Regents Chemistry Mid-year Exam

Periods 3 & 8 – Wednesday, Jan. 20, 2010

Periods 1 & 6 – Thursday, Jan. 21, 2010

(Bring a scientific calculator)

SchoolIsland Bonus Quiz – due Tuesday, Jan. 19th, 8 AM.

After school review – 2:40 PM, Tuesday, 1/19/08

Study Guide

  1. Matter and Energy
  2. Matter is anything that has mass and volume
  3. Types of matter
  4. Pure substances (one set of properties – can’t be separated by physical means)
  5. Elements
  6. Compounds
  7. Mixtures (two or more sets of properties – can be separated by physical means)
  8. homogeneous – uniform throughout
  9. heterogeneous – not uniform, “chunky”
  10. Differences in properties and structures
  11. Changes
  12. Physical changes – do not change the identity of a substance (example: phase changes)
  13. Chemical changes – form a new substance
  14. Energy – the ability to do work
  15. Kinetic energy – the energy of motion
  16. measured by temperature
  17. Calculations (Heat = mass x sp. heat x change in temp)
  18. Potential energy – the energy of position
  19. energy involved in phase changes (molecules change their position relative to one another
  20. Calculations
  21. melting/freezing – Heat = mass x Heat of Fusion
  22. boiling/condensing – Heat = mass x Heat of Vaporization
  1. Atomic Structure
  2. Development of the Atomic Theory
  3. John Dalton
  4. JJ Thompson
  5. Experiments with cathode ray tubes
  6. Discovered electron
  7. plum pudding/chocolate chip cookie model of the atom
  8. Ernest Rutherford
  9. Gold foil experiment
  10. Discovered the nucleus
  11. Atom is mostly empty space
  12. Niels Bohr
  13. Explained bright line spectrum of hydrogen
  14. electrons are in definite energy levels/orbits in the outer part of the atom.
  15. Electrons can jump from a lower energy level (ground state) to a higher energy level (excited state) with the addition of energy.
  1. Electrons give off energy as photons of light as they fall from higher energy levels to lower ones.
  1. Quantum mechanical/wave mechanical model of the atom
  2. Electrons are located in orbitals rather than orbits.
  3. Orbitals are areas of high probability of locating an electron within the atom.
  1. Definitions
  2. Protons, neutrons, electrons, nucleus
  3. Mass number
  4. Atomic number
  5. Average atomic mass
  6. Isotopes
  7. Calculation of ave. atomic mass given the mass of the isotopes and their relative abundance in nature.
  8. Each radioactive isotope has a unique half-life and decay mode that can be found in Table N.
  1. Periodic Table
  2. History
  3. Developed by Mendeleev based on chemical/physical properties and atomic mass
  4. Modern table is based on chemical/physical properties and atomic number
  5. Parts of the Periodic Table
  6. Periods – rows in the Periodic Table
  7. Groups/Families – columns in the Periodic Table. The elements in the same group share similar chemical and physical properties.
  8. Alkali metals – Group 1
  9. Alkali earth metals – Group 2
  10. Halogens – Group 17
  11. Noble gases – Group 18
  12. Metals – located to the left of the “Mason-Dixon” line – lose electrons in chemical reactions
  13. Non-metals – located to the right of the “Mason-Dixon” line – gain electrons in chemical reactions
  14. Metalloids/Semi-metals – adjacent to the “Mason-Dixon” line – can gain or lose electrons in chemical reactions.
  15. Transition metals – Groups 3 – 12 (tend to form colored compounds)
  16. Rare earth elements
  17. Lanthanide and actinide series
  18. Periodic Table and electron structure
  19. Elements in the same group have the same number of valence electrons (electrons in the outermost energy level).
  20. Elements in the same period have their valence electrons in the same energy level.
  21. Trends in the Periodic Table
  22. Metallic character – decreases as you go across a Period and increases as you go down a Group.
  23. Ionization energy – (energy required to remove the outermost electron) increases as you go across a Period and decreases as you go down a Group.
  24. Atomic radius – decreases as you go across a Period and increases as you go down a Group.
  25. Electronegativity – (the measure of an atoms ability to attract electrons in a bond) increases as you go across a Period and decreases as you go down a Group.
  1. Bonding
  2. Energy is released when bonds form and energy is required to break bonds.
  3. Ionic
  4. Results from the transfer of electrons
  5. Metal atoms lose electrons to form positive ions (cations)
  6. The charge on the metal ion is determined by the number of electrons that are lost.
  7. Group 1 (alkali metals) lose 1 valence electron to form 1+ ions
  8. Group 2 (alkali earth metals) lose 2 valence electrons to form 2+ ions
  9. Transition metals tend to form more than 1 type of ion. The charge on the ion is indicated by a Roman numeral. (Iron (III) is Fe 3+)

  1. Non-metal atoms gain electrons to form negative ions (anions)
  2. Non-metal atoms tend to gain enough electrons to get 8 electrons in their valence shell (Example: Oxygen atoms, with 6 valence electrons, tend to gain 2 electrons and form ions with a charge of 2-)
  3. Positive and negative ions are held together by the electric force.
  1. Formula Writing
  2. The sum of the positive and negative charges on the ions in a compound must always add up to zero.
  3. Negative ions
  4. Monatomic – end in ide
  5. Polyatomic – end in ate or ite (Table E)
  6. Do NOT use prefixes.
  1. Covalent
  2. Usually occur between two non-metallic elements
  3. Result from the sharing of valence electrons.
  4. Result in the formation of molecules, rather than ions
  5. Usually follow octet rule (bond to get 8 electrons in outermost shell)
  6. Formula writing
  7. Utilize prefixes to indicate the number of atoms in a molecule (mono, di, tri, tetra )
  8. Lewis Diagrams (Electron Dot Diagrams)
  9. Draw diagrams for neutral atoms, ions, ionic compounds, covalent compounds.

List of Required Calculations on next page.
Required calculations/measurement concepts

  1. Significant figures
  2. Determining the number of sig figs in a measurement (Atlantic-Pacific Rule)
  3. Read a measurement device to the correct number of sig figs.
  4. Using sig figs in calculations
  5. Density = mass/volume
  6. Conversions
  7. liters to milliliters, milliliters to liters
  8. meters to centimeters, centimeters to meters
  9. Energy
  10. Change in temperature ( Heat = mass x sp. heat x change in temp)
  11. Change in phase
  12. melting/freezing – Heat = mass x Heat of Fusion
  13. boiling/condensing – Heat = mass x Heat of Vaporization
  14. Average atomic mass - Calculation of ave. atomic mass given the mass of the isotopes and their relative abundance in nature.
  15. Formula mass – the sum of the atomic masses of the atoms that make up the formula (example: the formula mass of H2O is 18 amu)
  16. Moles
  1. Determine molar mass
  2. Convert grams to moles
  3. Convert moles to grams
  1. Percent composition