Chapter 10 – States of Matter

1. Chapter 10 – States of Matter GHS Enriched Chemistry Mr. Barbo

2. Section One - The Kinetic-Molecular Theory of Matter • Developed in the late 19th century. • Explains the behavior of the atoms and molecules that make up matter and the forces that act between them; • Based on the idea that particles of matter are always in motion.

3. Main Idea:The KMT explains the constant motion of gas particles. • Ideal gas – a hypothetical gas that perfectly fits all the assumptions of the kinetic-molecular theory. • It is based on five assumptions: • Gases consist of large numbers of tiny particles that are far apart relative to their size. • Gases occupy a volume roughly 1000 times greater than the volume occupied by an equal number of particles in the liquid or solid state. • Much more space between molecules, which allows gases to be compressed.

4. Ideal Gases (continued) • Collisions between gas particles and between particles and container walls are elastic collision. • One in which there is no net loss of total kinetic energy. • Gas particles are in continuous, rapid, random motion. They therefore possess kinetic energy, which is energy in motion. • There are no forces of attraction between gas particles. • The temperature of a gas depends on the average kinetic energy of the particles of the gas.

5. Main Idea:The KMT explains the physical properties of gases. • Expansion – gases have no definite shape or volume • Fluidity • Gas particles slide past each other just like liquids do. • Because liquids and gases flow, they are both considered fluids. • Low Density – 1/1000th the density of liquid or solid • Compressibility – volume of a gas can be greatly decreased • Diffusion and Effusion • diffusion – spontaneous mixing of the particles of two substances caused by their random motion. • effusion – a process by which gas particles pass through a tiny opening

6. Main Idea:Real gases do not behave according to the KMT. • A real gas is a gas that does not behave completely according to the assumptions of the KMT. • At high pressures and low temperatures gases tend to behave like a non-ideal gas. • Also, nonpolar gas molecules tend to behave more like ideal gases than do polar gas molecules, which are attracted to each other.

7. Section 10-2Liquids • Liquids are the least common state of matter in the universe. • This is due to the fact that a substance can exist in the liquid state only within a relatively narrow range of temperatures and pressures.

8. The intermolecular forces of liquids determine their properties. • Liquids have a definite volume and take the shape of their container. • The particles in a liquid are closer together than the particles in a gas. • Therefore, the intermolecular forces exert a stronger pull on the particles. • Dipole-dipole forces, London dispersion forces, and hydrogen bonding • Liquids are more ordered than gases because of the stronger I.F. and the lower mobility of the liquid particles. • Fluid – a substance that can flow and therefore take the shape of its container.

9. Relatively High Density • At normal atmospheric pressure, most substances are hundreds of times denser in a liquid state than a gaseous state due to the close arrangement of liquid particles. • Most substances are slightly less dense as liquids than solids (about 10%) • Water is an exception. • As water cools, its density increases. At 4oC, however, the density of water starts to decrease. The result is that ice is less dense than liquid water. • This is a good thing, as ice provides a protective shield for the organisms that live in the water.

10. Relative Incompressibility • When liquid water at 20oC is compressed by a pressure of 1000 atm, its volume decreases by only 4%. • Compare this with gases, that can be compressed to 1/1000th their original volume. • Liquid particles are more closely packed together.

11. Ability to Diffuse • Constant, random motion of liquid particles causes them to diffuse throughout any other liquid in which it can dissolve. • Diffusion is slower in liquid than gases because the liquid particles are closer together and the attractive forces slow their movement. • Diffusion speeds up as temperature speeds up.

12. Surface Tension • A property common to all liquids is surface tension, a force that tends to pull adjacent parts of a liquid’s surface together, thereby decreasing surface are to the smallest possible size. • Water has a higher S.T. than most liquids due to hydrogen bonds. • S.T. is responsible for the formation of droplets. • Capillary action, the attraction of the surface of a liquid to the surface of a solid, is a property closely related to surface tension. • Responsible for the meniscus in a graduated cylinder and for the transportation of water from the roots of trees and plants to their leaves.

13. Surface tension

14. Evaporation and Boiling • Vaporization – process by which a liquid or solid changes to a gas. • Evaporation – occurs at the surface of a nonboiling liquid. • Occurs because particles with higher-than-average energies can overcome the intermolecular forces that bind them to the liquid. • They then escape into the gas state. • Evaporation is a vital process to the water cycle and evaporative cooling. • Boiling is the change of a liquid to bubbles of vapor that appear throughout the liquid. It is different than evaporation.

15. Formation of Solids • When the average K.E. of particles decreases enough, attractive forces pull the particles into an even more orderly arrangement called a solid. • Freezing is an example of this. • Water – 0oC • Paraffin – room temp • Ethanol – -114oC

16. Section 3 - Solids Main Idea – The particles in a solid hold relatively fixed positions. • Particles are more closely packed. • Intermolecular forces exert stronger effects than in liquids or gases. • Particles are held in fixed positions, with only vibrational movement around fixed points.

17. Main Idea – The particles in a solid hold relatively fixed positions. (continued) • There are two types of solids • Crystalline solids – consists of crystals, which are substances arranged in an orderly, geometric, repeating pattern. • Amorphous solids – particles are arranged randomly.

18. Definite Shape and Volume • Solids maintain shape without a container. • Crystalline solids have distinct geometric shapes • Amorphous solids maintain a shape, but lack the distinct geometric shapes of crystals. • The volume of a solid changes only slightly with a change in temperature or pressure.

