Chapter 2 Review (Sections 1-8)

  1. Law of Conservation of Mass
    1. For all changes in matter, the mass before and after is the same.  Matter cannot be created nor destroyed.  Priestley and Lavoisier (1774).  This was actually hard to figure out.  After all when you burn something it seems to disappear leaving only bits of ash which weigh less than what you initially had.  Well the other products of fire are gases that drifted away - so you didn't get their mass.  But the mass of burning something is the same before and after. 
    2. Chemical changes can be called chemical reactions.  What you start with are called reactants.  What you end with are called products.  For chemical reactions the reactant mass = product mass.
    3. Law of definite proportions - pure samples always have the same elemental proportions by mass.  Law of multiple proportions (Click on activities then multiple proportions movie) - pure samples always have the same elemental proportions by atoms. For example, the mass of H in water compared to the mass of O in water is always the same - 1 to 8. And water is always two H atoms per one O atom.  Compare to hydrogen peroxide H2O2.  The mass of H in hydrogen peroxide compared to the mass of O in water is always the same - 1 to 16. And hydrogen peroxide is always two H atoms per two O atoms.  Just because water and hydrogen peroxide are made of the same two elements does not mean they are the same or even similar.  Water and hydrogen peroxide are quite different.
  2. Dalton's Atomic Theory (1808)                  
    1. John Dalton's theory is still largely true today  (IMPORTANT)
      1. Elements are made of tiny atoms.
      2. All atoms of the same element have the same mass.
      3. Elements combine in whole number ratios to make molecules and compounds.
      4. The number and type of atoms before a reaction will be the same after the reaction - what changes is how the atoms are combined. (must follow the Law of Conservation of Mass after all  )
    2. Only one of the 4 statements above is now considered false.  Which one?
    3. It is the second one.  All atoms of the same element don't have the same mass because of isotopes.
    4. Look at Example 2.1.  Try Problem 2.1 in your text.
  3. Electrons
    1. Thomson discovered electrons in the 1890's in cathode ray tube experiments. (Click on cathode ray tube) If you connect an evacuated glass tube with two pieces of metal at each end to a battery eventually electrons will pass from the negative metal to the positive metal.  This electron beam must be negatively charged since it goes to the positive side. This is how a TV works. See figure 2.3 in your text. We also know the electron beam is charged because a magnet can deflect the beam from traveling in a straight line. 
    2. Millikan (Click on Milikan oil drop) later found that the charge on an electron must be -1.6 x 10-19 Coulombs and the mass of an electron is 9.11 x 10-28 grams.
  4. Protons and Neutrons
    1. Matter is neutral overall so if electrons in atoms are negative there must be a positive part also.
    2. Rutherford (Click on activities then Rutherford movie) knew that Ra, Po, and Rn all emit alpha particles that have a +2 charge.  So he set up an experiment where alpha particles would pass through gold foil.  He thought that all would pass but to his surprise some alpha particles were deflected.  He decided that an atom was mostly empty space except for a tiny nucleus that held a positive charge and most of the mass. See figure 2.5 in your book. 
    3. If the nucleus of an atom were a marble, the atom would be the size of the Cardinal's stadium. The protons and neutrons live inside the nucleus.  (IMPORTANT)
      1. protons have +1 charge and mass = 1.67 x 10-24 g which is ~ 1 amu (atomic mass unit)
      2. neutrons have 0 charge and mass = 1.67 x 10-24 g which is  ~ 1 amu
      3. electrons have -1 charge and mass = 9.11 x 10-28 g which is  ~ (1 / 2000) amu (negligible mass)
    4. Electrons go around the proton and neutron containing nucleus.  Mass is due to the protons and neutrons.  Size is due to the electrons.  Charge is due to the protons and electrons. Atoms are neutral so the number of protons must equal the electrons. (IMPORTANT)
    5. Example:  Tommy dropped a thermometer in lab.  The mercury made a drop with a diameter of 2.5 cm.  A mercury atom is about 2.98 x 10-10 m in diameter.  How many mercury atoms across is this drop? Try first before looking at the answer!!!
    6. Answer:  2.5 cm ( 1 m / 100 cm) = 0.025m across the drop.  How many atoms will fit?  0.025 m (1 atom / 2.98 x 10-10 m) = 8.4 x 107 Hg atoms. 
    7. Look at example 2.2.  Try problems 2.2 and 2.3 in your text.
  5. Atomic Number  
    1. What makes atoms different from each other?  Well it is their number of protons which we call the atomic number.  The # protons determines the identity of an atom, not electrons, not neutrons, not mass - just number of protons!!!  (IMPORTANT)
