More on Moles...

Well... I'm going to start off with a funny story... My friend VIVIAN CHENG thought the "moles" were the black dots that grow on your face... then she thought it was the ANIMAL mole..... until i explained to her...

Well...we're still learning about moles, and how to convert moles... as much as I dislike moles, we still gotta learn about MOLES!...
Keep in mind Avogadro's Number  is 6.022 x 10^23 --> particles/mole

Molar Mass

You can convert molar masses from either:

1. Particles to moles
or
2. Moles to particles

Converting molar masses is somewhat simple. It is just like doing unit conversions.

For example:
1) Convert 4.5 x 10^24 particles to moles
4.5 x 10^24 particles x 1 moles/6.022 x 10^23 particles = 7.5 moles




What I did, was I divided 4.5 x 10^24 by 6.022 x 10^23. The particles cancel out. And because there is only one mole, the answer is simply 7.5 moles.

*Remember that you have to count to the correct number of significant digits.

2) *(2 atoms) Convert 0.82 moles CO₂ to molecules

0.82 moles x 6.022 x 10^23 particles/ 1 moles =   4.9 x 10^23 molecules CO₂

The moles cancel out, so you simply multiply 0.82 x (6.022 x 10^23) particles . Because its  4.9 x 10^23/ 1, it works out to just be 4.9 x 10^23 molecules CO₂

There are two atoms, so there is one more step involved. Well...it's not anything special... but...here goes!

4.9 x 10^23 molecules  CO₂  x 2 atoms O/ 1 molecule CO2  = 9.8 x 10^23 atoms of Oxygen 

And again... the molecules cancel out, and you just multiply the numbers. 

This video may help explain some questions you  may have: 

Are you tired of converting moles yet? Well we're almost done, so bear with me!
Now.. we are going to learn how to convert moles to grams, and grams to moles. Get your periodic table for this!!! We need it !!  To find the molar mass of an element, look for the atomic mass of the element!

Example:

1) Convert  3.06 moles Fluorine to grams
The molar mass of Fluorine is 19.0g/mol

3.06 moles x 19.0g/1 mole = 58.1 grams of F

Here, the moles cancel out, and you just multiply 3.06 x 58.1 grams of F...and.. BAM, you get the answer!


2) To convert 7.65 grams of Fluoride to moles
The molar mass of Fluorine is 19.0g/mol

7.65grams x 1 mole/ 19.0grams = 0.403 moles F

                                                                Remember this guys! 


Oh...my... I am Finally done :) ok... well now that my suffering is over... your turn to suffer!!!!
Here are some practice sites:
http://misterguch.brinkster.net/conversionsworksheets.html
http://www.sciencegeek.net/Chemistry/taters/Unit4GramMoleVolume.htm

This is another video to reinforce the idea of converting moles..




OK DONE! BYEBYE! =]

The Mole

So today, we learned about moles. No, not the round thing on the skin....the Chemistry mole.


But first, let's understand a hypothesis of a great Italian scientist named Avogadro.
He proposed that equal volumes of different gases at the same temperature and pressure will have the same number of particles.


Now, let's define Atomic Mass.
This is the mass of 1 atom of the element in the units amu. (atomic mass units)
For example, if you look at the atomic mass of an element on the periodic table, it will give you how much 1 atom of that element weighs..


So, let's take Carbon. Carbon's atomic mass is 12.0 amu. This means that every carbon atom weighs 12 atomic mass units.


Now formula mass. This means you add up all the masses of the atoms in the formula of an ionic compound.


So.... let's say we have NaCl.


Na is sodium, and the atomic mass is 23.0.
Cl is chlorine, and the atomic mass is 35.5.


So by simple addition: 23.0 + 35.5 = 58.5 amu. This means that NaCl weighs 58.5 atomic mass units.


Molecular Mass: All the atoms of a formula in a covalent compound are added in amu.


So: Carbon Dioxide is
 C + O₂
12.0 + 16.0X2
CO<sub>2 = 44.0 amu


Molar Mass is the atomic, molecular, or formula mass of any pure substance, and the unit is in grams per mole.</sub>
For example, 1 mole of oxygen is just its atomic mass (16.0 amu, but in a different unit, g/mol.)



