December 4/09

In todays class we learned about Molecular Formula and we did a lab on dilution.

For the lab Mr.Doktor showed us 5 test tubes each that are different colors and we had to compare our test tube to the other 5 and pick one of them.

First we had to measure the amount of water we needed
Then do our calculations to get the amount of Copper (II) Chloride we needed
Then add our solution to the water and stir
Our last step was to pick the test tube that was the most similar to ours.

A Molecular formula is what the equation should actually be.

You will need to know the Empirical Formula before doing the molecular formula lets jut say our empirical formula is CH2O Molecular Formula which is getting the mass in the equation dividing it by the original molar mass and then just multiply the equation and use subscripts.

Example : CH2O has a molecular weight of 180g/mol find the molecular formula.
Original Molar mass = 30g/mol
Mass in equation = 180g/mol
180g/mol / 30g/mol = 6 C = C6 H2 = H12 O = O6

The molecular formula would be C6H12O6

Video showing Molecular formula:

December 2 2009

In todays class we learned about giving directions for dilution of solutions and learn about dilution of solutions itself.

To give directions to a experiment we must first find the information we need.
concentration > moles > mass

concentration = n/v
N represents number of moles
V represents the volume

Once you find all of that you would tell them how much volume to measure how much grams of the solution you would put in and finally add the solution itself.

Example: Son needs to make a 4M solution of CaCl2. If he needs 2 L what procedure will he use?
4 mol x 2 L = 8mol
4 mol x 111g = 444g
Now we have all of the information we need
1) first measure 2L of water in your test tube
2) weigh 444g of CaCl2
3) add the CaCl2 to the water and stir the solution

We also learned about the dilution of solutions which is when you add water to the concentration it decreases. If the volume is doubled then the concentration would be halved. If the volume is halved then the concentration would be doubled.

4 L = 2M <--- what were starting with
8L = 1M <--- multiply 2 volume divide the M by 2
2 L = 4M <------ divide the volume by 2 multiply the M by 2

to solve for dilution of solutions we would use C1V1 = C2V2
C1 = initial concentration
C2 = final concentration
V1 = initial volume
V2 = final volume

An example would be:
If 30mL of .67M of NaBr is diluted to a total volume of 60ml what is the final Molarity of the solution?

C1 = .67M
C2 = ?
V1 = 30ml
V2 = 60ml

C2 = (V1)(C1)/V2 C2 = (30ml)(.67M)/60ml C2 = .34M\

Video helping to solve dilution of solutions:


Heres a link on Dilution of Solutions:
http://www.emsb.qc.ca/laurenhill/science/cv.html

November 30 2009

Today in chemistry we learned about Concentration which is the amount of solute and amount of solvent some units used for this are g/ml g/l mg/l mg/ml

Solution: a homogeneous mixture Solute: the one present in the smaller amount Solvent : the one present in the bigger amount

The unit we use the most is mol/L. Also known as Molarity.
Molarity = Moles/Volume

Molarity = Moles/Volume only works for aqueous solutions and not gases.

Here are some Examples:

Calculate the Molarity of 2.5mol of HSO3 in 3L solution of [HSO3]
[HSO3] 2.5mol/3L = .83M

How many grams of NaCl are contained in 300ml of .0420M of NaCl solution?
.0420M x .0300 L = .0126 mol
.0126 mol x 58.5g = .737g of NaCl

Heres a video on how to calculate Molarity

November 13/09

Today we learned about empirical formulas and how to create them. Molecular formulas show the actual bond between an atom and a molecule.

Ex. Cl8O4

An Empirical formula would show the bond between an atom and a molecule in its simplist form.

Ex. Cl2O

To find an empirical formula there are many steps involved i will now demonstrate what i learned in chemistry class today.

A sample of an unknown compound is analyzed and it contains 9g of Mg 15g of O and 5g of C

find the empirical formula.

You would want to start out by geting the molar mass of all the elements

Molar Masses:

Mg - 24

O - 16

C - 12

Then you would want to get there mass which you is stated above in the question.

Masses of elements:

Mg - 9g

O - 15g

C - 5g

After you would want to find the number of moles in each element you would do that by geting the mass and dividing it by the atomic mass.