    2. Atomic mass - well what has mass in an atom???  Mainly protons and neutrons.  So atomic mass = # protons + neutrons.
    3. Isotopes = same element, so same # protons, but different mass, so different # neutrons. (IMPORTANT)
      1. 1H and 2H and 3H are isotopes of hydrogen.  What is the same?  one proton  What is different?  # neutrons
      2. 12C and 13C and 14C are isotopes of carbon.  What is the same?  6 protons  What is different?  # neutrons
      3. What element has 7 protons?  nitrogen
      4. What element has 10 neutrons?  we don't know, identity of an atom depends on # protons only
      5. What element has 8 electrons?  we don't know, identity of an atom depends on # protons only
      6. If I add one neutron to a carbon atoms what element to I have?   carbon (did I get you?)
      7. Practice problems
        1. How many protons, neutrons and electrons are in:  28Si, 14C and 11B?
        2. Fill in this table for ATOMS:
          Symbol 127Xe   12C   41Ca 1H  
          # protons   38         1
          # neutrons   49   8     1
          # electrons       6      
        3. answers to a above: 28Si - 14p, 14n, 14e, 14C - 6p, 8n, 6e and 11B - 5p, 6n, 5e
        4. answers for the table:
          Symbol 127I 87Sr 12C 14C 41Ca 1H 2H
          # protons 54 38 6 6 20 1 1
          # neutrons 73 49 6 8 21 0 1
          # electrons 54 38 6 6 20 1 1
      8. Look at Example 2.3 and 2.4.  Try Problems 2.4, 2.5, 2.6 in text.
  6. Atomic Mass
    1. The atomic mass and atomic number are given on the periodic table.  The units for mass are amu.  So why aren't the masses whole numbers?  You can't have "part" of a proton or neutron?  Well the masses are averages for all the isotopes.
    2. Carbon is 12.011 amu.  That is mostly carbon-12 plus a little bit of carbon-13. Thus the average is a little higher than 12.
    3. To calculate the atomic mass sum the mass of each isotope times the percentage of that isotope.  Look at Example 2.5.  Try Problems 2.7 and 2.8. 
  7. Compounds and Mixtures
    1. Definitions (IMPORTANT)
      1. Element - sample of one type of atom
      2. Molecule - when 2 or more atoms are bonded together
      3. Compound - when 2 or more different atoms are bonded together
      4. Pure - each species or unit in the sample is the same (every molecule is the same for example)
      5. Mixture - 2 or more different atoms or molecules, opposite of pure
        1. Homogeneous - uniformly mixed, can't see the separate species (salt water, coffee)
        2. Heterogeneous - not uniformly mixed, can see the separate components (granite, dirt)
    2. Questions to try
      1. What is air?  A mixture
      2. What is tap water?  a mixture
      3. What is DI water?  pure, compound,  molecules
      4. What is  O2, oxygen gas?  pure, element, molecule
      5. Look -
      6. Are the pictures above elements, compounds, pure, mixtures, or molecules.  More than one may apply.  (IMPORTANT)
        1. molecules, pure, element
        2. pure, compound, molecules
        3. mixture
        4. pure, element
        5. pure, compound, molecules
        6. mixture
        7. mixture
        8. pure, element
  8. Ions and Covalent Bonds
    1. Chemistry is all about making and breaking bonds in chemical reactions.  Bonding is all about electrons. (IMPORTANT)
    2. For main group elements, bonds are either covalent or ionic
      1. Covalent (IMPORTANT)
        1. electrons are shared between nonmetals and other nonmetals
        2. the nonmetals may be the same element or different elements
        3. molecules are 2 or more atoms held together by covalent bonds (H2O, CO, CO2, H2O2, O2, N2)
        4. Diatomic elements are molecules - these elements occur naturally paired up (H2, N2, O2, F2, Cl2, Br2, and I2) They will not be found in nature as individual atoms. (IMPORTANT)
        5. Look at Key Example 2.3 (typo - should be 2.7 in book).  Try Key Problems 2.10 and 2.11) 
      2. Ionic (IMPORTANT)
        1. electrons are transferred from metal atoms to nonmetal atoms
        2. the metals lose negatively charged electrons so they become positively charged cations
        3. the nonmetals gain negatively charged electrons so they become negatively charged anions
        4. Opposites attract, so the cations and anions hang out  and become ionic solids
        5. There is no real "molecule" here - the ions just keep packing into a nice pattern for the solid
      3. Polyatomic ions - several atoms held together by covalent bonds but they have an overall charge so as a group make an ion and can form ionic bonds with other atoms. More on them later.
      4. Ionic or Covalent?  H2O, CaCl2, KBr, N2, Cu, HCl, H2O2     Answer:   H2O, N2, HCl, H2O2  are covalent.  The rest are ionic except Cu which is neither since it is not bonded!!!
      5. Look at Example 2.8.  Try Problems 2.12 and 2.13.

 

Suggested Problems from the End of Chapter (answers in back of book): 
26, 28, 34, 38, 42, 48, 52, 54, 60, 62, 68, 70