1. Liquids and Solids
Chapter 12

2. Liquid Has a definite volume and indefinite shape
Particles are in constant motion
Closer together than gases
Less kinetic energy than gases
Greater attractive forces than gases
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3. Fluid
Substance that can flow and therefore take the shape of its container
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4. Fluid density
At normal pressure, most liquids are thousands of times denser than their gases.
Particles are closer together
Different liquids can vary greatly in density
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5. Density
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6. Incompressibility Liquids are much less compressible than gases
Particles are closer together
Liquids can transmit pressure throughout themselves
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7. Diffusion Liquids diffuse in other liquids in which they can dissolve
Much slower than gases
Particles closer together
Attraction between particles slows them down
Faster at higher temperatures
More kinetic energy
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8. Surface Tension
Force that pulls parts of a liquid’s surface together, causing it to have the smallest possible size
From attractive forces between molecules
Liquid droplets take a spherical shape
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9. Capillary action
The attraction of the surface of a liquid to the surface of a solid
Causes meniscus
Water can travel up paper
Water traveling up a plant
Drawing blood in capillary tube
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10. Vaporization Changing from a liquid to a gas Evaporation Boiling
Higher energy particles escape from the surface of a nonboiling liquid
perfume
Boiling
Bubbles of vapor that appear throughout the liquid and travel to the surface
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11. Freezing Changing a liquid to a solid by removing heat
Energy of particles decreases until they are pulled into a more orderly arrangement
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12. Discussion Describe the liquid state using kinetic molecular theory.
Explain why liquids in a test tube form a meniscus.
Compare and contrast vaporization and evaporation.
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13. Solid Has definite volume and definite shape
Particles are in constant motion
Much closer together than liquid or gas
Much stronger intermolecular forces
Held in relatively fixed position – only vibrate
Most ordered state of matter
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14. High density and incompressibility
Substances are generally the most dense in the solid state
Slightly denser than liquids, much denser than gases
Virtually incompressible
Sometimes we compress air pockets in the solids
Wood, cork, etc.
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15. Diffusion Very, very slow
A few atoms may diffuse if clamped together for a long time
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16. Melting Change of a solid to a liquid by addition of heat
Melting point – temperature at which something melts
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17. Crystalline solids Consist of crystals Fragments have geometric shapes
Particles are arranged in an orderly, geometric, repeating pattern
Fragments have geometric shapes
Have definite melting points
When the crystal structure breaks apart
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18. Crystal structure
Total three-dimensional arrangement of particles in a crystal
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19. Lattice
Coordinate system that represents the arrangement of particles in a crystal.
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20. Unit cell
Smallest portion of a crystal lattice that shows the 3D pattern of the entire lattice
Each crystal lattice contains many unit cells packed together
Has one of seven types of symmetry – see page 369
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21. Ionic crystals Positive and negative ions in a regular pattern
Hard and brittle
High melting points
Good insulators
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22. Covalent network crystals
Individual atoms connected by covalent bonds
Giant molecules
Diamond
Quartz
Very hard and brittle
Rather high melting points
Nonconductors or semiconductors
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23. Metallic crystals Metal atoms surrounded by a sea of electrons
High electrical conductivity
Varying melting points
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24. Covalent molecular crystals
Covalently bonded molecules held together by intermolecular forces
Low melting points
Easily vaporized
Relatively soft
Good insulators
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25. Amorphous solids Noncrystalline solids
The particles are arranged randomly
Glass
Plastics
Can be molded
Fragments have irregular shapes
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26. Amorphous solids
Made by cooling molten substances in a way that prevents crystallization
Also called supercooled fluids
Retain certain fluid characteristics even at temperatures at which they appear to be solid
Can flow over a wide range of temperatures
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27. Discuss Account for each of the following properties of solids:
Definite volume
Relatively high density
Extremely low rate of diffusion
What is the difference between an amorphous solid and a crystalline solid?
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28. Possible changes of state
Melting: solid to liquid
Sublimation: solid to gas
Freezing: liquid to solid
Vaporization: liquid to gas
Condensation: gas to liquid
Deposition: gas to solid
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29. Closed system Matter cannot enter or leave, but energy can
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30. Equilibrium
A dynamic condition in which two opposing changes occur at equal rates in a closed system.
The same number of particles are entering and leaving.
The total number stays the same.
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31. Equilibrium
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32. Phase
Any part of a system that has uniform composition and properties.
Liquid or gas
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33. An Equilibrium equation
When a substance changes state, it either absorbs or gives off energy, usually as heat.
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34. Le Châtelier’s Principle
A system remains in equilibrium until a stress occurs on the system.
Stress: change in concentration, pressure, or temperature
When a system is disturbed by a stress, it attains a new equilibrium position that minimizes the stress.
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35. Shifting equilibrium
Shifts to the right or left, depending on which part of the equation gains concentration.
See table 12-3 on page 375
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36. Equilibrium vapor pressure
The pressure exerted by a vapor in equilibrium with its liquid at a given temperature
Increases as temperature increases
But not directly
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37. Kinetic-molecular theory
Increasing the temperature increases the energy and speed of the liquid particles
This means more particles evaporate, leading to higher vapor pressure
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38. Caution Equilibrium vapor pressure depends only on temperature.
If the system is not in equilibrium, gas laws must be used.
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39. Volatile liquids Evaporate easily
Weak forces of attraction between particles
Ether, acetone
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40. Boiling
Conversion of a liquid to a vapor within the liquid as well as at the surface.
Occurs when the equilibrium vapor pressure equals atmospheric pressure.
All the heat absorbed goes to evaporate the liquid, so the temperature remains constant.
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41. Chemistry chapter 12

42. Cooking
If atmospheric pressure is lower (high altitudes), liquids boil at lower temperatures and food takes longer to cook.
If the pressure is increased (pressure cooker), liquids boil at higher temperatures and food cooks faster.
If the pressure is decreased, it boils at low enough temperatures to avoid scorching milk and sugar. (evaporated and sweetened condensed milk)
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43. Molar heat of vaporization
The amount of energy needed to vaporize one mole of liquid at its boiling point.
(or the amount of energy released when one mole of vapor condenses)
A measure of the attraction between particles.
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44. Normal freezing point
Temperature at which the solid and liquid are in equilibrium at 1 atm pressure.
When a liquid freezes, energy is lost and order is gained.
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45. Clarification Boiling point is the same as condensation point.
Freezing point is the same as melting point.
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46. Molar heat of fusion
The amount of heat energy required to melt one mole of solid at its melting point
(or the amount of energy released when one mole of a liquid freezes)
Depends on the attraction between particles.
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47. Phase diagram
A graph of pressure versus temperature that shows the conditions under which the phases of a substance exist
Reveals how the states of a system change with changes in temperature or pressure
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48. Water’s phase diagram
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49. Curves on diagram
AB shows where solid and vapor can exist at equilibrium
AC shows liquid and vapor at equilibrium
AD shows solid and liquid at equilibrium
Usually a positive slope, but water is different
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50. Triple point
Point A
Shows the temperature and pressure at which solid, liquid, and vapor can exist in equilibrium
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51. Critical temperature
Temperature above which the substance cannot exist in the liquid state.
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52. Critical pressure
The lowest pressure at which the substance can exist as a liquid at the critical temperature
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53. Critical point
Point C
Where critical pressure meets critical temperature
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54. Water’s phase diagram
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55. Carbon Dioxide’s phase diagram
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56. Discuss
Section review on page 382
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57. Water Bent molecule Hydrogen bonding in liquids and solids
Bond angle = 105°
Hydrogen bonding in liquids and solids
Usually 4 – 8 molecules per group in liquid water
Without them, water would be a gas at room temperature
Ice has hexagonal arrangement
Empty spaces lead to low density
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58. Water Has highest density at 3.98 °C
When it melts, molecules can crowd together
When it gets hotter, increased kinetic energy makes them spread apart
Needs a lot of energy to completely break hydrogen bonds and vaporize
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59. Water
Pure liquid water is transparent, odorless, tasteless, and almost colorless
Odors or tastes are caused by impurities
Molar heat of fusion: kJ/mol
Relatively large
Density of ice: g/cm3
Molar heat of vaporization: kJ/mol
Quite high
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60. Calculating heat energy
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61. Example
Find the mass of liquid water required to absorb 5.23 x 104 kJ of heat energy on boiling.
2.31 x 104 g
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62. You try How much heat energy is absorbed when 16.3 g of ice melts?
5.44 kJ
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63. You try
Calculate the quantity of heat energy released when 783 g of steam condenses.
1.77 x 103 kJ
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64. Specific Heat
The amount of energy required to raise the temperature of one gram of substance by one one kelvin. J/(g∙K)
Used to calculate the energy absorbed or released by a substance during a temperature change
No change of state
Example: warm water  cool water
Chemistry chapter 12