Number of Moles:

Mg- 9 / 24 = .375mol

O - 15 / 16 = .938mol

C - 5 / 12 = .42mol

Then you would get your moles and divide it by the SMALLEST number of moles you got in this case it would .375mol.

Moles / SMALLEST mol :

Mg - .375 / .375 = 1

O - .938 / .375 = 2.5

C - .42 / .375 = 1.1

Then round your answers

Rounded answers:

Mg - 1

O - 3

C - 1

The final step would be get those numbers for each of the elements and put them in as subscripts. So the empirical formula would be MgO3C.

If you still want more help and example on empirical formulas this website should help.

http://dl.clackamas.cc.or.us/ch104-04/empirica.htm

November 11/09

Today in chemistry 11 we learned how to calculate the mass of an elements in compounds.


Find the percentage of each of the following elements C2H2O

First you would have to get all the molar masses of the compound and add them together

2C - 24
2H - 2
1O - 16
total : 42

Then you choose 1 element lets say we go with carbon first so we would get the molar mass of carbon and divide it by the total mass of all the elements 24 / 42 and multiply it by 100 which gives you 57%. You would do the same for hydrogen 2 / 42 x 100 4% and do the same for oxygen 16 / 42 x 100 which gives you 38%.

Percent
C2 would be 57%
H2 would be 4%
O would be 38%


We also learned how to find a certain amount of an element in a certain amount of grams.

For example find the mass of magnessium contained in 50g sample of MgNa
First find the molar masses of all the elements

Magnessium = 24
Sodium = 22
Then add them both together which gives us 46

Second since were looking for the mass of magnessium we get the molar mass of magnessium 24 and divide it by the total mass of the two elements 24 / 46 which gives us .52

After we would get .52 and multiply it by the given grams in the samepl of MgNa which is 50 x .52 which gives us 26. This would tell us how many grams of magnessium are contained in a 50g sample of MgNa which is 26g.



Heres a video pretty much summing up todays lesson and giving us a head start on the next lesson empirical formulas and percentages.

November 4/09

In today's class we learned about density and there relationship between moles.
Density= Mass/Volume Mass= Volume*Density Volume= Mass/Density

Finding the density of gases at STP is easier then finding it at solids and liquids because we know that the volume of the gas will be 22.4L. We can also find the mass by checking our periodic tables. After you just plug in the numbers and solve for the density of the gas.

Molar Mass g / 22.4L/mol = ?g/L

density of oxygen =
Density = Mass / Volume

32g / 22.4L/mol = 1.43g/L at STP

Finding the density of solids and liquids are much harder you do not have all the information you need such as the volume. Luckily Mr.Doktor gave us a neat chart to help us remember how to get the information we need.
use the formula g / molar mass of ele. 6.02 x 10 to the power of 23 subscripts
Density ---> Mass ----> Moles ---> Molecules ----> Atoms

Example: the density of Boron (solid) is 2.34g/mL how many molecules are in a 60ml piece?
2.34 g/ml * 60.0 mL = 140.4g
140.4 g * 1mol / 10.8 g = 13mol
13mol * 6.02 * 10 to the power of 23 / 1mol = 7.84 * 10 to the power of 24

Neat website that can solve density volume or mass for you:
http://www.1728.com/density.htm

November 2/09

In todays class we did a lab to see if the volume at STP is really 22.4L.

1.The materials were:
2.100ml graduated cylinder
3.lighter with butane gas
3.a sink filled with water
4.scale



The procedure was:
1)fill the sink with water don't over flow it
2)put the graduated cylinder underwater until it is completely filled with water and there no gas bubbles
3)put your lighter underwater so that water gets in the lighter then go weigh it and record the 4)mass make sure you dry your lighter before weighing it
5)then put the lighter underwater at the entrance of the graduated cylinder and release 10ml of butane you should see gas starting to form and water starting to decrease
6)once you released 10ml of butane dry off your lighter and record its mass



Observations:
Mass of lighter : 16.5
Mass of lighter after butane released: 16.3
number of moles in the lighters mass: 0.3g

Percentage error

22.4 - 16.3 / 22.4 * 100 = 27%
some errors we had in this experiment was we had to redo the experiment twice because water got in our lighter and made it heavier even though we released butane gas and we released 20ml of butane gas instead of 10ml.

