1. CHEMISTRY REVIEW
The following is a BRIEF overview of some important topics from General Chemistry I (CHEM 161) at HCC.
In CHEM 161 here, we cover Chapters 1 – 12.
General Chemistry II (CHEM 162) assumes that you have a mastery of this material.
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2. CHEMISTRY REVIEW
Significant figures – you are expected to use them in calculations.
Most answers typically rounded to two or three places.
In lab work, it is often critical that you round answers properly.
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3. CHEMISTRY REVIEW
Dimensional Analysis – a system using units along with numbers. When used properly in calculations, all units cancel except desired units.
Example – convert 2.5 pints to Liters
2 pt = 1 qt, 1 qt = 946 mL, 1 L = 1000mL
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4. CHEMISTRY REVIEW Density problems D = m / V
Ex) A concrete block has dimensions of 12.0” x 12.0” x 4.0”. If the density of concrete is 3.0 g/cm3, then what is the mass of the block?
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5. CHEMISTRY REVIEW Periodic Table Groups Periods
Main Group elements and Transition Group elements
Metals and Non-metals
Periodic Trends for radii, ionization energy, and electron affinity
Electron Configurations
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6. CHEMISTRY REVIEW
Nomenclature of Inorganic compounds depends on type of compound – Molecular or Ionic
Molecular = two non-metals or metalloid with non-metal.
Ionic = metal with a non-metal, but metal may have a fixed or variable charge.
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7. CHEMISTRY REVIEW Naming Molecular Compounds
Name first element in formula first followed by second element name, but change to –ide. Then, use prefix to indicate how many of each (subscript #).
Ex) CO2 = carbon dioxide
Ex) PCl3 = phosphorus trichloride
Ex) N2O5 = dinitrogen pentoxide
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8. CHEMISTRY REVIEW Naming Ionic Compounds
Name metal first followed by name of non-metal, but change to –ide.
If metal has a fixed charge, then no further action is needed.
If metal has a variable charge, then put charge after metal in Roman Numerals.
Ex) NaBr = sodium bromide
Ex) FeCl3 = iron(III) chloride
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9. CHEMISTRY REVIEW
Polyatomic Ions – a group of elements bound together with a net charge.
These NEVER change names when using in nomenclature.
Ex) Na2SO4 = sodium sulfate
Ex) Ni(NO3)2 = nickel(II) nitrate
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10. CHEMISTRY REVIEW Ionic Formulas
Need to remember how these are written – subscripts balance charge.
Need to be able to split apart into correct number of ions.
Ex) CaCl2 = Ca Cl-1
Ex) Na2SO4 = 2 Na SO4-2
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11. CHEMISTRY REVIEW Balancing chemical equations – by inspection.
Ex) __H3PO3  __PH3 + __H3PO4
Ex) __C3H8 + __O2  __CO2 + __H2O
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12. CHEMISTRY REVIEW Basic Stoichiometry
1 mole of anything = Formula or Molecular Weight
Ex) 1 mole of Na2SO4 = 2(23.0g) + 1(32.1g) + 4(16.0g) = 142.1g
1 mole of anything = 6.02 x 1023 atoms or molecules or formula units
Ex) 1 mole of CH4 = 6.02 x 1023 molecules of CH4 = 4 x (6.02 x 1023 atoms of H)
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13. CHEMISTRY REVIEW
Coefficients in balanced chemical equations represent mole amounts.
Ex) N H2  2 NH mol of N2 plus 3 mol H2 yields 2 mol NH3
Allows for the conversion of mass of one into mass of another.
Ex) 2.50g of N2 = ?g of NH3
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14. CHEMISTRY REVIEW Solutions – concentration measured in Molarity (M)
Molarity = Moles of solute / Liters of Solution
Dilution problems – will use this a lot!
M1V1 = M2V2
5.0 mL of 0.10M NaOH and 15.0 mL of water are added together. What is the new molarity of the NaOH?
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15. CHEMISTRY REVIEW
Titration problems – allow us to calculate a specific quantity of an unknown solution.
Very first lab will involve titration – and many others will use as part of the lab.
Ex) What mass of Na2CO3 is present in 100.0mL of a solution if it required 28.5mL of a 0.105M HCl solution to reach the endpoint? The balanced reaction is:
Na2CO HCl  H2O + CO NaCl
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16. CHEMISTRY REVIEW Acids – produce H+ in solution.
Formula usually starts with an “H”
Strong acids – completely ionize in water
HCl, HBr, HI, HNO3, H2SO4, and HClO4
Weak Acids – partially ionize in water
HF, HC2H3O2, etc.
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17. CHEMISTRY REVIEW Bases – produce OH- in solution
Strong bases – completely ionize in water
Group 1A and 2A hydroxides – NaOH, KOH, etc.
Weak bases – partially ionize in water
Nitrogen containing compounds – NH3, CH3NH2, etc.
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18. CHEMISTRY REVIEW Electron Configurations of the Elements
Four types of orbitals – s, p, d, and f
s = 1 orbital, p = 3 orbitals, d = 5 orbitals, and f = 7 orbitals
Shorthand configuration of
C: [He] 2s2 2p2
S: [Ne] 3s2 3p4
Fe: [Ar] 4s2 3d6
Number of unpaired electrons for atoms can be found using an orbital diagram.
For ions, gain/lose electrons according to charge – for transition metals, lose s electrons first.
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19. Math Stuff
Many equations involve the use of either the natural log or log functions.
ln x or log x
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20. Math Stuff
Some problems in chapter 14 will require use of the quadratic equation.
Need to rearrange into:
ax2 + bx + c = 0
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21. Math Stuff
Example problem
16 = (2x)2 / [(0.5-x)(0.1-x)]
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