Empirical Formulas

(November 13, 2009)
Today, we learned that empirical formulas are the simplest formulas. It is like the lowest common denominator of a molecular formula. Therefore, the empirical formula is the reduced of a molecular formula. For example, the molecular formula of tricarbon hexaoxide is C3O6. This can be reduced to CO2 (reduced by 3), this is the empirical formula.

Empirical formulas are the simplest whole number ratios in a compound. Molecular Formulas show the actual atoms and bonds.

- Example - : A sample of an unknown compound is analyzed and found to contain 8.4 g of "C", 21.0 g of "H", and 5.1 g of "O". Find the empirical formula.

• To do this, you must fill in a chart:
• You must find the amount of moles of each element by dividing its molar mass by the given mass.
• Out of all the moles, find the smallest.
• Divide every mole with the smallest mole
• Round your quotient to the nearest one
• Stick your answers as subscripts to the element.

That is all. Mr. Doktor is the best!

-- Jael Lumba

Percentage Mass of Elements in Compounds & Finding Mass of an Element in a Sample.

(November 10, 2009)
Today we learned how to calculate the percentage of mass of elements in compounds. Recall that molar mass is grams over moles: g/mol and molar volume is liters over moles: l/mol.

To find the percentage mass, first find the total molar mass of the compound. Then, for each element in the compound, find its individual molar mass. Divide the element mass with the compound mass:

% of Element = (Element Mass / Compound Mass) x 100

- Example - Compound: Fe3N2

• In iron (II) nitride, it's molar mass is 195.4 g. *Note that iron has a mass of 55.8g and nitrogen has a mass of 14.0g: 3(55.8) + 2(14.0) = 195.4 g
• Next, the individual mass of iron is 167.4 g (3 x 55.8 g), nitride is 28.0 g (2 x 14.0 g)
• Now find the percentage mass by taking element's mass and divide it by the compound's mass.
  
• % of Fe = (167.4 g / 195.4 g) x 100 = 85.7%
• % of N = (28.0 g / 195.4 g) x 100 = 14.3 %
• To ensure the percentages are accurate, add them together. If their sum adds up to 100, it is accurate. If their sum adds up to a couple of decimals off, that is because of the rounding. Try to adjust the decimals so that your sum will equal 100.

To find the mass of an element in a given sample, first find the total molar mass of the compound. Then, for each element in the compound, find its individual molar mass. Divide the element mass with the compound mass. Lastly, multiply the quotient with the mass of the given sample:

- Example - Compound: 62.0 g sample of ZnO

• In zinc oxide, the molar mass is 81.4 g. *Note that zinc has a mass of 65.4 g and oxygen has a mass of 16.0 g: 1(65.4 g) + 1(16.0 g) = 81.4 g
• Next, the individual mass of zinc is 653.4 g (1 x 65.4 g) and oxygen is 16.0 g (1 x 16.0)
• Now divide the individual masses with the compound mass
• Zn = (65.4 g / 81.4 g) = .8034
• O = (16.0 g / 81.4 g) = .1966
• Then multiply the quotients with the mass of the given sample (62.0 g).
• Zn = (.8034)(62.0 g) = 49.8 g
• O = (.1966)(62.0 g) = 12.2 g
• Therefore, in a 62.0 g sample of ZnO, there are 49.8 g of zinc and 12.2 g of oxygen.
• To ensure masses are correct, add them together so that they equal the mass of the given sample. If they do not equal the mass of the sample, that may because of the rounding. Adjust the decimals so that the sum will equal the given mass of the sample.

Here is a tutorial clip I found:


That is all. Mr. Doktor is the best!

-- Jael Lumba

Density and Moles

Today's class (November 4), We learned how to calculate the density of gases at STP (Standard Temperature and Presssure)

-density equals the mass per unit volume
-an easy way to rememeber how to calculate Mass, Density and Volume is to picture it as a triangle with Mass on top and Density and Volume on the bottom
-to calculate the density of gases at STP we follow this equation

DensitySTP=Mass of 1 mole\Volume of 1 mole=Molar Mass\22.4 L=g per mol\L per mol


Examples:

