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Transcript Unit3_Stoichiometry_vs2

Unit 3 Stoichiometry
Topic 1
November 2014
Lesson 1
Review of Matter and Phases
Friday, October 24, 2014
Lesson 1 – Review of Stuff!
• Atoms of different elements combine in fixed ratios to form
compounds, which have different properties from their component
• Names and symbols of the elements are in the IB data booklet in
Section 5.
• Mixtures contain more than one element and/or compound that are
not chemically bonded together and so retain their individual
• Mixtures are either homogeneous or heterogeneous.
Lesson 1 – Ionic Compounds
Applications and Skills:
• Deduction of chemical equations when reactants and products are specified.
• Balancing of equations should include a variety of types of reactions.
• Application of the state symbols (s), (l), (g), and (aq) in equations.
• Explanation of observable changes in physical properties and temperature
during changes of state.
• Names of the changes of state – melting, freezing, vaporization (evaporation
and boiling), condensation, sublimation and deposition – should be covered.
• The term ‘latent heat’ is not required.
Language of Chemistry
• Elements are represented by a chemical symbol, consisting of
either an upper case letter by itself or an upper case letter and a
lower case letter
• E.g. Na, F, He, etc.
• Section 5 of your data booklet has a complete list of names of
all elements
• Some elements are found in nature in their native form, like
gold. Others are only found in chemical compounds which
contain a fixed proportion of elements and are held together by
chemical bonds
• We represent chemical compounds using symbols to represent
the elements in the compound with subscripts show how many
atoms are in one molecule of that compound
• H2O  2 atoms of Hydrogen for 1 atom of Oxygen (when there is
no subscript, there is only one atom)
Chemical Equations
• A chemical change can be represented by a chemical equation
where the left hand shows the reactants (what you start with)
and the right side shows the products (what you end up with!)
• Since matter is neither created nor destroyed, the total number
of atoms of each element needs to be the same on both sides of
a chemical equation!
• To change the number of atoms of an element and balance an
equation, we can use stoichiometry coefficients in front of the
compounds and elements
Balancing Equations
Let’s Practice
• Write an equation for the reaction of thermal decomposition of
sodium hydrogencarbonate (NaHCO3) into sodium carbonate
(Na2CO3), water (H2O), and carbon dioxide (CO2).
2 NaHCO3 → Na2CO3 + H2O + CO2
Atoms on the Left
Atoms on the Right
Let’s Practice
1 Write balanced chemical equations for the following reactions:
(a) The decomposition of copper carbonate (CuCO3) into copper(II) oxide (CuO) and
carbon dioxide (CO2).
(b) The combustion of magnesium (Mg) in oxygen (O2) to form magnesium oxide
(c) The neutralization of sulfuric acid (H2SO4) with sodium hydroxide (NaOH) to
form sodium sulfate (Na2SO4) and water (H2O).
(d) The synthesis of ammonia (NH3) from nitrogen (N2) and hydrogen (H2).
(e) The combustion of methane (CH4) to produce carbon dioxide (CO2) and water
2 Write balanced chemical equations for the following reactions:
(a) K + H2O → KOH + H2 (b) C2H5OH + O2 → CO2 + H2O (c) Cl2 + KI → KCl + I2
(d) CrO3 → Cr2O3 + O2
(e) Fe2O3 + C → CO + Fe
Let’s Practice
3. Use the same processes to balance the following examples:
C4H10 + O2 → CO2 + H2O
NH3 + O2 → NO + H2O
Cu + HNO3 → Cu(NO3)2 + NO + H2O
H2O2 + N2H4 → N2 + H2O + O2
C2H7N + O2 → CO2 + H2O + N2
Lesson 1 cont.
Monday, October 27, 2014
Lesson 1 cont.
• A pure substance includes all elements and compounds; matter
that has consistent composition
• An element is a pure substance that contains only one type of
• An atom is the smallest part of an element that can still be
recognized as that element
• A compound consists of two or more atoms of different
elements chemically combined in a fixed ratio (Law of Fixed
• Ions are charged species; atoms or compounds that have gained
or lost electrons
• Anions are negatively charged and have more electrons than
• Cations are positively charged and have more protons than
• Polyatomic ion contain two or more atoms of different elements
covalently bonded (sharing electrons) but with an overall charge
Chemical vs. Physical Properties
• Chemical properties dictate how something reacts in a chemical
• Physical properties are all the other properties and can be
observed without changing that substance – melting point,
boiling point, color, density, etc.
Chemical Changes
• Synthesis – smaller substances combining into larger, more
complex substances: C(s) + O2(g)  CO2(g)
• Decomposition – reactions involving a single reactant breaking
down into smaller parts: 2 H2O2(l)  2 H2O(l) + O2(g)
• Single Displacement – one element replaces another in a
compound based on reactivity: Mg(s) + 2HCl(aq)  MgCl2(aq)
+ H2(g)
• Double Displacement – two compounds that ionize in solution
swap partners to form one insoluble product: AgNO3(aq) +
CuCl(aq)  AgCl(s) + CuNO3(aq)
• Mixtures consist of elements and compounds physically
combined (but not chemically bound together)
• A homogeneous mixture is one where it has uniform
composition and properties throughout
• A heterogeneous mixture is one where it does not have a
uniform composition
• Mixtures can be separated in a number of different ways that do
not rely on chemical reactions
Let’s Practice
4. Classify the following mixtures as homogeneous or
(a) sand and water
(b) smoke
(c) sugar and water
(e) ethanol and water in wine
(d) salt and iron filings
(f) steel
Separating Mixtures
States of Matter
• Matter can exist in three states – solid, liquid and gas
• Notice how particles are closest together in the solid phase and
furthest apart in the gas phase.
