(DA) Nov. 22, 2010: Volume Conversion

• At a specific pressure and temperature one mole of any gas occupies the same volume
• At 0 Celsius and 101.3 kPa = 22.4 Litres.
• This temperature and pressure is called STP [(S)tandard (T)emperature and (P)ressure]
• 22.4 L mol is the molar volume of STP
  
  Example:

1. A container with a volume of 893 L contains how many moles of air is STP?
   893L x 1 mol = 39.9 mol
               22.4L

            Breakdown:
What you want to do is put write down 893L:

• 893L

After, figure out what the question asks for. If it asks for a mol, you need to cancel out the 
information (units) that you already have by multiplying your information with 1 mol divided by STP.

• 893L x 1 mol =
              22.4L
  After this, divide 893 by 22.4

• 893L x 1 mol = 39.8660714
             22.4L

This is not your final answer. You need to watch out for your significant digits.

• 893L x 1 mol = 40.0
              22.4L

fun video! (not a class summary) just extra**

(NR) Nov. 18, 2010: Molar Mass (Mass of Atom)

- the mass (in grams) of 1 mole of a substance is called the molar mass.
- it can be determined from the atomic mass on the periodic table.
- measured in g/mol

  
Molar Mass of Compounds 

*REMEMBER SIGNIFICANT DIGITS

Element                           work                             Molar Mass

H2O                                       2(1.0) + 16.0                                        18.0 g/mol
NO2                                       14 + 2(16.0)                                         46.0 g/mol
NaCl                                       23.0 + 35.5                                          58.5 g/mol
FeO                                         55.8 + 16.0                                         71.8 g/mol
NaNO3                              23.0 + 14 + 3(16.0)                                  85.0 g/mol



Mole Conversions (converting between grams and moles)

- To convert between moles and mass we use molar mass as the conversion factor.
- be sure to cancel the appropriate units.  

(TG) Nov 16: AVOGDRO'S NUMBER; how we count atoms

Avogadro's Number

6.02 x 10^23

• atoms and molechules are extremely small
• macroscopic objects contain too many atoms to count or weigh individually
• Amedeo Avogadro proposed that the number of atoms in 12.000000 grams of Carbon be equal to a constant (one mole of a Carbon)
•  So what is Avogadro's number? Well its 6, 020, 000, 000, 000, 000, 000, 000, 000
• 1 mole = 6.02 x 10^23 atoms
• one mole is simply a multiply of things for example:

1.  pair = 2
2. dozen = 12
3. century = 100
4. mole = 6.02 x 10^23 

The Mole in Perspective

6.02 x 10^23

• So how big is a mole? (in perspective)
• 1 mole of meters would cross the entire galaxy over 3000 times
• 1 mole of smarties would cover 250 planets similar to the size of earth a kilometer deep!
• 1 mole of seconds is 100,000 times greater than the age of the universe
• 1 mole of blood cells more than every human on the face of the earth

A mole is also used to measure the smallest unit of a quantity. For example there is a such thing as a mole of NaCl ions. You do not thing of Na and Cl as separately, but rather as ONE UNIT. 

EXAMPLES:

A sample of  carbon contains 2.4 x 10^25 atoms. How many moles is this?

2.4 x 10^25     x       1 mole

                                    6.02 x 10^23               = 39.9 moles

THINGS TO REMEMBER:

significant figures

the units you want to cancel are always opposite each other (ie. top and bottom)

(NR) Nov. 5, 2010: Hydrate Lab

Yesterday, we did a hydrate lab. It was our first class lab and it was fun! like OMG!!!


Hydrates are ionic compounds that contain an inorganic salt compound loosely bound to water. The purpose of this experiment is to determine the empirical formula of a hydrate. In the lab we determined the anhydrous (without water) mass of the hydrate. We compared it with the actual mass of the water that should be presented.

the materials we used were:
- Bunsen burner***
- test tube
- test tube rack
-test tube clamp
-weight scale
- Cobaltous chloride hexahydrate

***REMEMBER: BE AWARE OF BUNSEN BURNERS! You can't see the hot light blue flame... and if you accidentally touch it... PEACE TO YOU!

note for Mr. Doktor:
We should do more outdoor experiments with chemicals... We students want to see something big explode! lol

(DA) Nov. 3, 2010: Naming Compounds

Chemical Nomenclature

• Today the most common system IUPAC for most chemicals
  - Ions
  - Binary Ionic
  - Polyatomic ions
  - Molecular Compounds
  - Hydrates
  - Acids / Bases

Chemical Formulas

Be aware of the differences between ion and compound formulas

 - Zn^2+ <------------------------- Ion Charge

 - BaCl2 <------------------------ Number of Ions

Multivalent Ions

• Some elements can form more than one ion
  eg. Iron -> Fe^3+ or Fe^2+
        Copper -> Cu^2+ or Cu^1+
• IUPAC uses Roman Numerals in parenthesis to show the charge
• Classical systems use latin names of elements ans suffixes
  - ic (larger charge) and -ous (smaller charge)
  
  Example:
  - Ferric Oxide -------------> Iron (Fe)
   [ -ic refers to larger charge
     -ous refers to smaller charge]

• Ferr - Iron
• Cupp - Copper
• Mercur - Mercury
• Stann - Tin
• Aunn - Gold
• Plumb - Lead

Complex Ions

• Complex ions are larger groups of atoms that stay together during a chemical reactions
• Almost all are anions
• Write the metal name and the polyatomic ion

Hydrates

• Some compounds can form latices that bound to water molecules
  - Copper Sulphate
  - Sodium Sulfate
• These crystals contain water inside them which can be released by heating.

To name hydrates

1. Write the name of the chemical formula
2. Add a prefix indicating the number of water molecules (mono=1, di=2, tri=3 etc.)
3. Add hydrate after the prefix

ie. CuSO₄·5H₂O          Copper (II) Sulphate Penta Hydrate

      LiClO₄·3H₂O          Lithium Perchlorate Tri Hydrate

Naming Acids And Bases

• Hydrogen Compounds are acids
  - HCl ---> Hydrochloric Acid
  - H₂SO₄ --> Sulfuric Acid
• Hydrogen appears first in the formula unless it is part of a polyatomic group
  - CH₃COOH --> Acetic Acid

(TG) Nov: ELECTRONIC STRUCTURE

ELECTRONIC STRUCTURE

ELECTRON DOT DIAGRAMS: 

• the nucleus is represented by the atomic symbol
• for individual elements determine the number of valence electrons
• electrons are represented by dots around the symbol
• four orbitals (one on each side of the nucleus) ea holding a maximum of 2 electrons
• Each orbital gets one electron before they begin to pair up

examples:

CARBON

LEWIS DIAGRAMS FOR COMPOUNDS AND IONS

• In covalent compounds electrons are shared

1. Determine the number of valence electrons for each atom
2. Place atoms do the valence electron are shared to fill each orbital

example:

NF(little3)

DOUBLE AND TRIPLE BONDS:

• Sometimes the only way covalent compounds can fit all their valence levels is if they share more than one electron  

example: 

CO(little2)

IONIC COMPOUNDS

• in ionic compounds electron transfer from one element to another
• determine the number of valence electrons on the cation. Move these to the anion
• Draw [ ] around the metal and the nonmetal
• Write the charges outside the brackets

example: