Chapter 5 Gases,Liquids, and Solids

1. Chapter 5Gases,Liquids, and Solids

2. States of Matter • The state that a sample of matter exists in at room temperature is dependent several things, including the type of bonding between the atoms in substance. For Ionic Compounds • Ionic compounds tend to be solids with high melting points  due to strong electrostatic attraction between + and - ions in the solid For Molecular Compounds, • State depends on size of molecule and the type of attractive forces between the molecules.

3. Gases • Behavior explained by the kinetic molecular theory of gases

4. Kinetic Molecular Theory of Gases Assumptions of the kinetic molecular theory: • Gases consist of particles constantly moving through space in random directions and with various speeds. • Gas particles have no volume. • Gas particles experience no attractive forces between them. • The average kinetic energy (KE) of gas particles is proportional to the temperature in kelvins. • Molecular collisions are elastic; when molecules collide, they may exchange KE but the total KE remains constant. • Molecules collide with the walls of their container; these collisions constitute the pressure of the gas.

5. Gas Pressure Gas pressure: The pressure is force per unit area exerted against a surface. • Most commonly measured in millimeters of mercury (mm Hg), atmospheres (atm), and torr. • Atmospheric pressure is measured using a barometer (next screen). • The pressure of a confined gas is measured using a manometer.

6. Gas Pressure Figure 5.2 A mercury barometer.

7. Gas Pressure Figure 5.3 A manometer.

8. Gas Laws Boyle’s law: For a fixed mass of gas at a constant temperature, the volume is inversely proportional to the pressure. Pressure varies with 1/V Charles’s Law: For a fixed volume of gas at a constant pressure, the volume is directly proportional to the temperature in kelvins (K).

9. Gas Laws Gay-Lussac’s Law: For a fixed mass of gas at constant volume, the pressure is directly proportional to the temperature in kelvins (K). in summary:

10. Gas Laws Boyle’s law, Charles’s law and Gay-Lussac’s law can be combined into one law called the combined gas law.

11. Gas Laws Problem: A gas occupies 3.00 L at 2.00 atm. Calculate its volume when the pressure is 10.15 atm at the same temperature. Begin with the combined gas law and solve for V2. Because the temperature is constant T1 = T2

12. Gas Laws Avogadro’s law: Equal volumes of gas at the same temperature and pressure contain the same numbers of molecules. • The actual temperature and pressure at which we compare two or more gases does not matter. • For convenience in making comparisons, chemists have selected one pressure as a standard pressure, and one temperature as a standard temperature. • The standard temperature and pressure (STP) selected are 0°C (273 K) and 1 atm pressure.

13. Gas Laws • All gases at STP or any other combination of pressure and temperature contain the same number of molecules in a given volume. But how many molecules is that? • One mole contains 6.022 x 1023 formula units; what volume of gas at STP contains this many molecules? • This quantity has been measured and found to be 22.4 L. • Thus, one mole of any gas at STP occupies 22.4 L.

14. Ideal Gas Law Avogadro’s law allows us to write a gas law that is valid not only for any P, V, and T but also for any mass of gas. Ideal gas law: PV = nRT P = pressure of the gas in atmospheres (atm) V = volume of the gas in liters (L) n = moles of the gas (mol) T = temperature in kelvins (K) R = ideal gas constant (a constant for all gases)

15. Ideal Gas Law We find the value of R by using the fact that 1.00 mol of any gas at STP occupies 22.4 L. • Problem: 1.00 mol of CH4 gas occupies 20.0 L at 1.00 atm. What is the temperature of the gas in kelvins? • Solution: Solve the ideal gas law for T, plug in the given values, and do the math:

16. Dalton’s Law of Partial Pressures Dalton’s law of partial pressures: The total pressure, PT, of a mixture of gases is the sum of the partial pressures, P, of each individual gas: • Problem: To a tank containing N2 at 2.0 atm and O2 at 1.0 atm we add an unknown quantity of CO2 until the total pressure in the tank is 4.6 atm. What is the partial pressure of CO2?

17. Kinetic Molecular Theory of Gases Ideal gas: The six assumptions of the KMT give us an idealized picture of the particles of a gas and their interactions with one another. Real gases • Their atoms or molecules do occupy some volume. • There are forces of attraction between their atoms or molecules. In reality, no gases are ideal. • At pressures below 1 to 2 atm and temperatures well above their boiling points, most real gases behave in much the same way as predicted by the KMT.

18. Intermolecular Forces (IMFs) The strength of attractive forces between molecules determines whether any sample of matter is a gas, liquid, or solid. • At or near STP, the forces of attraction between molecules of most gases are so small that they can be ignored. • When T decreases or P increases or both, the forces of attraction become important to the point that they cause condensation (gas to liquid) and ultimately solidification (liquid to solid). • In order to understand the properties of liquids and solids, we must look at the nature of these intermolecular forces of attraction.

19. Intermolecular Forces We discuss three types of intermolecular forces. • Their origins are electrostatic, that is, the attraction between positive and negative charges. • The strengths of covalent bonds are shown for comparison.

