Chapter 14: Liquids and Solids

1. Chapter 14: Liquids and Solids

2. 14-1 Condensed States of Matter • Condensed State- substances in these states have substantially higher densities than they do in the gaseous state. • Kinetic-Molecular Theory applied to Liquids and Solids

3. Physical Properties of the States of Matter

4. According to the kinetic-molecular theory, the state of a substance at room temperature depends on the strength of the attractions between its particles. • Attractive forces between solids are higher than liquids and higher than gases.

5. Ex. Water below 0 it’s a solid why? Kinetic energy of the water molecules is too low to overcome the strong attraction between the water molecules. • Above 0, molecules have enough kinetic energy to get away from each other and flow. • At 100 C, average kinetic energy of the molecules is so high they can escape the container.

6. Intramolecular forces • 3 types of chemical bonds. • Ionic – metal + nonmetal transfer electrons (all ionic compounds are solids at room temp) • Metallic – share electrons (sea of electrons) most are solid at room temp. • Covalent – sharing of electrons. Holds liquids and gases together • Even though covalent bonds can be really strong, they are liquids and gases because of intramolecular forces- a force that exists inside the molecule.

7. All of the other forces are Intermolecular forces are forces between molecules. Much weaker than ionic, covalent or metallic bonds. • ~ 463KJ to break an O-H bond but only 6KJ to melt a mole of ice. • No bonds are being broken when ice is melted only intermolecular forces are involved in a change of state.

8. 3 types of intermolecular forces~ All intermolecular forces, increase boiling point. The higher the IM force, the higher the BP • Dispersion Forces- intermolecular force between induced dipoles. • What is a dipole? Most atoms have spherical shape

9. Dispersion • Only happens when there is a temporary pole. Opposites attract. • Since there are 2 poles, it will force its neighbor to attract to whatever side (+ or -) and this is known as induced dipole. • **Only type of intermolecular bond in noble gases • Molecules that are large, have larger dispersion forces (dispersion forces increase with increasing molecular size and mass)..The larger the dispersion force, the higher the boiling point.

10. Dipole-Dipole • Dipole-Dipole forces – attractions between polar molecules. • A permanent dipole is a result of a covalent bond between 2 different atoms. Since they have different electronegativities, the atom with the higher e- has a δ-

11. H bond • Hydrogen Bonding- Hydrogen has a low e- and F, N and O have high e-. Because of the large difference in e-, the bonds are very polar and there is an extremely strong dipole-dipole force. It’s so strong, it’s called a H bond.

12. 14-2 Properties of Liquids • The physical properties of liquids are determined mainly by the nature and strength of the intermolecular forces present between their molecules. • Viscosity – resistance of motion. Tells you how easily something is poured. • The stronger the attraction between the molecules, the greater the resistance to flow, the higher the viscosity. • ~H bonds are so strong that they increase the viscosity. Water has a relatively high viscosity compared to rubbing alcohol and gasoline. • ~Viscosity increases as temp. decreases.

13. Surface Tension – since the molecules at the surface of a glass can only attract down (and out) they act like there is a film on top of the glass. • Water Why it’s so good • Unusually high bp (HF and NH3 are corrosive gases at room temp) • it can absorb or release heat without large changes in temp. • density of ice is less than liquid water • high surface tension. Allows plants to carry water from root to tip • has a high heat of vaporization (sweating) • is the universal solvent bc it’s polar

14. 14-3 The Nature of Solids • Crystalline solids- a solid with a highly ordered, repeating pattern. • Unit cells – the smallest possible repeating pattern of a crystal. • Crystal Formation and Water of hydration • Many chemical reactions for ionic substances occur in water. If water gets trapped in the bonds, it is called a hydrate and the water it was in is known as water of hydration

15. Amorphous solids- appears to be a solid but does not behave like a solid. Glass, rubber, and plastic. Think of them as a super-cooled liquid where the viscosity has become really high. • Bonding in solids – the physical properties of solids, such as hardness, electrical conductivity, and melting point, depend on the kind of particles that make up the solid and the strength of the attractive forces between them.

16. Metals – sea of electrons, good conductors because their valence electrons are free to move. (wires) Malleable (sheets) Ductile (wires) electrons can just change shape to adjust • Molecular solids – low mp because only the weak intermolecular forces have to be broken. Do not conduct b/c no free electrons • Ionic solids – make a crystal to maximize attraction between + and -‘s and minimize repulsion.

17. Covalent-Network solids – Atoms are bonded together with strong covalent atoms without forming molecules. Instead they form a network throughout the crystal.

18. 14-4 Changes of state – aka phase change, is the conversion of a substance from one of the three physical states of matter to another. It always involves a change of energy. • Energy and change of state • When a substances goes from solid to liquid or liquid to gas, it must overcome the attractive forces holding it together in the more condensed state. Gas has the most potential energy this means, gases have to absorb energy to jump up a phase.

19. Vaporization and Condensation • Vaporization = Liquid to gas • Condensation = gas to liquid • Evaporation = going from a liquid to a gas • Vaporization = evaporation and boiling

20. Evaporation • When molecules have enough energy to fly out of their containers. The higher the temp, the more evaporation • ~ volatile liquid – one that evaporate easily because the molecules are not strongly attracted to each other. (rubbing alcohol) • ~Evaporative cooling – the hot fast molecules leave right away and the slower cooler molecules remain.

21. Liquid-Vapor Equilibrium – • If you have a closed container with water in it, the water will evaporate at a constant rate, then it will condense and eventually will be in a dynamic equilibrium (a continually changing balance) • Equilibrium Vapor Pressure – As you increase the temp, there will be a greater number of vapor molecules.

22. Boiling point – as water is heated, bubbles of vapor are formed in it, they will collapse as long as the atmospheric pressure is greater. When the vapor pressure is equal to the atmospheric pressure, boiling will occur. The higher the atm pressure, the higher the boiling point. (why they change temps at high altitude) • Heat of Vaporization- the amount of heat needed to vaporize a certain amount of liquid

23. Freezing and Melting – they are the same! • The heat needed to convert solid to liquid is the heat of fusion • The heat releases during freezing is the opposite sign of heat of fusion • Sublimation and Deposition • Solid  gas sublimation • Gas  solid deposition • Solids with high vapor pressure sublime easily. Molecular solids will sublime the easiest.