Chemistry October 29/09

Atoms and Molecules

For monotomic elements: a molecule=an atom
For diatomic elements: a molecule Cl2 an atom CL

Molecules of Compounds

O

H H =H2O

2H atoms
1O atom

= 1 molecule

Ammonium Carbonate: NH3 + CO3= (NH4)2 CO3=14 atoms
Converting moles to molecules and vice versa:

6.02x10^23 molecules/1mole
OR
1mol/6.02x10^23 molecules

Examples:
How many molecules are in a 0.25 mol sample of CO2?
0.25 mol x 6.02x10^23 molecules/1mol = 1.51x10^23 molecules

There are 1.51x10^23 carbon atoms
and there are 3.02x10^23 oxygen atoms because its O(2)

5.1772x10^24 molecules of water = ? moles
5.1772x10^24 molecules x 1 mole/6.02x10^23 = 8.6 mol

Another example:
Find the number of 'H' atoms in 4.0 mol of ammonia.
NH3 convert: mol → molecules → atoms

4.0 mol x 6.02x10^23/1mol = 24.0x10^23 molecules
= 3(24.0x10^23)
= 7.22x10^24 H atoms

If you still don't get this here is a website that can help you:
http://www.chem1.com/acad/webtext/intro/MOL.html

Here's a video showing how to convert moles to atoms:

October 27/09 Chemistry 11

Today in chemistry 11 we did a lab, to determine the ratio of moles of copper produced to moles of iron consumed during the chemical reaction between Iron and Copper (II) Chloride.



The procedure for the lab is :

1)Weigh your empty 100ml beaker and record its mass.
2)Add around 8g of copper(II) chloride to the beaker and record its mass
3)Add 50ml of distilled water to the beaker and stir until the copper(II) chloride dissolves
4)Get two nails that are clean and dry record the two nails masses and record it
5)Put the nails in your beaker that contains the copper(II) chloride for 20 minutes you should start to see copper form at the bottom of your beaker
6)Use the tongs to pick up each nail and then use the scupula to scrape off the remaining copper on the nails into the beaker
7)Dry the nails in the drying oven and record its mass once its dry
8)Decant the liquid from the solid, pour the liquid into another beaker (just in case you lose some of the copper you can still retrieve it from the other beaker)
9)Rinse the solid with 25ml of distilled water then decant again into the beaker you decanted into before
10)Wash the copper with with 25ml of hydrochloric acid then decant the hydrochloric acid and was the copper with 25ml of distilled water decant again after
11)Put your beaker with the copper in it in the drying oven to dry
12)Let the copper with the beaker dry and record the mass of the beaker with the copper in it
13)Wash your hands clean up your lab station put away your materials and lab aprons and safety goggles

After that we recorded our information:

Empty beaker: 92.5g
Beaker with copper (II) chloride: 97.67
Nails: 5.54
Nails after the solution: 5.2
Beaker with copper: 94.24

After we managed to determine the amount of iron used which was .34 and amount of copper produced which was 1.74

We ended up with a ratio of 3:1 the correct answer was 3:2 we had a 20% percentage error.
Some errors where we went wrong in the lab was we did not let our materials completely dry before weighing them and we also did not let our nails sit for the full amount of time in our solutions.

Chemistry 11 October 22/09

In Chemistry 11 we learned about gases and moles.
the volume occupied by a certain gas depends on the temperature and pressure.
STP the standard temperature and pressure that we are in right now is T = 0°C and the standard pressure is 101.3 kPa

The volume of any gas at STP is 22.4L for every mole

Ex. find the volume of L occupied by 0.060mol of CO2 at STP

0.060 * 22.4l/1mol = 1.3L

find the volume occupied by 3.6kg of fluorine at STP

F2 = 38g

3.6kg * 1000 = 3600g

3600g/38 = 94.737mol

94.737mol * 22.4 = 2121g

If you are still struggling and having problems here is a link showing more examples
http://answers.yahoo.com/question/index?qid=20090804233941AAAvZHL

Chemistry 11 October 20/09

Today in Chemistry 11 we learned about Atomic Mass.
-the mass of 1 mole of atoms of an element
Ex.
mass of 1 mole of "C" atoms is 12.0g
mass of 1 mole of "Ca" atoms is 40.1g

