Lab 4C- Formula of a Hydrate

OK, so last Tuesday, all we did was have our quiz on percent composition, empirical formula and molecular formula.
Then last Thursday, we did Lab 4C- Formula of a Hydrate. This lab was done to determine the percentage of water in an unknown hydrate and to determine the moles of water presnt in this unknown hydrate.

First of all, we heated the crucible with a bunsen burner for a couple of minutes to make sure it was dry.

Then, we let the crucible cool down and weighed the empty crucible. We then added the hyrdate to the crucible and weighed it again. After, we heated the crucible until it was dull red on the bottom and waited for it to cool down and weighed it again. We repeated this step again- the second reading of the mass hydrate should have been within 0.03 grams of the first reading.

Finally, we added a few drops of water into the crucible and observed the changes.

Calculating the empirical Formula of Organic Compounds!

Yummmm, organic beets!

The empirical formula of an organic compound can be found by COMBUSTING the compound (reacting it with oxygen). The mass of the products can then be measured to find how much of each reactant was present. Recall: The Law of Conservation of Mass: the mass of the reactants will always equal the mass of the products!

Let's try an easy example: 

A 7.30 gram sample of a hydrocarbon is burned to give 23.8 grams of CO₂ and 7.30 grams of H₂O. What is the empirical formula?  

_{Notice that all of the C's and all of the H's went into making the carbon dioxide and water!} 

Step 1: Convert the grams of carbon dioxide and water into moles

Mol CO2 = 23.8g CO2 x (1Mol CO2)/(44.0g CO2)* = 0.541 mol CO2

* This is the molar mass of CO2

Mol H2O =7.30g H2O x (1Mol H2O)/(18.0g H2O)* = 0.406 mol H2O

* This is the molar mass of H2O

After this step, you can conclude that 0.541 moles of CO2 and 0.406 moles of H2O were produced.

Step 2: We want to isolate the carbon in CO2 and the hydrogen in H2O, these are the elements that make up organic compounds.

Mol C = 0.541 mole CO2 x (1 mole C)*/(1 mole CO2) = 0.541 mol
*Number of moles of C in 1 mole of CO2

Mol H = 0.406 mole H2O x (2 mole H)*/(1 mole CO2) = 0.812 mol

*Number of moles of H in 1 mole of H2O

After this step, you can conclude 0.541 moles of C and 0.812 moles of O were in the original organic substance.

Step 3: Recall how to find the empirical mass.

Divide both number of moles by the smallest molar amount (in this case it's C > 0.541)

C> 0.541/0.541 = 1 (Now read this: We have to multiply this because we multiplied 1.5)

H> 0.812/0.541 = 1.5 (Read this first: We have to multiply by 2 to make this into a whole number)

So, the empirical formula would be C2H3

Step 4: Check your answers!

Convert the moles of C and the moles of H to grams - they should add up to 7.30g

0.541 moles C x (12.0g C)/(1 mole C) = 6.49g C

0.812 moles H x (1.0g H)/(1 mole H) = 0.812g H

6.49g C + 0.812g H = 7.30g

If after checking your answer, you find that your masses do not add up, realize there must be a component of oxygen present in the compound. 

The mass of O = Mass of the compound - Mass of C + Mass of H

The mass of oxygen can then be converted to moles. Reapply step 3.

Some easy review for the quiz - http://lhs2.lps.org/staff/sputnam/practice/UnitV_EmpForm.htm

Here is a step by step tutorial on how to find empirical formula and molecular formula or the "real formula."

 

Empirical & Molecular Formula

Empirical formula: is the lowest ratio of atoms/moles in a formula.

Note: All ionic compounds are in its empirical formula

An example would be...

H2F10 is a molecular formula which reduces into HF5 which is the empirical formula!

For example, we were given  58.5% of carbon, 7.3% of hydrogen, and 34.1% nitrogen. What is the empirical formula?

