Chemistry 11 - Eckert - homeMole+Notes... · Web viewDivide the mass of each element by its molar mass to find the number of moles. Find the simplest whole number ratio. Often it

Chemistry 11: The Mole

Introduction

What is a mole? Yes it is a rat like creature with bad eyesight that tunnels

through the ground. But in chemistry, it means something more…much more.

Before we tackle that, let’s get some back ground info.

Pair : 1 pair of shoelaces = 2 shoelaces

Dozen : 1 dozen oranges = 12 oranges

Gross : 1 gross of pencils = 144 pencils

Ream : 1 ream of paper = 500 sheets of paper

Mole : 1 mole of cars = 6.02x1023 cars

(602000000000000000000000 cars)

What is a mole?

It is an amount, defined as the number of carbon atoms in exactly 12

grams of carbon-12.

1 mole = 6.02 x 1023 of the representative particles. (cars, donuts, atoms)

Treat it like a very large dozen

6.02 x 1023 is called: Avogadro’s number

mole of marbles would cover the earth with a layer 3 miles thick.

1 mol cars or 6.02 x 1023 cars

XDM Prodns (in conjunction with Simmer Down Inc.) 1

Molar Mass

Molar mass is the mass of one mole, of each element.

The units of molar mass are g/mol.

The molar mass of iron is different from the molar mass of oxygen, just

like the mass of a dozen donuts is different from a dozen eggs.

The molar mass of an element is the atomic mass shown on the periodic

table, expressed in grams.

For example:1 mole of lead has a mass of 207.2 g

207.2 g of Pb or 1 mole 1 mole 207.2 g of Pb

Write the molar mass of boron.

Write the mass of one mole of Lithium.

Write the mass of one mole of Calcium.

Write the mass of one mole of Iodine.


Molar mass in a compound is found by adding the atomic masses in grams

of all the elements in the compound.

For example: Find the molar mass of water, H2O

1 mole of O = 16.0g2 moles of H = +(2 x 1.0g)

18.0g of H2O/1 mole H2O

Example: What is the molar mass of Na2SO4

2 moles of Na = (2 x 23.0g) 46.0g4 moles of O = (4 x 16.0g) 64.0g1 mole of S = (1 x 32.1g) 32.1g

142.1g

142.1g of Na2SO4

1 mole Na2SO4


Try these ones…

1. C12H22O11

2. (NH4)2SO4

3. Cu(NO3)26H2O

4. Al2(SO4)3


Mass, molar mass and the mole. How are they connected?

You can calculate to following using the unit conversions below:

Moles to molecules and vice versaMoles to mass (grams) and vice versa

Mass Mole Molecules

XDM Prodns (in conjunction with Simmer Down Inc.)

1 molx grams

x grams1 mol

6.02 x 10 23 molecules 1 mol

1 mol6.02 x 1023molecules

5

Examples: 1.

2.

Examples 1. How many moles of Aluminum are in 54.0g of Al Mind map: 54.0g Al ? mol Al

1 mol Al 27.0 g Al

Ans: 54.0 g Al x 1 mol Al = 2.00 mol Al 27.0 g Al

Example 2. How many molecules in 4.25 mol of water?Mind map: 4.25 mol H2O ? molecules H2O

6.02 x 10 23 molecules H 2O 1 mol H2O

Ans: 4.25 mol H2O x 6.02 x 10 23 molecules H 2O = 2.59 x 1024 molecules H2O 1 mol H2O


(Unit Factor!!!!!!!!!!!!!!!!!!!!!)

(Unit Factor!!!!!!!!!!!!!!!!!!!!!)

6

Try these ones…

1. How many grams in 4.20 mol C12H22O11?

2. How many grams in 0.125 mol (NH4)2SO4?

3. How many moles in 239.4g Cu(NO3)26H2O

4. How many moles in 1.02g Al2(SO4)3?


Moles and The Volume of a gas!!!!

Avogadro’s hypothesis…Equal volumes of different gases, at the same temperature and pressure, contain the same number of particles.

The molar volume of a gas is the volume occupied by one mole of the gas.

Standard Temperature and Pressure (STP) = 0°C and 101.3 kPa

Avogadro’s Hypothesis is interpreted to mean:

1 mole of ANY GAS at STP has a volume of 22.4L

In other words,

1 mol gas or 22.4L gas 22.4L gas 1 mol gas

ExamplesFind the volume at STP1. 4.39 moles of C3H8

2. 1.00 mole of NO3

Find the mass given the volume at STP

3. 65.4 g of N2

4. 76.0 g of F2


Mass Mole Molecules

Volume


1 molx grams

x grams1 mol


6.02 x 10 23 molecules x mol

1 mol22.4 L (@STP

22.4 L @ STP1 mol

Gases only!!!

