Module 2 Notes Name:

Naming Inorganic Compounds:

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

Let’s 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 ions2) An anion is an ion with a negative charge. Eg. NO3- , Cl - 3) A cation is an ion with a positive charge. Eg. Al 3+ , NH 4+

4) Ne, Li+, He are monatomic species with ONE atom5) O2, NO, Hg22+, are diatomic species with TWO atoms6) A triatomic species include such molecules as O3, NOCl, I3- with THREE atoms7) A polyatomic species is made up of many atoms such as H2PO4- , NH 3. 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:1

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, ammoniumD) 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 ion2. “criss-cross” the charges on the ions3. 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 2SnO2

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

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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)2•2H2O 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 second2) The first nonmetal takes on the name directly from the periodic table3) The second nonmetal take the “ide” ending as usual4) Exception to every rule – if the first element only has one atom, don’t need to

use “mono” eg. CO2 carbon dioxide not monocarbon dioxide

Try: P2S3 CO CS2 N2O4

H) Diatomic moleculesN, O, F, Cl, Br, I and H are always diatomic moleculesIe. 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 includeHF Hydrofluoric acid used to frost glass

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HCl hydrochloric acid found in stomach acidH2SO3 suphurous acid acid rainH3PO4 phosphoric acid found in Coke HC2H3O2 or CH3COOH acetic acid vinegar (5%)HBr hydrobromic acidHI hydroiodic acidH2SO4 sulphuric acidHNO3 nitric acidHNO2 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 isHydrogen Name it as an acidA non-metal (and doesn’t contain NH4) Use the greek prefix naming systemIs 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 what’s made when two or more atoms join together

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Calculating Formula Masses The formula mass of a compound is calculated by adding the

atomic massesEg1. 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/molC – 12.0 g/molO – 16.0 g/mol

Combine the two to get the formula mass of the compound Na2CO3Na - (23.0 g/mol) x 2 atoms = 46.0 g/molC - (12.0 g/mol) x 1 atom = 12.0 g/molO - (16.0 g/mol) x 3 atoms = 48.0 g/mol +

Formula mass 100.0 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

H2OCO

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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. Avogadro’s 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? It’s just a number like a dozen means 12 donuts, molecules, nails, atoms, dogs or electrons. It just so happens N is a big number.

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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 formula2. Look up the atomic mass of each element in your Periodic Table.3. Multiply step one by step two for each element4. Add the results for step three together and round off as necessaryNote: Hydrates – Suppose you were asked to calculate the molecular weight of CuSO4·5H2O 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 nitrate213.00 grams of aluminum nitrate contains 6.02 x 1023 molecules of Al(NO3)3

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Conversion Factors and the Mole

Now it’s 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 molmol g

Eg1. What is the mass of 3.25 mol of CO2

1. First you have to find the molar mass of CO212.0 + 2(16.0) = 44g/mol

2. Pick the right conversion factor 44.0g or 1 mol1 mol 44.0 g

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

3.25 mol CO2 x 44.0g = 143 g CO21 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 SO32. Pick the conversion factor

80.0 g or 1 mol1 mol 80.0g

3. Find the mass of SO3 by dimensional analysis1.36 x 10-3 mol SO3 x 80.0g = 0.106 g SO3

1 molNotice, 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 N2This is backwards compared to the first 2 questions

1. Still find the molar mass of N214.0g x 2 = 28.0 g/mol

2. Pick the conversion factor28.0g or 1 mol1 mol 28.0g

Notice this time we pick the other conversion factor3. Find the number of moles by dimensional analysis

50.0g N2 x 1 mol = 1.79mol N228.0g

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

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2. Conversion factor is 1mol/32.0g3. Dimensional Analysis

0.250g CH3OH x 1 mol = 7.81 x 10-3 mol CH3OH32.0g

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 – we’ve been given grams and we’ve been given moles, so just divide them

Molar mass = g = 3.46g = 26.0 g/molmol 0.140mol

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Conversion Factors and Avogadro

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

Avogadro’s 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 10 23 entities 6.02 x 1023 entities 1 mol

Let’s try a few together.Eg1. How many pairs of socks are in half a dozen pairs?1. Start with 0.5 dozen pairs2. Conversion factor choices:

1 doz or 12 pairs12 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 socks2 socks 1 pair

.5 doz x 12 pairs x 2 socks = 12 socks1 doz 1 pair

Eg3. How many molecules in 3.2 moles?From above remember:

