MOLARITY!

Hi everyone! So last class we learned about the molar concentration aka molarity of solutions.

Recall:
- solution = a homogenous mixture where one substance is dissolved in another
- solvent = larger quantity; dissolves the solute
- solute = smaller quanitity; it's the amount of solute being dissolved in a certain amount of solution



Why do we need to know molarity?
ANSWER: To compare the amount of solute dissolved in a certain volume of solution!

"Molar Concentration" (aka molarity) is the number of moles of a solute in 1 litre of solution

• measures the concentration of solute in the solution
• "M" is the symbol we use to represent molarity
• "mol/L" is the unit

There are 3 different formulas (but as long as you know one, you can find the others):

1) MOLARITY = Moles/Litres
2) LITRES = Moles/ Molarity
3) MOLES = Molarity x Litres

Let's see some examples! 

1) Calculate the MOLARITY of a 1.5 L solution that contains 0.24 moles of HCL.

    Formula: M = mol/L
                  M = 0.24mol/1.5L
                  M = 0.16M of HCL

2) Calculate the concentration of 60 mL solution that contains 0.075 moles of NH4Cl. 

    Formulas: M = mol/L
                   M = 0.075mol/(0.060L)
                   M = 1 M  

3) What is the mass of solute for 125mL of 0.25 M Ba(NO3)2?

    Formula: mol = L x M
                  mol = 125mL x 1L/1000mL = 0.125 L
                  mol = 0.125L x 0.25 M
                  mol = 0.03125 mol Ba(No3)2 
                  0.031 mol x 261.3g of Ba(NO3)2/ 1mol = 8.2g of Ba(NO3)2 

 4) What volume in ml of a 0.25 mol/L NaOH solution is needed for 0.050mol of NaOH?


   Formula: L = mol/M
                 L = 0.050mol/0.25M
                 L = 0.20 L x 1000 = 2.0 x 10^2 mL



Here is an informative video on molarity to further educate you guys!

to be hydrated or ANhydrated? (LAB 4C)

So, over the course of two classes, we've once again did a lab!! This time, it was all about finding the empirical formula of a hydrate!

Now to give you a better understanding of this lab, here are some background information:

• A hydrate is a compound that contains a definite number of water molecules (H2O) 
• Without containing water, it becomes an anhydrous (without water) salt
• The formula of a hydrate is AB•xH2O where AB is the anhydrous salt and x is the number of moles of water per mole of the anhydrous salt

**always use crucible tongs and never touch or eat the substance!**


Here is the hydrate that we used:

Now in order to find the empirical formula, we have to find the mass of water that's in the hydrate by finding how much water was given off by burning the compound.
BUT WAIT! Before that, we must:
1. Find mass of empty crucible
2. Find mass of crucible + hydrate (before it's burned)
3. Subtract the mass in step 2 from the mass in step 1 to get mass of HYDRATE
Then, we heat the hydrate in the crucible until the substance changes to a white-ish colour. After letting it cool (SAFETY FIRST OF COURSE! - don't want to get your fingers burned) we weigh the crucible + anhydrous salt (as the water has been dried off). The mass should now be lighter than it was before. We do this process twice to make sure the water is completely gone! (the two masses should be within 0.03g)

(anhydrous salt of copper sulphate)

Next:
4. Using lowest mass of the crucible + anhydrous, subtract the mass of step 1 from this to get the mass of ANHYDROUS salt
5. Subtract the mass of the anhydrous salt from the mass of the hydrate to get mass of water that WAS in the compound


Given that the molar mass of the anhydrous salt (CuSO4) is 159.6g/mol, we can find the # of moles of the anhydrous.
mass of anhydrous salt x 1mol/159.6g = moles of anhydrous salt

Recall (from our quiz) that we have to find the ratio of moles of water to moles of anhydrous. Therefore, we find the moles of water from the molar mass of H2O:
mass of water given off x 1mol/18.0g = moles of H2O


Then, divide the two moles by the smallest molar amount (does this sound familiar now? :D)  and VOILA! You should get a whole-number ratio.


Thus, the empirical formula of this hydrate is CuSO4•5H2O (copper (II) sulphate pentahydrate)

Also, if you try adding water into the crucible with the anhydrous, an exothermic reaction will occur along with the changing of the color of the substance to the original hydrate color. Neat huh!

(teehee, goodluck!)