* Assume that there's 100.0g*

1) Convert grams --> moles

C: 58.5 g x 1 mol = 4.88 mol               
                     12.0 g
H: 7.3 g x 1 mol  = 7.3 mol
                   1.0 g
N: 34.1 g x 1 mol= 2.44 mol
                      14.0g
2) Divide all 3 by smallest molar amount.
C: 4.88/2.44= 2
H: 7.3/2.44= 2.99 --> 3 (Round it only if it's super close!)
N: 2.44/2.44= 1
Empirical Formula= C2H3N
3) Scale ratios to whole numbers if needed. (For this example, it isn't needed..but if for example you ended up with 1.5, multiple it by 2,3,4,5, etc until you get a whole number.)

Here's a nice little flow chart to show you the steps:

 

Molecular Formula (MF): multiples of an empirical formula & shows the actual number of atoms that combine to form a molecule.
To calculate the MF: 
                         n=      molar mass of compound       
                               molar mass of empirical formula 
 n= a whole number multiple of the empirical mass
So then... MF= N x empirical formula
Ex. A molecule has an empirical formula of HO and a molar mass 51.0 g. What's the molecular formula?
Mass of HO= 17.0 g
51.0 g= 3
17.0g                               
Using the above formula of MF=N x empirical formula --> 3x(HO)= H303
Ex. A compound contains 8.96 g of Nitrogen and 5.19 g of Oxygen. The molar mass of the compound is 88.0 g. What's the molecular mass?
N: 8.96g x 1 mol = 0.64 mol

                    14.0g

O: 5.19g x 1 mol  =0.32 mol
                    16.0g
Dividing by the smallest amount:
N: 0.64/0.32= 2                     Therefore the empirical formula is N2O.
O: 0.32/0.32= 1
Then using the molecular formula:
Molar mass of N2O= 44.0 g/mol
88.0 g/mol= 2                       MF= 2 x N20= N402
44.0 g/mol
Note: When writing down the molecular formula, if it doesn't form an ionic compound, write it alphabetically. 

Time for some random chemistry jokes? Kekekeeke.
Q: What is the chemical formula for the molecules in candy?
A: Carbon-Holmium-Cobalt-Lanthanum-Tellurium or CHoCoLaTe 
Ooooh..trickytricky. ;P
Q: What does a teary-eyed, joyful Santa say about chemistry?
A: HOH, HOH, HOH!
Q: What did the bartender say when oxygen, hydrogen, sulfur, sodium, and phosphorous walked into his bar?
A: OH SNaP!

Percent Composition

So what did we learn today? That's right! Percent Composition!!!

Percent Composition is the mass percentage of a element in a molecule. 

To calculate percent composition it's as easy as 1 2 3! There's no formula to memorize... you wanna know the secret to it? 

There's no secret! In fact you do it all the time to calculate the percent for your marks. But if you do insist a formula here it is: % Composition = mass of element      X 100%

                                                            mass of compound  

Now you know this not so  secret formula, let's try it out!

EX 1. Calculate % composition of C2H4O2

 Total MM = 60 g/mol (the mass of C2H4O2)

% of C = 24 g/mol X 100 = 40%

                60 g/mol

           

% of H = 4 g/mol X 100% = 6.7%

               60 g/mol

% of O = 32 g/mol X 100% = 53.3%

                60 g/mol

And that's it! It's that simple! Bet you can't wait to do more of these eh? Well, you're in luck here's a few more examples!

1) Calculate the % composition of FeO.

2) Calculate the % composition of H2O

3)Calculate the % composition of Co(NO3)2

4)What is the % composition of the underlined portion of this compound C8H9NO2

5) A compound has a mass of 51.2 g. It contains 20g of O, 19.2 g of S and a mystery amount of N. What is the % composition of N?

