CHEMISTRY SHENANIGANS

In today's Chem class, we had to clean up the cupboards and lab areas for the winter break. Also, we watched videos of parking fails and a swing fail. For homework, we were given a two page booklet about balancing equations.

If you're still a bit confused, take a look at these helpful hints!





And here's a video of a car reverse fail in case you missed out:


Have a Merry Christmas and enjoy the break!

In today's class, we discussed balancing chemical equations. Mr. Doktor reminded us of two key things: in order to balance a chemical equation, you must first have the correct unbalanced equation, and specify the state of matter, i.e. (g) for gas, (s) for solid.

To balance a chemical equation correctly, you must have the same number of atoms (or moles) of one element on one side as well as the other.

ex. (in-class) Mg_(s) + 0_{2 (g)} -> MgO
2Mg_(s) + 0_{2 (g)} -> 2MgO

*note that 2Mg + 0₂ are the reactants and 2MgO is the product.

ex. (in-class) AlBr₃ + SrCO₃ -> Al₂(Co₃)₃ + SrBr₂
2AlBr₃ + 3SrCO₃ -> Al₂(Co₃)₃ + 3SrBr₂

ex. CO₂ + H₂ -> CH₄ + H₂O
CO₂ + 4H₂ -> CH₄ + 2H₂O

ex. Mg + P₄ -> Mg₃P₂
6Mg + P₄ -> 2Mg₃P₂

If you want more practice on balancing equations, click here.

Also, keep in mind that in order to make the correct chemical equation, you will need your Peridoic Table of Elements. If you're feeling hungry, this might help.. it's the Periodic Table in cupcake form!!



And if you need some more help, here's a video!

Today, we had a first look into our new unit on chemical reactions. Mr. Doktor showed us 3 different demos on chemical reactions, one involving the bunsen burner and magnesium. We also learned a very important lab safety tip: never light a match when there is gas in the room. Towards the end of the class, Mr. Doktor told us to write out 6 ways in which we could tell a chemical reaction has occured. We came up with these: smell, bubbles, taste, colour, a precipitate is formed, gas, smoke and when new substance is formed.

Here is a video explaining the five major chemical reactions:

We had a chemistry test

Giving Directions

• find the mass you need
• con'c > moles > mass

Directions of Solutions

• when you add water, the con'c decreases
• if the volume is doubled, the con'c is halved
• when the water is heated and evaporates, the volume decreases and the con'c is doubled

Volume 2x = con'c 1/2 (when water is added)
Volume 1/2x = con'c 2x (when water is heated and evaporates)
Formula for both concentrations and volume:
C1 x V1 = C2 x V2
(concentration 1 muliplied by volume 1 = concentration 2 multiplied by volume 2)

• this formula is basically like math, isolating to get the variable by itself

Definitions
Solution: A homogeneous mixture
Solute: The one present in smaller amount
Solvent: The one present in greater amount
Concentration: Amount of Solute
Amount of Solvent

Some units for concentration
g, g, mg, mg
ml L L ml

The most common (and useful) units are
mol = Molarity = Molar Concentration
L

THE FOLLOWING ARE ONLY FOR AQUEOUS SOLUTIONS & DO NOT APPLY TO GASES
M= mol
L
mol= M(L)

L= mol
M

NOTE: Think triangle of density/mass/volume

and now an informational video..

Today was the last class before mid-term exams. In class, we discussed the empirical formula and went over what would be on the exam.

You need to know:

1. Binary Ionic
2. Multivalent
3. Polyatomic
4. Acids/Bases
5. Hydrates
6. Molecular Compounds
7. Classical Naming System

Mole Conversions

1. Mole Conversion Table
2. Mole to mass and volume (gases atSTP)
3. Density
4. Number of Molecules
5. Atoms

Significant Digits
SI system
Classification
Lab Safety & Labs

Here are some videos to help!