Organic Compounds - Formulablahblah

HAAII.

procrastinating for the win.

cats make me laugh so hard...

so last class we learned about empirical formulas, and this class we progressed to empirical formulas of ORGANIC compounds.  heres where things get tricky... :(

so a reminder: organic compounds MUST CONTAIN CARBON
-examples:  C
                    CH
                    CHO
CO is not one, because in an organic compound the hydrogen attaches to the carbon, and then the oxygen attaches to the hydrogen. NOT THE CARBON.

so to find the empirical formula of an organic compound, you need to first BURN the compound
second, you must collect and find the mass of the products (and convert to moles)

so say you have a mystery (organic) substance containing carbon and hydrogen, and when you burn it you end up with 10.0g of CO2 and 10.0g of H2O. Calculate the empirical formula.

STEPS:

1. Find the moles of the 2 compounds.

2. Divide each mole with the smallest molar mass.

3. If the resulting decimals are not even/close, multiply both substances until they are

(e.g. if you have 3.66mol H and 1 mol C, multiply both by 3 so you will have 11 mol H and 3 mol C)

4. If you are given the mass of the compound before it was burned, you should check the mass of the compound in the empirical formula to make sure it matches.  If not, then there might have been some oxygen in the original compound.

EXAMPLE

First Step: convert to moles

10.0g CO2 x 1/44 = .227mol

10.0g H20 x 1/18 = 0.556mol

Second Step: convert moles of compounds to C and H

0.227mol CO2 x 1mol C/1mol CO2 = 0.227mol C

0.556mol H2O x 2mol H/1mol H2O = 1.11mol H

and since there are 2H, then we multiply this by 2 to get a grand total of 1.11mol

Step 3: Now divide each by the smallest # of moles...

.227mol CO2 / .227 = 1

1.11mol H20 / .227 = 4.9, which rounds to 5 :)

so our empirical formula becomes CH5 

SUCCESS

Also remember to check your masses with the mass in the question (if stated) to check your answer! To do this, convert your moles of C and H back to grams and add them together. 

Total mass = grams of C + grams of H

If it's not equal (UH OH).. don't worry! There's just an extra step. Oxygen must be present, so the remainder of the grams is the grams of oxygen. Mass of O = mass of compound - mass of (C + H)

From there, convert the grams to moles and do step 3 again! Phew...

http://www.youtube.com/watch?v=9qGo9YTdnMg

^this is a video i found (and for the life of me, cant figure out how to insert...) and its kind of helpful especially if you want to isolate each element

as a reward, i has un autre chatton amusant <3

HAHAHAHAHAHAHAHAHAHAHAHAHAHAH TOOOO GOOOOOOOD

<3 Heather

Self Portrait:

sleep with one eye open...

Empirical & Molecular Formulas + Percent Composition!

Welcome back Beryllium Chemist followers! Today we will be talking about a few things. Firstly, the difference between empirical and molecular formulas. Then, we will move on to percent composition! (That's a lot!) I hope you'll survive through this blog!

In chemistry, there are two types of formulas that you MUST know: molecular and empirical

Empirical Formula

• shows the RATIO of the atoms or moles in LOWEST-TERMS 
• when you form ionic compounds on the periodic table, they are in empirical formulas since they're in lowest terms
• to get to empirical formula, divide molecular formula by its greatest factor
• you can calculate the empirical formula from percent composition and by mass
  
  eg. NaCl is an empirical formula but Na2Cl2 is not

Molecular Formula

• shows the total atoms each element are present in the molecule of the compound
• it is a MULTIPLE and not in lowest terms
• formula of the ionic and covalent compounds
  
  for example, C3H6 is a molecular formula because it's simplest form is CH2

Here is the key to converting:

Empirical Formula x whole number = Molecular Formula
Empirical Formula MASS x whole number = Molecular Formula MASS
Empirical Formula Mass (GRAMS) x whole number = mass of ONE MOLE of the compound


So if carbon and hydrogen are present in a compound in a ratio of 1:2, then the EF would be CH2. What is the molecular formula if 28.0g is the mass of one mole?

Step 1: Find the empirical formula mass.
12.0 + 2(1.0) = 14.0g/mol

Step 2: Divide the molar mass with the empirical formula mass.
(remember molecular formula mass = EF mass x whole number)

28.0 / 14.0 = whole number = 2

Step 3: Multiply whole number by the empirical formula.
Recall: Molecular formula = empirical formula x whole number

2(CH2) = C2H4


Now lets try some questions where we try to determine the EF with a given mass!
If you are given the question: Determine the empirical formula with the given 30.4g Nitrogen and 69.6g Oxygen. Follow through with these steps.