20. Attractive Forces Between Molecules • Must occur, otherwise gases wouldn’t condense Example: Liquid N2

21. 3 Main Types of IMFs • London Dispersion Forces • Dipole-Dipole Forces • Hydrogen Bonding

22. London Dispersion Forces • Making the Modern World - Intermolecular forces: Part 1

23. London Dispersion Forces • Arise from the motion of electrons within atoms and molecules • Movement of e- from one region to another results in e- imbalance in molecule, which causes e- in adjacent molecule to move away from higher electroneg. Region of first molecule • Result = temporary dipole, lead to short-lived electrostatic attractions between molecules or atoms

24. Dispersion Forces Analogy • “The wave” at a sports event

25. Dispersion Forces Analogy • The movement of one person influences and changes the behavior of another, and another, and so on • Long after original person has sat down, the result of his motions can still be seen through the crowd • So it is with electrons • London forces increase with increasing mass (# of electrons) of atom or molecule

26. London Dispersion Forces London dispersion forces are the attraction between temporary induced dipoles.

27. London Dispersion Forces • London dispersion forces exist between all atoms and molecules. • They are the only forces of attraction between atoms and nonpolar molecules. • In general, their strength increases as the mass and number of electrons in a molecule increases. • Even though these forces are very weak, they contribute significantly to the attractive forces between large molecules because they act over large surface areas.

28. Dipole-Dipole Interactions Dipole-dipole interactions: the electrostatic attraction between positive and negative dipoles. • Consider butane and acetone, compounds of similar molecular weight. • Butane is a nonpolar molecule. The only interactions between butane molecules are London forces. • Acetone is a polar molecule. Its molecules are held together in the liquid state by dipole-dipole interactions.

29. Hydrogen Bonding • Responsible for water’s • high boiling point • High heat of vap -lots of energy needed to pull molecules apart from liquid to gas • High surface tension • ice being less dense than water, to maximize h-bonds, expand into hexagons

30. Hydrogen Bonds • The strength of hydrogen bonds ranges from 2 to 10 kcal/mol. • The strength in water is approximately 5.0 kcal/mol. • By comparison, the strength of an O-H covalent bond in a water molecule is 119 kcal/mol. • Nonetheless, hydrogen bonding in liquid water has an important effect on the physical properties of water. • The relatively high boiling point of water is due to hydrogen bonding between water molecules. Extra energy is required to separate a water molecule from its neighbors. • Hydrogen bonds are not restricted to water; they form whenever there are O-H or N-H groups.

31. Intermolecular Attractive Forces (IMF) London Dispersion Forces • Occur between all molecules and all noble gas atoms Dipole-Dipole Forces • Occur between polar molecules Hydrogen Bonding • Occurs between molecules made from hydrogen and N, O, or F

33. Implications of IMFs “Like dissolves like” Rule • Polar substances will only dissolve other polar substances • Reason oil and water don’t mix • Oils contain long chain hydrocarbons, nonpolar oil nonpolar water

34. Implications for Biomolecules DNA – hydrogen bonding holds the double helix together, allows for replication

35. Fats • Made from fatty acids and glycerol • Fatty acids are long chain hydrocarbons, hydrocarbon region is nonpolar Fatty acid in beef fat Nonpolar hydrocarbon tail

36. Saturated vs. Unsaturated Fats • Unsaturated fats contain fatty acids in which there is at least 1 carbon-carbon double bond in the molecule. • This “kinks” the molecule, it can’t line up flat with other molecules, liquids at room temp “kink”

37. Saturated vs. Unsaturated Fats • Saturated fats contain fatty acids in which there are only single bonds between carbon atoms • Result is long, straight chains that stack nicely, forming solids at room temp

38. Cell Membranes Outside cell Inside cell Nonpolar hydrocarbon tails

39. Proteins • The 3-D shape of proteins is due to several types of intermolecular attractive forces

40. Liquids • As pressure increases in a real gas, its molecules come closer and closer with the result that attractions between molecules become important. • When distances decrease so that almost all molecules touch or almost touch, a gas condenses to a liquid. • In liquids, there is very little space between molecules; consequently, liquids are difficult to compress. • The density of liquids is much greater than that of gases because the same mass of molecules occupies a much smaller volume in the liquid state. • The position of molecules in a liquid is random and there is irregular space between them into which other molecules can slide; this causes liquids to be fluid.

41. Surface Tension Surface tension: The layer on the surface of a liquid produced by uneven intermolecular attractions at its surface: • Molecules in the interior of a liquid have equal intermolecular forces in all directions. • Molecules at the liquid-gas interface experience a greater attraction toward the interior of the liquid than toward the gas phase above it. • Therefore, there is a preferential pull of molecules on the surface toward the interior of the liquid. • This preferential pull crowds the molecules on the surface, and creates a thin elastic skin-like layer. • Surface tension is directly related to strength of the intermolecular attraction between molecules.

42. Surface Tension Figure 5.11 Surface tension.

43. Chapter 5 Gases, Liquids, Solids End Chapter 5