Moles to Volume Conversion (Brian)

-At a specific pressure and temperature one mole of any gas occupies the same volume
-At 0 Celsius and 101.3 Kilo pascals = 22.4 liters
-This temperature and pressure is called STP
-22.4 L/mol is the molar volume at STP

Example:
A certain gas is found to occupy 11.6L at STP. How many moles of gas are there?
11.6L x 1mol = .518 mol
              22.4L

Avogadro's Number (Brian)

-Atoms and molecules are extremely small
-Macroscopic objects contain to many to count or weigh individually
-Amedo Avogadro proposed that the number of atoms in 12.00000g
 of Carbon are equal to a constant (equal to 1 mol of Carbon)
-This value is now called Avogadro's number and forms the basis
 of all quantitative chem

* 1 mol of meters would cross the entire galaxy over 3000 times(1 mol repersents a huge number of   particles)

Nomenclature (Zac)

Naming Compounds
-Today the most common system is IUPAC for most chemicals (eg. Ions, Binary Ionic, Molecular Compounds, Hydrates, Acids/Bases
-Be aware of the differences between ions and compound formulas
Zn2+ =ion charge,  BaCl2 = # of ions
Multivalent Ions
-Some elements can form more than one ion (eg. Iron has a +3 and +2 charge, Copper has a +2 and +1 charge)
-The top number on the periodic table is more commonly used
-IUPAC uses roman numerals in parentheses to show the charge (eg copper II sulphide)
-Classical systems use latin names of elements and the suffixs -ic (larger charge) and -ous (smaller charge)
Complex Ions
-Complex ions are large groups of atoms that stay together during a chemical reaction
-Almost all are anions (negative) (eg. NaNO3, FeSO3)
Hydrates
-Some compounds can form lattices that bond to water molecules (eg. Cu(SO4) * 5H20(s) = Copper(II) Sulphate Pentahydrate)
-To name hydrates, write the name of the chemical formula, add a prefix indicating the number of water molecules (eg. mono, di, tri... etc), add hydrate after the prefix
Naming Acids & Bases
-Hydrogen compounds are acids
-HCl (aq) - Hydrochloric Acid
-H2SO4(aq) - Sulfuric Acid
- Hydrogen appears first in the formula unless it is part of a polyatomic group
(eg. CH3 COOH(aq) - Acetic Acid)

November 18, 2010: Mole Conversions

Converting Between Moles and Mass:
- To convert between moles and mass we use molar mass as the conversion factor
- Be sure to cancel the appropriate units.
Example 1: How many grams are there in 1.5 mol of H2?

Step 1: 1.5 mol. x [(2(1.0 g.)) / (1 mol.)] = ?
Step 2: 1.5 mol. x [(2 g.) / (1 mol.)] = ?
Step 3: 1.5 mol. x [(2 g.) / (1 mol.)] = 3 g.

There are 3 grams of H2 in 1.5 moles.

Example 2: A sample of HCl contains 0.72 mol. How many grams of HCl are there?

Step 1: 0.72 mol. x [(1(1.0 g.)) / (1 mol.)] = ?
Step 2: 0.72 mol. x [(35.5 g.) / (1 mol..)] = ?
Step 3: 0.72 mol. x [(35.5 g.) / (1 mol.)] = 26 g.

There are 26 grams of HCL in 0.72 mol.

Example 3: How many moles are there in 110 grams of Fe2O3?

Step 1: 110 g. x [(1 mol.) / (2(55.8 g.) + 3(16.0 g.)] = ?
Step 2: 110 g. x [(1 mol.) / (159.6 g.)] = ?
Step 3: 110 g. x [(1 mol.) / (159.6 g.)] = 0.689 mol.

In 110 grams of Fe2O3 there is 0.689 mol.

November 18, 2010: Molar Mass (Angelo)

Mass of Atoms:
- The mass (in grams) of 1 mole of a substance is the molar mass
- It can be determined from the atomic mass on the periodic table measured in grams per mole (g/mol).
Example:
What is the mass of Copper (Cu)?
63.5 grams  

Molar Mass of Compounds:
- To determine the molar mass of a compound, add the mass of all the atoms together.
Example: 
What is the molar mass of H2O? 
Write out the the mass of each atom and multiply that by its frequency in the compound and calculate: 2(1.0 g) + 1(16.0 g) = 18.0 g/mol.

NOTE: Because we are adding instead of multiplying or dividing figures, we round our final answer to the highest number of significant digits in our equation.

November 5, 2010: Hydrate Lab (Angelo)

Hydrate Lab:

Discussion:
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. Examples are: magnesium sulfate heptahydrate (epsom salts) and sodium carbonate decahydrate (washing soda). The formulas for substances are MgSO4*7H2O and Na2Co3*10H2O. They can also be represented as MgSO4(H2O)7 and Na2CO3(H2O)10. In this lab you will be dtermining the anhydrous (without water) mass of the hydrate. You will compare this with the actual mass of water that should be present.

