1. IB1 Chemistry Quantitative chemistry 1
1.1 The mole concept and Avogadro's constant
1.2 Formulas
1.3 Chemical Equations

2. Topic 1: Quantitative chemistry
1.1 The mole concept and Avogadro’s constant Apply the mole concept to substances Determine the number of particles and the amount of substance (in moles). 1.2 Formulas Define the terms relative atomic mass (Ar) and relative molecular mass (Mr) Calculate the mass of one mole of a species from its formula Solve problems involving the relationship between the amount of substance in moles, mass and molar mass Distinguish between the terms empirical formula and molecular formula Determine the empirical formula from the percentage composition or from other experimental data Determine the molecular formula when given both the empirical formula and experimental data. 1.3 Chemical equations Deduce chemical equations when all reactants and products are given Identify the mole ratio of any two species in a chemical equation Apply the state symbols (s), (l), (g) and (aq). 1.4 Mass and gaseous volume relationships in chemical reactions Calculate theoretical yields from chemical equations Determine the limiting reactant and the reactant in excess when quantities of reacting substances are given Solve problems involving theoretical, experimental and percentage yield Apply Avogadro’s law to calculate reacting volumes of gases Apply the concept of molar volume at standard temperature and pressure in calculations Solve problems involving the relationship between temperature, pressure and volume for a fixed mass of an ideal gas Solve problems using the ideal gas equation, PV = nRT Analyse graphs relating to the ideal gas equation. 1.5 Solutions Distinguish between the terms solute, solvent, solution and concentration (g dm–3 and mol dm–3) Solve problems involving concentration, amount of solute and volume of solution.

3. (Almost) everything is made of atoms
Everything is made of atoms, although ideas about what an atom is have changed. Originally indestructible and not made from anything else, we now know they are not elementary particles. Nevertheless, chemists are interested in atoms as the basic units of matter in most chemical changes. Chemical changes includes all the millions of reactions in living and non-living things across the world where new substances are made. Organic and inorganic.
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4. The periodic table lists the elements in order of atomic number
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5. Atoms are rearranged to make new substances in chemical reactions.
Atoms are not made or destroyed, just rearranged, and this makes substances with completely different properties. Hydrogen explodes and water is made. Highly exothermic reaction. Not all chemical reactions are as fast or as dramatic, but all make new substances.
Hydrogen + Oxygen  Water
2 H2 + O2  2 H2O
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6. Kinetic theory: atoms in solids, liquids and gases:
Particles (atoms, molecules, ions, etc) in constant random motion. Changes of phase do not create new substance. Some properties can be explained using particle theory. PhET states of matter simulations quick run through solids liquids and gases noting monatomic gases, diatomic oxygen and compound water.

7. 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.
A grain of sand is 10-3m and a planet is about 107m so a grain of sand is about 1010 times smaller. A human is around 1m.
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8. Chemists need to count atoms...
... because atoms join together in definite ratios
2H always joins with 1O to make water. if there are 2Os it is hydrogen peroxide. Ratios written in chemical formulas. Difficult because atoms are so small- 1g of hydrogen has about 600 thousand million million million. Use the mole to count atoms where 1 mole = 6.02x1023 atoms

9. Definition of a mole n = amount of substance in units of mol 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

10. How many in 1 mole? 1 mol of anything = 6.02×1023 units of that thing
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
but what is the unit? a mole of atoms or of molecules
In 1 mol of H2-molecules there is 2 mol Hydrogen atoms = 2
In 1 mol of H2O molecules there is 3 mol of atoms; 2 mol of H-atoms and 1 mol of O-atoms = 3 =2 =1

11. 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
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12. 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?

13. How heavy is an atom?
12g of carbon is made of 600 thousand million million atoms (6 x 1023)
Use of different units

14. How many atoms is the Earth made of?
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15. Ammonia (NH3)
1.2L of ammonia gas (NH3) contains 3.01×1022 molecules. Calculate the number of moles of hydrogen in 12L of ammonia.

