Today we had another lab which featured finding and calculating the mass of various objects such as anhydrous salt, empty crucible, water given off, etc. To begin the lab we had to wear the awesome looking safety goggles. Then ms. chen previewed over what we should do, and everyone i mean everyone ran to get all the supplies. Each group of two acquired a pipestem triangle, iron ring, stand, and bunsen burner. We set up the equipment so that the bunsen burner is directly underneath the crucible which is on the pipestem triangle that is held midair by the iron ring locked onto the stand. Before starting the lab we had to turn on the bunsen burner by connecting it to gas and lighting it up in order to dry out the crucible that is placed on the triangle. Once dry we weigh the crucible and record it. Shortly after we obtain some anhydrous salt equal to 1/3 of the max volume the crucible can hold and weighed and recorded them down again. Then comes the fun part........heating stuff up
anhydrous salt
Now we had to heat the anhydrous salt for 5 min. During the heating you can see that the anhydrous salt which WAS BLUE TURNED GREENISH WHITE. After the 5min and seeing cool colour changing, we let it cool and recorded its mass. This process was repeated to see if the mass were the same. Then we had to do the last step in our procedure, add water to the dehydrated salt. As soon as we added a few drops of water.......IT TURNED BACK TOO NORMAL OMGWTH!! NO WAI!!! Then we recorded down this amazing discovery and dumped it in the garbage. Then we filled in the rest of the information we needed to fill in with the data from the lab. As soon as we were done we handed our lab filled with all our data and observation to ms.chen. It was fun
Does that mean that the compound is twice as expensive and has absolutely no difference whatsoever? (distastefully biased plug on organic foods) No, it means that the compound contains carbon. BUT it is noteworthy that certain compounds containing carbon are indeed inorganic, like carbides and carbonates. In this lesson we learn how to calculate the empirical formula of such compounds with chemistry's favorite furry little critter, MOLES.
The smell of vinegar is caused by acetic acid (CH3COOH). It is an organic compound
We know that the empirical formula will contain the elements C and H because they are common in all organic compounds. What we do not know, are the amounts of each. If we let x represent the number of carbon and y represent the number of hydrogen, we can use a method to solve for each. THERE ARE 5 STEPS THAT ONE MUST TAKE TO ACHIEVE THIS GLORY. Actually 6, but the 6th is situaltional. It is also noteworthy that when an organic compound combusts, the compounds carbon dioxide and water are ALWAYS formed.
1st STEP:
Convert the mass of the CO2 and H2O
2nd STEP:
Find the moles of C and H from the moles of CO2 and H2O
There is 1 carbon in CO2 and 2 hydrogens and H2O so basically multiply the value of H2O in moles obtained from the previous step by 2
3rd STEP:
Divide both your values by the smallest molar amount
4th STEP:
Change the ratios until a whole number for both is reached by multiplying by 1 (2/2, 3/3, 4/4) ETC
The empirical formula will be represented by the fraction obtained
5th STEP:
Convert moles back to grams and add them to Check your work
Confused? EXAMPLES HO!
A 3.79 g sample of an organic compound is burned to yield 6.61 g of CO2 and 3.59 g of H2O. Find the empirical formula.
STEP 1:
(6.61g CO2)(1 mol CO2/44.0g CO2) = 0.150 mol CO2
(3.59g H2O)(1 mol H2O/18.02g H2O) = 0.199 mol H2O
STEP 2:
(0.150 mol CO2)(1 mol C/1 mol CO2) = 0.150 mol C
(0.199 mol H2O)(2 mol H/1 mol H2O) = 0.398 mol H
STEP 3:
0.150 mol C/0.150 mol C = 1
0.398 mol H/0.150 mol C = 2.65
STEP 4:
(2.65/1)(3/3) = 8/3 ----> 8 = y 3=x
C3H8
STEP 5:
(0.150 mol C)(12 g/1 mol)= 1.80g C
(0.398 mol H)(1 g/1 mol) = 0.398g H
1.80 + 0.398 = 2.198
BUT WAIT! 2.198 != 3.97.