1. Calculate the density of O2 @ STP

Dstp=32 g per mol\22.4 L per mol=1.43 g per L

2. The density of a gaseous mixture of C8H18 @STP=?

D=M\V
D=114\22.4
D=5.1 g per L

3. Calculate the density of Dihydrogen Phosphate

D=M\V
D=97\22.4
D=4.3 g per L



D&V-mass=d x v-MASS-molar mass-MOLE-6.02x10 to the 23-MOLECULES-subscripts-ATOMS

Molar Volume Lab

Todays class we did a lab that focussed on calculating molar mass and volume. The procedure for the lab is as follows:

1. Fill the sink with 3\4 of water and place the lighter in so no air remains.

2. Dry the lighter and weigh it

3. Place the gradulated cylinder into the sink and once again make sure there is no air.

4. Place the lighter in the water under the gradulated cylinder and press the button. Add 10 milileters of gas.

5. Record the volume of the gas and the mass of the lighter.

Moles of Iron and Copper

During class on October 27, 2009, the whole class participated in a lab. Through this experiment, it's main objectives were to:

Determine the number of moles of copper produced in the reaction of iron and copper(II) chloride

• Determine the number of moles of iron used up in the reaction of iron and copper(II) chloride
  

• Determine the ratio of moles of iron to moles of copper


• Determine the number of atoms and formula units involved in the reaction

The procedure of this particular lab consisted of:

1. Finding the mass of a clean, empty, dry 250 mL beaker. Record the mass to the nearest 0.01g.
2. Add aprrox. 8 grams of copper (II) cloride crystals to the beaker. Find the mass and record it in your notebook.
3. Add 50mL of distilled water to the beaker. Swirl th ebeaker around to dissolve all the copper(ii) chloride crystals.
4. Obtain two clean, dry nails. Find the mass of the nails and record it in your notebook.
5. Place the nails into the copper(II) chloride solution. Leave them undisturbed for approx. 20 minutes.
6. Use the tongs to carefully pick up the nails, one at a time. Use distilled water in a wash bottle to rinse off any remaining copper from the nails before removing them completely from the beaker.
7. After the nails are completely dry, find the mass of the nails and record it in your notebook.
8. Decan the liquid from the solid.
9. Rinse the solid again with about 25 mL of distilled water. Decant again. Repeat this step three or four times.
10. Wash the solid with about 25 mL of hydrochloric acid. Decant again; then, once more, clean the solid with 25 mL of distilled water.
11. After final washing, place the copper in a drying oven to dry.
12. Find the mass of the beaker plus the copper and record it in your notebook.
    

--Krizia Umali

Gases & Moles

(Thursday, Oct 22, 2009)

continuation of the mole..



This class we learned that the volume is occupied by a certain mass depends on the temperature and pressure. The Standard Temperature and Pressure (STP) is temperature = 0 degrees Celsius and pressure = 101.3 kPa. The volume of any gas at STP is 22.4 L for every mole.

22.4 L / 1 mol OR 1 mol / 22.4 L

Examples:

- Finding the volume (L) occupied by 0.0606 mol of CO2 at STP
*Multiply by 22.4 L / 1 mol
0.060 mol x (22.4 L / 1 mol) = 1.3 L

- Finding the mass of a 200.0 mL sample of NO2 at STP
*Multiply by 1 mol / 22.4 L, but mL must be converted to L. Note that NO2 has 46 g / mol.
200.0 mL x (1 L / 1000 mL) x (1 mol/ 22.4 L) x (46 g / 1 mol) = 0.41 g

Here's a website I found that can calculate STP: http://www.1728.com/stp.htm

That is all. Mr. Doktor is the best!

-- Jael Lumba

The Molecular Mass!

(Tuesday, Oct 20, 09)

Today we learned about the mole! A mole is the number of atoms f that element equal to the number of atoms in exactly 12.0 grams of carbon-12.

- For example, 1 mole of Carbon (c) atoms is 12.01 g,

- and 1 mole of Calcium (Ca) atoms is 40.1 g.

For the diatomic molecules, plus phosphorus and sulphur, are however different when not joint into a compound. The diatomic seven (hydrogen, oxygen, nitrogen, chloride, bromine, iodide, and fluoride) have 2.

- For example, hydrogen is 2, therefore 2(1.0 g) = 2.0 mol.

- Another example is oxygen. Oxygen is 2, therefore 2(16.0 g) = 32.0 mol.

