheric Pressure Is the pressure exerted by a column of air from the top of the atmosphere to the surface of the Earth.

Units for Measuring Gas Pressure SI Unit is the pascal (Pa) Millimeters mercury (mmHg) Atmosphere (atm) lbs/in 2 1atm= 760mmHg=101.3kPa

Barometers A glass tube filled with Hg placed in a dish of Hg At sea level, Hg rises to height of 760mm 1atm= 760mm Hg

The Meaning of Temperature Temperature is a measure of the average kinetic energy of the molecules in a sample. –Not all molecules have same kinetic energy Kinetic energy is directly proportional to the Kelvin temperature. –Average speed of molecules increases as the temperature increases

Kinetic Energy and Temperature What happens when a substance is heated? –The particles of the substance absorb energy –Some energy is stored (potential energy) –Some energy speeds up the particles (average kinetic energy increases) –Increasing average kinetic energy results in an increase in temperature –Kelvin temperature scale reflects relationship between temperature and KE Absolute 0 is when motion of all particles cease (-273ºC)

Temperature and Kinetic Energy The temperature of a gas is a measure of the KE of the gas particles For a given gas, the average KE is dependent on Temp.

Temperature and Kinetic Energy

Kinetic Energy and Mass of Gas Particles KE = 1/2mv 2 Lighter particles move faster Heavier particles move slower

Diffusion of Gases

Liquids Objectives –Describe nature of liquids in terms of attractive forces between particles –Differentiate between evaporation and boiling of a liquid using kinetic theory (liquid  gas)

Nature of Liquids Liquid particles have a greater range of motion than solids but less than gases. Particles spin, vibrate and can slide past one another Particles are attracted to one another, unlike gas particles These attractive forces are called “intermolecular attractive forces”

Attractive forces holding particles together Intramolecular forces –Bonds between atoms within compound Ionic bond Covalent bond Metallic bond –These types of bonds are strong. Intermolecular forces –Attractive forces between neighboring molecules –Have a wide range of strengths, but generally weaker than intramolecular forces (bonds)

Intermolecular Forces Reduce space between particles in a liquid –Liquids are more dense than gases –Increasing pressure on liquids has no effect on volume Affect ability of liquid particles to “escape” the liquid state

Intermolecular Attractive Forces Relate to properties which involve breaking the forces of attraction between particles m.p., b.p. Solubility Relate to 3 dimensional structure of molecule macromolecules such as DNA, proteins Intermolecular forces are electrostatic – Attractions between polar molecules Types of Intermolecular Attractive Force –Dipole-dipole –Dipole-induced dipole –London Dispersion Forces

Hydrogen bonds between water molecules

Liquid  Gas Conversion of a liquid to a vapor or gas is called vaporization E vaporation is when this conversion occurs at the surface of a liquid that is not boiling

Evaporation (open container) Only the particles on surface of a liquid with a certain minimum amount of KE (kinetic energy) can “escape” or break away from the surface.

Evaporation

If heat is added, KE of the particles increases –More particles have enough KE to overcome attractive forces and “escape” into the vapor state Evaporation is a cooling process –Particles with highest KE escape first –Particles left will have lower KE, thus a lower temperature

Evaporation (closed container) Some particles have enough KE to “escape” the liquid state When the space above liquid becomes saturated, some of the vapor particles are “captured” and return to the liquid state

Dynamic Equilibrium rate of evaporation of liquid = the rate of condensation of vapor H 2 O (l) H 2 O (g) Particles continue to evaporate and condense No net change in the number of vapor or liquid particles

Vapor Pressure The particles that have entered the gaseous (vapor) state collide with walls of container and exert a pressure (vapor pressure)

Vapor Pressure is a function of temperature As temperature increases, the number of molecules in vapor state increases, therefore the vapor pressure increases

Vapor Pressure (cont.) Liquid boils when its vapor pressure = atmospheric pressure –Normal boiling point –Raising external pressure raises boiling point, & vice-versa

Boiling Point

Normal Freezing and Boiling Points for water Freezes at 0°C (normal freezing point) –Pure water, 1.00 atm pressure Boils at 100°C (normal boiling point) –Pure water, 1.00 atm pressure –Boiling point increases as atmospheric pressure increases

Vapor Pressure Curve 1.What is the normal boiling point of pentane? What is the normal boiling point of water? Why is there a difference? 2. Which of the liquids has the greatest amount of attraction between the molecules? 3. If the atmospheric pressure in the mountains is 600 mm Hg, at what temperature will water boil?

Solids Objectives –Describe how organization of particles distinguishes solids form liquids and gases –Explain how allotropes of an element are different

Characteristics of solids Solid particles are not free to move Vibrate around a fixed point Particles are arranged in a highly organized pattern (crystalline shape) Dense and not compressible Do not flow or take shape of container

Melting Point Heat a solid, particles have more KE Organization of particles breaks down –Disruptive vibrations strong enough to overcome forces holding particles in fixed positions mp = When an solid becomes a liquid Most ionic compounds have a high mp NaCl has mp = 801ºC Most covalent compounds have a relatively low mp HCl = -112ºC

Crystal Structure Crystal lattice- orderly, repeating 3D pattern Shape of crystal is related to arrangement of particles A crystal has sides or faces –Angles of faces are characteristic of that substance Smallest group of particles within crystal that retains geometric shape is called a unit cell

Unit Cells Simple cubic Body centered cubic Face centered cubic

Allotropes solid substances that can exist in more than one form –Carbon Grahite Diamond buckminsterfullerene

Amorphous solids Solids which lack an ordered internal structure –Rubber, plastic, asphalt –Glasses Supercooled liquids Cooled to a rigid state without crystallizing Irregular internal structure between crystalline solid and free flowing liquid Gradually soften when heated

Changes of State

Heating Curve As the substance warms, what is happening to: –Energy of the particles? –Motion of the particles? –Arrangement of the particles? –State of matter?

Cooling Curve As the substance cools, what is happening to: –Energy of the particles? –Motion of the particles? –Arrangement of the particles? –State of matter?

Heating/Cooling Curve for Water

Phase Diagram Graph which shows relationships of solid, liquid, gas in a sealed container at varying temperatures and pressures

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Phase Diagrams 3 regions indicate the 3 states: solid, liquid, gas Points: –Triple point- point at which all 3 states exist in equilibrium –Critical point- Only the gaseous state exists at temperatures and pressures above this point –Normal bp- temp at which vapor pressure equals atmospheric pressure –Normal mp- temp at which atmospheric pressure cuts through solid-liquid equilibrium line Lines –Solid vapor equilibrium line –Solid liquid equilibrium line –Vapor liquid equilibrium line

Why do Molecules Attract Each Other? Intermolecular attractions are due to attractive forces between opposite charges. –+ ion to - ion –+ end of polar molecule to - end of polar molecule Larger the charge, stronger the attraction Even non-polar molecules have attractions due to opposite charges

Intermolecular Attraction Much weaker than ionic or covalent bonds There are several types: London dispersion forces, dipole-dipole attraction, and hydrogen bonding

Review Vocabulary allotrope amorphous solid atmospheric pressure barometer boiling point critical point crystal equilibrium evaporation gas pressure heating curve glasses kinetic energy Kinetic Theory melting, boiling point normal m.p., b.p. pascal phase diagram sublimation triple point unit cell vacuum vaporization vapor pressure