Loading in 5 sec....

Ch 10 Test Hints – ChemistryPowerPoint Presentation

Ch 10 Test Hints – Chemistry

- 50 Views
- Uploaded on
- Presentation posted in: General

Ch 10 Test Hints – Chemistry

Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author.While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server.

- - - - - - - - - - - - - - - - - - - - - - - - - - E N D - - - - - - - - - - - - - - - - - - - - - - - - - -

Ch 10 Test Hints – Chemistry

Chemical Quantities

Define mole

We’re talking representative particles here.

That could mean people, pickles, atoms, ions, molecules, etc.

(Psssst….You should KNOW how many that is without us having to say so)

2) How to determine/calculate the number of atoms of an element in a given number of molecules of a compound

So say we have 3 molecules of caffeine, C8 H10 N4 O2 in our coffee cup. (as if 3 molecules would be enough when we’re studying for a chemistry test!) Now, how many hydrogen atoms would we have in there?

Hmmmm…. Each molecule of caffeine has 10 whole hydrogen atoms in it, so that’s , C8 H10 N4 O2 plus , C8 H10 N4 O2 plus , C8 H10 N4 O2 . Altogether that would be…?

3) The various quantities that Avogadro’s number represents.

Check back to hint #1 and the examples of what a mole’s-worth of particles stands for. The key to this question is knowing the correct chemical formulas for atoms, molecules (even the diatomic ones), formula units, and ions.

So, Avogadro’s number is the same for Na, Na+, NaCl, and Cl2 (to name a few) [And you’d still better know exactly what that number is!!]

4) How to convert atoms of an element to moles of an element using Avogadro’s number.

If we said you have 4.3 x 1021 atoms of gold, nope, make that platinum. Would you be rich? How many moles is that?

Look closely at the exponent. Is this quantity more than 1 moles-worth or less than 1 moles-worth? Once you figure that out, you will know if you need to multiply or divide by Avogadro’s number. [seriously! You need to know what that number is.] Remember, that a MOLE is a group that is a certain size.

5) How to convert moles of an element to atoms of an element using Avogadro’s number

This is just the opposite of number 4. If you have 1.50 moles (groups) of platinum, how many atoms is that?

Mathematically, you do the opposite process compared to what you did in number 4. Think about it. This is more than one group, so it must be bigger than Avogadro’s number!

6) How the atomic masses of 2 different elements are related

The atomic masses of elements have something in common. They represent the same number of particles!

So, if you have 1.00797 g Hydrogen or 15.9994 g Oxygen or 140.9077 g of Praseodymium (just ‘cuz that one’s fun to say), they all have exactly

_ . _ _ _ x _ _ _ _ atoms. Yea!

7) How to calculate the molar mass of a diatomic (halogen) gas

I know! I know! How about Fluorine gas? F2

This is actually a molecule like H2O is a molecule. So, you have to add up the atomic masses of both fluorine atoms to get the molar mass! Simple.

8) Define molar mass.

Uh… the name says it all.

9) How to calculate the molar mass of a given compound

Well, we’ve practiced this a bazillion times in class, in labs, on homework.

All you have to do is add up the atomic masses of the elements shown in the formula.

10) How to convert moles of a given compound to mass – molar mass will be given

If 1.00 mole of water has a mass of 18.01534 g (You know how we got that. We used it in all the hydrate labs.) , then how much would 2.00 moles of water be (by mass that is)?

Just double it, right? As in multiply by 2.00. So what if the question actually has some weird number in it with decimals, like 4.56 moles or something? The math is the same.

11) How to convert mass of a given compound to moles – molar mass will be given

This one is just the opposite of number 10. You should be able to set up a factor-label problem. Some of you seemed to find using ratios easier, so go ahead and do that. The math is the same either way.

If you are confused on how to start this problem, remember that one whole mole’s-worth of a compound is represented by the molar mass. If you are given a mass of something like 5.00 g H2O and a whole mole of water has a mass of 18.01534 g, you can see that the 5.00g is LESS than the mass of one mole. What mathematical operation is that then?

12) How to calculate the mass of a specific element in a given compound (% composition) – molar mass will be given.

Ok. Now we’re talkin’. This is a problem like #5 on the percent composition worksheet. Look back at the answers that you wrote down when we went over this in class. [Didn’t do that? Then take a look at the practice problem on page 307 to get you started. Hey, the online textbook even has a Checkpoint problem that you can try!]

Say you are asked to find the mass of Na in 7.80 g of NaCl. First, you have to find the % Na in NaCl and then take that percent of 7.80 g

PS. If you get an answer over 300 g, you forgot to move the decimal.