EXPLAINED: Some organic compounds may contain oxygen aswell. THUS, the sample mass - mass of C and H = mass of oxygen
As the title suggested this class was more than just cool IT WAS AWESOME!! You won't be surprised because every single class of ours is super awesome because we are the awesome JOHNNY!!! (the Johnny part was a joke). okay back to our class~. Sooo what's did we learn today? Is that avocado lover with his not so cute moles planning for the next destructive weapon again??? The question remains to be answered but the important part is that we learned something in CHEM AGAIN~~~.
Okay the serious stuff is comming Empirical Formula mean The simplest whole number ratio of atoms of each element present in a compound. (WOW DOESN"T THAT SOUND COOL?)
Because the Pure awesomeness of Ionic compounds. They are given in Empirical formula already. (please don't go and ask the atoms why, they are just that awesome). The Covalent compounds are the strange ones they come in the right ratio but NOT SIMPLIFIED in math class that would be 1 mark off... so if a ionic compound went and did a math test with covalent the covalent would just barely pass while ionic gets perfect omg><.
Formula = n = molar mass of the compound/molar mass of the empirical formula
Okay now lets have some examples
A compound have 47.25% copper and 52.75% chlorine.
Find the empirical formula for this compound.
Cu 47.25g Atomic mass 63.6 47.25/63.6 = 0.74
Cl 52.75g Atomic mass 35.5
52.75/35.5 = 1.49
Cu 0.74/0.74 = 1
Cl 1.49/0.74= 2.01 = 2
So the empirical formula for Cu Cl
is CuCl2
Ex 2~ what is the molar formula? molecular mass = 132.16 , empirical formula = C2H4O
n = 132.16/42 = 3.2
MM = 42g/mol
MF= 3(C2H4O) = C6H12O3
Ex3 empirical formula of a compound is CH and the molar mass is 104 g/mol, calculate the molecular formula.
mass of C = 12.0 g/mol
mass of H = 1.01 g/mol
empirical formula mass = 13.0 g/mol
CH = (104 g/mol)(1 mol/13.0 g) = 8.00
It is exactly what it sounds like. Percentage Composition.
This class we learned about Percentage Composition and it is quite simple.
Here is an example
What is the percentage composition of NaCl?
First of all you must calcuate the total Molar Mass(MM) of this compound
Total MM of NaCl: 23.0 + 35.5 = 58.5g/mol
MM of Na = 23.0 g/mol
MM of Cl = 35.5 g/mol
From these numbers you calculate the percent composition of Na and Cl. (Round to 1 decimal place if the question does not give any numbers)
% of Na = 23.0g/mol
------ x 100 = 39.3% % composition = mass of element
58.5g/mol ---------------- x100
mass of compunt
% of Cl = 35.5g/mol
------ x 100 = 60.7%
58.5g/mol
From these two numbers, the % composition should add up to 100%.
Here is a video that explains another example of percent composition.
Here are the five questions and solutions we had to make.
1) What is the percent composition of Potassium carbonate? (K 2 CO3 )
Total MM of K2CO3 is :
K x 2 = 39.1 x 2 = 78.2 g/mol
C x 1 = 12.0 x 1 = 12.0 g/mol
O x 3 = 16.0 x 3 = 48.0g/mol
= 138.2g/mol
% of K = 78.2g/mol / 138.2g/mol = 56.6%
% of C = 12.0g/mol / 138.2g/mol = 8.7%
% of O = 48.0g/mol / 138.2g/mol = 34.7%
2. Sodium Bicarbonate (sodium hydrogen carbonate) is basically baking soda. Its formula is NaHCO3. Find the mass percentages (mass %) of Na, H, C, and O in sodium hydrogen carbonate.
Total MM = 84.0 g/mol
mass % Na = 23.0 g / 84.0 g x 100 = 27.4 %
mass % H = 1.0 g / 84.0 g x 100 = 1.2 %
mass % C = 12.0 g / 84.0 g x 100 = 14.3 %
mass % 3 * O = 48.0 g / 84.0 g x 100 = 57.1 %
3. Cetylpyridinium chloride, a common compound found in mouthwash, has a formula of C21H38NCl. What is the percent composition?