16. Relative mass Mass of a 12C-atom =12 by definition
All other atomic or molecular masses relative to 12C
Masses of single atoms and single molecules:
Relative atomic mass, Ar
Relative molecular mass, Mr
Relative formula mass, Mr
Relative masses have no units in IB chemistry (but can be u or amu)

17. Relative atomic mass, Ar
weighted mean mass of the naturally occuring isotopes
Iron Ar =
mix of Fe-54, Fe-56, Fe-57, and Fe-60
but mass of 1 mole of H2 molecules? = 1.014
Hydrogen Ar = 1.007
mix of H-1, H-2 and H-3

18. 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

19. Molar mass in g mol-1 55.8 g mol-1 12.0 g mol-1 18.0 g mol-1

20. 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) × ×4 = 132 gmol-1 The Molar mass of CuSO4.5H2O M= ×4 + 5(2×1 + 16) = gmol-1

21. Down: divide Up: multiply
Relationship between the amount of substance in moles, mass and molar mass
Quantity
Symbol
Unit
Mass m g
Molar mass M
gmol-1
Number of moles n
mol
or use triangle or formula m=nM
Down: divide Up: multiply

22. Example mole calculations
Formula of compound  Molar mass
Draw table
Complete and calculate

23. 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

24. 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

25. Magnesium and hydrochloric acid
Observations and measurements
Explanations
Further Questions

26. Describing chemical changes
Liquid oxygen and liquid hydrogen fuel
Burn hydrogen in oxygen and show water condensed.
Discuss why knowing reaction masses is critical (for example as Liquid hydrogen and oxygen in Saturn V for Apollo 11)
Launch of Apollo 11 (from 6 min)
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27. 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

28. Balanced equations show the mole ratio of reactants and products
Atoms are not made or destroyed, just rearranged, and this makes substances with completely different properties. Hydrogen explodes and water is made. Highly exothermic reaction. Not all chemical reactions are as fast or as dramatic, but all make new substances.
Hydrogen + Oxygen  Water
2 H2 + O2  2 H2O
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29. 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)

30. Balancing a chemical equation
Propane + oxygen  carbon dioxide + water _ C3H8(g) + _ O2(g)  _ CO2(g) + _ H2O(l)

31. 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

32. Percentage composition by mass
Mr = 18 mass of H = 2×1 % H by mass = 2/18×100 = 11%

33. 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)
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34. Calculate the formula from experimental data
iron oxide: 1.12 g of iron burn in oxygen to give 1.44 g of iron oxide
zinc oxide: g of zinc react with 0.08 g of oxygen
copper oxide: 32 g of copper react to give 39 g of copper oxide
calcium oxide: 4.0 g of calcium reacts with 1.6g of oxygen

35. Empirical formula from percentage composition
Assume that you have 100g of the compound
Calculate number of moles
Compare Mole ratios to find the formula

36. Calculate the formula for the compound with 70. 58 % C, 5. 93 % H, 23
Calculate the formula for the compound with % C, 5.93 % H, % O C H O
mass%
70.58
5.93
23.49
% (=100%) m
g (=100g) M
12.01
1.01
16.00
g/mol n
5.88
5.87
1.47
mol

37. divide by the lowest to find the ratio
C H O
4 : : 1
Empirical formula
C4H4O

38. 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

39. 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

40. Measuring chemical quantities: solids
in grams
using a balance (precision ±0.01g or ±0.001g)
Precision and accuracy
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41. Measuring chemical quantities: liquids
in litres (L) or decimetres cubed (dm3) 1L = 1dm3
In mL or centimetres cubed (cm3) 1mL = 1cm3
1000 mL = 1L
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42. Converting mass to volume
volume = mass x density
volume of pure water is 1.0 gcm-3

43. 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)

44. Links Powers of ten http://www.powersof10.com/
hydrogen explosion
states of matter phet