Molecular

• P₄O₁₀
• C₁₀H₂₂
• C₆H₁₈O₃
• C₅H₁₂O
• 
  
  N₂O₄
  

• 
  
  P₂O₅
  
  
  
• 
  
  C₅H₁₁
  
  
  
• 
  
  C₂H_6O
  
  
  
• 
  
  C₅H₁₂O
  
  
  
• 
  
  NO₂
  
  
  

ex. (in-class) A sample of an unknown compound is analyzed and found to contain 8.4g of C and 2.1g of H, and 5.6g of O.
C_8.4H_2.1O_5.6  <--------- this is wrong because it is not in whole numbers!

Element	Mass (g)	Atomic Mass	Moles
C	8.4	12.0 g/mol	8.4g ÷ 12.0 g/mol = 0.7 mol
H	2.1	1.0 g/mol	2.1g ÷ 1.0 g/mol = 2.1 mol
O	5.6	16.0 g/mol	5.6g ÷ 16.0 g/mol = 0.35 mol

ex. A compound was analyzed and found to contain 13.5 g Ca, 10.8 g O, and 0.675 g H. What is the empirical formula of the compound?

Element	Mass (g)	Atomic Mass	Moles
Ca	13.5	40.1 g/mol	13.5g ÷ 40.1 g/mol = 0.337 mol
O	10.8	16.0 g/mol	10.8g ÷ 16.0 g/mol = 0.269 mol
H	0.675	1.0 g/mol	0.675g ÷ 1.0 g/mol = 0.675 mol

Here's a helpful hint for putting your ratios into whole numbers:
If the ratio ends in... 
~0.5 multiply by 2
~0.33 or ~0.66 multiply by 3
~0.25 or ~0.75 multiply by 4
~0.2, ~ 0.4, ~0.6, ~0.8 multiply by 5

In today's class, we went over the percent mass of elements in compounds and did a few examples.

The percent composition: the percent mass of each element in a compound

How to find the percent composition:
1. Find the total molar mass of the compound.
2. Next, find the molar mass of each element. (Remember, if there is a subscript, you must multiply the element's mass from the Periodic table by that subscript.)
3. Finally, divide each element's molar mass by the toal molar mass.
Note: Do not round right away to a whole number!

ex. (from in-class) Find the percent composition of K₂Cr₂O₇
step 1:  2(39.10) + 2(52.0) + 7(16.0) = 294.2 g/mol (the total molar mass)
step 2 & 3:
2(39.10) ÷ 294.2 g/mol = K 27%
2(52.0) ÷ 294.2 g/mol = Cr 35%
7(16.0) ÷ 294.2 g/mol = O 38%

ex. What is the percent of carbon in glucose, C₆H₁₂O₆?
step 1: 6(12.0) + 12(1.0) + 6(16.0) = 180.0 g/mol (the total molar mass)
step 2 & 3:
6(12.0) ÷ 180.0 g/mol = C 40%

Here's a video to recap! Enjoy!

Today in class, we reviewed the past topics of conversions to and from Moles, Volume, Molar Mass, Density, and Atoms and Molecules. Don't forget about Mr. Doktor's chart!

Density <----> Molar Mass <----> Mole <-- (conversion factor is 6.02 x 10²³)--> Atoms and Molecules (the subscripts)
                                                  
Mole <--(conversion factor is 22.4 L/mol or vice versa)--> Volume                                                   

Below, we have some examples..

Ex. (in class) 1.25 L of an unknown gas has a mass of 3.47 g. What is the molar mass if it is 22.4 L/mol?

1.25 L x 1 mol  =  0.0558 mol
              22.4 L

Molar Mass  =    3.47 g       =  62.2 g/mol
                       0.0558 mol

Ex. (in class) 250 mL of a gas which is known to contain one sulphur atom and an unknown number of fluorides has a mass of 1.63 g at STP. Find the molar mass, then the number of fluoride atoms. (SFx)

a) the Molar Mass
0.25 L x   1 mol   =  0.011 mol
                22.4L

1.63g / 0.011 mol = 146.048 g/mol

b) the number of fluoride atoms
1(32.1) + x(19) = 146
                  19x = 146-32.1
                      x = 6
Ex.Calculate the volume of 11.2 mol of HCN (g) at STP.
11.2 mol x 22.4 L  =  250.88 L = 251 L
                   1 mol

Ex. What is the volume of 1.3 g of NO₂ at STP?
1N + 2O --> 1 (14) + 2(16) = 46 g/mol
1.3 g  x  1 mol   x   22.4 L  = 0.633 L
                46 g       1 mol


Here are two videos to help you review!