ANSWERS AND SOLUTIONS:

1) Total MM = 71.8 g/mol

% of Fe = 55.8 g/mol / 71.8 g/mol = 77.7%

% of O = 16 g/mol / 71.8g/mol = 22.3%

2) Total MM = 18 g/mol

% of H = 2g/mol / 18g/mol = 11.1%

% of O = 16.0g/mol / 18g/mol =88.9% 

3) Total MM = 156.9 g/mol

% of NO3 = 124 g/mol / 156.9g/mol = 79.0%

% of Co = 58.9g/mol / 156.9g/mol = 37.5%

4) Total MM = 151g/mol

% of C = 96g/mol / 151g/mol = 63.6%

% of H = 9g/mol / 151g/mol = 5.96%

% of N = 14.0g/mol / 151g/mol = 9.27%

% of O = 32g/mol / 151g/mol = 21.1%

5)51.2- 20-19.2= 12 g of N 

Total MM = 51.2 g/mol

% of N = 12g/mol / 51.2g/mol = 23.4%

"To mole or not to mole, this is the question"

On Thursday's class, we had a substitute. First, he gave us the answers to the mole conversion worksheet. Then, he gave us a couple of minutes to review before writing our mole conversions quiz. After the quiz, we worked on two mole worksheets until class ended.

So lets hear some jokes, shall we?

Q:What did Avogadro teach his students in math class?
A: MOLE-tiplication!

Q: What was Avogadro's favorite Indian tribe?
A: The MOLE-hawks

Q: Why did Avogadro stop going to the chiropractor on October 24th?
A: Cause he was only tense to the 23rd!

Q: How does Avogadro write to his friends?
A: By E-mole!

Q: Which tooth did Avogadro have pulled out?
A: One of his molars!

teehee.

Uh oh, the mole conversions are getting complex :O

As mole conversions get more and more complex, remember to draw a MOLE MAP!

This is a perfect example of a mind map to use.
To get from...
Grams to moles - multiply by 1mole/MMg (molar mass in grams)
Moles to molecules - multiply by 6.022 x 10^23/1mole
Molecules to particular kind of atoms in a particle (not present in this mind map) - multiply by number of the specific atom/1molecule

Particular kind of atoms in a particle to molecule - multiply by 1molecule/number of the specific atom
Molecules to moles - multiply by 1mole/6.022 x 10^23
Moles to grams - multiply by MMg/1mole 

http://www.fordhamprep.org/gcurran/sho/sho/convert/molews1.htm offers some extra practice!

Now a bit more on Mole and Avogadro!

FUN FACTS!
1)6.022 x 10^23 Watermelon Seeds would be found inside a melon slightly larger than the moon!
2) Mole Day is an unofficial holiday celebrated among chemists in North America on October 23, between 6:02 AM and 6:02 PM.
3) There are 3 types of moles that live underground in North America: Eastern Mole, Hairy-Tailed Mole and Star-Nosed Mole.

DID YOU KNOW?

There is a lunar crater called Avogadro - named after this fine looking lad!

I'm watching you.

Moooooole Conversions!

Last time we learned about molar mass, so now it's time to use it in conversions. Remember: Avogrado's number = 6.022 x 10^23 particles/mol
Remember: Put answers in the correct number of sigfigs.

1) Converting between particles <--> moles. In this conversion, we use Avogrado's number.

    a) Particles --> moles
    # of particles x                1 mol               =  # of moles
                                 6.022x10^23particles

Ex.  If there are 3.01 x 10^24 C particles, how many moles are there?
              3.01 x 10^24 C particles x                 1 mol                   = 5.00 moles of C
                                                                6.022 x 10^23 particles
    b) Moles --> particles

Another example:

If there are 0.75 mole of CO2, how many molecules are there?

       0.75 moles CO2 x    6.022 x 10^23 particles  = 4.5 x 10^23 molecules of CO2

                                                          1 mole

Now that you know how many molecules are in CO2, how many atoms of Oxygen are there?

4.5 x 10^23 molecules CO2 x   2 atoms of Oxygen     = 9.0 x 10^23 atoms of O

                                                        1 molecule of CO2

2) Converting between moles <--> grams. Remember in this type of conversion we use the molar mass.

    a) Moles--> grams
     # of moles x molar mass = # of grams
                                1 mole

Note: To find the molar mass, look at the periodic table for the atomic mass.


Ex. If there are 2.04 moles of Carbon, what is the mass?
             2.04 moles x   12.0 g   = 24.5 g of Carbon
                                      1 mole

    b) Grams --> moles

      # of grams x      1 mole      = # of moles

                                molar mass        

Ex. If there is 3.45 g of Carbon, how man moles are present?

