Chapter 10: States of Matter

Content Outline: Kinetic-Molecular Theory of Matter

  1. The States of Matter Present in the Universe.
  1. There exist 4 known “states” of matter:
  1. Solid – such as ice (frozen water) or aluminum cans.
  2. Liquids – such as water or Iodine.
  3. Gas – such as clouds (water vapor…”vapor” means “gas”) or Helium.
  4. Plasma – such as found in the composition of stars and suns.
  1. Kinetic Theory of Matter
  1. This theory states that the “states” of matter are defined by the kinetic (moving) ability of the composing atoms. All atoms/molecules of all substances move to some degree at the molecular level.
  1. Colder or heavier atoms/molecules move slower.
  2. Warmer or lighter atoms/molecules move faster.
  3. So all movement (Kinetic Energy) is directly related to temperature and mass.
  1. The history of the Kinetic Theory of Matter.
  1. This theory was initially proposed by a Roman philosopher named Lucretius. (50 B.C.)
  1. He stated, “That all macroscopic bodies were composed on a small scale of rapidly moving atoms”.
  1. Robert Boyle (1665) proposes that all atoms/molecules stop moving completely at “primum frigidum”.
  1. Primum Frigidum becomes known as Absolute Zero.
  1. Absolute Zero is defined as the temperature at which all kinetic movement ceases within atoms/molecules.
  1. Daniel Bernoulli (1738) expands upon this theory in his bookHydrodynamica.
  1. He states that all atoms/molecules move, but adds the average movement (kinetic) of the atoms/molecules is directly related to heat.
  1. More movement = more heat; less movement = less heat.
  1. Johann Heinrich Lambert proposes that Absolute zero has a temperature of -273.15° C or -460° F.
  2. Lord Kelvin (a.k.a William Thomson) proposes the Kelvin scale for temperature measurement in his paper, On an Absolute Thermometric Scale in 1848.

a.Absolute zero is defined as 0 K.

b.Kelvins can be found by taking the ° C and adding 273 to it.

c.It is written as “K” and not “° K”.

  1. Defining a “state” of matter.
  1. In defining a “state” of matter the following physical properties must be considered:
  1. Expansion – can the atoms/molecules “expand” their given “space”.
  2. Fluidity – this refers to the attractive forces between atoms/molecules and the ease of movement past one another.
  3. Density – this refers to how close the atoms/molecules are to each other.
  1. Heavier atoms/molecules  move slower at lower temperatures, such as room temp.
  2. Lighter atoms/molecules  move faster at lower temperatures, such as room temp.
  1. Compressibility – can the atoms/molecules be moved closer to each other?
  2. Diffusion – the ability of atoms/molecules to ‘mix’ together over time.
  1. Solids mix at an extremely slow rate, if ever, most of the time.
  2. Liquids mix at a fairly quick rate.
  3. Gases mix at a very rapid rate.
  4. Effusion – the ability of a gas to move through a tiny pore/hole in another substance.

Chapter 10: States of Matter

Content Outline: Solids

  1. There are 2 types of solids based upon atom/molecule arrangement within the solid:
  1. Crystalline Solids (Crystals)
  1. These substances possess atoms/molecules that are arranged in a repeating, orderly, geometric pattern.

For example: diamonds, rubies, emeralds, quartz, and salts

  1. The types of crystals based on geometric patterns (a.k.a. lattices) are:
  1. Cubic - Has equal dimensional sides.
  2. Hexagonal - Looks like a 3-D Stop sign, but with six sides.
  3. Orthorhombic - An elongated rectangular rectangle with rectangular, angled ends.
  4. Monoclinic - Small Diamond…almost cubic.
  5. Tetragonal - Elongated rectangle with square ends.
  6. Triclinic - Tilted rectangle with pointed ends.
  7. Trigonal - Diamond shaped cube with pointed ends.
  1. Amorphous Solids (“a” means “without”; “Morph” means “shape”.)
  1. These substances possess atoms/molecules that are randomly arranged.