States of Matter
• particles close packed
• inter-particle forces
• strong, particles
vibrate in position
• fixed shape
• fixed volume
• particles more spaced
• inter-particle forces
weaker, particles can
slide over each other
• no fixed shape fixed
• particles spread out
• inter-particle forces
• negligible, particles
move freely
• no fixed shape
• no fixed volume
• Liquids and gases are referred to as fluids, which refers to their
ability to flow
• Fluids are able to undergo diffusion, the process by which the
particles of a substance become evenly distributed, as a result of
their random movements.
• Diffusion rates are related to the average kinetic energy of the
particle: KE = (1/2)mv2
• Substances with lower mass diffuse more quickly than those with
higher mass.
Kinetic Theory
• The average kinetic energy of a system is directly related to the
temperature of that system in Kelvin
• The state of matter a substance is in depends on temperature,
pressure and the strength of its intermolecular (inter-particle)
• The SI Unit for temperature is the Kelvin; we also use Celsius
which has the same separation between degrees as Kelvin but a
different scale
• Absolute Zero on the Kelvin scale is the temperature where all
movement of particles stops!
• K = C + 273
Let’s Practice
• Which of the following has the highest average kinetic energy?
A. He at 100 °C
B. H2 at 200 °C
C. O2 at 300 °C
D. H2O at 400 °C
Let’s Practice
5 Write balanced equations for the following reactions and apply state symbols to all
reactants and products, assuming room temperature and pressure unless stated
otherwise. If you are not familiar with the aqueous solubilities of some of these
substances, you may have to look them up.
• (a) KNO3 → KNO2 + O2 (when heated, 500°C)
• (b) CaCO3 + H2SO4 → CaSO4 + CO2 + H2O
• (c) Li + H2O → LiOH + H2
• (d) Pb(NO3)2 + NaCl → PbCl2 + NaNO3 (all reactants are in aqueous
• (e) C3H6 + O2 → CO2 + H2O (combustion reaction)
6 A mixture of two gases, X and Y, which both have strong but distinct smells, is
released. From across the room the smell of X is detected more quickly than the
smell of Y. What can you deduce about X and Y?
7 Ice floats on water. Comment on why this is not what you would expect from the
kinetic theory of matter.
6 X has diffused more quickly, so it must be a lighter gas. Its particles
have greater velocity than the particles of Y at the same temperature.
(Note though that they will both have the same value for average kinetic
7 From the kinetic molecular theory we would expect a solid to be more
dense than its liquid, and therefore that ice would sink in water.
Phase Changes
• A phase change is a physical change in the state of matter
Energy of Phase Changes
Phase Change Diagram
• a–b As the solid is heated, the vibrational energy of its particles increases and
so the temperature increases.
• b–c This is the melting point. The vibrations are sufficiently energetic for the
molecules to move away from their fixed positions and form liquid. Energy
added during this stage is used to break the inter-particle forces, not to raise
the kinetic energy, so the temperature remains constant.
• c–d As the liquid is heated, the particles gain kinetic energy and so the
temperature increases.
• d–e This is the boiling point. There is now sufficient energy to break all of the
inter-particle forces and form gas. Note that this state change needs more
energy than melting, as all the inter-particle forces must be broken. The
temperature remains constant as the kinetic energy does not increase during
this stage. Bubbles of gas are visible throughout the volume of the liquid.
• e–f As the gas is heated under pressure, the kinetic energy of its particles
continues to rise, and so does the temperature.
Let’s Practice
8 A closed flask contains a pure
substance, a brown liquid that is at its
boiling point. Explain what you are
likely to observe in the flask, and
distinguish between the inter-particle
distances and the average speeds of the
particles in the two states present.
I movement of the particles increases
II distance between the particles
I only
II only
Both I and II
Neither I nor II
9 During very cold weather, snow often
gradually disappears without melting. • 12 You are given a liquid substance
Explain how this is possible.
at 80 degree C and told it has a
melting point of 35 degrees C. You
10 Explain why a burn to the skin
are asked to take its temperature at
caused by steam is more serious than a
regular time intervals while it cools
burn caused by the same amount of
to room temperature (25 °C). Sketch
boiling water at the same temperature.
the cooling curve that you would
expect to obtain.
11 Which of the following occurs at the
melting point when solid sulfur is
converted to its liquid form?
Other Helpful Info
Lesson 2
Monday, October 27, 2014
Lesson 2: The Mole Concept
The mole is a fixed number of particles and refers to the amount, n, of substance.
The value of the Avogadro’s constant (L or NA) is given in the data booklet in
section 2 and will be given for Paper 1 questions.
Masses of atoms are compared on a scale relative to 12C and are expressed as
relative atomic mass (Ar) and relative formula/molecular mass (Mr).
Molar mass (M) has the units g mol–1. Guidance
The generally used unit of molar mass (g mol–1) is a derived SI unit.
The empirical and molecular formula of a compound give the simplest ratio and
the actual number of atoms present in a molecule respectively.
Lesson 2: The Mole Concept
Applications and skills:
• Calculation of the molar masses of atoms, ions, molecules, formula
• Solution of problems involving the relationship between the number
of particles, the amount of substance in moles, and the mass in grams.
• Interconversion of the percentage composition by mass and the
empirical formula.
• Determination of the molecular formula of a compound from its
empirical formula and molar mass.
• Obtaining and using experimental data for deriving empirical formulas
from reactions involving mass changes.