-For now, all bases will be aqueous solutions of ionic hydroxides -NaOH
-Ba(OH)
2

-Use the cation name followed by hydroxide
-Sodium Hydroxide
-Barium Hydroxide
Example: Write the name of the following
-HI(aq)- Hydrogen Iodide
-H PO (aq)- Phosphoric Acid
3 4
-H PO(aq)- Phosphorous Acid
3 3
-HNO (aq)- Nitric Acid
3
-Mg(OH) (aq)- Magnesium Hydroxide
2
-HBr(aq)- Hydrobromic Acid

The Following are Oxalic Acids
-HOOCCOOH(aq)
2-
OOCCOO
2-
C O
2 4

-For monoatomic elements, a molecule= an element
-Diatomic element: Molecule (Cl2) Element (Cl)
-Molecules of Compounds H-O-H
--> molecule 2'H' atoms in 1 molecule, 1 'O' atom in 1 molecule
Example
Write ammonium carbonate (NH4)2CO3
N:2
H:8
O:3
C:1
Example
-moles<---->molecules
6.02x10^23 molec/1mol or 1mol/6.02x10^23molec

How many molecules are in 0.25 mol of CO2?
-0.25mol x 6.02x10^23molec/1mol = 1.505x10^23

5.1772x10^24 molecules of H2O = ? moles
-5.1772x10^24 molecules (1mol/6.02x10^23)=8.6mol

Find the number of 'H' atoms in 4.0mol of ammonia?(NH3)
- moles-->molecules--> H atoms
-4.0mol x 6.02x10^23/1mol= 2.41x10^24
-molecules x 3 = 7.22 x 10^24 'H" atoms

2Fe + 3CuCl₂ ---> 3Cu + 2FeCl₃

In our lab groups we dissolved copper chloride into water, added two iron nails into the solution, removed the two iron nails from the solution and dried them, and weighed the mass of the copper chloride, solution, and two iron nails each time a condition or characteristic of the material changed. Later, we discussed our recorded results with one another and handed in a lab report as a group.

Atomic Mass:
-The mass of 1 mole of atoms in an element
-The mass of 1.0 mol of 'C' atoms is 12.0g
-The mass of 1.0mol of 'Ca' atoms is 40.1g

Molecular Mass:
-The mass of 1.0 mole of atoms of an element or compound
N2, O2, H2, Br2, Cl2, F2, I2, P4, S8
-Assume all the rest are monoatomic
Element Symbol Formula Atomic mass Molar mass
Iodine I I2 126.9 g/mol 253.8g/mol
Silicon Si Si 28.1g/mol 28.1g/mol
Hydrogen H H2 1.0g/mol 2.0g/mol
Iron Fe Fe 55.8g/mol 55.8g/mol
Neon Ne Ne 20.2g/mol 20.2g/mol

Finding the molar mass of compounds:
-H2o
----> 2 H = 2(1.0) =2.0
----> 1 O = 1(16.0) =16.o
---->=18.0g/mol
-Find the molar mass of ammonium phosphate (NH4)3PO4
---->3N=2(14.0)
---->12H=12(1.0)
---->1P=1(31.0)
---->4O=4(16.0)
---->=149g/mol

Converting mass <----> moles:
-conversion factor g/mol or mol/g
-Find the mass of 2.5 mol of water
H20----> 18.og/mol 1mol/18.0g X 2.5mol =1/45g=45g
-Find the number of moles in 391g sample of nitrogen dioxide
NO2----> 1N=1(14.0), 2O= 2(16.0), =46g/mol
391 X 1mol/46g=8.5mol



In case you still don't understand, here's a short video:


And this is video is just for laughs:

1 Mole = {6.02 x10²³} known as Avogadro's number
= 602 000 000 000 000 000 000 000
.. therefore, 1 mole is just a really big number!