For example: glass, lava, and plastics

  1. Kinetic Theory of Matter Characteristics:
  1. Expansion (Can atoms/molecules “expand” their given “space”?)
  1. Solids generally cannot expand without a change in energy (heat or cooling).
  2. Solids have definite shapes without a container.
  3. Solids have definite volumes.
  1. Fluidity (The easeof movement past one atom/molecule by another because of attractive forces between atoms/molecules.)
  1. Solids possess atoms/molecules that are not fluid (able to move).
  1. Solids atoms/molecules vibrate in a fixed position.
  2. This inability to “move” is directly related to the strong intermolecular attractive forces between atoms/molecules.
  1. Intermolecular Attractive (Binding) forces of Crystals.
  1. Covalent Network(Examples Diamonds or Quartz –SiO2)

i. There are only very strong covalent bonds between all atoms/molecules.

ii. These are extremely hard substances; but brittle (able to be fractured into smaller

pieces).

iii. Have very highmelting and boiling points.

  • Melting – going from solid to liquid with absorbing energy (heat).
  • Boiling – going from liquid to gas with absorbing energy (heat).
  • Freezing – going from liquid to solid with loss of energy (heat).
  • Condensation – going from gas to liquid with loss of energy (heat).
  1. Covalent Polar Molecule(For example: Ice)

i. Individual molecules possess polar covalent bonds.

ii. The polarity of the individual molecules allows for the attraction between molecules by

London Dispersion Forces (a.k.a. Van der Waals Interactions), Dipole – Dipole

Interactions, or Hydrogen Bonds.

iii. These tend to be very soft (flexible)crystals.

iv. These tend to have low (negative) melting and boiling points.

α. The low melting and boiling points are due to the very weak intermolecular

attractive forces between individual molecules.

  1. Covalent Non-polar Molecules

(Examples: solid Oxygen, solid Hydrogen, solid Methane- CH4)

i. Individual molecules possess non-polar covalent bonds.

ii. Each individual molecule is “held” to other molecules by very weak London Dispersion

Forces (a.k.a. Van der Waals Interactions).

iii. These are extremely soft crystals.

iv. They have extremely low melting and boiling points.

α. The extremely low melting and boiling points are due to the very weak

intermolecular attractive forces between individual molecules.

  1. Ionic Crystals(Example: Salt)

i. These crystals are composed of positively charged ions that are strongly attracted to

negatively charged ions in an alternating pattern.

α. Positive Ions (Cations) – metals from Groups 1 or 2.

β. Negative Ions (Anions) – non-metals from Groups 16 or 17.

ii. They are hard solids, but very brittle.

iii. They have high melting and boiling points.

α. The high melting and boiling points are due to the very strong

intermolecular attractive forces between individual ions.

  1. Metallic Crystals (Examples: Copper wire, aluminum cooking pan, iron pan)

i. These solids are composed of individual metal cationssurrounded by a “sea of

delocalized(flowing) negatively charged electrons.

α. The delocalized (flowing) electrons make these solids great conductors of

electricity and heat. Remember, that electrons possess Kinetic Energy as they are

orbiting the atoms nucleus.

ii. They have very high melting and boiling points, hence why we can cook with them.

α. The high melting and boiling points are due to the very strong

intermolecular attractive forces between individual ions.

  1. Attractive (Binding) Forces of Amorphous Solids.
  1. These solids are held together by a variety of combined intermolecular forces.

i. Due to this combination of attractive forces, these solids do not have defined melting or

boiling points.

ii. This is because they are composed of various different atoms/molecules.

iii. These can sometimes be classified as supercooled liquids.

α. These substances look solid, but can flow.

(Examples: lava, molten steel, molten glass)

  1. Density (How compact/close are the atoms/molecules to each other.)
  1. Solids generally have very high densities due to the atoms/molecules being very compact already due to the intermolecular attractive forces. (See above in B.)
  1. Compressibility (Can the atoms/molecules be moved closer to each other?)
  1. Solids are generally not compressible do to the high density of the atoms/molecules.
  1. There exists very little vacant space between the atoms/molecules.
  1. Diffusion (The ability of atoms/molecules to “mix” together over time.)
  1. Some slight diffusion may occur, but rarely does, and is extremely slow.