SI Units
Systeme Internationale d’Unites (SI)
Base units for all quantities of measurements in the
physical sciences
Symbol for unit
cubic meter
Pa or N m-2
Amount of
Are these units feasible for
• Other units used in chemistry:
Symbol for unit
degrees Celsius
cubic centimeter
Amounts of Substance
• We measure mass and volume directly in the lab, but do they
measure the amount of a substance?
The Mole
The Mole
• Avogadro’s constant (Na) is the number of atoms of an element
in the mass of the element, in grams, equal to its atomic mass in
atomic mass units (amu).
• It is defined as 6.02 * 1023 mol-1
• Example:
Atomic Mass
Molar Mass
# of Atoms
6.02 * 1023
6.02 * 1023
6.02 * 1023
Fun Factoids
• If 6.02 x 1023 pennies were distributed to everyone currently alive, they could
all spend two million pounds every hour, day and night, of their lives.
• 6.02 x 1023 grains of sand would cover a city the size of Los Angeles to a
height of 600 m.
• 6.02 x 1023 soft drink cans would cover the surface of the Earth to a height
of over 300 km. This is the number of carbon atoms in a tablespoon of
• 6.02 × 1023 pennies distributed equally to everyone alive would make
everyone on Earth a dollar trillionaire;
• 6.02 × 1023 pencil erasers would cover the Earth to a depth of about 500 m;
• • 6.02 × 1023 drops of water would fill all the oceans of the Earth many
times over.
Let’s Practice
• A tablespoon holds 0.500 moles of water. How many molecules
of water are present?
• A solution of water and ammonia contains 2.10 × 1023
molecules of H2O and 8.00 × 1021 molecules of NH3. How
many moles of hydrogen atoms are present?
More Practice!
• Calculate how many hydrogen atoms are present in : (a) 0.020
moles of C2H5OH (b) 2.50 moles of H2O (c) 0.10 moles of
• Propane has the formula C3H8. If a sample of propane contains
0.20 moles of C, how many moles of H are present?
• Calculate the amount of sulfuric acid, H2SO4, which contains
6.02 × 1023 atoms of oxygen.
Relative Atomic Mass
• Isotopes are atoms of the same element that have the same
number of protons but different numbers of neutrons
• The relative abundance of each isotope is a measure of the % of
each isotope of an element; we can use this to calculate the
relative atomic mass (Ar) of an atom; the weighted average
• We can then use the relative atomic mass (Ar) of different
atoms to calculate the relative molecular or formula mass (Mr)
of compounds!
• We use this weighted average when performing stoichiometry
So many terms!
• The mass of one mole of a species is called the molar mass
(M). It is the relative mass expressed in g and has units of g
• The molar mass of an element which exists as atoms is the
relative atomic mass (Ar), expressed in g.
• The relative molecular mass (Mr) is defined as the sum of the
relative atomic masses of the atoms in the molecular formula.
• The molar mass (M) of a compound is the relative molecular
mass expressed in g mol-1
Relative Molecular
• In a chemical compound, the subscripts represent the ratio of
atoms of each element in one molecule
• Therefore, they also represent the relative number of moles!
• Example: The relative molar mass of methane is 12.01 (Ar of Carbon)
+ 4*1.01(Ar of H) = 16.05 (takes into consideration all isotopes!)
Chart Method
• What is the molar mass of CuSO4*5H2O?
Atomic Mass
# of Atoms
Mass (g)
TOTAL = 249.72g mol-1
Let’s Practice
• Calculate the relative molecular mass for the following using
Section 6 in your data booklet:
1. Cl2
2. NH4NO3
3. Al2(SO4)3
Molar Mass
• The Avogadro number is defined so that the mass of one mole
of a substance is exactly equal to the substance’s relative atomic
mass expressed in grams. This is known as the molar mass and is
given the symbol M with the unit g mol–1, which is a derived SI
How To Solve
• I like to use Factor Label (Dimensional Analysis) but you can
solve however you’d like as long as you show work!
Number of moles (n) = Mass of Substance/Molar Mass
Helpful Equations
Mass of one molecule = molar mass/Avogadro’s Number
# of Moles = Mass of substance/Molar Mass
However! It is more useful to understand rather than memorize
especially as these problems get more difficult!
Let’s Practice
What is the mass of the following?
• (a) 6.50 moles of NaCl
• (b) 0.10 moles of OH– ions
What is the amount in moles of the following?
• (a) 32.50 g (NH4)2SO4
• (b) 273.45 g N2O5
• 380g
• 1.7g
• 0.2456 mol
• 2.532 mol
More Practice
Lesson 4
Empirical and Molecular Formulas
Tuesday, October 28, 2014
Percent Composition
• The percent by mass of each element present in a compound can
be worked out using the following formula:
% Comp = (# of atoms of element * relative atomic mass of
element)/relative molecular mass
• I find it is easier to remember that a % is always the part over the
whole. The percent composition by mass of an element in a
compound is the total mass of that element dividing by the total
mass of the molecule!
Percent Composition Example
• What is the percent composition by mass of carbon in
Step 1: Find relative molecular mass
Step 2: Find total mass of carbon
6 * 12.01 = 72.06
Step 3: Divide part by whole!
72.06/123.12 * 100 = 58.53%
Let’s Practice
• What is the percentage by mass of N, H, and O in the
compound ammonium nitrate, NH4NO3?
Harder Example
• What is the mass of oxygen present in 2.20g of CO2?
Step 1: Relative molecular mass of CO2 = 44.01
Step 2: Total mass of oxygen is 32.00
Step 3: % comp of oxygen is 32.00/44.01* 100 = 72.71%
Step 4: 0.7271 * 2.20g = 1.60g
Empirical Formulas
• A molecular formula is the actual number of atoms present in
one molecule
• An empirical formula is the lowest whole number ratio of
atoms present in one molecule
Molecular Formula
Empirical Formula
Hydrogen Peroxide
Let’s Practice
• Which of the following are empirical formulas?