                 3.45 g x     1 mole      = 0.288 mole of Carbon

                                     12.0 g

Another example: If there are 6.2 g of MgCl2, how many moles are present?

6.2 g x     1 mole      = 0.065 mole of Mg Cl2

                  95.3 g

Whewww, that's a lot of info..time for a chem joke!

What does Avogrado put in his hot chocolate?

Marsh-mole-ows!

KEKEKE. ;P

It's time for the MOLE :D

What you need to know is...
- equal volumes of different gases has a constant ratio.

Do you know who Amedeo Avogadro is?

It's this guy right here!

His hypothesis is that equal volume of different gases at the same temperature and pressure have the same number of particles! 

Like this! 

So if the particles are the same then the mass ratio is the mass of the particles.

ATOMIC MASS:

- The units for atomic mass are amu , u or daltons.

FORMULA MASS:

- All the units of all atoms in ionic compound are in amu.

MOLECULAR MASS:

- All the units of all atoms in covalent compounds are in amu.

MOLAR MASS:

- molar mass contains atomic/molecular/formula mass and are in grams per mole (g/mol)

       ex. 1 mol of potassium = 39.1

Any pure substances have the same numbers of particles.

This is Avogadro's number: 6.022× 10^{23  PARTICLES} 

                                                                     ^MOLE

FUN FACTS:

Speaking of MOLES... did you know that there's something called a MOLE DAY! 

There's also a "anthem" for that day too! 

Check it out!!!

OHHHH NOOOOOO.. test next class D;

Today in class, we received a review sheet to review for our test next Monday. The material in the test includes: sig figs, uncertainty, measurement, graphing, density, scientific notation and unit conversions.

We worked on the review sheet for a while and spent the last half hour in the computer lab making three graphs on excel and completing a worksheet about the graphs.

Chem Joke of the day!

Teacher: What is the formula for water?

Student: H, I, J, K, L, M, N, O.

Teacher: Thats not what I taught you..
Student: But you said the formula for water was H to O...

LAB 2E QUIZ and GRAPHING FUN!

Today there was a Lab 2E Quiz. It covered the experiment we finished the previous day, in which we determined the thickness of aluminum foil.

Some important formulas to remember:

Density = mass/volume

Volume = length x width x height (which also happens to be thickness)


Don't forget that 1cm^3 = 1mL

You can also refer to the previous post for the Percent Experiment Error formula.

NOW ON TO GRAPHING!!!

One of the fundamentals in sciences and physics is learning how to express your data. Today, we were able to create graphs using Microsoft Excel.

Here are some instructions accompanied by screen shots that I took:

1) Enter your x values.
2) Under the next column, start by entering an equation that relates x and y. In my case I stated that y=A3(representing the x values) + 2.

 
3) Enter. The y value in accordance to the formula should appear.

4) Next, click on the first y value box. Hover your cursor over the bottom right hand corner. A small plus sign should appear.


5) Drag the plus sign down so that the rest of the y value column is covered.

 6) Enter again. All the y values should appear.

7) Now, highlight all the y values and the x values and click on the graphing icon outlined in green.

8) Customize the graph as you would like it. Keep clicking next, fill in the title, and click finish.

 

9) To add a linear trendline, right click on one of the points and click "Add Trendline." Go to options and click on the box that says "Display equation on chart" if you wish to display the equation.

10) To customize the chart, you can right click anywhere and choose "Format Chart Area."

11) Do a little editing and TADA! You can make your graph beautiful :)

Another great tool that can be used to create graphs or do another neat things is GEOMETER'S SKETCHPAD.

Watch to learn how to use it!

It's...LAB TIME...again..

For some reason, I end up always doing the blog on the day we do a lab...
Anyways, the lab was determining aluminum foil thickness. For this lab, the thickness will be an indirect measurement because we can't directly measure the thickness since it's so thin. So to determine the thickness, we used 2 formulas.