13) What information in general is needed to calculate the % composition of a compound.

If we asked you to find the % composition of water, we mean find the % H and the % O. The whole % composition worksheet was about this.

Can you describe how to calculate this?

14) How to calculate the % composition of a binary compound when given the masses of each element.

In this problem you will be given actual masses of elements. See page 306 #32 & 33 for examples. There is a sample problem in box 10.9.

You will need to find the total mass of the compound first and then do a basic % of element example.

Calculate the % of N.

9.03 g Mg 3.48 g N

15) How to calculate the % of a given element in a specific compound – molar mass will be given.

See your percent composition worksheet again. Example 10.10 on page 307.

Something like what is the % N in NH4NO3

16) How to set up the correct equation for finding % composition of an element.

If you can calculate the % of an element like in #15, this one will be super easy. All you have to do is identify which problem is set up correctly.

See, this is why we make you show your work all the time!

17) Define empirical formula.

Hey, look in your glossary, or the notes, or on page 309 in your book. It’s even on the online textbook.

You know this! We’ve asked you about it 100,000 times.

[Oh, another way to say it is ‘simplest formula’]

18) How to identify the molecular formula in a set of mostly empirical formulas.

Check Table 10.3 on page 311. Mr. Sustin say’s it’s wonderful!

Is the formula all nonmetals? Can the subscripts be reduced? If so, it’s a molecular formula and not an empirical formula.

Hey, I’ll bet there are podcasts posted on Mrs. Gregory’s Moodle page that you ALL have access to.

Practice problem #39 on page 312 rocks!

19) How to calculate the empirical formula of a compound when given the mass % of each element.

Well now, this is the type of problem that involves the most math.

On page 310, example problem 10.11 and two practice problems show you EXACTLY how to do this. Hey, the whole page is devoted to just this type of problem.

In your notes, there is exactly the same type of example for Acetic Acid.

Assume 100g sample – convert to moles – do a mole ratio.

20) How to distinguish between empirical and molecular formulas in general –basic characteristics/definitions of each.

This is one of those “Which of the following is NOT…” questions. Make sure you really, really know the definitions of empirical formula and molecular formula.

You have GOT to read all of the answer choices – twice, at least. How many times have you said to us during test corrections that you just misread the question? Well, don’t do that.

21) How to calculate the number of molecules in a given mass of a compound – molar mass will be given.

You will be given a mass like 5.67 g H2O and then you need to figure out how many particles. So GO TO MOLES! Use the mole map that is on page 303, figure 10.12, and then convert to particles. Yes, you actually get to use Avogadro’s Number this time.

KNOW THIS CHART!

(x Avogadro’s #) (x Molar Mass)

# of “particles”←Mole→ Mass in grams

22) How to calculate the % by mass of an element in a given compound – molar mass will be given.

This is just like #15, but it is a short answer problem. You did a bunch of these on the % composition worksheet. Yea!

23) How to calculate the molecular formula of a compound when given the empirical formulas and molecular masses of each.

If the empirical formula is HO and the molecular mass is 34 g/mole, what is the molecular formula?

Well, add up the mass of H & O (1 + 16 as a quick example) = 17

17 is not 34 because 34 is twice as big. So just double the subscripts to get the molecular formula H2O2.

Yep, it’s that easy.

24) How to calculate the number of moles of water in a hydrated compound when given the mass data (like the BaCl2 lab and the MgSO4 lab).

This is a problem just like the hydrate labs. You will be given data for mass of the crucible, mass of crucible with chemical before heating and mass of crucible with chemical after heating.

You need to know how to calculate the mass of water lost and the mass of the anhydrous (dry) chemical. Then, convert both of those to moles and do a mole ratio.

You are filling in the question mark in a formula like MgSO4 • ? H2O

Mrs. Gregory has a sample problem in her notes if you need that.

25) How to calculate the empirical AND molecular formulas when given the mass percentages of each element and the molecular mass of the compound.

This problem brings it all home. You will be given the %s of each element in a molecular compound. So, you need to convert each % to mass by assuming a 100.0 g sample. Change each mass to moles using molar mass. Calculate the mole ratio between the elements – always divide by the smallest number of moles so you get a whole number ratio.

But! There is one more step. Once you get the mole ration, which is the subscripts for the empirical formula, you have to add up the molar mass.

If that mass matches what is given as the molar mass, then the empirical formula = molecular formula. Chances are though that the molecular formula will be some multiple of the empirical (e.g. double, triple, etc.) so increase the subscripts by that much.

50.7% C, 4.2% H and 45.1% O with a molar mass of 110.0 g = ??