Chapter 10 States of Matter 1
Section 10.1 The Nature of Gases Objectives: Describe the assumptions of the kinetic theory as it applies to gases. Interpret gas pressure in terms of kinetic theory. Define the relationship between Kelvin temperature and average kinetic energy. 2
Kinetic Theory The word kinetic refers to motion The energy an object has because of it s motion is called kinetic energy The kinetic theory states that the tiny particles in all forms of matter are in constant motion! 3
Kinetic Theory & Gases Three basic assumptions of the kinetic theory as it applies to gases: Assumption #1: Gas is composed of particlesusually molecules or atoms Behave as small, hard spheres Each particle has insignificant volume; Relatively far apart from each other (when compared to solids or liquids) No attraction or repulsion between particles (particles are more free to move around than in a liquid) *Remember our discussion of intermolecular forces! 4
Kinetic Theory & Gases Assumption #2: Particles in a gas move rapidly in constant random motion Move in straight paths, changing direction only when colliding with one another or other objects Average speed of O 2 in air at 20 o C is an amazing 1700 km/h! 5
Kinetic Theory & Gases Assumption #3: All collisions between particles in a gas are perfectly elastic When large objects (like cars) collide some of the kinetic energy is changed into other forms of energy (such as heat from friction) Before After E Car1 = 10 E Car2 = 10 E car1 + E Car2 = 20 E Car1 + E Car2 = 15 + thermal energy (heat) = 5 6
Kinetic Theory & Gases When microscopic particles (like gas particles) collide kinetic energy is transferred from one particle to another- the total kinetic energy remains constant Before After E atom = 10 E atom = 10 E atom + E atom = 20 E atom + E atom = 20 When gas particles collide, no energy is lost we call this a perfectly elastic collision 7
Gas Pressure Gas Pressure is defined as the force exerted by a gas per unit surface area of an object 8
Gas Pressure Gas exerts pressure because moving particles exert a force when they collide The greater that force, the greater the pressure Less collisions More collisions The greater the number of collisions, the greater the pressure No particles present? Then there cannot be any collisions, and thus no pressure called a vacuum Vacuum 9
Measuring Gas Pressure The SI unit of pressure is the pascal (Pa) Older units of pressure: millimeters of mercury (mm Hg), and atmospheres (atm) Atmospheric pressure: pressure exerted on objects due to the gas particles in the atmosphere. Results from the collisions of air molecules with objects Decreases as you climb a mountain because the air thins as elevation increases air density decreases 10
Standard Temperature and Pressure For gases, it is important to relate measured values to a standard. This gives us a frame of reference. Standard conditions are defined as a temperature of 0 o C and a pressure of 101.3 kpa, or 1 atm TABLE A This is called Standard Temperature and Pressure, or STP 11
Section 10.2 The Nature of Liquids Objectives: Identify factors that determine physical properties of a liquid. Define evaporation in terms of kinetic energy. Describe the equilibrium between a liquid and its vapor. Identify the conditions at which boiling occurs. 12
The Nature of Liquids Liquid particles are also in motion. Liquid particles are free to slide past one another Gases and liquids can both FLOW However, liquid particles are attracted to each other, whereas gases are not Intermolecular attractions reduce the amount of space between particles of a liquid Thus, liquids are more dense than gases Increasing pressure on liquids has hardly any effect on it s volume 13
KE of Liquid Particles Particles of a liquid spin and vibrate while they move, thus contributing to their average kinetic energy But, most of the particles do not have enough energy to escape into the gaseous state they have to overcome their intermolecular attractions with other particles to be a gas Remember Water & its Hydrogen bonds! 14
Phase Changes (Liquid - Gas) The conversion of a liquid to a gas or vapor is called vaporization Let s look at two situations involving vaporization: 1. Vaporization of liquid that is NOT boiling 2. Vaporization of a liquid that IS boiling 15
1. Evaporation If vaporization occurs in a liquid that is not boiling, the process is called evaporation Evaporation occurs at the surface of the liquid Some of the particles have more kinetic energy (KE) than others. The particles with enough KE to overcome intermolecular forces break away and enter the gas or vapor state 16
Effect of Liquid Temperature on Evaporation When you raise a liquid s temperature by adding heat, you are increasing the average kinetic energy of the particles. A higher average kinetic energy means more particles are likely to have enough kinetic energy to overcome attractive forces between particles. So, liquids with higher temperatures will evaporate faster. 17
Effect of Evaporation on Liquid Temperature Particles left behind in liquid are the ones with lower average kinetic energies; they didn t have enough energy to escape Liquid left behind after evaporation is lower in temperature, b/c particles with highest energy left the system 18
Evaporation: A Cooling Process Evaporation is often called a cooling process. On a hot day we sweat. Explain how evaporation plays a role in helping to cool our skin. Body heat is transferred to liquid sweat. As sweat evaporates, high energy particles escape into air. Particles left behind have lower average kinetic energy. Lower kinetic energy on skin lower temperature. Why would this be less effective on a humid day? Diffusion: high concentration low concentration 19