Volume: V=LWH (Length x Width x Height)
Density: D=m/V (mass/volume)

The density formula is first used because we know the density of aluminum and mass. This gives us the volume. Next, is to figure out the height (thickness). We can do that by using the volume formula since we know the volume and length and width. An example would be:
If you're given this information...density of the aluminum foil is 2.70 g/cm^3, length= 15.57 cm, width= 13.56 cm, and mass is 0.85g, what is the thickness?

If D=M/V then it would be... 2.70g/cm^3= 0.85/15.57 x 13.56 x H

570.04884H=0.85

H= 0.0014911

= 1.49 x 10^-3 cm

Once, the thickness was determined for all 3 sheets of aluminum foil, we found the average. We gave it in proper scientific notation and with the right amount of sigfigs. Then we compared it to the accepted value by comparing how accurate we were and how precise. To find the percent experimental error...

 

On the other hand, here's a little cartoon to lighten the mood? =P

Guess What?? IT'S DENSITY TIMEE!!!!

Ready for some Density? OKay! 

First of all, the formula for density is.. Density =Mass

                                                                  volume

Other ways to write the formula are...

V = m       or       m = DV

       D

REMEMBER:
solid is g/cm3
liquid is g/mL
1 cm3 of water = 1mL

Density of water is 1000 g/L or 1.0 g/mL

If the d objeect < dliquid then it's floating

If the dobject is >dliquid then it's sinking

Here's an example question! 

In a balloon, helium occupies 3.5 L with a mass of 4.0 g. What is the density?

Ans: D = m/v = 4.0/3.5 = 1.1g/L

Now that you've done all these density things, it's time for something fun! Try this if you want! 

I'm Uncertain of Uncertainties..

Measurements:
No measurement is exact; each measurement is only an estimate which means that there is some sort of uncertainty involved. An exception to this is when you are able to physically count the objects. Counting people would be a good example of that.

Uncertainty:
There are two types of uncertainties; Absolute Uncertainty and Relative Uncertainty.

Absolute Uncertainty:
   -expressed in units of measurement and not in ratios
First Method of Absolute Uncertainty:
1) Cross out unreasonable data
2) Calculate the average of the other measurements
3) The absolute uncertainty is the largest difference between the average plus/minus the lowest or highest reasonable measurement
Example: The following measurements are in cm:
20.0, 20.2, 20.3, 20.5, 23.0
First of all, you want to cancel out the 23.0 because it has a big difference from the other measurements.
Then, you calculate the average from the rest of the numbers which turns out to be 20.25
Now the measurement that will give you the biggest difference will be 20.5 so 20.5-20.25 = 0.25
The answer is 20.25 ± 0.25

For example, since '4' is the uncertain digit and is in the second decimal place, the uncertainty will be in the second decimal place as well.

Second Method of Absolute Uncertainty:
1) Determine how many increments the instrument go up by
2) Determine one tenth of that measurement
Relative Uncertainty:
-Relative Uncertainty can be expressed as:
     a) percentage (%)
     b) by using significant figures
                                                          Absolute Uncertainty
    Relative Uncertainty =           --------------------------------------
                                                        Estimated Measurement

SIGFIGS! Kekeke..

An Introduction

Significant figures are all the number of digits reported in a measurement, including all certain digits in a measurement plus one uncertain digit (always the last digit).

The more precise a measurement is, the more significant digits it will have, but the last digit is always the one uncertainty, it is always a measurement.

Rules:
1)All numbers (excluding zero) are ALWAYS significant. e.g. 1234 has 4 sigfigs
2)Zeros at the beginning of a number are NEVER significant, they merely indicate the location of the decimal point. e.g 0.0000012 has 2 sigfigs
3)Zeros at the end of a number (BEFORE the decimal point) are NEVER significant. e.g. 50000 has 1 sigfig
4)Zeros at the end of a number (AFTER the decimal point) are ALWAYS significant. e.g. 5.000 has 4 sigfigs
5)Digits between two significant digits are ALWAYS significant. e.g. 5000.3 has 5 sigfigs

Some quantities are exact and require no rounding - especially quantities that pertain to real life examples: number of sheep, number of coins, number of students, etc.