19. Definite Melting Point • Melting – the physical change of a solid to a liquid by the addition of energy as heat. • The temperature at which a solid becomes a liquid is its melting point. • At this temperature, the kinetic energies of the particles within the solid overcome the attractive forces holding them together • Crystalline solids have a definite melting point, while amorphous solids have no definite melting point.

20. High Density and Incompressibility • This results from the fact that the particles of a solid are more closely packed than those of a liquid or gas. • Solid hydrogen is the least dense solid (0.086 g/cm3), while osmium is the most dense (22.6 g/mL). • Solids that seem compressible (cork, Styrofoam) actually have air pockets in them.

21. Low Rate of Diffusion • Rate of diffusion between solids is millions of times slower than in liquids or gases.

22. Main Idea – The particles in amorphous solids are not arranged in a regular pattern. • The word amorphous comes from the Greek for “without shape.” • Unlike crystals, particles that make up amorphous solids are not arranged in a regular pattern. • Glass and plastics are examples of amorphous solids.

23. Section Four – Changes of State Main Idea – Substances in equilibrium change back and forth between states at equal rates. • In a closed bottle of liquid, gas molecules under the cap strike the liquid surface and reenter the liquid phase through condensation. • A gas in contact with its liquid or solid phase is often called a vapor.

24. Equilibrium Vapor Pressure • Evaporation has a different outcome in an open system, such as a lake or an open container, than in a closed system, such as a sealed container. • In a closed system, the molecules collect as a vapor above the liquid. Some condense back into a liquid. • In an open system, molecules that evaporate can escape from the system.

25. When a partially filled container of liquid is sealed, some of the particles at the surface of the liquid vaporize. • These particles collide with the walls of the sealed container, producing pressure. • vapor pressure: a measure of the pressure/force exerted by a gas above a liquid(at equilibrium)

26. Equilibrium Vapor Pressure of a Liquid • The pressure exerted by a vapor in equilibrium with its corresponding liquid at a given temperature is called the equilibrium vapor pressure of the liquid. • Every liquid has a specific equilibrium vapor pressure at a given temperature. • The stronger the attractive forces are, the smaller the percentage of liquid particles that can evaporate, leading to a low equilibrium vapor pressure. • Volatile liquids, which are liquids that evaporate readily, have weak forces of attraction between their particles.

27. Vapor Pressure and Temperature Change • An increase in the temperature of a contained liquid increases the vapor pressure. • But, Vapor pressure is not affected by surface area of the liquid.

28. The vapor pressure data indicates how volatile a given liquid is, or how easily it evaporates. • Of the three liquids shown, diethyl ether is the most volatile and water is the least volatile.

29. 12.2 mm Hg or 1.63 kPa • 43.9 mm Hg or 5.85 kPa • Mercury • Mercury • Mercury • Air • Ethanol • Ethanol • Air at standard temperature and pressure • Ethanol at 0°C • Ethanol at room temperature (20°C) Vapor Pressure Measurements • The vapor pressure of a liquid can be determined with a manometer. The vapor pressure is equal to the difference in height of the mercury in the two arms of the U-tube.

30. So what does this have to do with boiling? • Boiling – the conversion of a liquid to a vapor within the liquid as well as at its surface. • boiling point - temperature at which the vapor pressure of the liquid is just equal to the external pressure(atmospheric pressure)on the liquid. • When a liquid is heated to a temperature at which particles throughout the liquid have enough kinetic energy to vaporize, the liquid begins to boil. • The lower the atmospheric pressure the lower the boiling point is.

31. 100°C • 101.3 kPa • 101.3 kPa • 34 kPa • 70°C • 70°C • Sea Level • Sea Level • Atop Mount Everest • Atmospheric pressure at the surface of water at 70°C is greater than its vapor pressure. Bubbles of vapor cannot form in the water, and it does not boil. • At the boiling point, the vapor pressure is equal to the atmospheric pressure. Bubbles of vapor form in the water, and it boils. • At higher altitudes, the atmospheric pressure is lower than it is at sea level. Thus, the water boils at a lower temperature. Boiling Point and Pressure Changes

32. Section 5 - Water What’s so great about water? • About 75% of the Earth is covered by water as rivers, oceans, lakes and glaciers. • Living organisms are 70 – 90% water. • Most chemical reactions take place in water. • In short, water is incredibly vital to our planet and everything living on it!

33. Section 5 - Water Main Idea – The properties of water in all phases are determined by its structure. • Water molecules consist of two hydrogen atoms and one oxygen atom united by polar-covalent bonds. • The angle between the two hydrogen atoms is about 105o.

34. Hydrogen Bonds • Water molecules are linked by hydrogen bonds. • The number of hydrogen bonds decreases with increasing temp and increase with decreasing temp. • If it weren’t for hydrogen bonds, water would be a gas at room temp • As liquid water is warmed from 0oC, water molecules pack closer together. • Water molecules are most closely packed together at 3.98oC.

35. Main Idea – The molar enthalpy of water determines many of its physical characteristics. • Because its molecules form an open rigid structure, ice at 0oC has a density of about .917 g/cm3, while liquid water is 0.99984 g/cm3. • Ice is therefore less dense than liquid water, which is vitally important. • The MEF of ice is 6.009 kJ/mol, which is relatively large in comparison. • Water’s boiling point is 100oC and its MEV is 40.79. • Both are quite high compared to nonpolar substances of similar molecular mass due to water’s hydrogen bonding. • Radiators/evaporative cooling

36. Evaporativecooling