I C6H6
IV Pb(NO3)2
% Composition into Empirical
• We can use the % composition of a compound to calculate the
empirical formula
Step 1: Divide the % composition of each element by its relative
atomic mass to get the # of moles
Step 2: Look at the mole ratio of all elements in the molecule and
turn it into a whole number ratio
• A sample of urea contains 1.120 g N, 0.161 g H, 0.480 g C, and
0.640 g O. What is the empirical formula of urea?
Using Experiments To Determine
Empirical Formulas
• We can use experiments to make new compounds which can
help us determine % composition and therefore empirical
formula for compounds.
• Example: A compound containing only copper and oxygen are burned
in air. When the reaction is done, all the oxygen has combined with
hydrogen and has evaporated away as water vapor. Given the following
measurements, what was the empirical formula of the original copper
Mass of Empty Dish (g)
Mass of Dish + Copper oxide
Mass of copper at end (g)
• Mass of copper oxide at start = 5.54g
• Mass of copper ending (and at start!) = 4.42g
• Mass of oxygen at start = 5.54-4.42 = 1.12g
• Moles of copper at start = 4.42/63.55 = 0.0696mol
• Mole of oxygen at start = 1.12g/16.00g/mol = 0.0700 mol
• CuO
Let’s Practice
• The mineral celestine consists mostly of a compound of
strontium, sulfur, and oxygen. It is found by combustion analysis
to have the composition 47.70% by mass Sr, 17.46% S, and the
remainder is O. What is its empirical formula?
Empirical Formula
• You took solid magnesium in the lab and burnt in with excess
oxygen. What is the empirical formula of the product?
Mass/g (±0.001)
empty crucible
crucible with
magnesium before
crucible with solid after
Empirical Formula
Mass/g (±0.002)
25.050 – 25.000 = 0.050
0.050/34.31 = 0.0021
25.084 – 25.050 = 0.034
= 0.034/16.00 = 0.0021
What is the ratio of magnesium to oxygen
Steps to Calculating Empirical
1. Calculate total mass of each substance present.
2. Convert each species to moles.
3. Determine the smallest ratio of atoms present.
Analyzing a Hydrate
• Heating a hydrated crystal drives off the water.
• This is called an anhydrous compound
• How could you determine the mass of water that
evaporated from this anhydrous compound?
Hydrated Salts
• A modification of this type of question is to analyze the
composition of a hydrated salt. These are compounds that
contain a fixed ratio of water molecules, known as water of
crystallization, within the crystalline structure of the compound.
• The water of crystallization can be driven off by heating, and
the change in mass used to calculate the ratio of water molecules
to the anhydrous salt.
• The formula of the hydrated salt is shown with a dot before the
number of molecules of water, for example CaCl2.4H2O.
Combustion Analysis
• We can also use combustion – reaction with oxygen – to get the
percent composition of elements in order to determine the
empirical formulas
Challenge Question!
• A 0.5438 g sample of a compound known to contain only
carbon, hydrogen, and oxygen was burned completely in oxygen.
The products were 1.0390 g CO2 and 0.6369 g H2O. Determine
the empirical formula of the compound.
Molecular Formula
• We can use the empirical formula and the relative formula mass
to find the molecular formula of a compound using the
following equation:
x (mass of empirical formula) = M, where x is an integer
Let’s Practice
• Calomel is a compound once used in the treatment of syphilis. It
has the empirical formula HgCl and a molar mass of 472.08 g
mol–1. What is its molecular formula?
First calculate the mass of the empirical formula: mass(HgCl) =
200.59 + 35.45 = 236.04 g mol–1
(236.04) × x = M = 472.08 ∴x = 2
Molecular formula = Hg2Cl2
Let’s Practice
24 Give the empirical formulas of the following compounds:
(a) ethyne, C2H2
(b) glucose, C6H12O6
(c) sucrose, C12H22O11
(d) octane, C8H18
(e) oct-1-yne, C8H14
(f) ethanoic acid, CH3COOH
25 A sample of a compound contains only the elements sodium, sulfur, and oxygen. It is found
by analysis to contain 0.979 g Na, 1.365 g S, and 1.021 g O. Determine its empirical formula.
26 A sample of a hydrated compound was analysed and found to contain 2.10 g Co, 1.14 g S,
2.28 g O, and 4.50 g H2O. Determine its empirical formula.
27 A street drug has the following composition: 83.89% C, 10.35% H, 5.76% N. Determine its
empirical formula.
28 The following compounds are used in the production of fertilizers. Determine which has the
highest percentage by mass of nitrogen: NH3, CO(NH2)2, (NH4)2SO4.
Let’s Practice
29 A compound has a formula M3N where M is a metal element and N is
nitrogen. It contains 0.673 g of N per gram of the metal M. Determine the
relative atomic mass of M and so its identity.
30 Compounds of cadmium are used in the construction of photocells. Deduce
which of the following has the highest percentage by mass of cadmium: CdS,
CdSe, CdTe.
31 Benzene is a hydrocarbon, a compound of carbon and hydrogen only. It is
found to contain 7.74% H by mass. Its molar mass is 78.10 g mol–1. Determine
its empirical and molecular formulas.
32 A weak acid has a molar mass of 162 g mol–1. Analysis of a 0.8821 g sample
showed the composition by mass is 0.0220 g H, 0.3374 g P, and the remainder
was O. Determine its empirical and molecular formulas.
Let’s Practice
33 ATP is an important molecule in living cells. A sample with a mass of
0.8138 g was analysed and found to contain 0.1927 g C, 0.02590 g H,
0.1124 g N, and 0.1491 g P. The remainder was O. Determine
the empirical formula of ATP. Its formula mass was found to be 507 g
mol–1. Determine its molecular formula.
34 A 0.30 g sample of a compound that contains only carbon, hydrogen,
and oxygen was burned in excess oxygen. The products were 0.66 g of
carbon dioxide and 0.36 g of water. Determine the empirical formula of
the compound.
35 You are asked to write your name on a suitable surface, using a piece
of chalk that is pure calcium carbonate, CaCO3. How could you calculate
the number of carbon atoms in your signature?
Lesson 4
Mole Ratio
Wednesday, October 29, 2014
• The mineral celestine consists mostly of a compound of
strontium, sulfur, and oxygen. It is found by combustion analysis
to have the composition 47.70% by mass Sr, 17.46% S, and the
remainder is O. What is its empirical formula?
Mole Ratio - Understandings
Reactants can be either limiting or excess.
The experimental yield can be different from the theoretical yield.
Avogadro’s law enables the mole ratio of reacting gases to be determined from volumes of
the gases.
The molar volume of an ideal gas is a constant at specified temperature and pressure.
Values for the molar volume of an ideal gas are given in the IB data booklet in Section 2.
The molar concentration of a solution is determined by the amount of solute and the
volume of solution.
The use of square brackets to denote molar concentration is required.
A standard solution is one of known concentration.
Mole Ratio – Applications and Skills
• Solution of problems relating to reacting quantities, limiting and
excess reactants, and theoretical, experimental, and percentage
• Calculation of reacting volumes of gases using Avogadro’s law.
• Solution of problems and analysis of graphs involving the
relationship between temperature, pressure, and volume for a
fixed mass of an ideal gas.
• Solution of problems relating to the ideal gas equation.
Mole Ratio - Guidance
The ideal gas equation, PV = nRT, and the value of the gas constant (R), are
given in the IB data booklet in Sections 1 and 2.
Explanation of the deviation of real gases from ideal behaviour at low
temperature and high pressure.
Obtaining and using experimental values to calculate the molar mass of a gas
from the ideal gas equation.
Solution of problems involving molar concentration, amount of solute, and
volume of solution. Guidance
Units of concentration to include: g dm–3, mol dm–3, and parts per million
Use of the experimental method of titration to calculate the concentration of a
solution by reference to a standard solution.
Mole Ratio
• The coefficients in a balanced chemical equation can be used as
a mole ratio to determine the amount of reactants or products
Let’s Practice
• Calculate the mass of carbon dioxide produced from the
complete combustion of 1.00 g of methane.
Step 1: Write the balanced chemical equation
CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)
Step 2: Determine the molar masses for all the compounds involved
Methane = 16.05g/mol and Carbon Dioxide = 44.01g/mol
Step 3: Use the coefficients to set up a mole ratio and solve
1.00g CH4(1 mol CH4/44.01g)(1mol CO2/1 mol
CH4)(16.05gCO2/1mol CO2)
All questions on reacting ratios involve a variation of this
• Write the balanced equation;
• Work out the mole ratio for the species identified in the
• Work out the reacting ratio by mass for these species, using m =
n M;
• Insert the data from the question and solve the ratio.
Let’s Practice
• Iodine chloride, ICl, can be made by the following reaction:
2I2 + KIO3 + 6HCl → 5ICl + KCl + 3H2O
• Calculate the mass of iodine, I2, needed to prepare 28.60 g of
ICl by this reaction.
• 17.88g
More Practice
36 Iron ore can be reduced to iron by the following reaction:
Fe O (s) + 3H (g) → 2Fe + 3H O(l)
(a) How many moles of Fe can be made from 1.25 moles of Fe O ?
(b) How many moles of H are needed to make 3.75 moles of Fe?
(c) If the reaction yields 12.50 moles of H O, what mass of Fe O was used up?
37 Lighters commonly use butane, C H , as the fuel. (a) Formulate the equation for the
combustion of butane. (b) Determine the mass of butane that burned when 2.46 g water were
38 Booster rockets for the space shuttle use the following reaction: 3Al(s) + 3NH ClO (s) →
Al O (s) + AlCl (s) + 3NO(g) + 6H O(g) . Calculate the mass of NH ClO that should be added
to this fuel mixture to react completely with every kilogram of Al.
39 Limestone is mostly calcium carbonate, CaCO , but also contains other minerals. When
heated, the CaCO decomposes into CaO and CO . A 1.605 g sample of limestone was heated
and gave off 0.657 g of CO . (a) Formulate the equation for the thermal decomposition of
calcium carbonate. (b) Determine the percentage mass of CaCO in the limestone.
(c) State the assumptions that you are making in this calculation.
Commence IA Review
Lesson 5
Limiting Reactant
• Today – Limiting Reactant
• Friday – Intro to Gas Laws
Monday – Lab
• Monday – Lab
• Tuesday – Kinetic Theory
Tuesday – Don’t see you
• Wednesday – Real Gases/Intro to Solutions
Wednesday – Practice
• Thursday – Solution Stoichiometry
• Friday – QUIZ/Dilutions
Monday – Lab/Review
• Monday – Double Review and exam
Tuesday - Exam
Chemical Reactions
Coefficients As Mole Ratios
• In a balanced chemical equation, the coefficients of the
reactants and the products can represent the number of moles
needed for the reaction
2 Na + Cl2  2 NaCl
Let’s Practice
• Calculate the mass of carbon dioxide produced from the
complete combustion of 1.00 g of methane.
• Write the balanced equation and deduce the mole ratio as above. Then pick
out from the question the terms that we need to analyse, here they are marked
in red; these are the species where we need to convert moles to grams.
Baking Brownies
6 ounces bittersweet chocolate (not unsweetened)
8 tablespoons unsalted butter (1 stick)
2 large eggs, at room temperature
1 cup granulated sugar
1 tablespoon unsweetened cocoa powder
1 teaspoon vanilla extract
1 teaspoon brewed espresso (optional)
1/4 teaspoon plus 1/8 teaspoon fine salt
1 cup all-purpose flour
You look in your pantry and have unlimited amounts of every ingredient but only 2
cups of sugar left. How many batches can you make?
Limiting Reactant
• The limiting reactant is the reactant in a chemical reaction that
will run out first and limit how much of the product will be
• REMEMBER: We are concerned with the mole ratio here; the
limiting reactant may actually have the highest mass but the
balanced chemical equation tells us how many moles of reactant
and product we need!
Helpful Vocabulary
• If one reactant limits the amount of product formed, the other
reactants will often have leftover amounts; they are said to be in excess
• The theoretical yield is the maximum amount of product obtainable
from a stoichiometric perspective. Often we do not actually get 100%
yield and can calculate a percent yield
Percent (%) Yield = Actual Yield/Theoretical Yield * 100
In college labs you are sometimes graded on your % yield as a measure of
lab skills!
Let’s Try One!
• Nitrogen gas (N2) can be prepared from this reaction:
2NH3(g) + 3CuO(s) → N2(g) + 3Cu(s) + 3H2O(g)
• If 18.1 g NH3 are reacted with 90.40 g CuO, determine the mass
of N2 that can be formed.
How do we start? What are we looking for?
1. Find the amount of product produced with each different
amount of reactants
2. The reactant that produces the lease amount of product is the
limiting reactant; that is the maximum that can be produced!
Let’s see what happens!
18.1g NH3 1 mol NH3 1 mol N2
28.03 g N2
14.9g N2
17.04g NH3 2 mol NH3 1 mol N2
90.40g CuO 1 mol CuO 1 mol N2
28.03 g N2
79.55g CuO 3 mol CuO 1 mol N2
CuO is the limiting reactant!!
10.6g N2
• How much NH3 will be left over when the reaction is done?
90.40g CuO 1 mol CuO 2 mol NH3 17.04g NH3 12.9g NH3
79.55g CuO 3 mol CuO 1 mol N2
Since we started with 18.1g and only used 12.9g, we would be
left over with 5.2g NH3
Experimental and % Yield
• Imagine we performed the reaction in the previous example and
only 8.2g of N2 were formed (experimental yield) instead of the
expected 10.6g (theoretical yield). What would the % yield be?
= 8.2g/10.6g * 100%
Solving Problem
All questions on reacting ratios involve a variation of this
1. Write the balanced equation;
2. Work out the mole ratio for the species identified in the
3. Work out the reacting ratio by mass for these species, using m =
n M;
4. Insert the data from the question and solve the ratio.
Let’s Practice
Let’s Practice
Lesson 6
Introduction to Gas Laws and Mole Ratio
Warm-Up Review
These four balloons are all filled
with 1 dm3 of gas. At 25 °C and
100 kPa, which balloon is the
Avogadro’s Law
• Equal volumes of all gases, when measured at the same
temperature and pressure, contain an equal number of particles.
• We can use this information to relate volumes of gases at the
same temperature and pressure to the mole ratio.
40 cm3 of carbon monoxide are reacted with 40 cm3 of oxygen in
the reaction: 2CO(g) + O2(g) → 2CO2(g)
What volume of carbon dioxide is produced? (Assume all volumes
are measured at the same temperature and pressure.)
2CO(g) + O2(g) → 2CO2(g)
2 moles 1 mole 2 moles
40cm3 O2
2 cm3 CO2
80 cm3 CO2
1 cm3 O2
40 cm3 CO
2 cm3 CO2
2 cm3 CO
40 cm3 CO2
You Try! Harder Problem!
• When 10 cm3 of a gaseous hydrocarbon (a compound
containing only carbon and hydrogen) is burned in excess
oxygen, the products consist of 30 cm3 of carbon dioxide and 30
cm3 of water vapour, measured under the same conditions of
temperature and pressure. Determine the molecular formula of
the hydrocarbon.
Molar Volume
• The molar volume for any gas under the conditions of STP
(standard temperature and pressure – 0 Celsius and (273K) and
100kPA (1atm) is: 22.7 dm3 mol-1 or 2.27 × 10–2 m3 mol–1
• LAST YEAR we used 22.4 L/mol. This year, remember that we
are using 22.7 L/mol (1 dm3 = 1L)
Let’s Practice
• Calculate the volume occupied by 0.0200 g of He at standard
temperature and pressure.
0.0200g He
1 mol He
4.00g He
1 mol He
0.114 dm3
Let’s Practice
• What volume of oxygen at standard temperature and pressure
would be needed to completely burn 1 mole of butane, C4H10?
1 mol C4H10 13 mol O2
2 mol C4H10
22.7 dm3 O2 147.6 dm3
1 mol O2
Volume Converting
Let’s Practice
Lesson 7
Kinetic Theory and Gas Laws
Kinetic Molecular Theory
• All matter consists of particles (atoms or molecules) in constant,
random motion.
• As the temperature increases, the movement of particles
• Particles have negligible volume.
• Particles have no intermolecular forces.