21 C = 252 u 38 H = 38.0 u 1 N = 14.0 u 1 Cl = 35.5 u
Total MM = 339.5 g/mol
mass % 21 C = (252/339.5)100 = 74.2% mass % 38 H = (38.0/339.5)100 = 11.2% mass % 1 N = (14.0/339.5)100 = 4.12% mass % 1 Cl = (35.5/339.5)100 = 10.5%
4. C9H11N2O4S is penicillin. Find the percentage composition for each Element.
We got quite a long time to study for the quiz. As soon as we walked in the cool teacher guy gave us time to study for our mole conversion quiz after we went over the homework. After studying we had about 1 hr of class left so that means we had 1 hr to do our quiz. About 20% finished quickly while the 80% including me finished near the bell. When we handed in our work, we got THE WONDERFUL WORLD OF MOLES YUPPY and other work. THE END to the super PRODUCTIVE day. I learned a lot yay.
Well, it is not so much harder as it is... more things to do.
We start, with a MOLE MAP
Made it myself. This is pretty much the entire lesson. If you find yourself nodding in satisfaction while reading it, you may stop here. If, however, you find yourself with either your left or right palm on your forehead, please proceed.
From the last post, you can learn how to apply these equations to convert your values. It is the basic conversion method. * Conversion factor refers to the value you must multiply by to obtain your initial number in converted units. It always has a division line with a numerator and denominator
1) Write your initial measurement
2) write the conversion factor so that the unit which you want to eliminate is on the opposite side. Ex if you wish to eliminate the mole unit, and it is in the numerator of your initial measurement, but the mole unit on the denominator of your conversion factor.
3) On the conversion factor, write the desired unit on the unfilled side of the division line.
4) On the conversion factor, put a 1 beside the larger of the two units. Beside the other unit, write the number of those units that go into the larger unit. Ex. 6.022 * 10^23 particles are in 1 mole, so put 6.022 * 10^23 beside the particles unit.
5) Cancel out the units that can be cancelled out, this leaves you with the desired unit
6) Multiply/Divide the numbers
7) Obtain new number in desired unit
However, new in this lesson, we have number of atoms in a particle.
To do this, count of number of X atoms in the formula and plug it in to where it says "# of atoms"
Ex. Convert 4.53 * 10^14 molecules of CO2 to amount of oxygen atoms
You count 2 atoms of O in CO2. Multiply 4.53*10^14 molecules by 2 atoms of oxygen/1 molecule.
4.53*2 = 9.06 ---> 9
9 * 10^14 atoms of oxygen
IT IS IMPORTANT TO KNOW THAT THERE IS NO SHORTCUT
In other words, there is no way to immediately jump from grams to particles. You must first convert to moles then to particles.
Ex. Find the mass of 3.67 * 10 ^11 molecules of phospate (PO4)
First convert to moles:
(3.67*10^11 molecules)(1 mole/6.022*10^23 molecules) = 6.0943208 *10^-13 moles
Now convert to grams:
(6.0943208 *10^-13 moles)(95 grams/1 mole) = 579*10^-13 ---> 5.79*10^-11 grams
wow... okay I just deleted everything I did.... fail..
So I'll start over again~
YES MOLES!!! THOSE LOVELY LITTLE ANIMAL!!! WITH WEIRD FACE(the once with the star ....)
ALSO THIS MUST BE REMINDED IT IS AVOCADO'S NUMBER BECAUSE THE PERSON WHO FOUND THIS NUMBER IS SO IN LOVED WITH AVOCADO AND HE NAME IT AFTER IT
Last class we studied converting moles to particles or grams
It was A REALLY FUNNNNN CLASS Since we learned something new~
yea since I deleted my intro before... I'll keep this one short
Okay so first convert particles to moles
If i have 5.0*10^12 particles of anything then how many moles would I have? The equation for this is (number of particles) * (1mole/6.022*10^23particles)
5.0*10^12 particles
1mole
6.022*10^23particles
So the answer would be 8.3*10^-12
Now let do it backwards from moles to particles
If I have 15 moles of carbon then how many particles would i have?
15moles
6.022*10^23particles
1mole
So I would have 9.0*10^24 particles for 15 moles of carbon
Now let convert Mole into grams
Since what ever the atomic mass is is the molar mass so carbon with an atomic mass of 12 would have a molar mass of 12mol/g
So lets have an example question
If I have 91 moles of Uuq then how many grams would it be?