What if the liquid is in a closed container? When some particles do vaporize, these collide with the walls of the container producing vapor pressure Some of the particles will return to the liquid, or condense After a while, the number of particles evaporating will equal the number condensing- the space above the liquid is now saturated with vapor A dynamic equilibrium exists Rate of evaporation = rate of condensation Increasing the temperature increases the KE of the particles increases the vapor pressure TABLE H 20
2. Vaporization of a Boiling Liquid We now know the rate of evaporation from an open container increases as heat is added The heating allows larger numbers of particles at the liquid s surface to overcome the attractive forces Heating allows the average kinetic energy of all particles to increase But what is happening when enough heat is added to make the liquid boil? 21
Boiling a Liquid The temperature at which a liquid will boil is called its boiling point (bp). When we boil a liquid we add enough heat for particles beneath the surface to have enough energy to overcome their attractive forces. When gas forms below the surface a bubble forms. Below the boiling point atmospheric pressure is greater than the vapor pressure of the liquid, this prevents bubbles from forming. At boiling temperature the vapor pressure of the liquid is just equal to the external pressure on the liquid (atmospheric pressure in an open container) 22
Boiling: A Cooling Process Our experiences with boiling are usually related to cooking and involve a continuous source of heat intended to maintain a boiling temperature. So, although we don t think of it as one, boiling is a cooling process just as evaporation is Those particles with highest KE escape first Turning down the source of external heat drops the liquid s temperature below the boiling point Supplying more heat allows particles to acquire enough KE to escape- the temperature does not go above the boiling point, the liquid only boils at a faster rate 23
Effect of Pressure on Boiling Point Since the boiling point is where the vapor pressure equals external pressure, the bp changes if the external pressure changes Normal boiling point- is defined as the boiling point of a liquid at a pressure of 101.3 kpa (or standard pressure) What if atmospheric pressure drops? 24
Effect of Pressure on Boiling Point Montrose: Normal bp of water = 100 o C Denver: Normal bp of water = 95 o C, Denver is 1600 m above sea level and average atmospheric pressure is about 85.3 kpa How pressure cookers work: Sealed container allows pressure to increase beyond atmospheric pressure Higher pressure higher boiling point Temperature rises above 100 o C b/c boiling point is above 100 o C 25
Section 10.3 The Nature of Solids Objectives: Evaluate how the way particles are organized explains the properties of solids. Identify the factors that determine the shape of a crystal. Explain how allotropes of an element are different. 26
The Nature of Solids Most solids have particles packed against one another in a highly organized pattern Tend to be dense and incompressible Do not flow, nor take the shape of their container Are still able to move, unless they would reach absolute zero Images from www.shorstmeyer.com 27
Movement of Particles in a Solids Particles in a solid are not free to move around the way particles in a liquid are Solid particles tend to vibrate about fixed points, rather than sliding from place to place 28
Heating Solids We said particles in a solid vibrate. When a solid is heated, kinetic energy of the particles increases. More kinetic energy means the particles vibrate more rapidly and more vigorously As a result, solids tend to expand when heated 29
Melting Solids If enough heat is added to a solid and the kinetic energy becomes high enough a solid will melt. The melting point (mp) is the temperature at which a solid turns to a liquid At the melting point, the particle vibrations are strong enough to overcome the interactions holding them in a fixed position Melting can be reversed by cooling the liquid so it freezes Solid liquid 30
Melting Points of Solids (REMINDER!) Ionic Compounds: Generally, have high melting points, due to the relatively strong forces holding them together Sodium chloride (an ionic compound) has a melting point = 801 o C Molecular Compounds: Generally, have relatively low melting points Not all solids melt- wood and cane sugar tend to decompose when heated Decompose = chemical change, not reversible 31
Microscopic Structure of Solids When distinguishing b/w solid, liquid and gas, we often represent solids as a group of particles, tightly packed but in no particular order In reality, most solid substances are crystalline in structure In a crystal, the particles (atoms, ions, or molecules) are arranged in a orderly, repeating, three-dimensional pattern called a crystal lattice All crystals have a regular shape, which reflects their arrangement The smallest group of particles within a crystal that retains the geometric shape of the crystal is known as a unit cell 32
Carbon Solids Carbon is a good example of the significance the crystal structure plays in the compound characteristics. Diamond, Graphite, Buckminsterfullerene are three different substances formed from pure carbon. Differences in how the C atoms bond to each determine the crystal structure and the substance properties 33
Carbon Solids Allotropes are two or more different molecular forms of the same element in the same physical state Diamond, Graphite & Buckminsterfullerene are called allotropes of carbon, because all are made of pure carbon only, and all are solid Few elements have allotropes Phosphorus, sulfur and oxygen Boron and antimony 34