Solids, Liquids and Gases• Solids (s): Molecules are held in place by intermolecular

interactions.• Liquids (l): Molecules are held next to one another by

intermolecular interactions, however, these interactions are not strong enough to prevent the molecules from flowingpast one another.

• Gases (s): The intermolecular interactions are too weak to hold the molecules next to one another, so the molecules wander off on their own.

DiffusionDiffusion is the spontaneous spreading out of a substance, and isdue to the natural movement of its particles.

Ink being diffused in water

Cotton wool soaked with

ammonia solution

Cotton wool soaked with

hydrochloric acid Ammonium chloride

Diffusion of ammonia and hydrogen chloride gases

The mass of atoms are very small

• Therefore scientists cannot count individual atoms or molecules directly

• Instead a scientist count there particles by weighing a certain number of them

• Same as a bank cashier, by knowing the weight of one coin they can calculate the number of coins by dividing the total mass by the mass of one coin.

A. What is the Mole?

• A counting number (like a dozen)

• Avogadro’s number (NA)

• 1 mol = 6.02 1023 items

A large amount!!!!

The Mole

• The mole is simply a chemist’s measurement for the amount of a substance.

• It is a very large number (6 x 1023) as it ‘counts’ very small ‘objects’ such as atoms and molecules which have a very small masses.

• This rather strange number was not selected deliberately – instead, it is the number of atoms in a sample of any element that has a mass in grams that is numerically equal to the elements atomic mass.

• It allows chemists to work with masses in the lab which are feasible.

The Mole Definition

A mole of any substance is defined as the amount of the substance that contains as many

particles as 12 g of carbon – 12.

B. Molar Mass

• Mass of 1 mole of an element or compound.

• Atomic mass tells the...– That 6.01 *1023 atoms will weigh this

amount– 6.01 *1023 atoms of carbons will weigh 12

grams– grams per mole (g/mol)– 1 mole of carbon is 12 grams per litre

• Round to 2 decimal places

B. Molar Mass Examples

• carbon

• aluminum

• zinc

12.01 g/mol

26.98 g/mol

65.39 g/mol

B. Molar Mass Examples

• water

• sodium chloride

– 2(1.01) + 16.00 = 18.02 g/mol

– H2O

– NaCl– 22.99 + 35.45 = 58.44 g/mol

B. Molar Mass Examples

• sodium bicarbonate

• sucrose

– NaHCO3

– 22.99 + 1.01 + 12.01 + 3(16.00)

= 84.01 g/mol

– C12H22O11

– 12(12.01) + 22(1.01) + 11(16.00)

= 342.34 g/mol

If a magnesium atom is twice as heavy as an atom of carbon

Then 100 moles of magnesiumAre twice as heavy as 100 moles Of carbonTherefore if we have a piece ofMagnesium that is twice as heavy As a piece of carbon they have an Equal number of atoms

• Atom of carbon weighs

• Mr=12

• Atom of hydrogen weighs

• Mr=1One gram of hydrogen atoms and twelve grams of

Carbon atoms both have the same number of atoms

• Atom of silver weighs

• Ar=108

• Atom of carbon weighs

• Ar=12

If we have 12g of carbon how do we calulate the mass of silver if it has the same number of atoms

12g 24g 27g 40g 56g 63.5 108g

carbon magnesium

aluminium

calcium iron copper

silver

All these masses contain the same number of atoms

Avogadro`s number = 6 X 1023 Symbol L

It is the number of atoms of carbon in 12g of the carbon-12 isotope

The amount of a substance which contains 6 X 1023 particles is called a mole of that

substance

Symbol mol

• One mole of carbon weighs 12grams

• 48g of Mg (one mole weighs 24 g)

• One mole of carbon contains 6 X 1023

atoms

• Is 2 moles of Mg

Converting moles to grams

• Mass of one mole of an element= relative atomic mass in grams

Element Relative atomic mass

H 1

C 12

N 14

Mg 24

Cu 63.5

Fe 56

I 127

The mass of

one mole of

an element

=RelativeAtomic Massin grams

The relative molecular mass of a compound is the sum of the relative atomic masses of all the atoms in the compound

H2O = 2(1) + (16) = 18 (relative molecular mass of H2O)

CuSO4.H2O =(63.5) + (32) +4(16) +5 (18)=249.5

Converting Grams to moles

• 12g of carbon = 1 mole• 1g of carbon = 1/12 moles• 36g of carbon= 36/12=3 moles

How many moles of sulfuric acid (H2SO4) are

there in 12.25 g of sulfuric acid

• Relative mol mass =2(1) + (32) + 4(16) = 98

• Total mass/mass of one mole• 12.25/98=0.125 mole

Converting moles to number of atoms or molecules

• How many atoms of sodium are present in 0.25 moles of a metal

• 1 mole of Na contained 6x1023 atoms• 0.25 x 6x1023 atoms= 1.5 x1023

atoms

232 2

How many molecules are there in 0.5 mols of chlorine gas?How many atoms are there in 2 mols of water?How many electrons are there in 1.5 mols of Calcium?

1 mol of Cl 6 10 molecules [Cl

(i)(ii)(iii)

(i)23 23

2

232 2

23 242

is composed of molecules]

0.5 mols of Cl 0.5 6 10 3 10 molecules

1 mol of H 0 6 10 molecules [H O is composed of molecules]

2mols of H 0 2 6 10 1.2 10 moleculesEach molecule of water contai

(ii)

24 24 24

23

23 232

ns 3 atoms

1.2 10 molecules contains 3 1.2 10 3.6 10 atoms.

1 mol of Ca 6 10 atoms [Ca is composed of atoms]

1.5 mols of H 0 1.5 6 10 9 10 atomsEach atom of Calcium contains 20 electrons

9

(ii)

23 24 2510 atoms contain 23 9 10 2.07 10 electrons.

Molar Volume

• One mole of any gas should occupy the same volume as one mole of any other gas at the same conditions of temperature and pressure.

• Gases are often compared at STP, standard temperature and pressure.Standard pressure = 101325 PaStandard temperature = 273 K

• Molar Volume = 22.4 L = 22,400 cm3 = 2.24 x 10–2 m3

4

3

34

34

What is the volume in litres at STP of 0.01 mols of methane (CH ) gas?

How many mols are there in 280 cm of fluorine gas at STP?

1 mol of CH 22400 cm

0.01mols of CH 0.01 22400 224 cm

1 mol

(i)

(ii)

(i)

(ii) 32

32

32

32

of F 22400 cm1

mols of F = 1cm22400

1280 mols of F = 280cm

224000.0125 molsof F 280cm

Relative Molecular Mass

The relative molecular mass is the average mass of a molecule relative to one twelfth the mass of the carbon 12 atom.The relative molecular mass has no units as it is the ratio of two masses.

2 4 2

2 4

2

Calculate the relative molecular mass of sulfuric acid, (H SO ) Oxygen, (O ).

H SO 2(1) 32 4(16) 98

O 2(16) 32

(i) (a) (b)

(i) (a)

(b)

Molar Mass• The molar mass of a substance is the mass in grams

of one mole of the substance.• The molar mass has the same numerical value as its

relative molecular mass, but its units are grams (g).

2 4

6 12 6

2

2

Sulfuric aicd, H SO 98 98 gGlucose, C H O 180 180 g

Oxygen, O 32 32gSulfur dioxide, SO 64 64 g

Substance Relative molecular mass Molar mass

1 mol of carbon, 12 g

2

2

2

How many moles are there in 990 g of cabon dioxide (CO )The daily intake of calcium for an adult is 800 mg, how many moles of calcium is this?

1 mol of CO 44 g1

mols of CO 1g441

990 mols o44

(i)(ii)

(i)

2

2

3 3 3

3

f CO 990 g

22.5 mols of CO 990 g

1 mol of Ca 40 g1

mols of Ca 1g401

800 10 mols of Ca 800 10 g [800mg 800 10 g]400.02 mols of Ca 800 10 g

(ii)

More Mole Calculations

Moles of X

Volume of X(if a gas at STP) in litres at STP

Number of particles of X

Mass of X in g

÷ 6 x 10 23x 6 x 10 23

÷ 22.4

x 22.4

÷ molar massx molar mass

Notice that there is no direct link from particles to grams, you must first convert to moles. Likewise going from volume to mass.

23

23

2323

What is the mass of one atom of calcium?

1 mol of Ca 40 g

1 mol of Ca 6 10 atoms

6 10 atoms 40 g40

1 atom 6.67 10 g6 10

2

2

2

The Nissan Micra 1.5 Diesel SVE is quoted to have a CO emission figure of 120 g / km.For a 40 km round trip to work calculate:

the mass of CO produced.the number of moles of CO produced.the

(i)(ii)(iii) 2 volume of CO produced at room temperature and pressure.

[Molar volume at room temperature and pressure 24.0 litres]

120 40 4800 km

actual massno. of mols

molecular mass4800

no. of mols mols44

no. of m

(i)

(ii)

2

2

2

ols 109.09 mols

1mol of CO 24.0 L109.09 mols of CO 109.09 24L109.09 mols of CO 2594.16 L

(iii)

How many iron atoms should be consumed daily to meet the recommended daily intake of iron in the diet of 0.014 g?

acutalmassno. of mols

molecular mass0.014

no. of mols56

no. of mols 2.5 10

[H2007, Q4 (e)]

4

3

4 4 3

4 0

mols

1mol 6 10 atoms

2.5 10 mols 2.5 10 6 10 atoms

2.5 10 mols 1.5 10 atoms

Chemical Formulas

Empirical Formula

• The empirical formula of a compound is the formula that gives that simplest whole number ratio in which the atoms of the elements in the compound are present.

• The empirical formula of glucose, C6H12O6

(ratio of atoms is 6:12:6) is CH2O, as this is the lowest ratio (1:2:1) of the atoms C, H and O.

A sample of a brown coloured gas that is a major pollutant is found to contain 2.34 g of Nand 5.34 g of O. What is the simplest formula of the compound?

2.34mols of N atoms 0.167

145.34

mols of O atoms

2

0.33416

Ratio of Nitrogen atoms to Oxygen atoms 1: 2NO

A dry - cleaning fluid composed of carbon and chlorine was found to have the composition 14.5% C, 85.5% Cl. What is the empirical formula of this compound?

C 14r

Element Percentage Percentage SimplestA Ratio

2

.5% 14.5 11.208

12Cl 85.5% 85.5 2

2.40835.5

CCl

A 1.28 g sample of sulfur was allowed to react with an excess of chlorine to produce 4.12 g of a product that contains only sulfur and chlorine. What is the empirical formula of the compound?

Mass of s

2

ulfur consumed = 1.28 gMass of chlorine consumed = 4.12 1.28 2.84 g

1.28Mols of sulfur consumed = 0.04

322.84

Mols of chlorine comsumed = 0.0835.5

Ratio of sulfur atoms to chlorine atoms = 1:2

SCl

Molecular Formula

The molecular formula of a compound is the formula that gives the actual number of atoms of each element present in a molecule of the compound.

2

2

A colourless liquid used in rocket engines, whose empirical formula is NO ,has molecularmass of 92. What is the molecular formula?

Formula mass 46Molecular mass 92

92No. of NO in formula = 2

46Molecular for

2 2 42mula NO N O

Percentage Composition by Mass

If the empirical formula of a compound is known, the percentage by mass of each element present can be calculated. It can be useful to calculate the percentage of a certain element in a compound, for example the percentage of nitrogen in fertilisers.

2 2

r

Calculate the percentage by mass of nitrogen present in urea, CO(NH ) .

mass of nitrogen% of N 100

M of Urea2(14)

% of N 10012 16 2[14 2(1)]28

% of N 100 46.66%60

Structural Formulas

The structural formula of a compound shows the arrangement of the atoms within a molecule of the compound.

Some Examples:

You will come across many more structural formulas in the organic section of the course.

C

H

H

H H

Methane, CH4

C C

H

H

H

H

Ethene, C2H4

C C OH

H

H

H

H

H

Ethanol, C2H5OH

Chemical Equations

Chemical Equations• A chemical equation is a way of representing a chemical change.• It shows reactants and products.• To balance an equation means to change the numbers of each

molecule involved, so that the same number of atoms of each element appear on the reactants side and on the products side.

• Chemical equations balance on an atomic level, not molecular.• You cannot change the formula of a substance, i.e. if the

equation has NH3 you cannot change this you can only put a

number in front of it, 2NH3, increasing the number of N’s and the number of H’s.

• Never change the subscripts (small numbers).• It is possible to write balanced equations for lots of reactions but

that does not mean that the reaction actually takes place.

Formation of Water

(Oxygen does not exist as single atoms)

(This reaction does not take place)

Balancing an Equation

4 2 2 2

Methane reacts with oxygen, forming carbon dioxide and water vapour only.Write a balanced chemical equation for this reaction.

CH O CO H OExamine each element in turn, to have the same no. o

Solution

2

f of each on either side.(Leave the O until last as it appears the most often)Examining C: 1 on LHS and 1 on RHS, leave as they are.

Examining H: 4 on LHS and 2 on RHS, put a 2 in front of the H O, o

atoms

4 2 2 2

2

4 2 2 2

n the RHS.

CH O CO H O

Examining O: 2 on LHS and 4 on RHS, put a 2 in front of the O , on the LHS.

CH O CO2 H O

2

2

Balance the following2 2

2 2

2 2 2

2 2 2

4 10 2 2 2

4 10 2 2 2

2 2 3

2 2 3

Fe HCl FeCl HFe HCl FeCl H

H O H O

H O H O

C H O CO H O

C H O CO H O

Al O Al O

Al O A

2

l O

12

132

32

4 5

1

2

4.

2

.

.

3.

Balancing Redox Equations

• In working out what is oxidised and what is reduced in a reaction, it is important to remember that oxidation numbers are not a charge.

• Write the oxidation numbers below each atom to which it applies, as shown in the examples.

• Oxidation is an increase in oxidation number.• Reduction is a decrease in oxidation number.

1 2 6 22

2

2

12 7H O Cr O OH

2 2 3 32 7 2

2 2 3 3

6 2 2 1 3 3 1 2

6 3 36

2 7 2

3 32

3 2 32

2

3

3

2 3

2

2

Balance the following redox equation

Cr O Fe H Cr Fe H O

Cr O Fe H Cr Fe H O

Cr 3e Cr Cr 6e 2Cr

Fe e Fe Fe e Fe

Cr 6e 2CrF6 e

1.

2.

3.

27 2

3

2 3 32

2 3 32

2 2 3 32 7 2

e FeCr 6Fe 2Cr 6Fe

Cr 6Fe 2C

6 6

O H Hr 6Fe

Cr O 6Fe H 2Cr 6

O

Fe H14 7 O

4.

5.

Assign oxidation numbers to all the atoms in the equation.

Write down half equations of what is oxidised and reduced.Attach subscripts to atoms oxidised and reducedand balance the half equations.

Balance the half equations.

Attach species that were attached to the oxidised and reduced atoms.

Include all the original species and complete the balancing by inspection.

Calculations using Chemical Equations• A balanced equation for a chemical reaction

gives the relative amounts of each reactant and each product involved in the reaction.

• If the amount of one substance is known, based on the molar ratios in the equation, the amounts (masses, particles or volume)of other substances can be calculated.

• Make sure to always work in moles.

A dry - cleaning fluid composed of carbon and chlorine was found to have the composition 14.5% C, 85.5% Cl. What is the empirical formula of this compound?

C 14r

Element Percentage Percentage SimplestA Ratio

2

.5% 14.5 11.208

12Cl 85.5% 85.5 2

2.40835.5

CCl