Molar Mass
- the mass in grams of 1 mole of molecules of an element or compound is the molar mass for the diatomics you would times there mass by two if there by themselves H2, Cl2, F2, I2, O2, N2, Br2, and same goes for the molecular ones P4, S8

Element Symbol Formula Atmoic Mass Molar Mass
Bromine Br Br2 79.9 159.8
Neon Ne Ne 20.2 20.2
Silicon Si Si 28.1 28.1
Hydrogen H H2 1 2
Iron Fe Fe 55.8 55.8

For compounds add the molar masses of all the atoms
Ex.
H2O
2H - 2
1O - 16
H2O = 18g/mol

Ca(NO3)2
1Ca - 40.1
2N - 28
6O - 96
Ca(NO3)2 = 164.1g/mol

When you are converting Grams to Moles you would want it to be g/mol or mol/g
Ex.
Find the mass of 2.5moles of water
2.5 mol of H2o * 18.0g/1mol= 45g


Find the number of moles in 391g sample of nitrogen dioxide

381g NO2 / 46g*1mol = 8.5 mols

62.6mg of C12H22O11 = ? moles
C12 = 144
H22 = 22
O11 = 176
C12H22O11 = 342

.0626g of C12H22O11 / 342g * 1mol = 1.83 * 10 to the negative 4

Down here is a video showing how to convert moles to grams

Chemistry 11 October 15/09

Today in chemistry we learned about moles.

1 mole is equal to 6.02 x 10²³

this number was created by Amedeo Avogadro. A pea of moles can cover all of the buildings in canada. Mr.Doktor also explained to us how to use the mole to balance equations.

Ex.
2 H2 + O2 = 2 H2O

2 molecules + 1 molecule = 2 molecules

12.04 x 10 power 23 + 6.02 x 10 power 23 = 12.04 x 10 power 23

2 moles + 1 mole = 2 moles

Avogadro also had a hypothesis and that was:
Equal volumes of any gas at the same temperature and pressure contain equal numbers of molecules.

We also learned about a man named John Dalton and how he used masses of chemicals to look for a patter in chemical reactions.

What he came up with were :

11.1g of H2 reacts with 88.9g of O2

46.7g of N2 reacts with 53.3g of O2

42.9g of C reacts with 57.1g of O2

Which there was no pattern.

And we learned about a guy named Joseph Gay-Lussac he also was looking for a pattern in chemical reactions by using simple ratios.

He came up with:

1L of H2 reacts with 1L of Cl2 = 2L of HCl

1L of N2 reacts with 3L of H2 = 2L of NH3

2L of CO reacts with 1L of O2 = 2L of CO2

Here's a youtube video explaining a concept of the mole:

Chemistry 11 October 13/09

We had our chemistry test today on matter, ionic compounds, molecular compounds, and naming chemical equations.

Chemistry October 9/09

Dear Mr.Doktor we were both sick on this day so none of us could do a blog for it =(.

Chemistry October 7/09

Acids and Bases and How to Name Them

Acids

-Solid, Liquid, or Gas SATP (Standard, Ambient, Temperature, Pressure)
T = 25oC p = 100KPA
-form conducting aqueous solutions
-turn blue litmus red
-dissolve in water to produce H+
-taste sour

Bases

-turn red litmus blue
-slippery
-non conductive
-dissolve in water to produce OH-

Naming Acids

-acids are aqueous (dissolve in water)
-hydrogen compounds are acids
Ex.
-HCl(aq) - Hydrochloric acid
-H2SO4(aq) - Sulfuric acid
-Hydrogen appears first in the formula unless it is part of polyatomic group
Ex.
CH3COOH(aq) - Acetic Acid
-classical rules use suffix ic or prefix hydro
Ex.
Hydrochloric acid
Sulfuric acid
IUPAC uses the aqueous hydrogen compound
Ex.
HCL(aq) - Aqueous Hydrogen chloride

Naming Bases

-for now all bases will be aqueous
-solutions of ionic hydroxides
Ex.
NaOH - sodium hydroxide
BaOH - barium hydroxide

Examples of naming acids and bases:

H3PO4(aq) - phosphoric acid
HLVO3(aq) - Nitric acid
HNO2(aq) - Nitrous acid
Mg(OH)2(aq) - Magnesium hydroxide
HBr(aq)- Hydro Bromic acid
HOOCCOOH - Oxalic acid

Here's a website that covers everything and more about today's lesson

http://chemistry.about.com/od/acidsbases/Acids_Bases_and_pH.htm

Here's a song about acids and bases

Chemistry October 5/09

Today in chemistry 11 we learned about hydrates, how to name them, molecular compounds, and how to name them.