Step 1: Convert grams to moles.
Nitrogen: 30.4g x 1mole/14g N = 2.1714 mol
Oxygen: 69.6g x 1mole/16g = 4.35 mol


Step 2: Divide each molar mass by the smallest molar amount.
Nitrogen: 2.1714/2.1714 = 1
Oxygen: 4.35/2.1714 = 2.003
And voila! Your final empirical formula is NO2.


BUT, if after step 2 your answers don't round easily to a whole number, you have to multiply them by a whole number.
For example: If you have 1.7344 Carbon and 4.2332 Nitrogen

Multiply by 2, and keeping going till you get a number close enough to a whole number
Carbon: 1.7344 x 2 = 3.4688. 1.7344 x 3 = 5.2032. 1.7344 x 4 = 6.9376
Nitrogen: 4.2332 x 2 = 8.4664. 4.2332 x 4 = 12.6996. 4.2332 x 4 = 16.9328Therefore your empirical formula will be C7N17


Hope this helps guys:)
-Kimberly

MORE CONVERSIONS! (2-Steps!)

Are you overwhelmed with moles yet? Maybe this might brighten your day..

One day on the Tonight Show, Jay Leno showed a classified ad that read: "Do you have mole problems? If so, call Avogrado at 602-1023!" ha-ha! get it? 6.02x10²³ :p

Back to business..

Here's a mole map.. treat it as your BEST FRIEND!! (very helpful)

Today we learned how to do two-step conversions! Fun eh?
How would you convert particles to grams? Let's refer to the mole map!

you would have to go from particles → moles → grams

Eg. What's the mass of 4.35 x 10²¹ atoms of Au (Gold)?

4.35 x 10²¹atoms  x  1mol/6.022x10²³atoms  x  197.0g/1mol = 1.42g  of Au

** Make sure that you find the formula/molecular mass beforehand if it's a compound or it has more than one atom**


How would you convert grams to particles? Let's refer to the mole map! (again)
this time, you'd have to go from grams → moles → particles (opposite as the first!)

Eg. How many atoms of Au (Gold) are there in 5.72g of Au?

5.72g  x  1mol/197.0g  x  6.022x10²³atoms/1mol = 1.75 x 10²² atoms of Au

---------------------------------------------------------------------------------------------------------
Here's some extra practice for you!
1) What's the mass of 7.7 x 10²² atoms of Carbon?
2) How many formula units of potassium sulphide are there in 20.1g of potassium sulphide?
3) What would be the mass of 6.32 x 10²¹ molecules of 
Ca(NO3)2?
4) How many molecules of BF3 are there 3.3g of boron trifluoride?

Answers:

1) 1.5g
2) 1.10x10²³ f.u
3) 1.72g
4) 2.9x10²² molecules

And voila! You're an expert in mole conversions! (What a relief!) ;)

CONVERSIONS WITH THE MOLE >:(

What can i say? its just that good.

:D

So, what if you have this mole.

And you, being the nosy person you are, want to know how much it weighs.

So you put it on this scale.  (After you kick the sheep off, of course!)

So then you got the mole's weight in grams!! Yay!
...But wait. Teachers aren't that nice.  So picky picky, they want the weight in moles!! How ever will you do that?!?!

Well, with a handy-dandy calculator and the supreme knowldege of the mole conversion chart, you should be just fine.

Since I'm technologically challenged, this chart isn't very pretty.  Boo hoo.

Grams to moles: x (1mol / grams)
Moles  to grams: x (grams / 1 mol)

Moles to Formula Units: x (6.022x(10 to the 23rd) / 1 mol )   [hahaha my intimite knowlegde of blogger is clearly displayed here]
Formula Units to Mole: x(1 mol / 6.022x(10 to the 23rd))

F.U. (hah...) to atoms: x (#of atoms / 1 molecule)
atoms to F.U.: x (1 molecule / # of atoms)


*molar unit is g/mol
*label all your conversions incase you forget what you're doing and don't want to start ALLLL OVEERRR AGAAAAIIINNN

