Solids, Liquids and Gases

Georgia Perimeter College

Objectives

  1. Compare and contrast solids, liquids, and gases.
  2. Compare and contrast the movement of particles in solids, liquids, and gases.
  3. Compare properties such as surface tension, boiling point, freezing point/melting point, viscosity, vapor pressure and capillary action.

This section addresses, in whole or in part, the following Georgia GPS standard(s):
  • S8P1 Students will examine the scientific view of the nature of matter.
    c. Describe the movement of particles in solids, liquids, gases, and plasma states.
    d. Distinguish between physical and chemical properties of matter as physical (i.e., density, melting point, boiling point) or chemical (i.e., reactivity, combustibility).
     

This section addresses, in whole or in part, the following Benchmarks for Scientific Literacy:
  • Atoms and molecules are perpetually in motion. Increased temperature means greater average energy of motion, so most substances expand when heated. In solids, the atoms are closely locked in position and can only vibrate. In liquids, the atoms or molecules have higher energy, are more loosely connected, and can slide past one another; some molecules may get enough energy to escape into a gas. In gases, the atoms or molecules have still more energy and are free of one another except during occasional collisions.

This section addresses, in whole or in part, the following National Science Education Standards:
  • A substance has characteristic properties, such as density, a boiling point, and solubility, all of which are independent of the amount of the sample. A mixture of substances often can be separated into the original substances using one or more of the characteristic properties.

Phases of Matter

There are three phases (or states) of matter.  There are actually four phases, if you count plasma.
  1. Solids have a fixed shape and volume.  Their molecules are locked in place.  The molecules constantly vibrate, but they can't move or switch places with other molecules.  As a result, a solid retains its shape and size.
     
  2. Liquids - The molecules in a liquid move faster than those in a solid.  They can slip out of position and move over and around one another.  Liquids flow because the molecules can move.  The liquid has a definite volume, but no definite physical shape.  It conforms to the shape of whatever container it is poured into.
     
  3. Gases - In a gas, the molecules are widely spaced and separated from one another.  They can move freely  and randomly.  Gas does not have a definite size or shape.
     
  4. Plasma - A gas which has been heated to the point where it has been ionized.  This means that one or more of the electrons have been stripped off the gas molecule.  For most materials, stripping of electrons requires temperatures in excess of 10,000oC.
    The ionized particles  of plasma interact by long-range electromagnetic forces associated with their charges and motion. 
    We see plasmas in lightning, the aurora borealis (northern lights), and in neon and fluorescent lights.  (In the case of neon and fluorescent lights, they are not heated; instead electrical current is passed through gases, which strips off the electrons.) 
    The Sun and other stars are made of plasma.  Planets with magnetic fields (like Earth) are surrounded by plasmas, consisting of particles of the solar wind interacting with particles in the Earth's upper atmosphere

     


Phase Changes

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

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
 

1. Water drops are spherical because of intermolecular forces of attraction between water molecules (e.g., water "beads" on waxed car)

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

 

B. Capillary Action

Capillary action is the spontaneous rising of a liquid in a narrow tube

1. Capillary action is due to:
 
a) adhesive forces - attraction of liquid for molecules of the tube if the tube has bonds of the same polarity as the liquid (i.e., nonpolar liquids are attracted to nonpolar tubes; polar liquids attracted to polar tubes)

b) cohesive forces - attraction of liquid molecules for each other
 

2. meniscus is related to capillary action of liquids
 
a) meniscus is concave if adhesive forces are greater than the cohesive forces (e.g., polar water molecules in polar glass tube); water rises in glass tube higher than water outside tube due to capillary action

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)
 

C. Viscosity - a measure of a liquid's resistance to flow
 
1. liquids with large intermolecular forces have high viscosity

2. examples: syrup has higher viscosity than water; glass is actually a liquid even at room temperature but is extremely viscous

 

Crystalline Solids
 
A. Lattice: the positions of atoms/ions in a crystalline solid represented in a 3-dimensional picture
 
1. unit cell: the smallest repeating pattern of the lattice

2. some examples of types of unit cells:

a) simple cubic
b) body-centered cubic
c) face centered cubic
 

 

B. Examples of crystalline solids
 
1. atomic crystalline solid: e.g., diamond structure and graphite structure -- both are crystalline structures of carbon atoms; diamond and graphite are both pure carbon but they have very different properties due to their different crystal structures

2. ionic crystalline solid: e.g., NaCl (simple cubic unit cell)

3. molecular crystalline solid: e.g., H2O (ice)

 


Optional content - beyond the standards

Intermolecular Forces

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.

 

A. Dipole-Dipole Forces
  1. molecules with polar bonds exhibit dipole moments
  2. polar molecules exhibit a net dipole moment (recall: not all molecules with polar bonds are polar)
  3. oppositely charged ends of molecules with dipole moments are attracted to each other
  4. Hydrogen bonding - a special kind of dipole-dipole attraction; involves molecules with a H atom bonded to a very electronegative element (N, O or F); hydrogen bonding is a stronger intermolecular force than the dipole-dipole forces of attraction (not true bonds)

  5. Dipole-dipole forces are more noticeable in the solid phase than in the liquid phase and are much more significant in the liquid phase than in the gas phase. This is because molecules are closer in solid and liquid phases than in the gas phase.
  6. Hydrogen-bonding causes boiling points of molecules such as H2O, NH3 and HF to be much higher than expected; it takes more energy than expected to move the molecules apart to go into the vapor phase from the liquid phase due to the extra intermolecular forces of H-bonding.
B. London Dispersion Forces

 - 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.

 

1. Instantaneously (only for an instant), electrons in a nonpolar molecule might be arranged so that there is temporarily a dipole moment

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,
Clarkston, GA

Page created December 20, 2006
Modified June 24, 2007