Chapter 14

1. Chapter 14 Condensed States of Matter

2. I. States of Matter • Three common states • Solid • Gas • Liquid • Changes to solid if temp lowered below freezing point • Changes to gas if temp is raised about boiling point • These changes are called phase changes • During this process, the substance is not chemically changed, it is physically changed

3. A. Kinetic Molecular Theory: Liquids and Solids • Kinetic molecular theory of gases used assumptions that simplified and allowed us to predict behavior • Particles of an ideal gas are in constant, random, straight-line motion • Collisions between gas particles are perfectly elastic • Average kinetic energy of gas particles depends upon temperature • Gas particles exert no attractive or repulsive forces on one another

4. Why do these assumptions work? • Gas particles are relatively far apart (as compared to particles of a liquid or solid) • Since they are so far apart, the forces of attraction are very weak • At moderate temperatures and pressures, the attraction can be ignored • These assumptions do not work for liquids and solids because the particles are closer together

5. Intermolecular Forces • Covalent bonds are the force that holds individual atoms of a molecule together • So, what holds one molecule close to another in a liquid or solid? • Intermolecular forces: forces of attraction between two or more molecules • Which state of matter has the weakest intermolecular forces? • Gases Why?

6. Effect of Temperature • Average kinetic energy of a gas increases with temperature • At higher temperatures, if two particles get close enough to be attracted, the average kinetic energy is great enough to overcome this attraction • At lower temperatures, the average kinetic energy is lower, and therefore the attraction cannot be overcome and the particles cannot break away from each other • They group together and form a liquid • This process is known as condensation • As intermolecular forces increase, condensation occurs at higher temperatures

7. Effect of Pressure • Increasing pressure forces gas particles closer together • The closer they are, the stronger the intermolecular forces • If pressure is increased enough, particles are close enough together to condense and form a liquid

8. B. Types of Intermolecular Forces • Van der Waals Forces: The attractions between neutral particles • London dispersion forces • Occur between nonpolar molecules • Due to instantaneous dipoles (temporary) which arises from random movements of electrons • This can cause an induced dipole in another molecule A second atom or molecule, in turn, can be distorted by the appearance of the dipole in the first atom or molecule (because electrons repel one another) which leads to an electrostatic attraction between the two atoms or molecules. Because of the constant motion of the electrons, an atom or molecule can develop a temporary (instantaneous) dipole when its electrons are distributed unsymmetrically about the nucleus

9. Dipole- dipole forces • Occur between (neutral) polar molecules • Electrostatic connection between positive end of one dipole and the negative end of an adjacent dipole

10. Hydrogen bonding • Dipole-dipole interactions that involve a hydrogen atom in a polar bond (F, O, & N)

11. Ion-Dipole Forces: The attraction between an ion and the charged end of a polar molecule • Common in solutions containing soluble ionic salts

12. II. Change of State • A. Phase Changes • Evaporation: liquid to vapor • Rate is dependent upon temperature • At any given temperature, the rate is constant • More and more particle enter the vapor phase • This increases the number of gas particles hitting the liquid surface and some particles condense • So as the number of gas particles increases (over time), the rate of condensation increases • Eventually, rate of vaporization and rate of condensation become equal • Explain the demonstration in terms of relative strength of intermolecular forces

13. Vapor Pressure: partial pressure of the vapor above a liquid • Since partial pressure of a gas increases as number of gas particles increases, evaporation increases the vapor pressure • At each temperature, vapor pressure is constant (since rate of evaporation is constant) • So, as temperature increases, vapor pressure increases due to increased evaporation • Is related to strength of intermolecular forces • Stronger forces equal lower vapor pressure • Molecules that readily evaporate (due to weak IM forces) are referred to as being volatile

14. Boiling point: temperature at which the vapor pressure equals the external pressure acting on a liquid • How was this definition arrived at? • At sea level (where atmospheric pressure is 1 atm), water boils at 100°C • In Denver (where altitude is 1600 m and atmospheric pressure is 0.83 atm), water boils at 95°C • The increased pressure at sea level compresses the particles, increasing the intermolecular forces • This means it takes more energy (or higher temperature) to break apart those forces • Explain the demonstration in terms of external pressure and intermolecular forces

15. Freezing point • Change from liquid to solid • Reverse process is called melting or fusion • The stronger the intermolecular forces, the higher the melting point • Sublimation • Change from solid directly to gas • Reverse process is called deposition

16. B. Heating Curves • Temperature v Time graph • Heat of fusion • Energy it takes to melt a substance • Molar heat of fusion • Energy required to change 1 mole of ice into 1 mole of water at 0°C • Heat of vaporization • Energy to boil a substance • Molar heat of vaporization • Energy required to change 1 mole of water into 1 mole of water vapor

17. III. Structure and Properties of Liquids • A. Surface Tension • When the surface of a liquid acts like a weak, elastic skin • The attractive force that occur between molecules in a single (same) substance is called cohesion • This force pulls equally in each direction on individual particles (and therefore no net force) in the interior of a liquid • On the surface, however, particles are attracted to the interior so there is a net inward pull

18. Strength is related to the strength of intermolecular forces • The stronger the intermolecular forces, the greater the cohesion, thus the greater the surface tension • Prevents water from being a good wetting agent • Force of attraction between two different substances is called adhesion • “wetting ability” depends on relative strength of cohesion and adhesion • Explain why the addition of soap caused the needle to sink in the demonstration

19. B. Water: A Unique Liquid • Densities of water and ice • Hydrogen bonding occurs in the solid and liquid phases of water • When water freezes, the hydrogen bonds define the three-dimensional shape of the crystal • In this case, the molecules are held farther apart in the solid phase • This results in ice being less dense than liquid water, so ice floats • Explain how this property (expansion in solid form) is beneficial to plant and animal life in lakes.

20. Water’s heat capacity • Heat capacity: the thermal energy required to raise the temperature of one kilogram of a substance by one degree Kelvin • Water has much higher heat capacity compared to other liquids due to strong hydrogen bonding • While other liquids also have hydrogen bonding, the HB in water is 3-Dimensional due to two unbonded pairs of electrons on each oxygen

21. IV. Structure and Properties of Solids • Solids are classified as either crystalline or amorphous (without shape) • A. Metallic and Covalent Solids • Metallic solids • Unit cell: the smallest unit from which the larger crystal can be built (i.e. the repeating pattern in smallest form) • Seven different types found in all crystalline solids; four of them in metallic solids • Three are cube-shaped; one is hexagonal • Held together by nondirectional metallic bonds • Covalent network solids • Held together by extensive network of directional covalent bonds

22. B. Molecular and Ionic Solids • Molecular solids • Molecules held in place by relatively weak intermolecular forces • Ionic Solids • Three-dimensional arrangements of positive and negative ions • Held together by electrostatic attractions • Arrangement depends on: • Sizes of ions • Ratio of cations to anions