Module 2 Notes

Module 2 Notes

Name:

Naming Inorganic Compounds:

This should be a review of a Grade 10 chemistry with a little bit of added information.

Lets start.

Some background information:

First letter is capital, second letter lower case eg. Ge, Br, Li

When first two letters have been used jump to the third letter eg. Ar (argon), As (arsenic), At (astinine)

Some element take on Latin names eg. Cu (cuprum), Au (aurum), Pb (plumbum)

Some elements just use the first letter eg. P, O, S, U

On your periodic table use a highlighter to highlight the nonmetals (these are all elements left of the metalloids C, N, P, O, S, Se, F, Cl, Br, I, At, He, Ne, Ar, Kr, Xe, Rn)

The metals are found to the right of the metalloid staircase

Indicate the charges:

Above Column 1 (above H hydrogen) write +1

Above Column 2 (above Be beryllium) write +2

Above Column 13 (above B boron) write +3

Above Column 14 (above C carbon) write + 4 (acts like a metal and a non-metal)

Above Column 15 (above N nitrogen) write -3

Above Column 16 (above O oxygen) write -2

Above Column 17 (above F Fluorine) write -1

Each element in these columns have the indicated charges when not in a chemical bond. For example, Na is always +1, or Na1+, Ca is always +2 or Ca2+, Se is always -2 or Se2-.

In addition, we will use some of the transition metals. As there are no general rules here, use the Common anions and cations box on page 2 as a reference.

Definitions:

1) Metals form positive ions while nonmetals form negative ions

2) An anion is an ion with a negative charge. Eg. NO3-, Cl-

3) A cation is an ion with a positive charge. Eg. Al3+, NH4+

4) Ne, Li+, He are monatomic species with ONE atom

5) O2, NO, Hg22+, are diatomic species with TWO atoms

6) A triatomic species include such molecules as O3, NOCl, I3- with THREE atoms

7) A polyatomic species is made up of many atoms such as H2PO4-, NH3. The polyatomic ions are listed on the back of your PT.

Try: Circle the correct Answer.

OH is a neutral/cationic/anionic, monatomic/diatomic/triatomic/polyatomic species

H20 is a neutral/cationic/anionic, monatomic/diatomic/triatomic/polyatomic species

What about chromate, acetate, iodine ion, ammonium, iodine gas?

A) Naming metal ions:

Use the name of the metal and add ion

Eg.Aluminum metal forms aluminum ion (Al3+)

If there is more than one possibility, put the charge in roman numerals immediately after the name eg. Lead metal forms lead (II) ion or lead (IV) ion. Sometimes the Latin name is also used. Eg. Lead (II) = plumbous, Lead (IV) = plumbic

Try: Name Cu+_______________, Cr3+___________________, W6+___________________

Write the formula for nickel (II)ion_____________, vanadium (V) ion ________________

B) Naming non-metal ions:

Take off the ending you see on the periodic table or the Element List and replace with ide

Eg. Fluorine element (F) forms a fluoride ion (F-)

Try: Bromine (Br) _________, Iodine(I)_________, Oxygen (O)__________, Sulphur (S)___________, Nitrogen (N)___________, Phosphorus (P) _____________

C) Naming polyatomic ions is beyond the scope of this course use the back of your PT

Eg. Nitrate, hydroxide, ammonium

D) Ionic Compound Name to Formula:

An Ionic compound is a compound made of ions; one cation and one anion

Compound are always NEUTRAL. # positive charges = # negative charges

There are 3 steps:

1. write the formula for the positive ion followed by the negative ion

2. criss-cross the charges on the ions

3. Clean-up : divided subscripts by common number, omit - and +, omit subscript if its a 1

Eg.Tin (IV) oxide

1.Sn4+ O2-

2.Sn 4+ O 2-

Sn2O4

3.Sn2O4 divide subscripts by 2

SnO2

Try: give formula for calcium phosphide, iron (II) phosphate, ammonium sulphate

Naming Compounds Part II

E) Ionic Compound Formula to name:

1)If the cation only has one possible charge, write the names of the ions one after the other without the word ion. Remember for the anion, replace the ine ending with ide

Eg. ZnCl2 zinc chloride not zinc chlorine

2)If there are a few possibilities for the cation, uncriss-cross and double check.