Materials:
- Bunsen burner
- Test tube
- Test tube rack
- Test tube clamp
- Weight scale
- Cobaltous chloride hexahydrate

Procedure:
1. Fill a test tube about 1 cm. with the hydrate.
2. Carefully place the test tube on the scale and record the mass of the hydrate and test tube.
3. Using extreme caution, connect andl ight your Bunsen burner. Adjust the gas flow until the flame is about five cm. tall.
4. Pick up the test tube with the clamps and carefully hold it in above the Bunsen burner.
5. Gently heat the test tube by moving the test tube in and out of the flame for about 5 minutes or until all the water has boiled away.
6. Carefully re-weight the test tube ensuring none of the chemicals inside spill.

Observation:
Test Tube Weight - 18.78 grams
Mass Before Heating - 19.38 grams
Mass After Heating - 19.09 grams

Analysis:
1. Determine how much water was released during the heating? 0.29 grams
2. What percent of the hydrate was water? 48%

Conclusion:
1. The actual percent water in this hydrate is 45%. Determine your percent error for part 2.
Write out the formula: [(measured - accepted) / (accepted)] x 100 
Plug in the numbers: [(48% - 45%) / (45%)] x 100
Calculate: 6.7%

In this lab conducted, the percent error was 6.7 percent.

October 28th,2012:Trends on the Periodic table(Brian)

Elements close to each other on the periodic table display similar characteristics.
The are 7 important trends:
1) Reactivity
     - metals and non metals show different trends 
     - the most reactive metal is Francium and the most reactive non-metal is Fluorine
2) Ion Charge
     - an elements ion charge depends on their group(column)
3) Melting Point
     - elements in the center of the table have the highest melting points
     - Noble gases have the lowest melting points
     - starting from the left and moving right, melting points increase(until the middle of the table)
4) Atomic Radius
     - the radius decreases going up and right on the table
     - Helium has the snallest atomic radius and Francium has the largest atomic radius
5) Ionization Energy
     - is energy needed to completely remove an electron from an atom
     - it increases going up and right
     - all noble gases have high ionization energy
     - Helium has the highest and Francium has the lowest
     - opposite trend of the Atomic Radius
6) Electronegativity      - refers to how much an atom wants to gain electrons
     - same trend as ionization energy
7) Density
   

October 25th, 2010: Isotopes & Atoms (Zac)

Atomic Number
- Atomic # - # of protons
- Atomic mass - atomic number = # of Neutrons
- Isotopes - same atom but different mass
- For example, there are 3 types of chlorine atoms (35H, 36H, and 37H)

Mass Spectrometers
- Are used to determine the abundance ad mass of the isotopes of the element
- A device known as a mass spectrometer can be used to determine the "relative abudance" and the "mass" of  the "isotopes" of the elements

Oct.21,2010:Quantum Mechanics (brian)

Bohr Theory
-electrons are particles that must be in the orbital of an atom
-Quantum Theory
-an electron is like a cloud of negative energy/wave
-orbitals are areas in a 3D space where the electrons most probably are
-energy of the electron is in its vibrational modes
-photons are produced when high energy modes change to lower energy modes

1. S orbitals
   - hold 2 electrons 
2. P orbitals
   -have 3 suborbitals
   -each contain 2 electrons
   -total electrons = 6  
3. D orbitals
   -have 5 suborbitals
   -each contain 2 electrons
   -total electrons = 10
4. F orbitals
   -have 7 suborbitals
   -each contrain 2 electrons
   -total electrons = 14

                                

                  

October 15th, 2010: Bohr's Model (Zac)

- Bohr (1920's) based his model on the energy (light) emitted by different atoms
- Each atom has a spectra of light

- To explain this emmission spectra, Bohr suggested that electrons occupy shells or orbitals
BOHR'S THEORY
- Electrons exist in orbitals
- When they absorb energy the move to a higher orbital
- As they fall from a higher orbital to a lower one they release energy as a photon of light

October 13, 2010: Atomic Theory (Angelo)

Atomic Theory:
- Many theories have been made to explain atoms. Not all of them are true today.

Aristotle (384 B.C. to 322 B.C.):
- Invented the four elements theory. (Water, Earth, Wind, and Fire)
- The four elements theory lasted for about 2000 years.
- It is not a scientific theory because it could not be tested against observation.

Democritus (460 B.C. to 370 B.C.):
- In 300 B.C., Democritus said atoms were indivisible particles.
- This was the first mention of atoms (atomus).
- Not a testable theory, only a conceptual model.
- No mention of any atomic nucleus or its consituents.
- Cannot be used to explain chemical reactions.

