IB1 ChemistryQuantitative chemistry 1

1.1 The mole concept and Avogadro's constant1.2 Formulas

1.3 Chemical Equations

Topic 1: Quantitative chemistry1.1 The mole concept and Avogadro’s constant1.1.1 Apply the mole concept to substances.1.1.2 Determine the number of particles and the amount of substance (in moles).

1.2 Formulas

1.2.1 Define the terms relative atomic mass (Ar) and relative molecular mass (Mr).

1.2.2 Calculate the mass of one mole of a species from its formula.1.2.3 Solve problems involving the relationship between the amount of substance in moles, mass and molar mass.1.2.4 Distinguish between the terms empirical formula and molecular formula.1.2.5 Determine the empirical formula from the percentage composition or from other experimental data.1.2.6 Determine the molecular formula when given both the empirical formula and experimental data.

1.3 Chemical equations1.3.1 Deduce chemical equations when all reactants and products are given.1.3.2 Identify the mole ratio of any two species in a chemical equation.1.3.3 Apply the state symbols (s), (l), (g) and (aq).

1.4 Mass and gaseous volume relationships in chemical reactions1.4.1 Calculate theoretical yields from chemical equations.1.4.2 Determine the limiting reactant and the reactant in excess when quantities of reacting substances are given.1.4.3 Solve problems involving theoretical, experimental and percentage yield.1.4.4Apply Avogadro’s law to calculate reacting volumes of gases.1.4.5 Apply the concept of molar volume at standard temperature and pressure in calculations.1.4.6 Solve problems involving the relationship between temperature, pressure and volume for a fixed mass of an ideal gas.1.4.7 Solve problems using the ideal gas equation, PV = nRT.1.4.8 Analyse graphs relating to the ideal gas equation.

1.5 Solutions1.5.1 Distinguish between the terms solute, solvent, solution and concentration (g dm–3 and mol dm–3).1.5.2 Solve problems involving concentration, amount of solute and volume of solution.

(Almost) everything is made of atoms

Images:http://en.wikipedia.org/wiki/Rock_(geology), http://en.wikipedia.org/wiki/Tree , http://en.wikipedia.org/wiki/Human, http://en.wikipedia.org/wiki/Star

The periodic table lists the elements in order of atomic number

Image: http://en.wikipedia.org/wiki/File:Periodic_table.svg

Atoms are rearranged to make new substances in chemical reactions.

Hydrogen + Oxygen Water

2 H2 + O2 2 H2O

Image: http://labspace.open.ac.uk/mod/resource/view.php?id=438900

Kinetic theory:atoms in solids, liquids and gases:

How big is an atom? about 10-

10m

if an atom was the size of a grain of sand, humans would be the size of a planet.

Image: http://en.wikipedia.org/wiki/Earth, http://en.wikipedia.org/wiki/Sand

Chemists need to count atoms...

... because atoms join together in definite ratios

Definition of a mole

1mol =

Avogadro’s constant=

L= NA=

6.02×1023

number of atoms in 12 grams of pure carbon-12

n = amount of substance in units of mol

How many in 1 mole?

1 mol equals: 6.02×1023 Hydrogen atoms, H 6.02×1023 Hydrogen molecules, H2

6.02×1023 Water molecules, H20

6.02×1023 formula units of Sodium Chloride, NaCl

1 mol of anything = 6.02×1023 units of that thing

A mole of atoms is like a box of eggs...

an eggbox contains 6 eggs

a mole contains 600 thousand million million million atoms

23 grams of sodium is 1 mole

Image:http://commons.wikimedia.org/wiki/File:Activated_Carbon.jpg, http://en.wikipedia.org/wiki/Sodium

Eggs and moles

1. a) How many eggs in 2 boxes?

b) How many in 0.5 boxes?

c) How many eggs in 20 boxes?

2. a) How many atoms in 2 moles?

b) How many atoms in 0.5 moles?

c) How many atoms in 20 moles?

How heavy is an atom?

12g of carbon is made of 600 thousand million million atoms (6 x 1023)

How many atoms is the Earth made of?

Image: http://en.wikipedia.org/wiki/Earth

Ammonia (NH3)

1.2L of ammonia gas (NH3) contains 3.01×1022 molecules.

Calculate the number of moles of hydrogen in 12L of ammonia.

Relative mass

Mass of a 12C-atom =12 by definition

All other atomic or molecular masses relative to 12C

Þ Masses of single atoms and single molecules:• Relative atomic mass, Ar

• Relative molecular mass, Mr

• Relative formula mass, MrRelative masses have no units in IB chemistry (but can be u or amu)

Relative atomic mass, Ar

weighted mean mass of the naturally occuring isotopes

Hydrogen Ar = 1.007mix of H-1, H-2 and H-3

Iron Ar = 55.845mix of Fe-54, Fe-56, Fe-57, and Fe-60

Relative molecular mass, Mr

Relative formula mass, Mr

The relative molecular mass is the relative mass of the atoms in one molecule.

The formula mass is the sum of the relative mass of atoms in the formula for an ionic coumpound

Molar mass in g mol-1

55.8 g mol-1

http://en.wikipedia.org/wiki/Iron, http://en.wikipedia.org/wiki/File:Hydrogen_discharge_tube.jpg, http://en.wikipedia.org/wiki/Graphite

12.0 g mol-1

1.0 g mol-1

18.0 g mol-1

Calculating molar mass from Periodic table

The Molar mass of water , H2O

M = 2×1+16 = 18gmol -1

The Molar mass of (NH4)2SO4

M= (14+4×1) × 2 + 32 + 16×4 = 132 gmol-1

The Molar mass of CuSO4.5H2O

M= 63.5 + 32 + 16×4 + 5(2×1 + 16) = 249.5 gmol-1

Relationship between the amount of substance in moles, mass and molar mass

Down: divide

Up: multiply

Quantity Symbol Unit

Mass m g

Molar mass

M gmol-1

Number of moles

n mol

Example mole calculations

1. Formula of compound Molar mass

2. Draw table

3. Complete and calculate

Calculate the number of moles in 34 g of Ammonia.

Quantity Symbol Unit

Mass m g

Molar mass

M gmol-1

Number of moles

n mol

Calculate the mass of 0.50 mol of NaCl.

Quantity Symbol Unit

Mass m g

Molar mass

M gmol-1

Number of moles

n mol

Magnesium and hydrochloric acid

Observations and measurements

Explanations Further Questions

Describing chemical changes

Liquid oxygen and liquid hydrogen fuel

Image: http://en.wikipedia.org/wiki/Saturn_V

Launch of Apollo 11 (from 6 min)

Chemical equations describe a chemical change where reactants change into products

Word equations

Hydrogen + Oxygen Water

Chemical equations (symbol equations)

H2 (g) + O2 (g) H2O (l)

reactants products

state symbols show solid, liquid, gas or aqueous

Balanced equations show the mole ratio of reactants and products

Hydrogen + Oxygen Water

2 H2 + O2 2 H2O

Image: http://labspace.open.ac.uk/mod/resource/view.php?id=438900

2 H2 + O2 2 H2O

green numbers = subscripts

cannot be changed (a compound has one formula- changing the formula changes the compound)

red numbers = coefficients

changes to balance the reaction (coefficients valid only for a specific reaction)

Balancing a chemical equation

Propane + oxygen carbon dioxide + water

_ C3H8(g) + _ O2(g) _ CO2(g) + _ H2O(l)

C3H8(g) + 5 O2(g) 3 CO2(g) + 4 H2O(l)

What does an equation tell us about molecules in the reaction?

you need 5 oxygen molecules for 1 molecule of propane

1 propane molecule will produce 3 carbon dioxide molecules and 4 Water molecules

2 molecules of propane produce 6 molecules of carbon dioxide

1 mole of propane produces 3 moles of carbon dioxide

Percentage composition by mass

Mr = 18

mass of H = 2×1

% H by mass = 2/18×100

= 11%

Empirical and Molecular formula

Molecular formula: shows the actual number of each atom/element in a compound, e.g.

Ethane C2H6

Glucose C6H12O6

Empirical formula: Shows only the ratio of the elements in a compound, e.g.

Ethane CH3

Glucose CH2O

(Formulas of salts are empirical formulas)Image: http://en.wikipedia.org/wiki/Ethane, http://en.wikipedia.org/wiki/Glucose

Calculate the formula from experimental data

a) iron oxide: 1.12 g of iron burn in oxygen to give 1.44 g of iron oxide

b) zinc oxide: 3.25 g of zinc react with 0.08 g of oxygen

c) copper oxide: 32 g of copper react to give 39 g of copper oxide

d) calcium oxide: 4.0 g of calcium reacts with 1.6g of oxygen

Empirical formula from percentage composition

1. Assume that you have 100g of the compound

2. Calculate number of moles

3. Compare Mole ratios to find the formula

Calculate the formula for the compound with 70.58 % C, 5.93 % H, 23.49 % O

C H O

mass% 70.58 5.93 23.49 % (=100%)

m 70.58 5.93 23.49 g (=100g)

M 12.01 1.01 16.00 g/mol

n 5.88 5.87 1.47 mol

divide by the lowest to find the ratio

C H O

5.88 5.87 1.47 1.47 1.47 1.47

4 : 4 : 1Empirical formula

C4H4O

Molecular formula

with Empirical formula, C4H4O, and Molar mass 136 g/mol

you can calculate the Molecular formula.

C4H4O M = 68 g/mol Too Low

C8H8O2 M = 136 g/mol Correct

C12H12O3 M = 204 g/mol Too High

Measurement of mass and volume

Apparatus Quantity measured

Units Range Scale uncertainty

Precision

25mL Measuring cylinder

volume mL or cm3

0 – 25 ±0.25 ±0.5

10mL Pipette

25mL Burette

Centigram balance

Milligram balance

Volumetric flask

Measuring chemical quantities: solids

in grams

using a balance (precision ±0.01g or ±0.001g)

Image: http://en.wikipedia.org/wiki/Weighing_scale

Measuring chemical quantities: liquids

in litres (L) or decimetres cubed (dm3) 1L = 1dm3

In mL or centimetres cubed (cm3) 1mL = 1cm3

1000 mL = 1L

Images: http://commons.wikimedia.org/wiki/Category:Laboratory_glassware

Converting mass to volume

volume = mass x density

volume of pure water is 1.0 gcm-3

Solutions

a solute dissolved in a solvent gives a solution

units of concentration grams per litre (gL-1)

or moles per decimetre cubed (moldm-3)

Links

Powers of ten http://www.powersof10.com/

hydrogen explosion http://www.youtube.com/watch?v=DjcztiNGg_8

states of matter phet http://phet.colorado.edu/en/simulation/states-of-matter