Memorising that chart would probably be helpful.  My advice would be to draw that chart at the top of the test paper, and refer to that throughout the test to avoid that awkward moment when you begin to doubt yourself and then you change all your answers and then get everything WRONNNGGG. Not that that has ever happened to mee... anyways if you have that memorised the conversions are pretty basic.

here are some examples for your pleasure...

convert 5.0g of carbon to moles.
                mass of carbon (according to periodic table) = 12.0
                5.0g C x (1/12) = .416666666666666666666666666666666666666666666
                and there were 2 sigfigs in the question... so your answer would be .42mol
easy peasy shampoo squeezy

lets try the other way:
20.0mol nitrogen to formula units
                              20.0g N x (6.022 x (10 to the 23rd)) = 1.20 x 10 to the 25th

comprendez-vous?
if not, ask someone else.

funneh kitteh.

no, not morbid at all.

-A

ha just kidding its heather. im just dying to finish this blog so i can watch pretty little liars hahahaha

All About Moles!

The Mole

.. and no, I'm not talking about the "beauty" mark or the animal!!!


Then what is it?!

it is a unit of measurement! (a counting unit) that allows chemists to count large quantities such as atoms and molecules by weighing its mass
- mole has same number of particles found in 12g of carbon-12




- Volumes that are equal of different gases have a constant ratio, 
 
  For example:   Oxygen : Hydrogen    16:1
             
when they have the same # of particles, the ratio of oxygen to hydrogen will always be 16:1      
 
Avogadro's Hypothesis

 
Avogadro's hypothesis states that equal volumes of different gasses, at the same temperature and pressure, contain the same number of particles.
 
- If they have the same number of particles, the mass ratio is due to the mass of particles.
- principle for the relative masses of all atoms are on the periodic table of elements


To understand the hypothesis better, imagine two boxes with oranges in one and apples in the other. You believe that there are equal number of fruits in each box. Not counting the masses of the boxes, you can find the mass of each of the fruit and determine the ratio between the oranges and apples. This is the relative mass since it is a comparison. 

Relative Mass
 
- Relative masses are expressed when comparing it mathematically to the mass of another object.
- Currently, we compare it to the mass of carbon-12 (1 atom is one twelfth the mass of carbon)


 
Molar Mass
 
- The molar mass is the mass of ONE mole of particles.
- The molar mass of an element is the mass shown on the periodic table which is expresssed in grams. (same value as molecular mass, formula mass, or atomic mass)
 
  For example:    1 mole of Oxygen = 16.0 g/mol
                         1 mole of Sodium = 23.0 g/mol
                          1 mole of CO = 28.0 g/mol

 
* All of these have the same number of particles.
 
- Molar atomic mass is the mass of 1 mole of that element.
- "Grams per mole" is the unit for molar mass.
 
Atomic Mass
 
- The mass of 1 atom of the element in amu (u)


eg. atomic mass of Iron (Fe) is 55.8u



Formula Mass
- Total mass of all atoms in the formula of an ionic compound
- in atomic mass units (u)
 
  For example:    Manganese Oxide
                               Mn           O
                            54.9    +   16.0
                             MnO     =   70.9 u
 
Moledular Mass
 
- All atoms of a formula in covalent compound.
 
  For example:    Nitrogen Oxide
                              N       O 
                           14.0  +  16.0
                           NO  =  30.0 u
 
Avogadro's Number
 
Do you want Avogadro's number?! 
 
- Well it's....
 
6.022 × 10^{23    =}  particles per mole   * memorize it!
 
- This is the number of particles in 1 mole of any amount of substances.

How big is it?


Have Avogadro's number of pennies placed in a rectangular stack of about 6m by 6m at the base and it would stretch 9.4 x 10¹²km and extend outside of our solar system. It'd take light approximately a year to travel from one end of the stack to the other!! That's a LOT of pennies!!

 

Now that you learned about mole, you should learn the mole song!
You will love it...

LAB DAYY! (DENSITY)

Our lab today was determining aluminum foil thickness.

The objectives for this lab was:

• to calculate the thickness of a sheet of aluminum foil and express the answer in the terms of scientific notation and significant figures.

All supplies that were used were 3 sheets of aluminum foil, 15cm by 15cm large, a metric ruler, and a centigram balance.

The steps to doing this lab correctly is very simple.