91moles
289grams
mole
2.6*10^4 grams of Uuq if I had 91 moles of it
Okay now the other way around
If i have 5846 grams of Americium than how many moles would it be?
5846grams
mole
243grams
So I would end up with 24.1moles of Americium~
http://www.fordhamprep.org/gcurran/sho/sho/convert/molecalc.htm this is an extremely good source if you didn't get the lesson
This is a little game that I found on the internet that is about moles the site is below
http://nobel.scas.bcit.ca/chemed2005/tradingPost/TUPM_S2_4_15ChemFunGames.pdf
The Great Mole Relay Race
Purpose: Students will work as a team in a relay race format in order to solve 1-step and
multi-step problems involving mass, moles, and representative particles.
Materials: Whiteboards mounted on wall, dry erase markers, slips of paper with different
problems printed on each
Set-up: You will want to separate your class into teams of about 3-6 people each. Be
strategic in your team formation so that no one team has a lop-sided advantage or
disadvantage. Each team will use a different section of the whiteboard.
Place the problems which are on the different slips of paper in cups labeled for each
team. Each cup will contain the same five or six problems, but students will choose the
slips at random, so that each team will probably be working on a different problem at any
particular time.
As teams work out the problems you will want to make sure you have a clear answer key
already written out so that you can check their work.
Game Play:
1) On your signal, the first student from each team will pick a slip of paper from his
team’s cup. He will then write the problem on the board in any form he chooses
so long as the rest of his team understands the problem (note: no one else is
allowed to look at the slip of paper). Player 1 then sits down.
2) Player 2 then heads to the board and begins the problem, proceeding through the
first conversion factor.
3) Player 3 then heads to the board and continues the problem by writing the next
necessary conversion factor.
4) Player 4 will write the next conversion factor, or if there is no need for another
conversion factor, she must use her calculator to correctly compute the answer
with units.
5) Check the answer. If it is correct, the next person may begin the next problem. If
it is wrong the next person must go to the board and figure out what is wrong and
fix it. This requires each person to be engaged in the whole problem. Require
each person to write each problem on his or her own paper. Collect all of their
work at the end of the game.
6) The winning team is the one that finishes all of the problems the fastest.
Say What? Ewww who would want to learn about those big brown things.
NO! This is chem class.
Okay this guy is Avogadro. This guy basically invented the mole.
One of the most important things to note is his hypothesis and it is named after him
Avogadro's Hypothesis
Equal volumes of different gases at the same temperature and pressure have the same number of particles
Enough with this charming looking dude.
You are still wondering what a mole is right? Right, because I have not told you yet.
The Mole allows chemistst to count atoms and molecules.
Avaogadro's Number
Yes children, the mole has a number. And this number is just 1 mole
Sorry but i cannot tell you how big it is. It is big. Very Big.
Back to Avogadro's hypothesis if there are the same number of particles in atom , the mass ratio is due to the mass of the particles. (used for the relative masses of all atoms on the periodic table.)
The mass of 1 atom of the element is counted in atomic mass units or amu.
Eg. Fluorine = 19.0 amu (because its atomic mass is 19)
When you want to find the mass of an ionic compound, you use the formula mass.
E.g. Potassium(K)+Fluoride(F)
39.1 + 19.0 amu = 58.1 amu
For covalent compounds, you use the Molecular mass.
E.g. Carbon Dioxide
C O2
12 + 16.0x2
C02 = 44.0 amu
For pure substances, you use the atomic/molecular/formula mass in grams per mole)
Ex. 1 mole of oxygen is 16.0g/mol
Here is a song that we had in chem class talks about moles. Its fun.
But if you really want to learn about moles, watch this.
Today we did quite a lot! We worked VERY VERY VERY hard, we were given a review sheet on sci notation, sig figs and measuremants and uncertainty. Totally rad. The purpose was to help us get ready for the test. Btw Joe had the worst accuracy and jennifer has the worst precision TROLOLOL. Anyway we worked very hard on the review sheet VERY HARD LIKE GOOD STUDENTS and got most of the work done. Occasionally we got a wtf r u doin and l2work look but otherwise it was all groovy.