Rounding Rules, oooooohhhhh...
1) To round, always look at the digit to the right of the one you wish to round
2) If that digit is greater than 5, round up. If that digit is less than 5, round down
3) If that digit IS 5, and there are nonzero digits after it (symbolizing that it is in fact MORE than half), round up
4) if that digit IS 5, and ends at five, round so that the last digit is even (either up or down depending on the situation) e.g. 1.235 rounded to the nearest hundredth is 1.24

Adding and Subtracting
1) Round the answer to the fewest number of decimal places. e.g. 1.234 + 567.98=569.21 < We can only be accurate to the hundredth position

Multiplying and Dividing
1) Round the answer to the fewest number of sigfigs. e.g. 1.3 x 15462.2 = 2.0 x 10^4 < When we tried to round 20 100.86 to two sigfigs, our result was 20 000. This only has one sigfig, so the best way to fix this would be to express 20 000 in scientific notation, thereby making sure that two sigfigs (2.0) are represented.

KEKEKE.

Separation of a Mixture Through Paper Chromatography

From the previous post, you can tell that there are many ways of separating a mixture...but today we focused on paper chromatography. =)
So what happens is that the food dye (solute) was spotted onto a piece of filter paper ("stationary phase"). It was then put into water (solvent). The solute was then moved by the solvent which acts as a moving carrier("moving phase"). After approx 20 min or so, the spot is spread out on the piece of filter paper in bands.  The different bands or spots on the filter paper are the separated substances. We then calculated the Rf values.


Calculations we used today:
Rf= d1/d2
Rf --> "Ratio of fronts" (ratio of distance traveled by the solute to the distance traveled by the solvent) Varies from 0-1.
d1= distance of solute
d2= distance of solvent

This is an example of how the lab would of looked like. It starts off with the spot (in this case, it's the ink spot). Over time, the water travels up the filter paper and spreads the solute and will show the different components making up the original spot. Then we would calculate the Rf values to determine the identities of those components. 

Chromatography can be used  in detection & measurement of pesticides in foods, separating alcohol, amino acids, and sugars in plants, and others. It can even separate complex mixtures such as drugs or plastics. Even if the sample size is small, it can analyze it and will still be accurate and precise.

Separating Mixtures :D

Basis for separation: different components , different properties
Strategy: Divide a process that decreases between different properties and different components.

Separation:
- The more similar properties are it is difficult to separate them
ex. oil and water

Other Techniques of Separation:
Filtration: select components by the size of their particles
Floatation: select components by density
Crystallization and Extraction: select components by solutions
Distillation: select components by their boiling points

Hand Separation and Evaporation:
- Hand separation (solid and solid)
- Evaporation (solid dissolves in liquid)
- Boiling the liquid the solid remains as it is

Filtration:
- solids no dissolve in liquids

Crystallization:

- Solid and Liquid

- Solids are separated by filtration

- Solution is a solid

- Solid is cooled and creates pure crystals

- Crystals are then filtered again

Gravity Separation:

- solid based

- separation allowing denser components to settle

Solven Extraction:

- better if a solvent mixture dissolve in only component.

- MECHANICAL MIXTURES use liquids to dissolve solid

- Solution: the solvent will dissolve substances and will leave unwanted substances behind. 

Distillation:  

- Liquid to Liquid

- Low boiling points can cause a mixture to vaporize

- The liquid with the lowest boiling temperature will boil first

Chromatography:

- Different speed components can flow over the material at different speed.

- Are able to separate very complex mixtures (drugs, plastic, foods etc.

Sheet Chromatography:

1) Paper Chromatography:

        - Is in a stationary phase

        - appear as spots separated on a paper after it dries

2) Thin Layer Chromatography:

        - In a stationary phase that absorbs (AL2, O3, SiO2) [covering it with glass will absorb more quickly]

        - appear as spots on a sheet

                

                       

Naming ACIDS!