1. Take the sheets of aluminum foil and number them 1,2 and 3.
2. Use the metric ruler and measure the length and width of each sheet.
   * REMEMBER TO RECORD ALL YOUR DATA*
3. Then use the centigram balance to find out the mass of each sheet.
   *always use significant digits!*

Your table should have these labels: Sheet; Length(cm); Width(cm); Mass(g), and te numbers 1,2,3, for the aluminum foil.

To calculate for thickness, use the following formulas:

Volume = Length x Width x Height

Volume = Mass/Density

ex. 

• Take your information for sheet #1: length-17.11cm; width-17.91cm; mass-1.27g
• With the given density, find out the volume with the formula highlighted blue with green font.
• Then with the volume you just calculated, plug it into the formula in above in pink!
• After all calculations are done, voila! You have your height aka thickness.

To find the experimental error, here is another colorful the formula:

Experimental Error = |(your average thickness) - (given average thickness)| / (given thickness)

hope this helped guys:)

GRAPHING TIME! and Lab next day!

Fun comic of the day!

Ha-Ha! Hopefully that wont happen on the lab the next day...



Today, we went to the computer lab to... *drum rolls*


MAKE GRAPHS!
and interpret them...

Anywho, we made graphs using Microsoft Excel and it was FUN! We got to decorate and customize it however way we want! :) 
Why do we make graphs? Graphs are one of the key tools to present our data! It helps us visualize the relationship between the x and y values.

How do you make a graph on MS Excel, you ask?
Well here is a visual step-by-step procedure that may help!

1) On the second row of column B and C, make a table of values.

2) Click "Insert" at the top and then click "Scatter". From there, click on the first diagram.

 3) Now you will have a graph! Make a title and make sure to label your axis! Click on "Text Box" on the top to create your labels.

 4) Right click on one of the points on the graph then click on "Add Trendline". Examine the graph and determine what type of graph it is. In this case, it is "Linear". Also, click on "Display Equation on chart".

 5) You can play around with the graph by customizing it with shadows and colors. Right click on the chart or the plot (actual graph), then click on "Format". From there, play around with the designs and make your graph beautiful!

AND YOUR DONE! Wasn't that fun?!
Now to calculate the slope, just look at the equation! It's that easy :)

A REMINDER TO ALL THAT WE HAVE A LAB THE NEXT DAY! (YAY FOR LABS!)
Don't forget to do your PRE-LAB and also dress appropriately for the lab!!



Hopefully, nothing goes wrong and all will go well! ..... Hopefully.....

DENSITAAYYY

You are soo dense. ha.
Just thought I'd get that in there.



Density is a physical property of matter that measures mass per unit volume.
Anyways here are some formulas you maayy want to store in the old memory bank.  They may come in handy sometime. You never know.

Density = Mass/Volume

Volume = Mass/Density

Mass = (Volume)(Density)


Density helps determine different components in a mixture as the denser component will sink.
If the density of an object is more than the density of the liquid, it SINKS. If it doesnt, it FLOATS.
It is also handy to know that 1cm³ = 1mL so 1000cm³ = 1L (in case you need to convert!)
The density of water is 1.0g/mL or if you convert, it would be 1000g/L (1.0g/mL x 1000mL/L)

So here is an example of how you may see or find density in life. It is also the best experiment you will ever do. 

You're gonna need some OREOS (noooww i have your attention...), some milk (i used 75mL), something to put the milk in, a calculator, and a pen and paper.

So the first thing you need to calculate the density of an OREO is it's mass.  You can't really tell from the picture, but it says that 2 OREOS = 24g.  According to my calculator, that means each OREO has a mass of 12g.  Well that was easy.

The next thing you need is to calculate the volume.  I started with 75 mL of milk in the measuring cup shown below.  Again, stupid flash makes it difficult to read.  The smallest division on the cup was 25mL, I measured to the nearest 2.5mL.   I dropped in the cookie...

And voila!  From what I saw with my own eyes (not the camera) the cookie plus milk had a total volume of 97.5mL. 
volume of milk = 75
total volume = 97.5
97.5 - 75 = 22.5mL (volume of OREO)

Alright now that we have all that figured out...

Notice how I rounded to 2 sig figs? Smarty pants.

Oh that's weird... where did the oreo go? ;)

You also need to know how to calculate the expression of error. This is done by taking your three measurements, and calculating the average (using the correct number of SigFigs, of course).  You then find out what the accepted measurement is, and use the below equation to find the percent of error:

[(your measurement - accepted measurement) / accepted measurement] x 100

ta da!