Hydrogen Bomb Equation
2H₂ + O₂ -> 2H₂0
is actually..
(2 H₂ molecules) and (1 O₂ molecule) = (2 molecules of H₂0)
is also..
12.04 x 10³ + 6.02 x 10³ --> 12.04 x 10²³
is also..
2 mol of H₂ + 1 mol of O₂ --> 2 mol of water

How big is Avogadro's number?
if 1 mole = 6.02 x 10³ was represented in dollars and divided among every person in the world..
about $1.0 x 10¹⁴ -> each person would have $100 trillion!!!!

How Gases Combine
- John Dalton looked at the masses of gases such as H₂, N₂, and C in comparison to O₂ but unfortunately, discovered there was no pattern.
- Joseph Gay-Lussac combined gases based on volume.
ex. 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
He discovered that gases combine in simple, whole number ratios.

Avogadro's Hypothesis
- he came to believe that equal volumes of any gas at a constant temperature and pressure contain equal numbers of molecules (but have different mass)
ex. [(*) to represent one molecule]
H₂ (*)(*)(*)(*)
O₂ (*)(*)(*)(*)

And in case you still don't understand the mole, here is a video to help:

Acids
-solid, liquid or gas at SATP (25°, 100kPa)
-form conducting aqueous solutions
-turns blue litmus red
-dissolve in water to produce H+
-taste sour

-turn red litmus paper blue
-slippery
-nonconductive
-dissolve in water to produce OH-

- acids are aqueous (dissolved in water)
- hydrogen compounds are acids
-HCL_(aq) ---> Hydrochloric Acid
-H₂SO_4(aq) ---> Sulfuric Acid
-Hydogen appears first in the formula unless it is part of a polyatiomic group
-CH₃COOH_(aq) ---> Acetic Acid
-classical rules use the suffix ic and/or the prefix hydro-
ex. Hydrochloric acid
-IUPAC systme uses the aqueous hydrogen compound
ex. HCL_(aq) ---> Aqueous Hydrogen Chloride

-for now, all bases will be aqueous solutions of ionic hydroxide
-NaOH
-Ba(OH)₂
-use the cation name followed by hydroxide
-sodium hydroxide
-barium hydroxide

EXAMPLES:
-HI_(aq) hdroiodic acid
-H₃PO_4(aq) phospheric acid
-H₃PO_3(aq) phosphorous acid
-HNO_3(aq) nitric acid
-HNO_2(aq) nitrous acid
-Mg(OH)_2(aq) magnesium hyfroxide
-HBr_(aq) hydrobromic acid
-HOOCCOOH_(aq) oxalic acid

In additon to the notes we took today, Mr. Doktor also showed us a demo for acids and bases. We had baking soda ready to neurtalize the extremely, strong smell it produced.

Heres a video on how to make a ph indicator with red cabbage:

In today's class, we talked about hydrates and molecular compounds, as well as how to name them and the IUPAC formulas.

Hydrates: are compounds that form lattices which bond to water molecules, and the crystals that contain water inside can be released by heating

If a compound does not have water, it is often preceded by "anhydrous" (i.e. copper sulfate anhydrous)

How to name hydrates:

1. Write the name of the chemical formula
2. Add a prefix indicating the number of water molecules
3. Add "hydrate" after the prefix
    

ex. Cu(SO₄) · 5H₂O _(s) = copper II sulfate pentahydrate
ex. Nickel II sulfate hexahydrate = Ni(SO₄) · 6H₂O _(s)

The Prefixes in order to name the number of molecules:
Mono =1
Di = 2
Tri = 3
Tetra = 4
Penta =5
Hexa = 6
Hepta = 7
Octa = 8
Nona = 9
Deca = 10



Molecular compounds: are composed of two or more non-metals; they have a low melting point and boiling point; they share electrons; usually end in -gen, or -ine

• Diatomic molecules: (when 2 of the elements are the same) H₂, N₂, O₂, F₂, Br₂, I₂
• Polyatomic molecules: (S₈, P₄)

How to name molecular compounds:

1. Write the formula for the least electronegative ion first, then the formula for the most electronegative ion.
2. Criss-cross the charges, moving the numbers below.
3. Reduce the ion numbers to the lowest common multiples. Note: You do not need to write the subscript '1' or the ionic charges.