Cool story bro, nice review
After spending 4/6 of the class working on the review sheet, we started working on some real work. Plotting graphs on excel. We had to plot 3 graphs each on a different experiment. One included volume of gas while other used density, etc. We spent the whole time working very hard and plotting the graphs and answering questions on a hand out she gave us. Turns out the worksheet was not even collected. QQ i actually finished it.
This is basicly what we did but we chose scatter graph instead of column.
Monday, November 8, 2010
Last class, we did a lab as you may or may not have deduced from our previous post, which by the way, was the first lab blog not to include the words "super" and "fun" with the latter following the former. This was mainly due to the fact that we have banned Cole from making lab blogs.
ANYWAYS. We had a lab last class... BUT there was no lab report! WHAT A TWIST! Going all M. Knight Shyamalamadingdong on this blog, yo. INSTEAD, we had a lab quiz. BUT HOW?! How does one combine these two seemingly incompatible elements? By ingeniously transferring electrons from the metal element lab or La2+ to the gaseous element Quiz or Qu3-, the ionic compound Lab Quiz or La3Qu2! Actually, it was just a lab report disguised as a quiz, so I may have over hyped that a little.
The quiz consisted of hypothetical measurements and calculations with said measurements. Also, accuracy and precision was also discussed and a few hypothetical follow up questions that applied your knowledge of density and volume calculations. All of them must have incorporated scientific notation and significant digits, which was pretty much the underlying theme of the whole lab. It was more to test your sig-fig knowhow than to test your ability to make density calculations and weigh aluminum foil.
Apres le quiz, were taken to ze laBORaTORy of ze computers. There we made graphs from given data on a spreadsheet. The data was of the volume and mass of hot and cold water. Which this we found the density of both and the hot water was found to have slightly lower density than cold water. I concluded that this was because of the higher level of kinetic energy within the hot water, resulting in faster movement on more space between particles.
cool density experiment you can try AT HOME, except you probably don't have all this stuff
Today we had a really fun day since we just had another LAB.~ But other than the lab we didn't really do much. So late me explain what happened during the lab. At first we were told what to do. Get 3 pieces of alumium foils and measure them. Then we are suppose to put each aluminum foil on the centigram and find out the mass of the the aluminum foil. After we found out the mass the volume of the foil were stated in the lab book and a simple calculation was able to show us the thickness. We repeated the calculations 3 to find out the 3 aluminum foils thickness. In the end we all had to calculate the percentage of error. (estimate - actual) / actual * 100, in absolute value. And for the rest of the class we basically didn't do anything yay~ free block~.
So you're wondering, what is density and why is it important? Well since my vocbalulary is too intense, ill dumb it down a bit and sum it up in a few words. Density is a physical property of matter specified as mass per unit volume. It is important because if you have two objects that have the same size, you can differentiate the two by calculating density from the following pyramid.
Mass = Density x Volume
Density = Mass / Volume
Volume = Mass / Density
Being a student the stress with memorzing formulas, please consider this pyramid because it can save your life. It also works in physics.
Some ways to represent density:
For a solid : g/cmᵌ
For a liquid: g/mL
1cmᵌ of water = 1mL
*If you want to know if an object floats or sinks, look at this equation below.
ᵈobjects > ᵈliquid = sink
ᵈobjects < ᵈliquid =float
Now for an example to help you with your problems.
Ex. Calculate the density of a solid gold house that has a mass of 133337g and a volume of 137.7L
Today we learned about a very very interesting things called.....PRECISION, ACCURACY, AND ABSOLUTE + RELATIVE UNCERTAINTY. Notice it is in CAPS, it means that its very important.
Precision = is how reproducible a measurement is compared to other similar measurements
Accuracy = Is how close the measurement comes to the accepted or real value
(If you could use definitions then so can I)
For example if you the real value is 10 cm. If you measured 9.7cm then it is accurate however if it is like 50cm then its not very accurate. If you repeatedly measure and get the results of 9.8cm, 9.8cm, 9.9cm, 9.8cm then it is precise since the results can be reproduced. BUT if you are very bright and smarter than the average lad and measure 4cm, 9.7cm, over 9000cm then it is not precise but 9.7 is considered accurate unless you take all the numbers into account for average measurement then it is neither accurate or precise. If you dont understand read it over 50 times.