When finding the molar mass in compounds, the molar mass of all atoms must be added:

- H2o (water) has 2 hydrogen atoms and 1 oxygen atom. The mass of hydrogen is 2(1.0 g) = 2.0 mol and oxygen is 1(16.0 g) = 16.0 mol. Together, the molar mass of the compound is 18 mol

To find the mass in moles of a compound we use g/mol. To find the number of moles in a compound we use mol/g. For example:

- To find the mass of 2.5 moles of water..
2.5 mol of H2O x (18.0 g / 1 mol) = 45 g
H = 2(1.0)
O = 1(16.0)
= 18.0 g / mol

- To find the number of moles in a 391 g sample of nitrogen dioxide..
391 g of NO2 x (1 mol / 46 g) = 8.5 mol
N = 1(14.0)
O = 2(16.0)
= 46 g / mol

Here's a handy clip in finding molar mass:



That is all. Mr. Doktor is the best!

-- Jael Lumba

2-4 Elements and Compounds

Today's class we did two main experiments. The pickle electroloysis and the sugar and sulfuric acid dehydration demonstration. We saw in the pickle electroloysis what happens when the compound splits apart into its constituent elements. The sugar and sulfuric acid dehydration was by far the highlight of the class. When Mr. Doktor first added the sulfuric acid and the sugar turned a brownish colour, it really wasn't all that exciting. Without warning however, the sugar rose up as a black solid and gave off a strong scent which had everyone coughing and laughing. Following the two demonstrations, I had to leave for volleyball so I wasnt there for the acids and bases notes. I'll be getting the notes from Jarren so those will be put up soon. Here is a quick video that re-enacts our dehydration demonstration. http://www.youtube.com/watch?v=uaB70TgLfqs

That is all. Mr Doktor is the best!

Michelle Haughian

2-4 Elements and Compounds

Todays class we focussed on four main concepts: Hydrates, Naming Hydrates, Molecular Compounds and Naming Molecular Compounds.

I. HYDRATES
1. Some compounds can form lattices that bond to water molecules; without water the compound is often preceded by 'anhydrous'.
Example: Copper Sulfate 2. Those crystals contain water inside them which can be released by heating.

II. NAMING HYDRATES

1. There are three simple steps to naming hydrates:

a. write the name of the chemical formula.

b. add a prefix indicating the number of water molecules.

c. add "hydrate" after the prefix.

Examples:

1. Cu(SO4) 5H2O= Copper (II) Sulfate pentahydrate

2. Li(CLO4) 3H20= Lithium Perchlorate trihydrate

3. Nickel (II) Sulfate hexahydrate=NiSO4 6H20

TABLE OF PREFIXES:

1-mono

2-di

3-tri

4-tetra

5-penta

6-hexa

7-hepta

8-octa

9-nona

10-deca

III. MOLECULAR COMPOUNDS

I couldn't get the video to embed as well...enjoy!

http://www.britannica.com/EBchecked/topic-video/108614/109656/Molecular-compounds-are-formed-when-molecules-such-as-those-of

1. they are covalent; 2 or more non metals

2. low melting and boiling points

3. they share (not exchange) electrons

4. elements usually end in -gen

5. 7 molecules are Diatomic

a. 2 of the same elements

b. H2, N2, O2, F2, Cl2, Br2, I2

6. 2 molecules are Polyatomic

a. P4, S8

IV. NAMING MOLECULAR COMPOUNDS

In class we all decided that we knew how to name molecular compounds so we just did a few examples on the board.

1. N2O4= Dinitrogen tetraoxide

2. CS2= Carbon disulfide

3. P4O10=Tetraphosphorous decaoxide

4. Nitrogen trichloride= NCl3

5. Sulphur dibromide= SBr2

6. Dihydrogen oxide= H2O

IUPAC NAME FORMULA

1. Water H2O

2. Hydrogen Peroxide H2O2

3. Ammonia NH3

4. Glucose C6H12O6

5. Sucrose C12H22O11

6. Methane CH4

7. Propane C3H8

8. Octane C8H18

9. Methanol CH3OH

10. Ethanol C2H5OH

11.Ethane C2H6

That is all. Mr. Doktor is the best!

