12.2 Forces of attraction

1. 12.2 Forces of attraction • Intermolecular forces (dispersion forces, dipole-dipole forces, and hydrogen bonds) determine a substance’s state at a given temperature

2. Intermolecular forces • Inter- means between or among • Intermolecular forces can hold together identical particles or two different types of particles • Weaker than intramolecular forces (bonds)

3. Dispersion Forces • Weak forces that result from temporary shifts in the density of electrons in electron clouds

4. Exist between all particles • Weak for small particles • Get stronger as the number of electrons involved increases • F2 • Cl2 • Br2 • I2

5. Dipole-dipole forces • Attraction between oppositely charged regions of polar molecules • Polar molecule = • Neighboring polar molecules orient themselves so that oppositely charged regions align

7. Hydrogen Bonds • Dipole-dipole attraction that occurs between molecules containing a hydrogen atom bonded to a flourine, oxygen, or nitrogen atom

8. Explain why water is liquid at room temperature while compounds of similar masses are gasses

9. 12.3 Liquids and Solids • The particles in solids and liquids have a limited range of motion and are not easily compressed.

10. Liquids • Kinetic molecular theory also applies to liquids and solids • Must take intermolecular forces into account to apply it

11. Density and compression • Much denser than gasses • Due to intermolecular forces holding particles together • Incompressible • Why can you compress a gas but not a liquid?

12. Fluidity – both gases and liquids are classified as fluids because they can flow and diffuse • Liquids diffuse more slowly because intermolecular attractions interfere with the flow

13. Viscosity - measure of the resistance of a liquid to flow • Attractive forces – stronger intermolecular forces = higher viscosity • Particle size – larger molecules = higher viscosity • Temperature – lower temperature = higher viscosity

14. Surface tension – the energy required to increase the surface area of a liquid by a given amount • Caused by intermolecular forces pulling down on the particles on the surface of a liquid which stretches it tight like a drum

15. Stronger the attraction between particles in a liquid = greater surface tension • Surfactant – lowers the surface tension of water by disrupting hydrogen bonds between water molecules

16. Cohesion – force of attraction between identical molecules • Adhesion – force of attraction between molecules that are different

17. Solids • Solid particles have as much kinetic energy as liquids or gasses but much stronger attractive forces between particles • Limit the motion of particles to vibrations

18. Density of solids – almost always greater than density of liquids Exception = water

19. Crystalline solids – solid whose atoms, ions, or molecules are arranged in an orderly, geometric structure • Unit cell = smallest arrangement of atoms in a crystalline solid that has the same shape as the whole crystal

21. Categories of crystalline solids • Classified based on the types of particles they contain and how they are bonded together

22. Molecular solids • Molecules are held together by dispersion forces, dipole-dipole forces or hydrogen bonds • Most are not solid at room temperature • Poor conductors

23. Covalent network solids • C or Si, can form multiple covalent bonds which allow it to take many forms • Allotrope – element that can exist in different forms at the same state

24. Ionic solids • Made of cation + anion • Each ion is surrounded by ions of the opposite charge • High melting point • Brittle

25. Metallic solids • Positive metal ions surrounded by a sea of mobile electrons

27. Amorphous solids • Particles are not arranged in a regular, repeating pattern • Does not contain crystals • Forms when molten material cools too quickly for crystals to form • Glass • Rubber • Some plastics

28. 12.4 Phase changes • Matter changes phases when energy is added or removed

29. Phase changes that require energy • Melting • Heat flows from an object at a higher temperature to an object at a lower temperature • Ice absorbs heat which does not raise temperature but is used to break hydrogen bonds • When hydrogen bonds are broken molecules can move further apart into the liquid phase

30. Melting point – temperature in which forces holding a solid together are broken and it becomes a liquid

31. Vaporization – process by which liquid changes to vapor • Vapor – gaseous state of a substance that is normally liquid at room temperature • Evaporation – when vaporization occurs only at the surface of a liquid • Vapor pressure – the pressure exerted by a vapor over a liquid

32. Boiling – temperature at which the vapor pressure of a liquid equals the atmospheric pressure • Energy being input causes molecules to move around more and vaporize

33. Sublimation – changing from solid to gas without becoming a liquid • Dry ice • Moth balls • Solid air fresheners

34. Phase changes that release energy • Freezing • Heat flows out of warmer object into cooler object • Molecules slow down & become less likely to flow past one another • Intermolecular forces cause the molecules to become fixed into set positions • Freezing point – temperature in which a liquid becomes a solid

35. Condensation – process by which a gas or vapor becomes a liquid • Deposition – substance changes from gas or vapor to solid without first becoming a liquid • frost

36. Phase Diagrams • Temperature and pressure both effect the phase of a substance • Have opposite effects • Phase diagram – graph of pressure vs temperature that shows which phase a substance will be in under different conditions.

37. Triple point = point at which all three phases exist at the same time