Then the next thing under CAPS is ABSOLUTE UNCERTAINTY. Im absolutely sure it would mean the same thing if absolute uncertainty is uncertainty but adding absolute made it sound more scientific. Anyway onto absolute uncertainty. Everything that is measured IS NEVER EVER EVER EVER NEVER EVER EVER NEVER EVER EVER exact ( except things you can count like human beings though you could be half of a person but that is debatable) There ALWAYS be a degree of some uncertainty. And we learned how to calculate the uncertainty below.
If you are not bored and didnt exit out of this site then read below then exit out of the site. HF
there are 2 ways to calculate uncertainty but know this, it is always expressed in the units of measurements such as M, cm, etc and never in a ratio.
METHOD 1: THE MATH METHOD
Make measurements (the more the better) and take out any data that looks off or your sure you measured wrong then calculate the avg. The absolute uncertainty is the greatest difference between the avg of the measurements and one of the measurements you made that are REASONABLE.
Trial Mass of a Swift Scout Holding 3 Mushrooms(kg)
1 20
2 19
3 20
4 over 9000
5 20
First you take the most unreasonable one. I would say the 20 but i am forced to take away over 9000 eventhough it seems most reasonable. Then you take the average of the numbers left over. The average is 19.75. Then you find the greatest difference with the reasonable measurement which would be 19.75-19 = 0.75. Then you write the answer like this. The mass of the swift scout holding 3 mushrooms is 19.75kg plus or minus 0.75kg. As long as the mushrooms dont blow up, your answer will be relatively correct.
METHOD 2: THE OTHER MATH WAY
uncertainty of instruments = Uncertianty of instruments that measure stuff
This is SUPER SUPER SUPER SUPER SUPER SUPER SUPER EASY. You take the smallest segment on your instrument and times it by 0.1 and that is your uncertainty for the instrument.
Let's say your're driving across the country in your Dodge Caravan, or Honda Civic if you want to promote the stereotype, and you come across a sign that says, "Entering Awesometown, population: 271,835". This figure is most likely COMPLETELY AND UTTERLY INCORRECT. But what if the sign said, "Entering Awesometown, population: 270,000". This figure is not entirely incorrect. By applying significant digits, this sign has become 68% more reliable.
A precise less precise number like 270,000 implies that the zeros are subject to change, whereas a number like 271,835 implies that the digits are there and cannot change and in most situations when dealing with inconsistent numbers like population, these precise numbers are incorrect.
HOWEVER, Significant digits or Significant figures or if you want to be super cool, Sig-Figs must be measured and have at least some certainty. The rule is that the last sig-fig in a number is uncertain while all other sig-figs are certain.
Ex. 5486; digits 5, 4 and 8 are certain while 6 is not
But how does one know which which digits are significant? Well, I have devised a set or rules, or commandments if you will for telling which digits are significant.
THE 5 COMMANDMENTS OF SIGNIFICANT DIGITS:
COMMANDMENT 1) All non-zero digits ARE significant.
ex. 7.342 has 4 Sig-Figs
COMMANDMENT 2) Zeros that trail after Sig-Figs and are before the decimal ARE NOT significant.
ex. 1340000 has 3 Sig-Figs
COMMANDMENT 3) Zeros that trail after the decimal point ARE significant.
ex. 54.00 has 4 Sig-Figs
COMMANDMENT 4) Zeros that come before Sig-Figs ARE NOT significant.
ex. 0.000078 has 2 Sig-Figs
COMMANDMENT 5) Zeros between other Sig-Figs are ARE significant.
ex. 5400.045 has 7 Sig-Figs
Exact Numbers
Some numbers have a set value assigned to certain things. One human has a set value of 1, if the subject is how many humans. We cannot say that there are 3.5 humans unless we were able to find 3 humans and half of one. If 3.5 humans were somehow calculated it would be rounded down to 3 humans.
Rounding
you will need to know your general place values, refer to this chart
If you have calculated an ugly answer such as 6542.5865 and you can only keep 2 significant digits, you must round your answer to accomplish this.