Lavoisier (1743 to 1794):
- Created Law of Conservation of Mass.
          - States that the mass of a system will remain constant.
- Created Law of Definite Proportions
          - Water is always 11% H (Hydrogen) and 89% O (Oxygen)

Proust (1754 to 1826):
- If a compound is broken down into its constituents, the products exist in the same ratio as in the compound.
- Proust experimentally proved Lavoisier's laws.

Dalton (1766 to 1844):
- Thought atoms are solid, indestructable spheres (like Billiard balls).
- Thought each element had different types of atoms (different color, shape, etc.).
- Based on the Law of Conservation of Mass.
- Have a molecule (atoms combine in simple whole number ratios) explains the Law of Constant Composition.
- If the atoms are not destroyed then the mass does not change.

J.J. Thompson (1856 to 1940):
- Raisin bun
- Solid, positive spheres, with negative particles embedded in them.
- First atomic theory to have positive (protons) and negative (electron) charges.
- Demonstrated the existence of all electrons using a cathode ray tube

Rutherford (1871 to 1937):
- Showed that atoms have a positive, dense centre with electrons outside it.
- Resulted in planetary model.
- Explains why electrons spin around nulceus.
- Suggest atoms are mostly empty space.

Sodium Chloride Lab (Brian)

Problem: What is the maximum amount of table salt that can be dissoloved in 200mL of water?

Observations:

   

Trial	Volume of Water (mL)	Mass of Salt (g)
1	10	1.01g
2	20	2.14g
3	40	4.02

Analysis:

    1. the units for the slope of our graph was g/mL

    2. the slope repersents how many g of salt in a certain amout of ml of water

    3. our best fit line prediction was 20.1 g

   


              

20.1 - 70.8 = -252
     20.1

September 30th, 2010 - Density & Graphing (Zac)

-The density of an object is its mass divided by its volume.
   eg. d= m/v
-It is usually expressed kg/L, kg/m3, or g/km3 (3 - to the power of 3)
-All Graphs must contain 5 important things
1. Labelled Axis
2. Appropriate Scale
3. Title
4. Data Points
5. Line of Best Fit
- There are 3 things you can do with a graph
1. Read it
2. Find the slope (rise/run)
3. Find the area under the graph

September 28, 2010: Dimensional Analysis (Angelo)

Dimensional Analysis:
- Want to know what 100 km/h (kilometres per hour) is in mi/h (miles per hour)?
- Just like converting between currencies, in chemistry, it is usually necessary to convert between units.
- This process is called Dimensional Analysis

Steps for the Dimensional Analysis process:

1. Find the unit equality
2. Find the conversion factor
3. Apply the conversion factor
4. Cancel Units

EXAMPLES:

1. How many miles are equal to 200 kilometres?

    Find the unit equality: 1 mile (mi.) = 1.6 kilometres (km.)
    Find the conversion factor: 1 = 1 mi. / 1.6 km.
    Apply the conversion factor: (200 km.) x (1 mi. / 1.6 km.)
    Cancel Units: 125 mi.

    Answer: 200 km. is equal to 125 mi.

2. What is 100 kilometres per hour in metres per second?

    Find the unit equality: 1000 metres (m.) = 1 kilometre (km.) and   
                                    3600 seconds (s.) = 1 hour (h.)
    Find the conversion factor: 1 = 1000 m. / 1 km. and 
                                           1 = 1 h. / 3600 s.
    Apply the conversion factor: (100 km/h) x (1000 m. / 1 km.) x (1 h. / 3600 s.)
    Cancel Units: 28 m./s.

    Answer: 100 km./h. is equal to 28 m./s.

Here's a helpful example of changing inches to feet using Dimensional Analysis:

September 23, 2010: Scientific Notation & S.D. (Brian)

Significant Digits

-Accuracy and precision is very important as well as communicating the accuracy carefully

-Calculators are not smart enough to decide what is and isn't precise

-Non zero digits are always significant

-If the zero is a place keeper its not significant


-Any numbers to the left of the decimal are significant
Example 6.004

-Zeros after another number are significant
Example 4.00

-When adding or subtracting round to the lest precise number
Example 8.1240 - 3.33 = 4.79 

-When multiplying or dividing round to the number with the fewest S.D.s
Example 4.83 x 9.1 = 43.953 = 44

-Constants on your data sheet have infinite S.D.'s

Scientific Notation
-Used to show really big numbers and really small number 

September 21st, 2010: Experimental Accuracy (Zac)

- In general, the maximum accuracy of any measurement is 1/2 of the smallest division of the measuring device
- A ruler with measurements of millimieters has a maximum accuracy of  plus or minus .5 mm.
Expressing Error
- Error is a fundamental part of science
- There are usually 3 reasons for error
1. Physical errors in the measuring device
2. 'Sloppy' measuring
3. Changing ambient conditions
Calculated Errors
- There are two different errors, absolute & percentage error 
Absolute Error
- Measured value minus accepted value 
Absolute Error = Measured - Accepted
Percent Error
- Most Common
- Eg.