This section addresses, in whole or in part, the following Georgia GPS standard(s):
This section addresses, in whole or in part, the following Benchmarks for Scientific Literacy:
This section addresses, in whole or in part, the following National Science Education Standards:
Phases of Matter
Melting point - the temperature at which solid and liquid phases can coexist with one another. The temperature of the freezing point is the same as the temperature of the melting point. Think about water. Ice melts at 0oC (32oF). Liquid water also freezes at 0oC (32oF).
Boiling point - The temperature at which a liquid becomes a gas. The molecules are heated and begin moving so fast that they escape out of the liquid state and become a gas. Liquid and vapor (or gas) phases are at equilibrium at the boiling point.
Different materials have different boiling points. The boiling point of water is 100oC (or 212 oF).
Evaporation - The change from a liquid to a gas. Molecules escape from the liquid and become a gas. As a result, the temperature drops. Think about when you step out of the shower and begin to feel cold. This is because water molecules are evaporating from your skin. It is the highest energy water molecules that escape from the liquid to become a gas. Cooler, lower energy molecules are left behind.
Condensation - The change from gas to liquid. The molecules in water vapor collect together to form small droplets. Clouds are large groups of water droplets. When the droplets become large enough, they will fall to Earth as rain. You can also see condensation on the outside a glass of ice water. The ice water cools the surrounding air. Water molecules in the air that are in a gaseous state are cooled (and slowed down) when they contact the outside of the glass, and they become liquid.
Sublimation - The transition from solid to gas, such as a transition from ice to water vapor. You may have noticed that old ice cubes in your freezer begin to shrink. They have not melted. The water has changed to a gas and escaped.
Density is a property that describes the relationship between mass and volume. In general terms, it is how heavy something feels for its size.
To determine density, you measure the mass of a solid object, and the volume of that object. Then you divide the mass by the volume.
Density = mass / volume
Density gives information about how tightly the atoms or molecules of a particular material are packed. The closer together the atoms are, the more dense the substance.
As a material is heated, the atoms begin to vibrate more, the volume expands,
and the density decreases.
Solids expand when heated.
Solids contract when they are cooled. The atoms vibrate less as
they become cooler.
Most materials are more dense in the solid phase than in the liquid phase. Water is a notable exception. Water expands when it freezes because of its crystal structure. Ice crystals contain a large amount of empty space. The atoms are farther apart than they are in liquid water. As a result, ice is less dense than water. So ice cubes float in your glass of water. And ice forms on the surface of a lake, rather than sinking to the bottom.
The Liquid State - Properties of Liquids
A. Surface Tension
2. Molecules on the outer surface are attracted inward toward the other water molecules due to intermolecular forces
3. In order to increase the surface area of a liquid, these intermolecular forces must be overcome since the molecules must be moved away from each other.
4. This causes water molecules to resist having their surface area increased; this resistance to increased surface area is called surface tension.
5. The stronger the intermolecular forces, the higher the surface tension
Capillary action is the spontaneous rising of a liquid in a narrow tube
b) cohesive forces
- attraction of liquid molecules for each other
b) meniscus is convex if cohesive forces are greater than the adhesive forces (e.g., Hg in a glass tube, since nonpolar Hg atoms are attracted to each other more than to polar glass)
polar liquid in polar
tube nonpolar liquid in polar tube
(adhesive > cohesive) (cohesive > adhesive)
2. examples: syrup has higher viscosity than water; glass is actually a liquid even at room temperature but is extremely viscous
2. some examples of types of unit cells:
2. ionic crystalline solid: e.g., NaCl (simple cubic unit cell)
3. molecular crystalline solid: e.g., H2O (ice)
Forces between 2 or more molecules; weaker than covalent or ionic bonding forces; affect changes of state and so affect properties like melting point and boiling point of a substance.
- very weak intermolecular forces; only significant in nonpolar molecules; all molecules should have these forces, but they are so weak compared to other forces that they are only significant in molecules that have no other intermolecular forces.
2. This temporary dipole moment induces a nearby molecule to have a dipole moment also. Then there is an attraction between the molecules for that instant that they both have dipole moments.
3. This is very brief, so London Dispersion forces are very weak
4. Noticeable in molecules like H2, CH4, CO2, and for noble gas atoms, all of which are nonpolar
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Content provided by Dr. Margaret Venable and Dr. Michael Denniston
Page created by Pamela J.W. Gore
Georgia Perimeter College,
Page created December 20, 2006
Modified June 24, 2007