To do so, follow the basic rounding rules:
Look to the digit to the right of the desired place value that you want to round your answer to. If that number is larger than 5 round up and vice-versa. Simple, right? BUT WHAT IF THE NUMBER IS 5?!?!?!?!
CALM DOWN, BUDDY. First look to see if there are numbers after the 5. If so, round up because that indicates that the number is larger. If not, and the number is 5, this means that the number is exactly half-way, in which case, rounding it up wouldn't really be more correct or more incorrect than if you rounded it down. However, if you have a set of digits, you wouldn't want your data to be rounded up more than down. Round to the nearest even, or in other words, if the digit to the left of 5 is odd, round up, if it is even, round down.
Operations with Significant Digits
Addition and Subtraction:
line up your numbers like you would if you were doing old-school math, pre-calculator style with the decimals lined together.
779.233
+ 68.4
710.833
Examine the numbers: 779.233 goes to the thousandths place while 68.4 goes the the tenths place. ALWAYS PICK THE LOWER ONE. Round your answer to the tenths place. 710.8. Same rule applies for subtraction.
Multiplication and Division
44.56 * 846651.243 = 37726779.38808
Again examine your numbers and again, choose the lower one. This time, notice the number of sig-figs in each number. 44.56 has 4 while 846651.243 has 9. Round your answer to 4 sig-figs. 3773. Same rules applies for division.
SO THE CONCEPT IS: Lower amount of sig-figs means more reliable, because it reduces the chances of being COMPLETELY AND UTTERELY INCORRECT.
TL;DR: Watch this vid of Mr.Coolteacherguy with the pro hair talk about Sig-Figs.
Today we did another lab to see how to separate solutions in full action. At first Ms Chen told us to get all the equiptments that we need then we all had to WEAR OUR SAFETY GOGGLES ( Always have safe even when we are dealing with food dye. There might be a chance for it to accidentally color our eyes). As we moved on with the lab we all had 3 pieces of paper and we had to cut the piece of paper in to a pencil shape. Then with a pencil we marked a line 4 cm from the tip. Then the fun part begins we filled the test tubes with 2 cm of water and got a drop of food dye with the color of our choice (red yellow blue, BLUE FTW!!!) from Ms Chen. We put the chromatograph with a dot from the food dye into the test tubes. When we are finish the food coloring that Ms Chen give us we had to record it on the board and get new colors (Unknown and green) from the back.
After 20 minutes the food dye finish separating and the end results is food coloring that have been separated into it original colors. After we pretty spent the rest of the time doing the lab questions...
Since components keep their identities, you can devise a process that discriminates between components with different properties. There properties include:
Filtration: select components by particle size through a filter
Floatation: by density
Distillation: by boiling points
Chromatography: by attraction for a stationary phase
Hand Separation: (solids and solids): A mechanical mixture/heterogenous misture can be separated by a magnet or a sieve(sifter)
Evaporation: (solid dissolved in a liquid solution): boils away the liquid and leaves the solid behind.
Here is how each process works:
Filtration:(undissolved solids and liquid): Passes a mixture that contains solid particles through a porous filter. If the pores are smaller than the particles, then the solid particles stay on the filter.
Crystallization: (solid and liquid) A saturated solution of a desired solid is used. Which is then turned into a precipitate through chemical or physical change. The solids are then separated by floatation or by filtration. The liquid is then evaporated or cooled and the solids come out as pure crystals.
Gravity Separation:(solids based on density) A centrifuge whirls a test tube at high speeds forcing denser materials to the bottom. This works best with small quantities.
Solvent Extraction: A component movies into a solvent shaken with the mixture. Works best with one component that dissolves.
Mechanical Mixture: (solid and solid): liquid to dissolve one solid but not the other, so the desired solid is lieft behind or dissolved
Solution: solvent is insoluble with solvent already present. The solvent dissolves one or more substances and leaves the unwanted behind.
Distillations(liquid in liquid solution): Distillation is collecting and condensing volatilized components. Heating a mixture can cause low-boiling components to vapourize. The component that boils evaporates and enters a condenster; the gas the ncools and condenses back into liquid dropping as a purified liquid.
exmaple of distillation
Chromatography: There are many different types of chromatography and they include both the mobile and stationary phase. It is highly accurate and precise with its analyses and it is used to separate very complex mixtures (for eg. drugs, plastics, foods, pesticides). Then the separated components can be collected individually.