Chapter 10: States of Matter

Content Outline: Liquids

  1. Liquid
  1. A “state” of matter where the substance can flow and takes on the shape of its container.
  1. The flow is usually in the direction of the force of gravity or channeling.
  2. The flow is possible because the atoms/molecules are not fixed in position by strong attractive forces, like a solid.
  3. The flow is possible also because the individual atoms/molecules possess more Kinetic Energy, than do solids.
  1. This greater Kinetic Energy per atom/molecules helps them overcome some of the restrictive attractive forces.
  1. Kinetic Theory of Matter Characteristics:
  1. Expansion (Can atoms/molecules “expand” their given “space”?)
  1. Liquids, while they can change shape, by changing the container, cannot expand their volumes.
  2. Liquids have defined volumes, like solids.
  1. Fluidity (The ease of movement past one atom/molecule by another atom/molecule because of attractive forces between atoms/molecules.)
  1. Liquids atoms/molecules have greater ease of motion due to the possession of greater amounts of Kinetic energy per atom/molecule.
  1. The greater Kinetic Energy makes it easier to overcome some of the restrictive (stronger) intermolecular attractions between atoms/molecules within the liquid, but some attractions still exist.
  1. These would be London Dispersion Forces (a.k.a. Van der Waals Interactions), Dipole-Dipole Interactions, and Hydrogen Bonds.
  2. These attractive forces are what give the liquid “order” or “organization”, just like solids.

Example: Water molecules attracted to water molecules by Hydrogen Bonding.

  • Cohesion – molecules binding to like molecules.
  • Adhesion – molecules binding to some other substance.
  1. At the surface of liquids, there exists a boundary between the liquid’s atoms/molecules and the atmosphere’s (air) atoms/molecules that are in a gaseous “state”.
  1. Surface Tension (“Tension” refers to “pulling strength”.)
  1. This is a cohesive attraction that tends to pull adjacent atoms/molecules of a liquids surface together into a more stable position with less surface area.

α. Greater attraction between atoms/molecules = greater tension.

b. Weaker attraction between atoms/molecules = weaker tension.

Remember, these stronger/more numerous attractions increase the boiling points of liquids, just like with solids.

  1. This is the attractive force that creates the concept of a “drop”.

α. Drop – the spherical (round) shape of a liquid in the air.

iii. Evaporation

α. This term refers to the “escaping” of surface atoms/molecules of a liquid to the

atmosphere.

b. The ability to “escape” from the liquid’s surface is because atoms/molecules

absorbed more Kinetic Energy from the surrounding environment.

For example, such as heating up a liquid, or putting a liquid in the sunshine, such as

you do when you sweat on a hot day.

c. This “escaping” of greater Kinetic energy containing atoms/molecules also takes

heat away from the liquid and thus providing a cooling effect known as

Evaporative Cooling.

  • Vaporization - turning a solid or liquid into a gas.
  • Vapor – a gas.
  1. For liquids, there also exists a boundary between the liquid’s atoms/molecules and the solid surface or the container.

For example, water in a glass beaker or plastic cup.

  1. Capillary Action (“Capillary” refers to a “narrow tube”.)
  1. This is an adhesive attraction that exists between a liquid’s atoms/molecules and a solid container’s surface.
  2. This attractive force helps to “pull” liquids up the surface of the container.

For example, seen in the meniscus of a graduated cylinder or water moving up through the xylem tissue (which are dead, hollow plant cells).

  1. The narrower the tube greater the adhesive climb.
  1. Density (How compact/close are the atoms/molecules to each other?)
  1. Most liquids are only slightlyless dense than their solid “state”.
  1. Water is a rare exception to the law. Water is denser than ice; therefore ice floats.
  1. The floating of ice is directly due to the Hydrogen bonds between water molecules, which the bond angle is fixed and maximized, in terms of distance… hence why it is less dense and floats.
  2. Water is most dense at 4OC… about 1g/cm3.
  3. Water gets less dense as it approaches the boiling point (100OC).
  4. Items with greater densities sink in water. For example: Glycerol
  5. Items with lesser densities floaton top of water. For example: Oils and Alcohols
  1. The density of a liquid can affect the boiling points of liquids.
  1. Typically, less dense liquids boil at lower temperatures.
  2. Typically, more dense liquids boil at higher temperatures.
  3. This is all due to the greater attractive forces between individual atoms/molecules within the liquids.
  1. Compressibility (Can the atoms/molecules be moved closer to each other?)
  1. Most liquids are only slightly compressible.
  2. It usually requires extremely massive amounts of pressure to achieve a minimal compression.
  1. Hydraulics
  1. This is the study of fluids under pressure to move or lift objects.
  1. Pressure diffuses through liquids in all directions evenly.
  1. Diffusion (The ability of atoms to “mix” together over time.)
  1. Most liquids can diffuse from into another liquid, but it depends upon 2 dissolvability factors:
  1. Polar versus non-polar – Like dissolves like; they do not “like” the other and keep separate.
  2. Temperature – colder liquids dissolver slower than warmer liquids.

c. Stirring liquids can increase the rate of diffusion.