Hydrates

-compounds that form lattices that bond to water molecules
-when there are no water compounds in a hydrate often preceded by anhydrous ex. copper sulfate
-crystals contain water inside, it can be released by heating



Naming Hydrates

-write name of the chemical formula
-add prefix indicating number of water molecules (mono, di, tri)
-add hydrate after prefix
-if something has more then one charge you can use ic at the end indicating a larger charge or ous indicating a smaller charge
Examples: cu(So4Li - 5H2O8 - cuppric sulfate pentahydrate
: Li(ClO4) - 3H2O - Lithium perchlorate trihydrate
: Nickel (II) sulfate hexahydrate - Ni(SO4) - 6 H2O

You will also need to know your prefixes for naming hydrates and molecular compounds
mono - 1
di - 2
tri - 3
tetra - 4
penta -5
hexa -6
hepta -7
octa - 8
nona - 9
deca- 10

Here's a video showing you how to name hydrates



Molecular Compounds

-composed of two or more non metals
-low melting and boiling points
-share (not exchange) electrons
-usually end in gen ex. oxygen hydrogen nitrogen
-7 molecules are diatomic (they are the same element)
Ex. H2, N2, F2, O2, Cl2, Br2, I2
-2 of them are polyatomic
Ex. P4, S8

To find out more information about naming molecular compounds check out this website http://www.ehow.com/how_2386009_name-molecular-compounds.html

Examples:

Name the Following Compounds:

N2O4 - Dinitrogen tetraoxide

CS2 - Carbon disulfide

P4O10 - Tetraphosphorous decaoxide

Write the Chemical Formulas for:

Nitrogen trichloride - NCl3

Sulphur dibromide - SBr2

Dihydrogen oxide - H2O

IUPAC Names 5 important ones only

Water - H2O
Hydrogen Peroxide - H2O2
Ammonia - NH3
Glucose - C6H12O6
Sucrose - C12H22O11

Chemistry October 1/09

Today in chemistry we learned about protons, electrons, neutrons, and ions. Ions are atoms or molecules that have an electric charge on it, cations would be a positive ion, and anions would be a negative ion.

Heres a video giving you a good and clear understanding of what a proton neutron and electron is.



Protons:
the number of protons would be the atomic number in an element. By adding or removing neutrons there would be no change in the element. Protons have a positive charge and can be located in the nuclei

Electrons:
electrons are 1800 times smaller then the protons. Chemical reactions occur between electrons in different atoms and compounds. Electrons have a negative charge and can be located outside of the nucleus

Neutrons:
By adding or removing neutrons there is no change in the element because neutrons and protons almost have the same mass. Neutrons have no charge and can be located inside of the nuclei with the protons.



We also learned about the periodic table today. Families form vertical columns in the periodic table, while periods are horizontal rows. Elements change from metal to non-metals going from left to the right. All the elements of a family have similar traits and characteristics to one another.

Another thing we learned today in chemistry 11 is that some elements have a latin and a english name examples would be:
copper - cuprum
gold - aurum
iron- ferrum
lead - plumbum
silver - argentum


We also learned about naming the formula for the compounds examples would be:
potassium iodide - KI
barium chloride - BCl2
lithium bromide - LiBr
sodium hypochlorite - NaCLO

Here is a link to find out more about how to write and name formulas for compounds. This neat little link also gives you a practice test once you have read and understand how to name formulas for compounds.
http://www.geocities.com/tjachem/symname.html

Chemistry Class September 29/09

Today in chemistry 11 we learned that matter can be divided into either a heterogeneous mixture or a homogeneous mixture. In a homogeneous mixture there is only 1 visible substance for example water and graphite are homogeneous substances. Heterogeneous substances contain more than one visible component, examples would be chocolate chip cookie and granite.