Sheet Chromatography: Paper Chromatography (PC) involves a strip of paper. The solvent is applied on the strip of paper and then the components appear as separate spots spread out on the paper.
OMG OMG OMG OMG OMG TODAY WE LEARNED ABSOLUTELY THE MOST FUN AND SUPER COOL THING EVER LEARNED IN CHEMISTRY....NAMING STUFF! OMG I KNOW!
First of all we learned what an acid is. An acid is formed when a compound composed of hydrogen ions and a negatively charged ion are dissolved in water. Ions separate when dissolved in water? I know so super cool.
ex. H + CL = HCL ---- HCL + H20-->H30 + CL H ion joins with H20 to form H30
Now the super super fun part and the main super super fun thing we learned....NAMING
First of all we learned how to name SIMPLE acids and they follow the following super cool rule:
1.Use "hydro" as the beginning
2.Last syllable of the non-metal is dropped and replaced with "-ic"
3.Add "acid" at the end
should be like this ____ide ---> hydro_____ic acid
Then the TOTALLY BEST part of class is that we got a worksheet. The worksheet TOTALLY DID NOT have a lot of questions and was TOTALLY FUN. SO FUN in fact that I am so hyped up talking about it. Who doesnt like doing 100 questions? Work is sooooooooooooooooooooooooooooo fun. This concludes my totally super super fun blog about this super super fun day......Yay...until next time...yay.
Tuesday, October 12, 2010
As you may recall, in our previous class, we did a lab. This class was basically split in to 2 sections. First we continued with our lab report and write up. My partner and I ran around the class looking for two other groups to compare with. Eventually, Ms. Chen stopped the absolute madness and commenced the lesson.
Writing and Naming Ionic and Covalent Bonds!
"Wait a minute! That's sounds somehow oddly familiar... didn't we cover the same subject in grade 9 and 10?"
"That's right Billy, and now we get to learn it AGAIN!!"
"Golly gee wilikers, what fun! I sure hope we get a whole bunch of redundant practice problems!
"197 of them, in fact!"
"HOOORRAYY ADVENTURE HO!!!!"
Ionic compounds
metal + non-metal (oppositely charged) electrons from metal transfer to non-metal Symbol: Criss cross apple sauce the charges, 1's need not write and simplify if possible
Ex.
Li1+ + O2---> Li2O
Zr4+ + S2- --> Zr2S4 --> ZrS2
Name: write out metal name, then non-metal name but with "ide".
Ex.
calcium + chlorine ---> calcium chloride
Usually metals near the middle of the periodic table. In this case, include the number of the charge being used in a roman numeral.
Ni2+ + Br --> nickel (II) bromide
Ni3+ + Br --> nickel (III) bromide
Back in the day, they would use "ic" to represent 2+ and ous to represent 1+ charge. For charges higher than 2, "ic" would represent the higher one while "ous" the lower one.
Ex.
Cupric = 2+
Cuprous = 1+
A common ionic compound, NaCl, sodium chloride or better known as salt.
Covalent Compounds
Unlike ionic compounds, covalent compounds share electrons and combine non-metals with non-metals Naming: Greek prefixes must be put in front of elements according to their charge.
1 - mono
2 - di
3 - tri
4 - tetra
5 - penta
6 - hexa
7 - hepta
8 - octa
9 - nona
10 - deca
Ex.
P4O3 --> tetraphosphorus trioxide
*Mono for the first element need not be written
CO2 --> carbon dioxide
Diatomics: certain elements can combine with another of themselves
H2 O2 F2 Br2 I2 N2 Cl2 Hofbrincl
Carbon dioxide can be found in our atmosphere. It is debatable whether it contributes to global warming.
Polyatomics
Groups of ions form together as one atom. Naming is simply writing the two compounds/elements involved Symbol: When criss-crossing, brackets must be put if there is an existing subscript.
Ex.
Calcium Nitrite --> Ca2+ + PO43- --> Ca3(PO4)2
For the rest of the class we finished our lab reports.
The polyatomic citric acid (C6H8O7) is used in orange juice