So 1 mole of oxygen = 16.0 g/mol
1 mole of carbon = 12.0 g/mol


They both have the same number of particles. Hence, the molar atomic mass of an element is the mass of 1 mole of that element.
It's just a unit that helps us and scientists count atoms and molecules, without actually counting every single atom.
And what Avogadro found, was that the number of particles in 1 mole of any amount of substance is always 6.022 X 10^23 particles per mole.



Now here's a video if you're still somewhat confused.







And here's a website to summarize the different types of masses to calculate.
http://www.mpcfaculty.net/mark_bishop/molar_mass_conversion_factors_help.htm

More Graphing

Well, next class (November 15, a long way I know. It's because of the Remembrance Day holiday coming up) will be our Chapter 3 Chemistry Test.

This test will focus on Significant Digits, Scientific Notation, Uncertainty, Density, the lab, and some conversions and graphing we learned last time.

We reviewed the unit, and then we went to the computer lab to do more graphing.

Please refer to the previous post for instructions on how you can make your own graph! :)

Anyways, got to go study now!

Graphing Analysis

In this Chemistry class, we went to the computer lab and made graphs that demonstrated the relationship between volume in hot and cold water.

Today, we learned how to make these graphs using OpenOffice; however, most of you have Excel on your computer, so that's the alternative.

First, you want to open up Microsoft Excel.

Next, you want to plug in a table of values into the cells.
Let's begin on A1, and labelling it as X values.
Then, put down some numbers for X. At the end you should get:

X values            
1
2
3
4
5

In the column beside (so starting with B1), label it as Y values.
In B2, enter an equation that can relate X and Y.
This equation must start with "=" then A_  (whatever cell number your values start at, in this case A2) then an operation.

So, we could have something in cell B2 like.....   =A2+5
Once the equation is set, you should see the number 6. (because 1+5 = 6)
So:

X values    Y values
1                6
2
3
4
5

Now, click on the cell with the y value 6. There should be a small box in the bottom right hand corner. Drag that all the way down until it covers the last y value.

You should see all the numbers plugged in by now.

X values    Y values
1                6
2                7
3                8
4                9
5                10

Now, highlight all the values, and click on the bar graph/chart option in the toolbar.
This will pop out a menu. In this menu, choose the scatter plot, and click next. Then, customize it to your liking, and fill in the x-axis and y-axis information, as well as a title.

After you click finish, a graph should pop out.



To add a linear trend line, right-click a point on the graph and click Add Trendline. You can even show the equation, which will give you the slope of this line you just graphed. (in y=mx+b form)

Adding a trendline:




The finished graph before you customize (i.e. changing colours, fonts):


Format your graph by changing the colours, fonts, etc.

AND You have now created a nice graph that shows your data clearly and effectively.

Another useful tool that can be used to make graphs or many cool geometric things, or even just having fun on the computer, is the Geometer's Sketchpad, introduced to me by my math teacher. :)

For more help on making a graph:

Lab 2E: Measuring the Thickness of Aluminum Foil

Lab 2E occurred on November 3, 2010. In this class we used basic math formulas to calculate the thickness of aluminum foil pieces. These formulas are. V=L*W*H,  D=M/V, or V=M/D. Basically each group weighed each piece of aluminum foil. We were then given the density of each piece, which is 2.70g/cm^3. Our next step was to measure the length and width of the aluminum with our ruler. After that we would find the volume by dividing the mass by the density. Once we got the voulme, we divided it by the length and width to get the height (aka Thickness).

We also used a formula to calculate if there was any errors with our test results.

experimental error = estimated measurement - accepted value x 100%

                       accepted value

Well that wraps about everything up. This lab was fun and easy. It was fun because we didn't have to wear any goggles. LOL, jokes.