Eg. PbO2 ( un criss cross Pb2- O1- now check charge on anion oxide must be a (2-) charge so we have to multiply both charges by 2.

So the ions are actually Pb4+ and O2- and the name is Lead (IV) oxide

Try:Ag2SO4

Cu2O

U(SO4)2

F) Naming Hydrates:

Molecules which include water molecules in their structures are called Hydrates

This often includes crystal structures such as copper sulphate (CuSO4 5H20)

Each water molecule is called a Hydrate and you number them according to the prefixes given:

Mono (1)

hexa (6)

Di (2)

hepta (7)

Tri (3)

octa (8)* please memorize prefixes*

Tetra (4)

nona (9)

Penta (5)

deca (10)

So copper sulphate is actually properly named

copper (II) sulphate pentahydrate

Try: Zn(CH3COO)22H2O

aluminum nitrate nonahydrate

G) Naming molecules with two nonmetals: (prefix naming system)

1)Use the number prefix above for each element whether it comes first or second

2)The first nonmetal takes on the name directly from the periodic table

3)The second nonmetal take the ide ending as usual

4)Exception to every rule if the first element only has one atom, dont need to use mono

eg. CO2 carbon dioxide not monocarbon dioxide

Try:P2S3

CO

CS2

N2O4

H) Diatomic molecules

N, O, F, Cl, Br, I and H are always diatomic molecules

Ie. N2, O2, F2, Cl2, Br2, I2 and H2

I) Naming Acids:

1)If the first element in a compound is Hydrogen (H) it is an acid use hydro as the first name and add the name of the ion followed by ic acid

2)If the acid involves a polyatomic you usually have two choices one ending in ate and one ending in ite. If the polyatomic ends in ate the acid takes on the ic ending. If the ployatomic ends in ite, the acid takes on the ous ending. Compare sulfurous and sulfuric acid on the next page

3)Some common acids include

HFHydrofluoric acid

used to frost glass

HClhydrochloric acid

found in stomach acid

H2SO3suphurous acid

acid rain

H3PO4phosphoric acid

found in Coke

HC2H3O2 or CH3COOH acetic acid

vinegar (5%)

HBr hydrobromic acid

HI hydroiodic acid

H2SO4 sulphuric acid

HNO3 nitric acid

HNO2 nitrous acid

*** Rather than try to remember a rule to which there are many exceptions, you may just choose to memorize this list of acids ***

Naming Compounds Summary:

If the first element or ion in the formula is

Hydrogen

Name it as an acid

A non-metal (and doesnt contain NH4)

Use the greek prefix naming system

Is not listed on the back of your periodic table

Use name of cation followed by anion (with ide ending)

Is listed on the back of your periodic table

Use Roman numerals for the metal, followed by name of anion

Pure Substances:

A quick reminder:

Elements

Are pure substances that cannot be chemically decomposed into anything simpler because they only consist of one type of atom

About 110 element are known and listed on the Periodic Table

Elements can be a solid, liquid or gas (depends on space between particles, not the particles themselves)

Examples: oxygen (O2), nitrogen (N2), gold (Au), copper (Cu)

Compounds

Are pure substances that can be chemically decomposed into elements.

Made up of more than one element (2 or more different elements)

Can be separated by chemical means

Atoms are the smallest particles which are chemically indivisible. The atoms of one element differ from the atoms of all other elements.

Molecules are whats made when two or more atoms join together

Calculating Formula Masses

The formula mass of a compound is calculated by adding the atomic masses

Eg1. Na2(CO3)

Na2(CO3) is composed of 2 sodium atoms, 1 carbon atom and 3 oxygen atoms.

Now use your periodic table to find out the mass each element.

Na 23.0 g/mol

C 12.0 g/mol

O 16.0 g/mol

Combine the two to get the formula mass of the compound Na2CO3

Na - (23.0 g/mol) x 2 atoms =46.0 g/mol

C - (12.0 g/mol) x 1 atom =12.0 g/mol

O - (16.0 g/mol) x 3 atoms=48.0 g/mol +

Formula mass100.0 g/mol

Try a few more to make sure you get the hand of this. In the following table, give the name and formula, calculate the formula mass, state whether the molecule is an element or compound; finally, count the number of atoms per molecule.