See you later alligater

In a while crocodile

See you soon big baboon

XOXO

Beryllium Chemists

Heather is over and out.

Measurement & Uncertainty

Today we learned about Measurement and Uncertainty! Yayyy ...

Measurement and Uncertainty

Okay, sooo when you make a measurement the results usaully specify in a range of values. To show the range of values of a measurement is......

measurement = best estimate ± uncertainty

- Every estimate has some degree of uncertainty because it is only the "best" estimate, no estimate is exact

For example!
For a measurement of 6.07g  ± 0.02g, it means that the "true value" is between the range of 6.05g and 6.09g.

- We can only get an exact number if we count

For example, you can see and count two test tubes in the above image. You can't have 1.75 of a test tube (unless it's broken of course, but you don't use broken test tubes!) so therefore 2 test tubes is an exact number.




Absolute Uncertainty

Absolute uncertainty is the uncertainty in a measurement which is expressed in the units of measurements.

If a measurement is 2.12 and the "true value" is 2.00, the absolute uncertainty would be 2.12 - 2.00 = 0.12.
If the mass of something is measured like 3 times with the values of 1.00g, 0.95g, and 1.05g, the absolute uncertainty would be ± 0.05g.

Method 1

When you make atleast 3 measurements, you can calculate the average. The greatest defference between the average and the highest or lowest reasonable value is the absolute uncertainty.

Method 2

When finding the uncertainty of instruments, always measure to the best precision. Estimate to a fraction of 1/10 of the smallest segment on scale of whatever instrument you're using.

For example, a ruler's smallest segment is millimeter which is 0.1cm. The uncertainty of the ruler would be 0.01 because 1/10th of 0.1 is 0.01.













Relative Uncertainty and Sig Figs

Relative Uncertainty = absolute uncertainty/estimated measurement

- It can be expressed in percentage.
- Using sig figs indicate the relative uncertanty, the last digit is uncertain.


~ ~






Hope this helped you at least a little bit!
- Melody L. =)

Precision, Accuracy & Significant Figures

Today in class, we learned about precision, accuracy, but mostly significant figures.
Here are some definitions...


Precision - how reproducible measurement is compared to other similar measurements.
Accuracy - how close the (average) measurement  comes to the accepted or real value.
Significant Figures - measured in meaningful digits and the more precise means the more significant digits.


A - Significant Digits
- the last digit in a measure is uncertain as it could be one digit higher or lower. 
[ example: 3.24L. The 3 & 2 are certain whereas the 4 is uncertain ]
- significant digits in the measurement include ALL of the certain and first uncertain digit.
[ example: 3.24 = 3 sig figs ]


B - Significant Figures: Digits
- leading zero are not counted.
[ example: 0.04 = 1 sig fig ]
- trailing zeros after the decimal point are counted.
[ example: 10.0230 = 6 sig figs ]
- trailing zeros without a decimal point are not counted.
[ example: 15100 = 3 sig figs ]


C - Exact Numbers
- some quantities are defined as exactly a certain amount and no rounding is required.
- have an infinite number of sig figs.
[example: 4 dogs can be expressed as 4.000 dogs]
[example: a "pair" of shoes = 2 shoes (not 1.853892 shoes rounded up to 2) ]


D - Rounding Rules
1. look at the digit after the position of rounding.
2. if the digit is <5, round down.
3. if the digit is >5, round up.
4. if the digit is =5, and there are more non-zero digits after the 5, round up.
5. if that digit is =5 and it ends at the 5, round to make the last digit an even number.


Here are some practice rounding questions: round to the nearest tenth.
a) 26.73 =
b) 26.79 =
c) 15.31 =
d) 83.298634 =
e) 62.4585 =


Answers:
a) 26.7
b) 26.8
c) 15.3
d) 83.3
e) 62.5


E - Math Rules: Adding & Subtracting
- when adding and subtracting, round to the fewest number of decimal places (least precise).
[example: 12.544g + 1.3g = 13.844 ---> 13.8 (1.3 only has one decimal place) ]
here's a visual example:













F - Math Rules: Multiplying & Dividing
- when multiplying and dividing, round to the fewest sig figs.
[example: 12.27g x 3.8g = 46.626 ---> 47g(squared) ]
[example: 24.845g(squared) / 1.3g = 19.11 ---> 19g ]
visual example:

















hope this helped!
submitted by Kimberly Leung:)