When writing a molecular compound back into words, remember to write the prefix in front in front of each element (mono, di, or tri, etc.). You MUST do this for both the first and second part of the compound. However, if there is only one of the first element, you do not write the prefix 'mono'. Also, do not forget to add the ending of 'ide' to the second ion.




Here's an example of a molecular compound. Look closely, and you'll see that the compound is bonded together due to the SHARING of electrons rather than the exchanging of electrons.



In addition, we learned some IUPAC formulas:

IUPAC NAME        FORMULA
Water*                         H₂0
Hydrogen Peroxide*     H₂0₂
Ammonia*                    NH₃
Glucose*                      C₆H₁₂0₆
Sucrose*                      C₁₂H₂₂0₁₁
Methane                       CH₄
Propane                       C₃H₈
Octane                         C₈H₁₈
Methanol                      CH₃0H
Ethanol                         C₂H₅0H

*important formulas to know for tests and future use


Here's a video about gas hydrates, talking about how ice can be used for fuel!

In today's class, we discussed atoms, ions, elements and compounds as well as how to name ionic binary compounds.

Things to know:

• An element is a substance that cannot be separated into simpler substances by a chemical change.
• A compound is a substance that contains two or more elements combined in a fixed proportion.
• A hydrate is any class of compound containing water. The procedure of how to name hydrates is in the following post.
• A multivalent element has more than one ionic charge. In the periodic table, the most common ionic charge will be placed first for the element.
• An atom is the smallest particle of an element that retains the chemical identity of the element; made up of negatively charged electrons, positively charged protons, and uncharged neutrons. Its structure:
• An ion is an atom or group of atoms that has a positive or negative charge because it has lost or gained electrons.
• The difference between an atom and an ion: an ion is an atom or group of atoms, normally electrically neutral, that has lost or gained one or more electrons.
• An ion consisting of a single atom is called a monatomic ion, and an ion consisting of multiple atoms is called a polyatomic ion. Larger ions containing many atoms are called molecular ions.

Remember, in chemical formulas, the ion charge is indicated by a superscript (a small number to the top right), and the number of ions is indicated by a subscript (a small number to the bottom right).
How to name ions:

• for metals, use the name of the element and write the ion charge
• for non-metals, remove the original ending and add 'ide' and write the ion charge
• for polyatomic ions, there are special names

How to name binary ionic compounds:

1. Write the formula for the cation first, and the formula for the anion.
2. Criss-cross the charges, moving the numbers below.
3. Reduce the ion numbers to the lowest common multiples. Note: You do not need to write the subscript '1' or the ionic charges.

Naming polyatomic compounds is quite similar to that of naming binary ionic compounds. The steps are exactly the same. The only difference is that there will be more than one element written in one ion. Don't get confused!

How to name multivalent ions:

• the more common charge will be placed on top in the periodic table
• remember to use roman numerals in parentheses to show the charge

The Periodic Table of Elements to help you name the compounds, which includes SOME of the elements' ionic charges:


Additional help on how to name ionic compounds:

In today's class, we learned about homogeneous and heterogeneous substances, as well as solutions, pure substances, and how to tell the difference, and mixtures and how to separate them.

• Homogeneous substances consist of only one visible component (i.e. distilled water, oxygen) 
• Heterogeneous substances contain more than one visible component (i.e. chocolate chip cookie) We also learned about what pure substances are. There are two type of pure substances, elements and compounds.
  