Name

Formula

Formula mass

Element or Compound?

# atoms in one molecule?

Nitrogen gas

N2

14.0g/mol x 2 = 28.0g/mol

Element

2

Carbon dioxide

CO2

12.0g/mol + 2(16.0g/mol) =

42.0g/mol

Compound

3

H2O

CO

The Mole & Molar Mass

The mole is the standard method in chemistry for communicating how much of a substance is present.

Here is how the International Union of Pure and Applied Chemistry (IUPAC) defines "mole":

The mole is the amount of substance of a system which contains as many elementary entities as there are atoms in 0.012 kilogram of carbon-12. When the mole is used, the elementary entities must be specified and may be atoms, molecules, ions, electrons, other particles, or specified groups of such particles.

This is the fundamental definition of what one mole is. One mole contains as many entities as there are in 12 grams of carbon-12 (or 0.012 kilogram).

In one mole, there are 6.02 x 1023 atoms. Here's another way: there are 6.02 x 1023 atoms of carbon in 12 grams of carbon-12.

one mole of ANYTHING contains 6.02 x 1023 entities.

The word "entities" is simply a generic word. For example, if we were discussing atoms, then we would use "atoms" and if molecules were the subject of discussion, the word entities would be replaced in actual use by "molecules."

6.02 x 1023 is such an important number, it has been named after the Italian chemist who in 1811, made a critical contribution to the measurement of atomic weights. Avogadro's Number (N) has been very carefully measured in a number of ways over many decades. Avogadros Number has a unit associated with it. It is mol-1, as in 6.02 x 1023 mol -1 , or 6.02x1023 per mol (6.02x1023/mol)

Why is there no unit in the numerator? There could be, but it would vary based on the entity involved. If we were discussing an element, we might write atoms/mol. If we were discussing a compound, we would say "molecules per mol." What is in the numerator depends on what "entity" (atom, molecule, ion, electron, etc.) is being used in the problem.

The symbol for mole is mol . (like the symbol of meter is m)

Here it is again: one mole of ANY specified entity contains 6.02 x 1023 of that entity. For example:

One mole of donuts contains 6.02 x 1023 donuts

One mole of H2O contains 6.02 x 1023 molecules

One mole of nails contains 6.02 x 1023 nails

One mole of Fe contains 6.02 x 1023 atoms

One mole of dogs contains 6.02 x 1023 dogs

One mole of electrons contains 6.02 x 1023 electrons

Get the idea? Its just a number like a dozen means 12 donuts, molecules, nails, atoms, dogs or electrons. It just so happens N is a big number.

Getting back to Avogadro's Number role in chemistry; please note that counting atoms or molecules is very difficult since they are so small. However, we can "count" atoms or molecules by weighing large amounts of them on a balance.

When we weigh one mole of a substance on a balance, this is called a "molar mass" and has the units g/mol (grams per mole). This idea is very critical because it is used all the time.

A molar mass is the weight in grams of one mole.

One mole contains 6.02 x 1023 entities.

Therefore, a molar mass is the mass in grams of 6.02 x 1023 entities.

OK. How does one calculate a molar mass? Get ready, because you already know how to calculate a molar mass.

The molar mass of a substance is the formula mass in grams.

All you need to do is calculate the formula mass (or molecular weight) and stick the unit "g/mol" after the number and that is the molar mass for the substance in question.

Four steps to calculating a substances molar mass:

1. Determine how many atoms of each different element are in the formula

2. Look up the atomic mass of each element in your Periodic Table.

3. Multiply step one by step two for each element

4. Add the results for step three together and round off as necessary

Note: Hydrates Suppose you were asked to calculate the molecular weight of CuSO45H2O remember does not mean multiply.

Eg. Calculate the molar mass of Al(NO3)3

(1 x 27.00) + (3 x 14.00) + (9 x 16.00) = 213.00 g/mol

213.00 grams is the mass of one mole of aluminum nitrate

213.00 grams of aluminum nitrate contains 6.02 x 1023 molecules of Al(NO3)3

Conversion Factors and the Mole

Now its time to piece it all together and pave the way for the rest of the course. Remember the conversion factor equivalences we have seen so far, from 60min = 1 hr to 1000mg = 1 g.

I have now introduced you to one of the most important conversion factors in Chemistry and that is the mole (( mass conversions.

The two conversion factors are:

g

or

mol

mol

g

Eg1.What is the mass of 3.25 mol of CO2

1. First you have to find the molar mass of CO2

12.0 + 2(16.0) = 44g/mol

2. Pick the right conversion factor

44.0g

or 1 mol

1 mol

44.0 g

3. Find the mass of CO2 by dimensional analysis. Remember start with what you know.

3.25 mol CO2x44.0g=143 g CO2

1 mol

Notice by choosing this conversion factor, the moles cancel out leaving us with the grams we want.

Eg2.What is the mass of 1.36 x 10-3 mol of SO3?

1. Molar mass of SO3

32.0g + 3 (16.0) = 80.0g/mol SO3

2. Pick the conversion factor

80.0 g

or

1 mol

1 mol

80.0g

3. Find the mass of SO3 by dimensional analysis

1.36 x 10-3 mol SO3 x 80.0g

=0.106 g SO3

1 mol

Notice, I have included units and the name of the compound in every step. I have also rounded to the correct number of sig figs (3)

Eg3.How many moles of N2 in 50.0 g of N2

This is backwards compared to the first 2 questions

1. Still find the molar mass of N2

14.0g x 2 = 28.0 g/mol

2. Pick the conversion factor

28.0g

or

1 mol

1 mol

28.0g

Notice this time we pick the other conversion factor

3. Find the number of moles by dimensional analysis

50.0g N2 x1 mol=1.79mol N2

28.0g

Grams cancel out leaving us with mols

Eg4.How many moles of CH3OH is there if I have 0.250g?

1. Molar mass of CH3OH is 32.0g/mol

2. Conversion factor is 1mol/32.0g

3. Dimensional Analysis

0.250g CH3OH x1 mol=7.81 x 10-3 mol CH3OH

32.0g

Eg5.If 0.140 mol of acetylene gas has a mass of 3.64g, what is the molar mass?

This question really checks for understand of this topic, rather than falling into the trap of memorizing the steps.

Recall the units for molar mass are g/mol weve been given grams and weve been given moles, so just divide them

Molar mass = g =3.46g

=26.0 g/mol

mol

0.140mol

Conversion Factors and Avogadro

So far the concept of the mole has given us the conversion factor relating mass (g) and moles (mol).

Avogadros number gives us another conversion factor to consider, namely the following two possibilities:

Recall:

1 mole = 6.02 x 1023

So

1 mol

or

6.02 x 1023 entities

6.02 x 1023 entities

1 mol

Lets try a few together.

Eg1.

How many pairs of socks are in half a dozen pairs?

1.Start with 0.5 dozen pairs

2.Conversion factor choices:

1 doz

or

12 pairs

12 pairs

1 doz

Final calculation:

0.5 doz x 12 pairs=6 pairs of socks

1 doz

Eg2.

How many individual socks are in half a dozen pairs?

Exactly the same as Eg1, only we have one more conversion factor to include:

1 pair

or

2 socks

2 socks

1 pair

.5 doz x 12 pairsx2 socks=12 socks

1 doz

1 pair

Eg3.

How many molecules in 3.2 moles?

From above remember:

1 mol

or

6.02 x 1023 entities

6.02 x 1023 entities

1 mol

So

3.2 molx6.02 x 1023 molecules=1.93 x 1024 molecules

1 mol

Eg4.

How many molecules in 4.3 moles of O2?

4.3 molx6.02 x 1023 molecules=2.6 x 1024 molecules O2

1 mol

Eg 5.

How many oxygen atoms in 3 moles of O2?

This is the same as Eg 5, except we go one step further. We need to include the fact that there are 2 atoms per molecule of O2 or 2 atoms = 1 molecule

4.3 mol x6.02 x 1023 molecules x2 atoms =

3.2 x 1024 atoms O2

1 mol

1 molecule

Eg 6.