Michelle Haughian

2-4 ELEMENTS AND COMPUNDS (cont.)

Because our previous class was shortened, we continued our lesson by learning:

There are many methods in seperating mixtures. Depending on the type of mix, possible methods of seperating mixtures include

• By hand


• Filtration


• Distillation


• Crystallization


• Chromatography

Some elements use Latin. But, there are a few exceptions. Some of those exeptions include:

• Copper= Cuprum


• Gold= Aurum


• Iron=Ferum


• Lead= Plumbum


• Silver= Argentum

Naming chemical compunds is a difficult task. But, the most common system is the IUPAC

• IONS


• BINARY IONIC


• POLYATOMIC IONS


• MOLECULAR COMPOUNDS


• ACIDS

When naming ions, remember that for metals: use the name of your element and add ion. For non-metals, remove the original ending and ade -ide.

Binary Ionic Compounds contain two elements (one metal and one non-metal)

• Metallic and non-metallic bond together
• Electrons trasnfer from metal to non-metal
• Net charge=0 ( total positive charge=total negative charge)

That is all. Mr. Doktor is the best!

BY: Krizia Umali

2-4 ELEMENTS AND COMPOUNDS

Although today's lesson was cut short, we learned that...

Matter could be divided into:

• homogeneous- consists of only ONE visible component
• heterogeneous- consists of MORE than one visible component

We also learned...
The differences between an element and compund is that:

• Elements cannot be broken down into simpler substances by chemical reactions
• Compunds are made up of two or more elementsby chemical reactions.

Finally, we learned...
the difficulty of differentiating an element and compound

• One method> Electrolysis(process which splits the substance into its constituent elements)

http://www.youtube.com/watch?v=WB0Kr8VA-74 - For some reason, I couldn't get the EMBED for this video. But, enjoy!

That is all. Mr. Doktor is the best!

BY: Krizia Umali

What Matters?

We know matter commonly as 3 forms: Solid, Liquid, and Gas... and for the more intelligent/Star Wars Fanatic, there is also Plasma, which is a state of matter thats like gas and liquid. We learned the different boiling points for water, and the effects of each change.

We also learned of three different types of change in matter:

- Physical Change
- Chemical Change
- Nuclear Change

In physical change: matter is changed in shape or form, but no new substances are created. Chemical changes result in a new substance being formed. We did not learn so much about Nuclear change... yet.

That is all. Mr. Doktor is the best.

--- by JBATO

Test Time E Block

So today instead of a usual lesson we had our highly anticipated Unit Test ... yay!
Our test consisted of safety symbols, lab materials, conversion, and calculating percentage error. It consisted of us knowing our:

- Lab symbols
- Significant Digits
- Percent Error
- and our Conversion Units !


That is all. Mr. Doktor is the best.

--- by JBATO

(Thursday, Sept 17, 2009)

None of us were present for this class.
That is all. Mr. Doktor is the best.

-- by Jael Lumba

In today's lesson (Wednesday, Sept 16, 2009), we learned about dimensional analysis and graphing. Dimensional analysis is the technique of converting between units. An example would be changing kilograms to milligrams:

120 kg x 1000g/1kg x 1000mg/1g = 120 000 000 mg = 1.2 x 10 to the power of 8

We learned that by doing an equation like this, it will help us cancel out units not needed in order to find the needed unit.

Here's a youtube video I found with lots of example on dimensional analysis:

We also learned about the basics of graphing:
1. Title
2. Labeled axis
3. Correct scaled axis
4. Data plotted correctly
5. Best-fit line

This is a perfect example of a graph done for votes in Iran:

That is all. Mr. Doktor is the best.

-- by Jael Lumba

(lesson on Monday, Sept 14, 2009)

This class' lesson was the introduction of the International System of Units aka SI. We learned that it consists of 7 metric units. These include metres for length, kilograms for mass and weight, and metres cubed for volumn. We also learned the metric prefixes, which is a syllable attached to the beginning of a metric unit to make it larger or smaller. Examples would be:

- kilo- (k) which means 1000 and is 10 to the power of positive 3. (to make a unit larger)

and..

- milli- (m) which means 0.001 and is 10 to the power of negative 3. (to make a unit smaller)



That is all. Mr. Doktor is the best.

-- by Jael Lumba