• There are two types of pure substances:
1. Elements: substances that cannot be broken down into simpler substances by chemical reactions (i.e. oxygen, iron, magnesium)
2. Compounds: substances that are made up of two or more elements and can be changed into elements or other compounds by chemical reactions (i.e. water, sugar)
• How to tell the difference between an element and compound:
1. The differences are really only visible on an atomic level
2. Electrolysis: to connect the substance to an electric current, and split the compound apart into its constituent elements
   
• Solution: a homogenous mixture of two or more substances that may not always involve liquids (i.e. fog, steel)
1. Solvent: the component present in a solution that is in greater amount
• water is the most common solvent
• the symbol (aq) is used when something is dissolved in water
2. Solute: the component present in a solution that is in lesser amount
   
• Mixtures: many are easy to identify, but others are easily confused as pure substances
  
1. In heterogeneous mixtures, the different parts are clearly visible (i.e. granite, sand, fog)
2. In homogenous mixtures, the different parts are not visible (i.e. salt water, air, brass)
3. Separating mixtures is a physical change. There are many methods to separate mixtures, depending on the type of mixture:
• by hand, filtration -> are for heterogenous mixtures only
• distillation, crystallization, chromatography

MATTER

What is Matter?
-Anything that has mass and occupies space

-Matter can exist in many different states, the most common are:
a)solid,liquid,gas
b)plasma, aqueous,amorphous
Solids: Holds one shape and has a definite volume (has strong bonds)
Liquid: Can change shape, but has a definite volume (has weak bonds)
Gas: Can change shape and volume (has no bonds)




CHANGES IN MASS
-Matter can undergo many changes
-Nearly all changes can be broken down into three catagories:
-Physical Changes
-Chemical Changes
-Nuclear Changes



Physical Change
-Involves changing shape or state of matter (Crushing, Tearing, etc.)
-No new substances are formed (Boiling water, Cutting wood, Smashing cars)

Phase Changes
-Changing from a solid to a gas can often be confused as a chemical change
-Chemcials remain the same
-During the melting process chemicals usually follow this path

Chemical Change
-New substances are formed
-Properties of the matter change
-Conductivity
-Acidity, colour etc.
-ex. Iron rusting, burning wood, digesting food

Conservation of Matter
-In physical and chemical changes, matter is neither created nor destroyed. Ever! PERIOD!
- This is called the Conservation of Matter



Thus ends the lesson of the class..

Today, we did our Unit 1 Test in class. For homework, we had to read a bit about the periodic table as a preview for the lesson.

In today's class, the labs were finished up by all groups. It was also the last class before the Unit 1 Test, so we were reminded of what to study and know for next class. This would include WHMIS symbols, safety in the laboratory, scientific notation, dimensional analysis, unit conversions, SI units and metric prefixes, inaccuracies in measurement, as well as graphing.

Below are some useful pieces to know and help study for the upcoming test:



A tidbit on scientific notation..




Some examples on conversions and dimensional analysis..





For useful study tools in SI units and metric prefixes, visit the past post titled Measurement & Chemistry. Good Luck!

Last class, we did our first lab.

Just a reminder of what NOT to do in the lab..


Problem: What is the maximum amount of Sodium Chloride you can dissolve in 200mL of
water?



We completed 4 trials one with 10mL, another with 20mL, one with 30mL and the last one with 40mL. After figuring out the maximum mass of salt we can dissolve in each amount of liquid, we graphed the points. From there we then tried to estimate the maximum amount of Sodium Chloride we can dissolve in 200mL of water from a line of best fit.

Our results:

Trial	Volume of Water (mL)	Mass of Salt (g)
1	10	0.70
2	20	1.83
3	30	2.67
4	40	3.62

Today in class, we learned about unit conversions. We also learned about scientific notation and looked through the equipment in our class. In addition, we did an experiment with a balloon blowing out fire when it gets heated with some elements. Just like converting currencies in Chemistry, it is usually necessary to convert between units. This process is called Dimensional Analysis and is demonstrated below.

Today in class, we were taught the importance of measurements in Chemistry, as well as a bit of background information on the common systems of measurement. While the Metric system was founded about 30 years ago in France, the most common system we use today is the SI system. In addition, we discovered the 7 fundamental units of Chemistry as well as several prefixes and SI units.

7 FUNDAMENTAL UNITS:


PREFIXES & SI UNITS: