ib1 chemistry quantitative chemistry 1 1.1 the mole concept and avogadro's constant 1.2...
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IB1 ChemistryQuantitative chemistry 1
1.1 The mole concept and Avogadro's constant1.2 Formulas
1.3 Chemical Equations
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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.
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(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
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The periodic table lists the elements in order of atomic number
Image: http://en.wikipedia.org/wiki/File:Periodic_table.svg
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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
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Kinetic theory:atoms in solids, liquids and gases:
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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
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Chemists need to count atoms...
... because atoms join together in definite ratios
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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
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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
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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
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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?
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How heavy is an atom?
12g of carbon is made of 600 thousand million million atoms (6 x 1023)
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How many atoms is the Earth made of?
Image: http://en.wikipedia.org/wiki/Earth
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Ammonia (NH3)
1.2L of ammonia gas (NH3) contains 3.01×1022 molecules.
Calculate the number of moles of hydrogen in 12L of ammonia.
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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)
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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
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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
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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
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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
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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
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Example mole calculations
1. Formula of compound Molar mass
2. Draw table
3. Complete and calculate
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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
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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
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Magnesium and hydrochloric acid
Observations and measurements
Explanations Further Questions
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Describing chemical changes
Liquid oxygen and liquid hydrogen fuel
Image: http://en.wikipedia.org/wiki/Saturn_V
Launch of Apollo 11 (from 6 min)
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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
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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
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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)
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Balancing a chemical equation
Propane + oxygen carbon dioxide + water
_ C3H8(g) + _ O2(g) _ CO2(g) + _ H2O(l)
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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
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Percentage composition by mass
Mr = 18
mass of H = 2×1
% H by mass = 2/18×100
= 11%
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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
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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
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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
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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
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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
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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
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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
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Measuring chemical quantities: solids
in grams
using a balance (precision ±0.01g or ±0.001g)
Image: http://en.wikipedia.org/wiki/Weighing_scale
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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
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Converting mass to volume
volume = mass x density
volume of pure water is 1.0 gcm-3
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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)
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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