9

Percent Composition

Percent composition: is the percentage (by mass) of the species in a

chemical formula.

The molar masses of the atoms can be used to calculate percentage

composition by mass.

The formula of a compound gives its elemental composition in terms of

moles of atoms. ex. 1 mole H2O = 2 moles H atoms and 1 mole O atom

To calculate the % composition, see the examples below


Practice Problems: Determine the percent composition for each of the elements in the compound.

1. aluminum sulfide

2. nickel(II) iodide

3. calcium cyanide, Ca(CN)2


Empirical Formulas / Molecular Formulas

Scientists use two kinds of formulas: molecular and empirical.

A molecular formula describes the molecule as a whole; it’s subscripts

show the actual number of each atoms in the molecule. (Example: H2O,

C6H6 (Benzene))

Empirical formula gives the simplest whole number ratio of the atoms in a

compound.

Examples: C6H6 is Benzene. Its empirical formula is CH.

CH2, C2H4, C3H6, C4H8 and C5H10 all have twice as many H’s as C’s. The

empirical formula for all is CH2

Finding the empirical formula is basically the opposite of percent

composition. Follow these 4 easy steps and you will be golden

1. Find the mass of each element in the sample. If % composition data

is given, assume the sample weighs 100 g. If one mass is a

percentage or missing, find it by taking the difference (Subtract )

2. Divide the mass of each element by its molar mass to find the

number of moles.

3. Find the simplest whole number ratio. Often it can be calculated by

inspections, if not, divide ALL the molar quantities by the smallest

one. If these quotients are not whole numbers, multiply then by 2, 3,

4 etc to make them a whole number ratio.

4. Write the formula. (The atom ratio will be the same as the mole ratio

as determined in step 3)


***Always carry out calculators to 3 or 4 digits and NEVER round off

intermediate values. Rounding will cause errors when trying to clear

the fraction***

Examples:


Molecular formula

The molar mass of benzene (C6H6) is 6 times the molar mass of CH, its

empirical formula. Hydrogen peroxide (H2O2) is twice the molar mass of

HO, its empirical formula. Notice anything?

The formula of a molecular compound is always a whole-number multiple

of its empirical formula.

This implies that, the molar mass of a molecular formula is always a

whole-number multiple of its empirical formula.

To find this multiple, we must divide the molar mass by the empirical

mass.

Note: empirical mass is the molar mass of the empirical formula

Multiple = N = molar mass Empirical mass

Since the molar mass is a multiple of the empirical mass, then the

molecular formula must be the same multiple of the empirical formula:

Molecular formula = N x empirical formula

Recap:1) Divide the molar mass of the compound by the empirical mass

to find the whole number multiple.2) Multiply the empirical formula by this multiple to get the

molecular formula.

Examples:


Molar Concentration - Molarity

This section is concerned with the idea of concentration. (Symbol = [ ])

The concentration of a solution is the amount of solute in a given quantity

of solvent or solution. It is a measure of the solution’s “strength”.

A concentrated solution has a high concentration.

A dilute solution has a low concentration.

Knowing the concentration of a solution provides a way to find how

much of a particle substance exists in a given volume of the solution.

In chemistry, the # of moles is usually more important than mass (grams),

so a concentration unit based on moles is better than one based on mass.

Remember the following…

Molarity (M) is defined as the number of moles of solute per litre of

solution:

Molarity = Number of moles of solute Volume of solution in LITRES

Molarity is sometimes called molar concentration.

The units for Molarity are mol/L.

mol/L = M which is written as molar.


Some helpful Diagrams

Remember:Molarity M = moles

L

mol

M vol

Molarity

Mass Mole Molecules

Volume


1 molx grams

x grams1 mol


6.02 x 10 23 molecules x mol

1 mol22.4 L (@STP

22.4 L @ STP1 mol

Gases only!!!

x molL

16

Examples Give the answers to the correct sig figs and in M (mol/L)

1. What is the [ ] of a solution if 3.25 mol are dissolved into 6.500 L?

2. What is the [ ] of a solution if 0.0256 g of silver nitrate are dissolved into

6.50 L of water?

3. Find the number of moles in 10.00 mL of 3.78 M NaCl solution.

4. What is the [ ] of a solution if 1.806x1024 molecules of an unknown

substance are dissolved into 3.00 L of water?


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Chemistry 11 - Eckert - homeMole+Notes... · Web viewDivide the mass of each element by its molar mass to find the number of moles. Find the simplest whole number ratio. Often it