PH

Density

Density is a material that is specified as mass per unit volume, or weight per unit volume. Now.. why is density important? Well, density is applied to many chemicals that are being tested. It can find out the buoyancy of an object or fluid. To calculate density, there is a formula...and here it is..

 

Density = mass/volume                       

           OR  

Volume = mass/density 

           OR  

Mass = Density x Volume

*Once you memorize one of them, you can move them around as you would mathematically to calculate density, volume or mass.

now... make sure you REMEMBER THE FORMULAS OF DENSITY!!!

To represent density:

solids - g/cmᵌ

liquid - g/mL 

If there is 1cmᵌ of water = 1 mL

then the density of water = 1.0g/mL

                                      = 1000g/L

Ex. Calculate the density of a liquid which has a volume of 28mL and a mass of 26.4g

density = mass ÷ volume

density = 26.4 ÷ 28 = 0.94 g/mL

Unless the fluids are mixed together, fluids that are less dense float on fluids that are more dense.

The Cup of fluids on the left have different density`s and as you see, the fluid on the top ( the clear fluid) is the least dense, the fluid on the bottom (the pink fluid) is the most dense, the fluid in the middle (the blue fluid) has density that is between the clear and pink fluid. The cup of fluid on the right can be mixed together with other fluid.

 

ᵈobjects > ᵈliquid = sink

ᵈobjects < ᵈliquid =float

For more practice, go to this site: http://serc.carleton.edu/mathyouneed/densitysp.html

watch this video for more help: 

What does it mean to be "uncertain"?

Before we move onto the answer to the question, let us define a few terms to better understand "uncertainty". 

Precision - this is how reproducible a measurement is compared to other similar measurements. In other words, the more decimal places you have, the more precise your number is. 

Accuracy - this is how close your measurement (or average measurement, in some cases) comes to the accepted/real value

Let's say you're shooting arrows at a target. If you shoot 5 arrows, and they all end up close to the bullseye, but all of them are far apart from each other, then the shots were accurate, but not precise.

If you shoot 5 arrows and they are far away from the bullseye, but are very close together, as if you shot all 5 of them at the same time into the same area, then your shots were very precise. 

If your 5 shots were both very close to the bullseye, and very close together (meaning you can hit the same area consistently), then your shots were both very accurate and very precise.

Now: 

Realize that no measurements are exact. Every measurement is just a best estimate, meaning there is room for error and "uncertainty". 

However, when you count a set of objects, it is exact. There are 5 humans. There can't be 4.34 humans.

Now that we know the background, what is uncertainty? 

A: Uncertainty is the margin of error, usually stated by giving a range of values that contains the real or true value.

There are 2 types of uncertainty: Absolute and Relative.

In absolute uncertainty, there are 2 methods. 

Method 1: First, discard any unreasonable data first. You must have at least 3 reasonable measurements  in order to use this method. Then, take the average of the measurements. Then, find the largest difference between the average and either the lowest or highest reasonable measurement. 

For example:

Trial 1 - 15.3 g

Trial 2 - 15.5 g

Trial 3 - 15.2 g 

Trial 4 - 11.9 g (remove it, because it looks very different from the rest of the data)

Average of the reasonable measurements: 15.3g

Difference between average and lowest number = 15.3g - 15.2 g = 0.1 g

Difference between average and highest number = 15.5g - 15.3g = 0.2 g

Absolute certainty based on average: 15.3 ± 0.2 g

Method 2: By using the uncertainty of each instrument

Measure to the best precision, then estimate to 0.1 of the smallest segment on instrument scale. 

For example, a ruler's smallest segment is 1 mm. The uncertainty would be the tenth of 1 mm, which is 0.1 mm. So the data recorded for a measurement with a ruler would be: 

 _____(measured value to the best precision) ± 0.1 mm

In relative uncertainty, it is the ratio of absolute uncertainty to the estimated measurement.

So for example, if the data was 39.3 ± 0.1 g, then the relative uncertainty would be 0.1 / 39.3 

This can be expressed in percentage, or in Significant Figures. 

The number of significant figures is the relative uncertainty. Refer to the previous blog post for info on Significant Figures.

Here's a video