Acids are formed when a compound composed of hydrogen ions and a negatively charged ion is dissolved in water. Also known as aqueous (aq)

Guidelines for Naming Simple Acids:
1) Has to have the word "Hydro" as the beginning
2) The last syllable of the non-metal is dropped and replaced with "ic"
3) Have to add the word "Acid" at the end
*Skeleton Form*
________ ide --> Hydro______ ic Acid

Ex.  Name the Acids.
1) HCl --> Hydrochloric Acid
2) H₂S--> Hydrosulphuric Acid
3) H₂Se--> Hydroselenic Acid
4) H₂O--> Water- which is neutral

Naming Complex Acids:
1) ______ ate turns into  ________ ic
_______ ite turns into _______ ous
2) Add the word "Acid" at the end of it
"We ate ic-y sushi and got
appendic ite ous"
A way of remembering that "-ate" goes with "-ic" and
"-ite" goes with "-ous".
Ex.
1) HCH₃COO--> Acetic Acid (Also known as Vinegar)
2) HNO₂--> Nitrous Acid
3) H₂Cr₂O₇--> Dichromic Acid
4) HClO--> Hypochlorous Acid
5) HCN --> trickk questionn! This is a Simple Acid.. and the name for it would be "Hydrocyanic Acid"

Writing and Naming Ionic and Covalent Compounds!

Today, we reviewed how to write the chemical formulas of compounds and how to name them as well.

Ionic compounds are compounds made up of at least two particles that form bonds by giving or receiving electrons. Ionic compounds are often made up of metals (which give up an electron) and non-metals (which do the receiving).

E.g. 

Calcium gives up its two valence electrons (now it is stable) to two chlorine atoms (now they are stable as well). The charges balance out and calcium chloride is formed.

NOTE: Multivalent metals (metals that can form ions in more than one way) can often have names ending in -ic (for the higher charge) and -ous (for the lower charge)

For example: Cupric is Copper (II) and Cuprous is Copper (I)


When naming these compounds, make sure that the metal is mentioned first and that the suffix of the non-metal is changed to -ide.

Covalent Compounds share electrons. This bond occurs between non-metal elements.

Eg.

The hydrogen atoms share electrons with the oxygen atom creating a water molecule.

When naming, take into account these prefixes:

Here is a link to some practice worksheets:
http://misterguch.brinkster.net/pra_namingwkshts.html

GOOD LUCK!

AHHH! It's freezing..oh wait, now it's melting?

Ring stand and test tube.

Today we did a lab on matter and its changes: heating and cooling curves of a pure substance. (supposedly it was dodecanoic acid?). The purpose? To investigate the heating and cooling rates and determine the melting and freezing points of the substance.

First off, we started with our goggles! SAFETY EQUIPMENT FIRST EVERYONE!!! Anyways, we began the cooling of our substance. Placing the thermometer in the test tube, we lowered it into a beaker filled with cool water. We recorded the temperature at each 30 s interval until it reached 25 degrees Celsius. Immediately, we moved onto the next part...

For the heating process, we placed the beaker on a hot plate and then lowered the test tube into it. Once again, we recorded the temperature in 30 s intervals until it reached around 50 degrees Celsius. Also, we recorded any other observations. Once we finished, we cleaned up, washed our hands (don't skip this step! It's important..cause you never know what might happen if there's still some of the substance on your hands), and began to work on the lab reports. Most important step of the day was to NOT pour the substance down the drain...otherwise it would solidify and nothing would be able to get through that drain!

Next class, we'll be finishing off our lab reports and graphs!

TIME FOR.... FINDING MORE ABOUT MATTER!!!!

This is a graph that shows how each pure substance reacts after going through each level: 


A-B = SOLID
B-C = MELTING
(C-B = FREEZING)
C-D = LIQUID
D-E = EVAPORATION
(E-D = CONDENSATION) 
E-F = GAS
AB-EF = DEPOSITION
EF-AB = SUBLIMATION