• All collisions between particles are completely elastic (no energy
gained or lost)
KMT Describes Ideal Gases
• The kinetic molecular theory describes ideal gases; no gas fits
this description since real gas particles do have volume,
intermolecular forces and more
• However, this description of an ideal gas is useful for predicting
the behavior of real gases, especially under the conditions of
high temperature and low pressure
• Today, we are going to review gas laws
• REMINDER: Just a reminder in case you forgot from last year,
gas laws only work when we use the Kelvin temperature scale.
• The Kelvin scale is an absolute scale, meaning at 0 Kelvin all
movement stops
• Last year, we often used atmospheres (atm) as the preferred unit
for pressure
• This year, the IB prefers to use the actual SI unit for pressure, the
Pascal which is equal to 1 N/m
Defining Pressure
• Pressure is the force
exerted per unit area.
• Gas molecules exert
pressure when they collide
with the walls of their
• As the number of gas
molecules in a container
increases, pressure
increases because there are
more collisions.
Boyle’s Law
• As pressure increase, while
temperature and number of
moles remains constant,
volume decreases
• P1V1 = P2V2
Boyle’s Law Graphs
• Know the graphed mathematical relationships!!
Charles’ Law
• As temperature increases, keeping the number of moles and
pressure constant, volume of the gas also increases
• V1/T1 = V2/T2
Gay-Lussac’s Law
• As temperature increases at constant volume and number of
moles, pressure also increases
• P1/T1 = P2/T2
Combined Gas Laws
Let’s Practice
• The molar volume of a gas at STP is 22.7 dm3 mol–1. Calculate
the molar volume at 25 °C at the same pressure.
Ideal Gas Law
Ideal Gas Law
• New for this year!
PV = nRT
• Can be used to find pressure, temperature, number of moles or
volume of a gas when given three other variables
• R is the universal gas constant and is given in Section 2 in your
data booklet
R = 8.31 N m K–1 mol–1 or 8.31 J K–1 mol–1
Tricky Tricky!
• Applying the ideal gas law is almost too easy for IB right?
• How they sometimes ask questions is to trick you by giving you
or asking for the density of the gas! What?! I don’t see density in
that equation
• All you have to realize is density = mass/volume sooo . . .
(n * molar mass)/Volume = Density
• We can have n and V in the Ideal Gas Equation and we can
always calculate molar mass!!
Tricky Trick *2
Helpful Hints!
• Pressure, P: must be in Pa (N m–2); if kPa are given, multiply by
• Volume, V: must be in m3; if dm3 are given, divide by 103, if cm3
are given divide by 106.
• Number of moles, n: this is often derived by application of n =
Mm (mass * molar mass).
• Temperature, T: must be in Kelvin; if °C is given, add 273.15.
Let’s Practice
• A helium party balloon has a volume of 18.0 dm3. At 25 °C the
internal pressure is 108 kPa. Calculate the mass of helium in the
First ensure all data are in SI units!
P = 108 kPa = 108 × 103 Pa
V = 18.0 dm3 = 18.0 × 10–3 m3
T = 25 °C = 298 K
• Then: PV=nRT  108 × 103 Pa × 18.0 × 10–3 m3 = n × 8.31 J
K–1 mol–1 × 298 K  n= 0.785 mol
• Now, transform moles = mass  0.785mol * 4.00g/mol = 3.14g
Two More
• A sample of gas has a volume of 445 cm3 and a mass of 1.500 g
at a pressure of 95 kPa and a temperature of 28 °C. Calculate its
molar mass.
• A gas has a density of 1.65 g dm–3 at 27 °C and 92.0 kPa.
Determine its molar mass.
Let’s Practice
Lesson 8
Gas Laws / Solutions
Ideal Gases
• An ideal gas obeys the Ideal Gas Law, PV=nRT under .all
• A graph of PV/RT for one mole of gas would equal 1 and look
like this:
Real Gases
• Real gases, however, do not always behave like ideal gases and a
graph of PV/RT for one mole of gas will tend to deviate, as
shown below
Real Gases
• Real gases behave most like an ideal gas at low pressure and
shows the greatest deviation at high pressure
• Real gases behave most like an ideal gas at high temperature and
shows the greatest deviation at low temperature.
• Why?!
Real Gases – Ideal Conditions
• At low pressure, the volume occupied by the gas particles
becomes negligible. Remember that the kinetic molecular theory
assumes that the volume of the particles is negligible
• At low pressure, particles are also so widely spread apart that the
intermolecular forces become negligible
• At low temperatures, particles have low kinetic energy and
cannot overcome the intermolecular forces
Van der Waals’ Equation
• Van der Waal’s equation is the Ideal Gas Law with adjustments
that allow it to more accurately predict the real behavior of gases
• However, for most conditions, we can simply use PV=nRT
Let’s Practice
61 (a) List the main features of the kinetic theory for ideal gases.
(b)Explain the reason for the difference in behavior between real and
ideal gases at low temperature.
62 Ammonia, NH3, forms a relatively strong type of intermolecular
attraction known as a hydrogen bond, whereas methane, CH4, does not.
Explain the relative deviation from ideal behavior that each gas is likely
to show.
63 Gases deviate from ideal gas behavior because their particles:
A have negligible volume
B have forces of attraction between them
C are polyatomic
D are not attracted to one another
61 (a) Particles are in constant random motion and collide with
each other and with the walls of the container in perfectly elastic
collisions. The kinetic energy of the particles increases with
temperature. There are no inter-particle forces and the volume of
the particles is negligible relative to the volume of the gas.
