1. Unit 11: States of Matter Ch. 13, Sections 2-4

2. ++ -- Types of Covalent Bonds Polar Covalent Bond –e - are shared unequally –asymmetrical e - density –results in partial charges (dipole)

3. Nonpolar Covalent Bond –e - are shared equally –symmetrical e - density –usually identical atoms Types of Covalent Bonds

4. If ΔEN is:Bond type is: < 0.4Nonpolar covalent 0.4 < Δ EN < 1.7Polar covalent > 1.7Ionic Given: Electronegativities of these elements H = 2.2C = 2.55N = 3.04O = 3.44F = 3.98 Na = 0.93K = 0.82P = 2.19S = 2.58Cl = 3.16 Determine bond type for the following bonds: H – H_____________________H – O________________________ H – C_____________________Na – Cl________________________ C – Cl _____________________F – F ________________________ K – F _____________________H – N ________________________ Take the absolute value of the difference!

5. Intermolecular Forces (IMF) Attractive forces between molecules. Much weaker than chemical bonds within molecules. a.k.a. van der Waals forces

6. Types of IMF London dispersion forces –Exist in all atoms and molecules (temporary dipoles) –Attraction between two instantaneous dipoles –Weakest –Larger the molecule, larger the dispersion force.

7. Types of IMF London Dispersion Forces View animation online.animation

8. Types of IMF -Dipole-Dipole forces -Attraction between two permanent dipoles -Polar molecules -Medium strength -Stronger when molecules are closer together

9. Types of IMF Dipole-Dipole Forces + + - - View animation online.animation

10. Types of IMF Hydrogen bonding –Special kind of dipole- dipole –Occurs between molecules that have an H bonded to either O, N, or F. –Extremely polar bonds –Strongest –Not chemical bonding

11. Types of IMF Hydrogen Bonding

12. Types of IMF The hydrogen bonds in water explain its relatively high boiling point, considering that it is a small molecule. The H-bonds hold the water molecules together as a liquid, so you have to heat it a lot before it will change to a gas. Compare boiling points of these molecules: MoleculeIMF (s) presentMolar Mass (g/mol) Boiling Point ( o C) CH 4 16.05- 164 HCl36.46- 85 H2OH2O18.02100 London Disp. Dipole-Dipole London Disp./Dipole- Dipole/Hydrogen Bonding

13. Properties of Liquids Density: –Higher in liquids than gases Intermolecular forces hold liquid particles together Gas particles are far apart and have minimal attraction between particles. Compressibility: –Only slightly in liquids Molecules already closely packed Gases have a lot of empty space between particles and can be compressed.

14. Liquid Properties Viscosity –Resistance to flow Viscosity Demo

15. Liquid Properties Surface Tension –attractive force between particles in a liquid that minimizes surface area

16. Action of detergents: Detergents are used to clean dirt/oil from clothing, dishes, etc. How does this work? – The detergent interacts with water in such a way that the hydrogen bonding is disrupted, and this breaks the surface tension of the water, so that the dirt can mix with the water. Detergents are called surfactants, because they are surface active agents. Detergents have a special structure that aids the cleaning process. A micelle is formed when detergents interact with dirt/oil and water. –Water is polar. –Dirt and oils are nonpolar. –Detergents are long molecules that have polar “heads” and nonpolar “tails”. –“like dissolves like” Because detergents have a dual nature, they can dissolve grease and carry it away in the water at the same time.

17. Cohesion and Adhesion Cohesion is the force of attraction between identical molecules in a liquid (cohesion is a result of intermolecular forces). Adhesion is the force of attraction between liquid molecules and a solid that is touching them.

18. Liquid Properties Capillary Action –attractive force between the surface of a liquid and the surface of a solid watermercury Meniscus of water in a glass tube is concave: adhesion > cohesion Meniscus of Hg in a glass tube is convex: cohesion > adhesion

19. PROPERTIES OF SOLIDS Density: In general, the particles of a solid are more closely packed than those of a liquid. So most solids are more dense than most liquids, but there are exceptions (wax, cork). When solid and liquid states of a substance coexist, the solid is more dense than the liquid, so the solid will sink in the liquid. Water is an important exception (ice floats in water). This is because the H 2 O molecules are less closely packed in ice than in liquid water.

20. Types of Solids Crystalline - repeating geometric pattern –covalent network –metallic –ionic –covalent molecular Amorphous – –no geometric pattern decreasing m.p.

21. Crystalline vs. Amorphous: A crystalline solid has particles which are arranged in an orderly, geometric, 3- D structure. – Examples: sodium chloride, ice, gems and minerals In an amorphous solid, the particles are not arranged in any particular pattern. –Examples: rubber, plastics.

22. Types of Solids Covalent Molecular (H 2 O) Covalent Network (SiO 2 - quartz) Amorphous (SiO 2 - glass)

23. Phase Changes Temperature E. Q=mCΔT D. Q= mH v T b  B. Q =mH f C. Q=mCΔT T m  A. Q= mCΔT Thermal Energy (Heat)

24. Thermal Energy (heat) Phase Changes: –B and D represent phase changes –Occur at constant temperature Temperature Changes –A,C, and E –Temp is changing –Sloping portion of the graph

25. Thermal Energy (heat ) What is happening at each part of the graph: A. Substance is a solid. Can heat it up or cool it down along this line. B. Phase change: solid-liquid. The temperature at B (Tm) is the melting (freezing) point. C. Substance is a liquid. Can heat it up or cool it down along this line. D. Phase change: liquid-gas. The temperature at D (Tb) is the boiling (condensation) point. E. Substance is a gas. Can heat it up or cool it down along this line.

26. Heat Calculations Q = m c ∆ T –Q = heat or thermal energy in Joules (J) or calories (cal) –m= mass in grams (g) –C = specific heat in J/(g o C) or cal/(g o C) –∆T= change in temperature in oC

27. Heat Calculations Q =mH f –Q = heat or thermal energy in Joules (J) or calories (cal) –m= mass in grams (g) –H f = Heat of fusion (J/g or cal/g); use for liquid- solid phase change

28. Heat Calculations Q =mH v –Q = heat or thermal energy in Joules (J) or calories (cal) –m= mass in grams (g) –Hv = Heat of vaporization (J/g or cal/g); use for gas-liquid phase change

29. Things to remember You can move from left to right or right to left along the curve. If you move left to right: – all the processes are endothermic (heat must be supplied). –Your answer for heat calculations will be positive.

30. Things to remember If you move right to left: – all the processes are exothermic (heat is removed/released). – Your answer for heat calculations will be negative. You need to hand-correct Hf and Hv so they are negative.