How many moles in 8.25 x 1025 atoms?

8.25 x 1025 atomsx1 mol

=137 mol

6.02 x 1023 atoms

Moles and Multiple Conversions:

Lets recall

On the first day of the mole, we did questions involving mole (( mass conversions like this. (Please complete and check with the answer given)

What is the mass of 3.7 moles of NaCl?

Ans: 220g NaCl

What is the number of moles in 48g of H2O?

Ans: 2.7 mol H2O

Then on the second day of the mole, we did questions involving mole (( molecule/atom conversions like this. (Please complete and check with the answer given)

How many moles of H2SO4 are there if I have 5.6 x 1024 molecules?

Ans; 9.3 mol H2SO4

How many atoms of oxygen are there in 5.6 moles of carbon dioxide molecules?

Ans: 6.7x1024 atoms O

Now the big step is to combine the two. This involves multiple conversions.

x # atoms

atoms molecule

x N x mol mass

mol molar mass

xmolecule Molecules

x molar mass density

#atoms

mol g/L

x mol MOLe x mol

N

22.4L

x 22.4L volume

mol

The squares indicate your start and end point and the arrow indicate conversion factors.

You can start and end at any point for example atoms ( volume, or molecules ( mass,

or volume ( molecules, the list is endless.

Remember N = Avogadros number = 6.02 x 1023

Lets try a couple.

1. What is the mass of 1.00x1012 atoms of Cl?

Step one: locate the start and end point on the chart above

Atoms ( moles( mass

This means this calculation with take 2 conversion factors

Step two: complete conversion factor calculation

1.00x1012 atoms Cl x 1mol x 35.5g = 5.90x10-11g

6.02x1023atoms mol

2. How many atoms in 5.0g of NaCl?

There are 3 conversion factors: mass (moles ( molecules ( atoms

5.0g NaCl x 1mol x 6.02x1023 molecules x 2 atoms = 1.0 x 10 23 atoms

58.5 g NaCl

1 mol

molecule

Avogadros Hypothesis - Moles and Gas Volumes:

First a little History Lesson:

Amedeo Avogadro (1776 1856)

Avogadro was born on August 9 of the year of the Declaration of Independence of the United States of America in Turin, a town in Northern Italy. As it was the custom in Italy at that time to give long names, he was given the name Lorenzo Romano Amedeo Caroo Avogadro de Quaregna e di Cerreto.

Amedeo was a very bright young man. He started college at 13 and graduated in law at the age of 16. By the time he was 20 he had his doctorate degree in ecclesiastical law. Amedeo practiced law for a few years, but become more and more interested in physics.

By the age of thirty he had given up law and was teaching mathematics and physics in a small local college. Later he has made Professor of Mathematical Physics at the University of Turin. He retired at the age of 74 and died at the age of 80.

Avogadro was a quiet man, well liked by his students, and quite well known for his mischievous humour. He was happily married and the father of 6 sons. During his university career he was involved in various committees. All in all, he appears to have been a good father, a good citizen and a conscientious thinker.

The highlight of Avogadros life was an article written in 1811, when he was 35 years old. In that article, Amedeo discussed an assumption which must be make regarding Gay-Lussacs Law of Combining Volumes in order to explain the law at the atomic and molecular level.

Gay-Lussac had shown that gases in a chemical reaction combine in volumes of fixed proportions. The fixed proportions were simple whole number ratios. Avogadro combined Gay-Lussacs discovery with John Daltons application of his Atom Theory to chemical equations. Dalton had proposed that in chemical reactions atoms or compound- atoms react in fixed simple-whole-number proportions. Avogadro saw that the gas-volume ratios of Gay-Lussac and the atom or compound-atom ratios of the Dalton were identical. This led to a hypothesis by Avogadro that equal volumes of gases at the same temperature and pressure contain the same number of molecules. If the gas volume ratio in a chemical reaction was to equal the atom or compound atom ratio this had to be true.

The hypothesis was not as straightforward nor as immediately acceptable as might appear. First, the chemical formulas for many substances were not known. The writing of balanced equation for many reactions was merely guesswork. As a result it was difficult to obtain the ratio of molecules which theoretically should combine.