We also learned that there are two types or pure substances. Elements is one of them which is a substance that can't be broken down to simpler substances by a chemical reaction. Examples would be oxygen, iron, magnesium. The other pure substance is a compound, which are substances that are made up of two or more elements and can
be changed into elements or other compounds by chemical reactions. Examples would be water, chocolate chip cookies, and sugar.
It is very difficult to know the difference between a compound and a element you can only see the difference on an atomic level. There is one method though, this method is to connect the substance to an electric current. This technique is called electrolysis it will split the compound apart into its consistent element.

Down Here is a video explaining Heterogeneous and Homogeneous mixtures.

Chemistry September 24/09

Today in Chemistry 11 we learned what matter was in general and in how many various states it can exist. ex. Solid, liquid, gas, plasma. We also learned 2 new forms, Aqueous and amorphous, where aqueous is something dissolved in water, and amorphous is something lacking definite form (shapeless). We also learned about the strength of the bonds between the forms and the 3 categories of changes:
-physical change ex. crushing tearing etc.
-chemical change ex. cooking rusting etc.
-nuclear change ex. no example, but assuming nuclear decay
We learned that in physical change, the shape or state of matter changes, and that changing from a solid to gas is often confused as a chemical reaction. We then learned that in a chemical change, a new substance is formed, the properties of matter change, and that it could change the conductivity, acidity etc. We also briefly learned the law of conservation of matter.


Quick summary:
-what is matter? Anything that occupies space and has mass
-Matter can exist in many forms: solid, liquid, gas, plasma, amorphous, aqueous
-bond strength from strongest to weakest: solid, liquid, gas, plasma
-Physical change is the change of shape or state of matter
-Chemical change is the change of the properties of matter and when new substances are formed
-Law of conservation of matter states that matter, like energy, is neither created or destroyed in any process

heres a neat explanation of a solid, liquid, gas website online
http://sciencecastle.com/sc/index.php/articles/view/2100
and heres one talking about chemical and physical changes
http://www.chem4kids.com/files/matter_chemphys.html

Chemistry September 22/09

Today in class we had our chapter 1 test. It covered:

percentage error
scientific notation
conversion
WHMIS
lab materials
significant digits

Chemistry September 18/09

In chemistry 11 today we did a lab. The lab was to find out how much salt can dissolve in 40ml of water. We had to put salt in bit by bit into the water and stir it, until we could not see anymore salt at the bottom of the water. When salt still remained at the bottom that means that the water could no longer dissolve anymore amount of salt. The amount of salt that could be dissolved in 40ml of water was 11.5grams of salt. The procedure for this lab can be found on the following website.
http://www.ehow.com/how_5074919_dissolve-salt-water.html

We also learned about the basics of graphing today. We would have to know this because Mr.Doktor had wanted us to graph our results on paper. He specifically said on your graph you must have at least 5 basic elements, those are:
1. Title
2. Labeled axis
3. Correct scaled axis
4. Data plotted correctly
5. Best-fit line

Chemistry Class Number Two September 16/09

Today in chemistry we learned about conversions(dimension analysis) and the SI unit system.
Example: 1L --> ?mL 1L = 1000ml
to find out more about conversion you can have a look at the following website.
http://www.purplemath.com/modules/units.htm

	SI unit system
Giga-	billion	G	109
Mega-	million	M	106
kilo-	thousand	k	103
centi-	hundredths of	c	10-2
milli-	thousandths of	m	10-3
micro-	millionths of	µ	10-6
nano-	billionths of	n	10-9
pico-	trillionths of	p	10-12

We also learned that we will be doing a lab next class on how much salt can dissolve in 40ml of water. We are learning more and more each chemistry class and things are geting harder to do. We will need to know how to do these conversions for our chapter test on tuesday.

Chemistry Septermber 14/09

In our first class of chemistry we learned about significant digits.
For example: 3.14= 3 significant digits
2.0= 2 significant digits
2 = 1 significant digit
We also learned about scientific notation.
For Example: 145,000,000,000 = 10<sup>11

To learn more about significant digits you can have a look at this website.
http://www.physics.uoguelph.ca/tutorials/sig_fig/SIG_dig.htm
To find out more about scientific notation you can check out this website.
http://www.nyu.edu/pages/mathmol/textbook/scinot.html</sup>