1 mol or 6.02 x 10 23 entities 6.02 x 1023 entities 1 mol

So…3.2 mol x 6.02 x 10 23 molecules = 1.93 x 1024 molecules

1 mol

Eg4. How many molecules in 4.3 moles of O2?4.3 mol x 6.02 x 10 23 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 molecule4.3 mol x 6.02 x 10 23 molecules x 2 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 atoms x 1 mol = 137 mol

6.02 x 1023 atomsMoles and Multiple Conversions:Let’s recall…

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

x # atomsatoms molecule x N x mol mass

mol molar mass x molecule 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 = Avogadro’s number = 6.02 x 1023

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Let’s 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 aboveAtoms moles mass

This means this calculation with take 2 conversion factorsStep two: complete conversion factor calculation

1.00x1012 atoms Cl x 1mol x 35.5g = 5.90x10-11g6.02x1023atoms 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 6.02x10 23 molecules x 2 atoms = 1.0 x 10 23 atoms

58.5 g NaCl 1 mol molecule

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Avogadro’s 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 Avogadro’s 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-Lussac’s 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-Lussac’s discovery with John Dalton’s 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 Avogadro’s Hypothesis was a difficult forward. It was not until after Avogadro’s death that his ideas were accepted. Stanisiao Cannizzaro championed Avogadro’s ideas 45 years after they were presented and finally saw their adoption.

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

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

What we take from this is Avogadro’s Hypothesis which states:

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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 volumeFor 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 22.4L22.4L 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?10.0L x 1 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)

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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 = MassVolume

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: g or mLmL g

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

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

Let’s look at a few examples.

Eg1. What is the volume occupied by 3.00 mol ethanal (CH3CH2OH) if it’s 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 factorMoles number of particles 6.02 x 10 23 particles or 1 mol .

1 mol 6.02 x 1023 particlesMoles 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 LMolecules atoms Atoms or 1 molecule

1 molecule atoms

Percent Composition and the Empirical Formula

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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.o 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 = 2.0gMass of Sulphur = 32.1 x 1 = 32.1gMass of Oxygen = 16.0 x 4 = 64.0gMass 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 x 100% = 2.0%98.1g

Percent of S = 32.1g x 100% = 32.7%98.1g

Percent of O = 64.0g x 100% = 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 CuSO4·5H2O?Water = 90.0g x 100% = 36.1% H2OCuSO4●5H2O 249.6g

Try a few on your own:

Find the percent composition of the blue group of atoms 1. Cr(NH3)6Cl3·H2O2. Cu(C2H3O2)2·2NH33. (NH4)2Sn(OH)64. 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

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Compare the ratio of C:H in the following: CH2, C2H4, C3H6, C4H8, C5H10They 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 C20.0% 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 = 16.67

N/A

20.0g H x 1 mol = 20.0 mol H 1.0g

20.0 = 36.67

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% N2. 81.8% C, 18.2% H3. 39.1% Si, 61.0% O

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Calculating Molecular FormulasThus 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 C3H6All 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 FormulaEmpirical 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 MassEmpirical 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 NN = molecular mass = 34.0 = 34.0g = 2

Empirical mass 1.0 + 16.0g 17.0g

Step 2: Now find the molecular formulaMolecular 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 NFill 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:

1.62g x 22.4L = 104g/mol = molecular mass0.350L mol

This can now be substituted into our first equation:

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N = molecular mass = ? = 104g = 1Empirical mass POF3 104g

Step 2: Find the Molecular FormulaMolecular 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/molMol 0.0275mol

So N = molecular mass = 62.2g = 62.2g = 2Empirical mass SiH3 31.1g

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

Now it’s 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 & DilutionsA) Molarity (aka Molar Concentration)Molarity is a measure of CONCENTRATION or how much solute (eg. Salt) is dissolved in solvent (water).

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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 molesLitre

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

5.12g NaCl x 1 mol x 1000mL = 0.350mol/L = 0.350M250.0mL 58.5g 1L [NaCl] = 0.350MEg2. 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 NaOH1L 1molEg. 3. What is the molatrity of pure sulphuric acid H2SO4 having a density of 1.839 g/mL?

1.839 g x 1 mol x 1000 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 CaCl2•6H2O to 500.0mL?

15.00g x 1mol x 1000mL = 0.1369mol/L500.0mL 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 x 200.0mL = 0.200M1L 500.0mL [NaCl] = 0.200MEg2. 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] = 0.156MEg3. 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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