Another related difficulty in obtaining the theoretical ratios (to compare to Gay-Lussacs gaseous volumes) was that there was not, at that time, a clear distinction between atoms and molecules. What Dalton has called compound atoms, Avagadro was the first to call molecules. In the context of the time Avogadros Hypothesis was a difficult forward. It was not until after Avogadros death that his ideas were accepted. Stanisiao Cannizzaro championed Avogadros ideas 45 years after they were presented and finally saw their adoption.

In Avogadros honor, 6.02x1023 is called Avogadros number. We now refer to 6.02x1023 or Avogadros number (N) as the mole.

http://www.bulldog.u-net.com/avogadro/avoga.html

What we take from this is Avogadros Hypothesis which states:

Equal volumes of different gases, at the same temperature and pressure contain the same number of particles.

This brings us to the idea of molar volume which is the volume of gas occupied by one mole of gas at standard temperature (OoC) and standard pressure (101.3kPa) -STP.

Implications At STP, equal mol of any gas occupy the same volume

For example if 2 mol of O2 gas takes up 44 L, 2 mol of CO2 also takes up 44L

It has been determined that

1 mol gas = 22.4 L

@ STP for gases

This of course gives us two more conversion factors to work with:

1 mol or 22.4L

22.4L

1 mol

These conversion factors have already been added to the flow chart from the last lesson

*** This is only true for Gases at STP, NOT solids or liquids***

Lets try a few examples.

Eg1.How many moles of gas are in a 10.0L balloon at STP?

10.0Lx1 mol=0.446 mol of gas

22.4L

Eg2.What is the volume?

a) 12.5 mol NH3(g)

b) 0.350 mol O2(g)

c) 4.25 mol HCl(g)

Eg3.Moles at STP?

a) 85.9L of H2(g)

b) 375mL of SO3(g)

Density as a Conversion Factor:

Density is obviously an integral part of Chemistry. Up until this point in your Science careers Density has been calculated using the formula:

Density = Mass

Volume

I would like you to shift your thinking a little so we can incorporate density calculations in our dimensional analysis.

So by focusing on the units

Density = Mass

=grams

(or kg)

Volume

mL

L

So we have two conversion factors:

gormL

mL

g

These conversion factors allow us to find the volume of solids and liquids

Remember for gases, at STP we would just use Avogadros hypothesis (22.4L/mol)

Lets look at a few examples.

Eg1.What is the volume occupied by 3.00 mol ethanal (CH3CH2OH) if its density is 0.790g/mL

Eg2.How many moles of Hg(l) are contained in 100mL of Hg(l) (d=13.6g/mL)

Eg3.Density of O2(g) at STP?

Eg4.2.50L bulb contains 4.91g of gas at STP. What is the molar mass?

Eg5.Al2O3(s) has a density of 3.97 g/mL. How many atoms of Al are in 100mL of Al2O3?

Moles, moles, moles

We have now covered all the conversion factors you could ever need regarding moles.

Conversion

Conversion factor

Moles ( ( number of particles

6.02 x 1023 particles or 1 mol .

1 mol 6.02 x 1023 particles

Moles ( ( mass

Molar mass (g) or 1 mol

1 mol molar mass (g)

Moles ( ( volume (gases at STP)

22.4 L or 1 mol

1 mol 22.4 L

Molecules ( ( atoms

Atoms or 1 molecule

1 molecule atoms

Percent Composition and the Empirical Formula

A) Percent Composition

Sometimes I might ask you to tell me what percentage of the total mass an element takes up in a compound.

An analogy to this would be to ask what is the percentage by mass of mayonnaise used in a tuna fish sandwich.

To figure this out you would divide the mass of mayo used by the mass of the total sandwich

Mass of mayo x100%= % by mass of mayo on sandwich

Mass of sandwich

Eg1.What is the percent composition of H2SO4?

