Introduction to the Particulate Nature of Matter

Elements, Compounds, and Mixtures

    Elements:

  • Described as the building blocks of matter (cannot be further broken down chemically)
  • Consist of only one type of atom
  • Can exist as individual atoms, or as atoms of the same element bonded together
    • E.g. Two hydrogen atoms are often found bonded together in the molecule H2

    Compounds:

  • Composed of two or more different elements chemically bonded in fixed ratios
  • The atoms in a compound cannot be separated using physical methods
  • Examples:
    • Water (H2O) is a compound with two hydrogen atoms and one oxygen atom per molecule
    • Table salt (NaCl) is a compound with equal amounts of sodium and chlorine atoms

    Mixtures:

  • Unlike elements and compounds, they are not classified as pure substances
  • Composed of two or more elements or compounds in varying ratios, which can be separated by physical methods (not chemically bonded)
  • The components of a mixture keep their individual properties, unlike the elements in a compound
  • Types:
    • Homogeneous mixtures, which have a uniform composition where the components are equally distributed and are in the same state
    • Heterogeneous mixtures, which have non-uniform compositions with visible phase differences or boundaries between components

Separation Techniques

    Filtration:

  • The separation of an insoluble solid from a liquid or solution
  • A filtrate is the substance needed to be separated, which passes through a filter
  • Residues are the insoluble components of the filtrate, which are collected in the filter
  • Example: using a paper filter to separate sand from a mixture of sand and water

    Evaporation:

  • The separation of a mixture which has a solute dissolved in a solvent (usually water)
  • Solutions (solute + solvent) are heated in an evaporating dish, where the solvent evaporates and the solute is left behind
  • Example: heating up salt water so the water evaporates, leaving just the separated salt

    Solvation:

  • The separation of a mixture containing two solids, using the difference in their solubilities
  • A solvent (often water) will dissolve the soluble molecules into a solution, leaving behind the insoluble component
  • Other techniques can be used afterwards to separate the components, such as filtration and evaporation
  • Example: using water to separate a mixture of sugar and sand, where the sugar will dissolve but not the sand

    Distillation:

  • The separation of a liquid mixture, using the difference in boiling points (or volatility) of the components
  • Mixtures are heated to a certain temperature, where the component with a lower boiling point (more volatile) will evaporate and be later cooled/condensed in a different container
  • Example: ethanol boils at 78°C, so continuously heating a mixture of water and ethanol will result in ethanol evaporating first, leaving behind the pure water

    Paper Chromatography:

  • The separation of a mixture of solutes in a solvent, using the solvation differences of the components
  • In the mobile phase, the mixture to be separated is dissolved in a certain solvent to create a solution
  • In the stationary phase, a piece of chromatography paper is placed in the solution, and the dissolved components travel up the paper at different speeds (dictated by their affinity for the stationary phase)
  • Example: an ink can be separated using paper chromatography, where the various pigments of the ink will travel different distances up the paper

    Recrystallization:

  • The separation of impurities in a solid, using the solubility differences found at different temperatures
  • Impure mixtures are first completely dissolved in a hot solvent, allowing for the filtration of insoluble components
  • The new solution then gets cooled (decreases solubility), where the desired substance crystallizes and leaves behind impurities to be filtered again – thus purifying the solid
  • Example: rock candy can be made by dissolving sugar in hot water, then cooling the solution until the pure sugar crystallizes (leaving out impurities)

States of Matter and Changes of State

    The Kinetic Molecular Theory (KMT):

  • All matter is made up of small particles
  • These particles all have kinetic energy (the energy of motion) which causes them to constantly move
  • The amount of kinetic energy is proportional to the temperature of the substance, where particles at lower temperatures have lesser motion and tend to vibrate instead of move in straight lines
  • Collisions between particles are elastic, which means there is no loss in kinetic energy

    Solids:

  • Particles are held in certain positions, but can vibrate around fixed points
  • Cannot be compressed, as the particles are already very close together with little space between them
  • Strong forces of attraction hold the particles together
  • Have a fixed shape and a fixed volume, so they cannot flow

    Liquids:

  • Particles are spread out, but are still attracted to each other
  • Cannot be compressed, as the particles are still close enough together
  • Medium forces of attraction between the particles, which are weaker than those in a solid
  • Do not have a fixed shape, and tend to take the same shape as the bottom of their container

    Gases:

  • Particles are very spread out, and move randomly in straight lines
  • Can be compressed, as the particles are far apart with lots of space between them
  • Weak forces of attraction between the particles (weakest among the three states)
  • Do not have a fixed shape nor a fixed volume, where they take the same shape as their container, and their volume depends on the temperature and pressure of the gas

    Density:

  • The mass of a substance per unit volume
  • Calculated by: density = mass / volume
  • General order of densities for different states of matter:
    • Solids > Liquids > Gases

    State/Phase Change:

  • When a substance changes from one physical state to another
  • Classified as a physical change, since the substance does not undergo any chemical changes
  • Changes absorbing heat:
    • Melting: solid → liquid
    • Evaporation: liquid → gas
    • Sublimation: solid → gas
  • Changes releasing heat:
    • Freezing: liquid → solid
    • Condensation: gas → liquid
    • Deposition: gas → solid

Temperature and Energy

    Temperature Scales

  • The Celsius scale is based on the freezing point of water (0°C) and the boiling point of water (100°C)
  • The Kelvin scale is based on absolute zero (0K), which is the lowest possible temperature where particles have zero kinetic energy
  • The two scales are related by: T(K) = T(°C) + 273.15 or T(°C) = T(K) - 273.15
    • Thus, a change of 1°C is equal to a change of 1K
    • Absolute zero (0K) is the same as -273.15°C

    Heating and Cooling Curves:

  • Graphs of temperature vs thermal energy input, which show how the state of matter changes as heat is added/removed
  • Heating curves start at the solid state and end at the gaseous state, while cooling curves go the other way
  • Generally the temperature increases/decreases when heat is added/removed, however during state changes the temperature stays constant until the entire substance transforms (graphed as a flat line)
    • This is because heat is used to overcome intermolecular forces during heating, and the opposite during cooling