A	-The solid state at any temperature below is below its melting point - The particles are packed closely together -The forces of the particles are so strong that the molecules can only vibrate a little!
A-B	- Heated molecules are then converted into kinetic energy. -This will cause the molecules to vibrate faster and therefore the temperature will increase.
B	- The molecules are still solid but will it will gradually start to melt. - The temperature will remain the same    = Liquid form
B-C	- Exist in both solid and liquid state - Temperature remains the same because it helps hold particles together. - This constant temperature is referred to the “melting point.” - Heat energy will absorb to overcome intermolecular forces, also referred as “Latent Heat of Fusion.”
C	- All molecules have been melted          SOLID à LIQUID!
C-D	-Molecules are still in liquid form but the temperature is increasing. - When the liquid molecules have heated it will move faster because kinetic energy is increasing.
D	- Exist in liquid state -Molecules have overcome the forces of attraction between particles in liquid. -Some molecules start to move freely -Liquid begins to turn into gas.
D-E	-Exist in liquid and gas -The temperature remains the same -Again, heat energy is absorbed to overcome intermolecular forces between particles of liquid rather then increasing in temperature.
E	-All liquid into gas           LIQUID à GAS!
F	-Gas has now absorbed the energy and the particles begin to move faster and freely! -The temperature rises and heating continues.

TIME FOR SOME TEXTBOOK READING!! YAY!..


Matter in the Macroscopic World:


Most of the time when we look at a painting or a artwork, we wouldn't think about the different strokes or dots this painting has. We just see it as it is... 
But what we don't see are actually patterns and dots and what the color affects are. 
In your Chemistry 11 textbook on page 25, look at figure 2-2. It's a artwork right?
So what do you see? 
Well, what we can see are plates, people eating, their face expressions and all that kind of stuff. But if we look closely to the picture with our magnifying glass or even a microscope, you'll see more then a cup or a chair. 
As scientist, we use these type of skills to OBSERVE and after we OBSERVE we EXPERIMENT.
But as Chemists, we ask ourselves "How is this different to another?" or "What does this have in common with the other one ?" We ask... Can this happen?, How is this different?, What does this have in common with..? etc.


What YOU know about Matter!


We live in a world where everything is made up of matter, like water is a liquid that can be stored in a solid container. 
Speaking of water... water is the most familiar kinda of matter but not all water is the same! Compare fresh water to salt water, muddy water or even rain! They taste, look boil differently! 


Purifying Matter:

Muddy water left in a cup for hours will separate into layers of dirt and clear water. So this water is not pure and could be defined as a mixture (--two kinds of matter that separate and show their identities... also said to be impure.)  How do we purify the water? How can we tell if certain water is clean? Tap water is an example, not all tap water are clean and how we purify that is by adding alum and line to the water and that will resolve into a jellylike substance material. There are mixtures in that formula that do not scatter light. Mixtures like salt/sugar water that look uniform, do not scatter light are called solutions. A process which you boil something until it's dry and crystal produce remains is called distillation. 

Characteristics of Pure Substances:

-Pure substance have a constant boiling point; mixtures do not.
-The temperature at which a liquid changes to a solid is called freezing point.
- The temperature at which a solid becomes a liquid is called melting point.


Chemical and Physical Changes:
-Chemical changes produces new substances/matter.

• This change is called decomposition because one kind of matter are separated into two or more kinds of matter.

-Physical changes creates no new substances/matter. ex. melting or cooking anything, it's irreversible!

Compounds and Elements:

-Electrolysis involves passing an electric current causing a substance to decompose to create more kinds of matter.
-It' is not practical to do a experiment in electrolysis (of sodium chloride) at school because the requirement of the temperature to melt the salt is too high and chlorine gas can produce toxin.
-Pure substances that can be decomposed into new kinds of matter are called compounds
-Pure substances that can not be decomposed are called elements

This is a video about mixtures and compounds!! Feel free to watch it :D

Compounds have a Definite Composition:

-An important difference between mixtures and compounds is that...

• Mixtures can have as many compositions has they please but compounds are only allowed to have a certain amount of composition.
• ex. Imagine your younger sibling got to have as many fish sticks s/he wanted and you could only have hlf a fish stick. How unfair is that?

This is called "LAW OF DEFINITE COMPOSITION"

However, there's a twist! 

There are some compounds that have two or more compounds with different proportions of the same elements are known as "LAW OD MULTIPLE PROPORTIONS."