*** A convenient way to attack these type of questions is to assume we have a mole of the compound this way we can use the formula masses right off the periodic table, and not actually with it***

Step 1: Find the masses of each element, and the entire compound:

Mass of Hydrogen =1.0 x 2 =2.0g

Mass of Sulphur=32.1 x 1=32.1g

Mass of Oxygen=16.0 x 4=64.0g

Mass of H2SO4

=98.1g

Step 2: Now find the percent composition. This is like finding the percent of anything. Do your best to keep sig figs in mind.

Percent of H =2.0 g x100%=2.0%

98.1g

Percent of S=32.1g x100%=32.7%

98.1g

Percent of O=64.0gx100%=65.2%

98.1g

Step 3: Make a final statement

The percent composition of H2SO4 is a follows: Hydrogen 2.0%; Sulphur 32.7%; Oxygen 65.2%.

Eg2.What is the percent of water in CuSO45H2O?

Water

=90.0gx100%=36.1% H2O

CuSO45H2O

249.6g

Try a few on your own:

Find the percent composition of the blue group of atoms

1. Cr(NH3)6Cl3H2O

2. Cu(C2H3O2)22NH3

3. (NH4)2Sn(OH)6

4. K3Fe(CN)6 ( each element

B The Empirical Formula:

This is essentially the opposite of find the percent composition

You are given the percentage of each element and you have to find the simplest possible formula

Compare the ratio of C:H in the following:CH2, C2H4, C3H6, C4H8, C5H10

They all have a C:H ratio of 1:2; in other words, they all have the Empirical Formula of CH2. However their molecular formulas are clearly different.

Eg1.What is the empirical formula of a compound consisting of 80.0% C and 20 % H?

***The assumption here is we have 100g of the compound, since we are dealing with percent***

Step 1: Change the percent to grams as per the above assumption

So 80.0% C ( 80.0 g C

20.0% H (20.0 g H

Step 2: Now convert the masses to moles so we can compare them directly to each other

g( mol

By smallest result from first column

Multiply to achieve whole number multiple if necessary

80.0g x 1mol = 6.67mol C

12.0g

6.67 = 1

6.67

N/A

20.0g H x 1 mol = 20.0 mol H

1.0g

20.0 = 3

6.67

N/A

Step 3: Determine empirical formula

Based on above calculations C H3 is the empirical formula give these percents.

Try a few on your own:

What are the Empirical Formulas for the following?

1. 58.5% C, 7.3% H and 34.1% N

2. 81.8% C, 18.2% H

3. 39.1% Si, 61.0% O

Calculating Molecular Formulas

Thus far we can calculate the empirical formula given the percent composition. We need to take it one step further in order to calculate the molecular formula.

Remember

CH2 C2H4 C3H6

All have the same empirical formula CH2, the molecular formulas being a multiple (x2, or x3) of this.

So in general

Molecular Formula = N x (empirical formula)

Where N = a whole number multiple

Keep in mind we can shuffle the equation so

N = Molecular Formula

Empirical Formula

Since N can be determine by the ratio of the molecular formula and the empirical formula it also works for the molecular masses, so

N = Molecular Mass

Empirical Mass

Eg1. A molecule has an empirical formula of HO and a molar mass of 34.0g. What is the molecular formula?

Step 1: First find N

N = molecular mass= 34.0 =34.0g=2

Empirical mass

1.0 + 16.0g

17.0g

Step 2: Now find the molecular formula

Molecular Formula = N x (Empirical Formula)

=2 (HO)

=H2O2

(hydrogen peroxide)

Eg2. A gas has empirical formula POF3. If 0.350L of gas at STP has a mass of 1.62g. What is the molecular formula?

Step 1: Find N

Fill in what we know:

N = molecular mass= ?

Empirical mass

POF3

We have not been given the molecular mass as such, but we do have some density information. Since this is at STP, we can convert the density to the molar mass:

1.62gx22.4L=104g/mol = molecular mass

0.350L

mol

This can now be substituted into our first equation:

N = molecular mass= ?

=104g = 1

Empirical mass

POF3

104g

Step 2: Find the Molecular Formula

Molecular Formula = N x (Empirical Formula)

=1 (POF3)

=POF3

Eg 3. Given an empirical formula of SiH3 and the fact that 0.0275 mol of this compound weighs 1.71g, what is the molecular formula?

Remember molar (molecular) mass = g=1.71 g=62.2g/mol

Mol

0.0275mol

So N = molecular mass=62.2g=62.2g=2

Empirical mass

SiH3

31.1g

Therefore MF = 2(EF) = 2(SiH3)=Si2H6

Now its time to try some on your own.

Calculating Molecular Formulas Practice:

1. A gas has the empirical formula CH2. If 0.850L of the gas at STP has a mass of 1.59g, what is the molecular formula?

2. A gas has the percentage composition: 30.4% N and 69.6% O. If the density of the gas is 4.11g/L at STP, what is the molecular formula of the gas?

3. A compound has an empirical formula C5H11. If 0.0275 mol of the compound has a mass of 3.91g. What is the molecular formula of the compound?

4. A gas has an empirical formula CH. If 450mL of the gas at STP has a mass of 0.522g, what is the molecular formula?

5. When a sample of nickel carbonyl is heated, 0.0600 mol of a ga containing carbon and oxygen is formed. The gas has amass of 1.68g and is 42.9%C. What is the molecular formula of the gas?

6. A gas sample is analyzed and found to contain 33.0% Si and 67.0% F. If the gas density is 7.60g/L at STP, what is the molecular formula of the gas?

7. A gas has the percentage composition: 78.3% B and 21.7% H. A sample bulb is filled with the unknown gas and weighed. The mass of unknown gas is found to be 0.986 times the mass of a sample of nitrogen gas, N2(g), in the same bulb under the same conditions of temperature and pressure. What is the molecular formula of the unknown gas?

8. A gas has an empirical formula CH2. If 0.500L of the gas at STP has a mass of 0.938g, what is the molecular formula of the compound?

A sample of gas has an empirical formula of O and has a molar mass which is 3 times that of CH4. What is the molecular formula of

Molarity & Dilutions

A) Molarity (aka Molar Concentration)

Molarity is a measure of CONCENTRATION or how much solute (eg. Salt) is dissolved in solvent (water).

While this could be measured as how many grams is put in a certain number of litres, Chemists choose to measure how many moles are in a certain number of litres. However we are not worried about this because we are now experts at converting grams to moles right??!!!

In other words Molarity (M) is moles

Litre

Eg1. What is the molar concentration of NaCl ([NaCl]) containing 5.12g of NaCl in 250.0mL of solution?

5.12g NaClx1 molx1000mL=0.350mol/L = 0.350M

250.0mL

58.5g

1L

[NaCl] = 0.350M

Eg2. What is the mass of NaOH contained in 3.50 L of 0.200 M NaOH?

0.200mol NaOH x 3.50L x 40.0g =28.0g NaOH

1L

1mol

Eg. 3. What is the molatrity of pure sulphuric acid H2SO4 having a density of 1.839 g/mL?

1.839 g x1 mol x1000 mL = 18.7 mol = 18.7 M

mL

98.1g

1L [H2SO4] = 18.7 M

Eg. 4. What is the molarity of CaCl2 in a solution made by dissolving and diluting 15.00g of CaCl26H2O to 500.0mL?

15.00g

x 1molx1000mL=0.1369mol/L

500.0mL

219.1g

1L [CaCl2] = 0.1369M

B) Dilutions:

The purpose of dilution calculations is to calculate how the concentration is changed, by adding more solvent, mixing solutions or if evaporation occurs.

Eg1.If 200.0mL of 0.500M NaCl is added to 300.0mL of water, what is the resulting [NaCl] in the mixture?

0.500mol NaCl x200.0mL= 0.200M

1L

500.0mL [NaCl] = 0.200M

Eg2.If 300.0mL of 0.250M of NaCl is added to 500.0 mL of 0.100 M NaCl, what is the resulting [NaCl] in the mixture?

(0.300L x 0.250mol NaCl+ 0.500L x 0.100mol NaCl ) x 1 = 0.156mol/L

L

L

0.800L

[NaCl] = 0.156M

Eg3.What volume of 6.00M HCl (stock solution) is used in making 2.00L of 0.125M HCl?

0.125mol HCl x 2.00L x L . = 0.04166 L = 41.7mL HCl

L

6.00mol

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