(b) At low temperature, the particles have lower kinetic energy,
which favours the formation of inter-particle forces and reduces gas
pressure. PV < 1 nRT
62 NH3 shows greater deviation than CH4 due to stronger
intermolecular attractions, especially at low temperature.
63 B
• A solution is a homogeneous mixture of two or more substances
that may be solids, liquids, gases or a combination
• The solvent is the component that is present in greater quantity
• The solute is the component that is present in smaller quantity;
dissolved in the solvent
• The concentration of a solution (c) is determined by the amount
of solute (n) and the volume of solution (V). It is usually
expressed as mol dm–3.
Why IB? Why?!
• From the IB:
“The term molarity, M, has been widely used to express amount
concentration, but it is falling out of common usage. It will not be
used in IB examination questions, so make sure you are fully familiar with
the terms mol dm–3 and g dm–3. (Note that M is used specifically to refer to
molar mass.)”
Let’s Practice
• Explain how you would prepare 100 cm3 of a 0.100 mol dm–3
solution of NaCl.
Let’s Practice
• Calculate the concentration of a 0.0400 mol dm–3 solution of
sodium carbonate, Na2CO3, in g dm–3.
Part Per Million
• A different unit of concentration can be used to describe very
dilute solutions where there is a very small amount of solute per
solvent – part per million
Other Concentrations
• % by mass = mass solute/total mass * 100
• % by volume = volume solute/total volume * 100
• Molality = moles solute/mass (in kg) of solvent (NOTE: NOT
Lesson 9
• In the lab, we often buy stock solutions, or solutions of a
specific concentration (previously referred to before IB decided it
was passé as Molarity or % by mass); these stock solutions are
far more concentrated than what we want to use
• To get the appropriate concentrate, we can perform a dilution by
adding more solvent (typically water)
• The key to performing dilutions is to realize that adding more
water (or solvent) does not change the number of moles of the
• We know that C = n/V (concentration = moles of
solute/volume total solution.
• Rearrange this to get n=cV. Since n remains constant we get:
Before, taken from
the stock solution.
c1V1 = c2V2
After more solvent
is added.
Let’s Practice
• Determine the final concentration of a 75 cm3 solution of HCl
of concentration 0.40 mol dm–3, which is diluted to a volume of
300 cm3.
• HINT: Watch your units!
• C1 = 0.40 mol/dm3, V1=75 cm3, V2 = 300 cm3, C2 = ???
• c1V1 = c2V2
• (0.40 mol dm–3) (75 cm3) = c2 (300 cm3)
• C2 = 0.10 mol dm–3
• Titration is a way to figure out the unknown concentration of
an acid or a base; it is also called volumetric analysis
• The basic reaction is neutralization – acid + base  salt +
• Dumb question? Can’t we just use a probe to measure pH and
then use pH to roughly determine concentration? Why yes we
can! Just go with it for now.
• When we study acids and bases we can see how the resulting
titration curve looking at the change in pH is actually useful in
determining the strength of acids and bases
What Are These?!
a. Erlenmeyer flask
b. Beaker
c. Measuring cylinder
d. Volumetric flask
e. Pipette
f. Burette
Titration - Steps
1. A pipette is used to measure a known volume of one of the
solutions and it is put into a flask
2. The other solution is put into a burette which can slowly add it
to the other solution drop by drop.
3. An indicator is placed into the flask which will change colors
when the solution goes from being acidic to basic or vice versa;
this is extremely important as we need to know when the
number of moles of acid exactly equal the number of moles of
the base!
4. The solution in the burette is slowly added to the solution in the
flask until the indicator changes color, indicating the end point
of the titration.
Now What?
• At the end point or equivalence point, we assume that the
number of moles of H+ equals number of moles of OH-.
• Therefore, we can use this equation:
CacidVacid = CbaseVbase
• NOTE: We are really looking at the concentration of the H+
and OH- so if we have acids that make more than one of these,
we have to adjust, i.e. H2SO4 (we would have to double the
concentration of the acid in general)
• We are going to spend more time on titration when we do acids
and bases.
Lab Skills
• IB expects you to have a working knowledge of basic laboratory
techniques such as serial dilution and titration in addition to
knowing what tools you should use!
Let’s Try One
• 25.00 cm3 of 0.100 mol dm–3 sodium hydrogencarbonate,
NaHCO3, solution were titrated with dilute sulfuric acid,
2NaHCO3(aq) + H2SO4(aq) → Na2SO4(aq) + 2H2O(l) + 2CO2(g)
• 15.20 cm3 of the acid were needed to neutralize the solution.
Calculate the concentration of the acid.
• Knowns: Cacid = ??, Cbase = 0.100 mol/dm2, Vbase = 25.00cm3,
Vacid = 15.20 cm3
(0.100 mol/dm3)(25.00cm3) = (15.20 cm3)(x)
X = 0.164 mol/dm3
BUT!! That is for concentration of H+ ions. So we need to divide
this by 2 (1 mol H2SO4 = 2 mol H+)
Concentration = 0.0822 mol/dm3
Back Titration
Sometimes, the end point of titration is difficult to determine or one of
the reactants is impure. In that case, we can perform a back titration
where a known excess of a reagent is added to react with the unknown
and the leftover is titrated to see how much reacted.
1. First write the equation for the reaction;
2. Look for the reactant whose volume and concentration are given and
calculate its number of moles from n = cV;
3. Use this answer and the mole ratio in the equation to determine the
number of moles of the other reactant;
4. Use the number of moles and volume of the second reactant to
calculate its concentration from c = n/V.
This is it, I promise!
• An antacid tablet with a mass of 0.300 g and containing
NaHCO3 was added to 25.00 cm3 of 0.125 mol dm–3
hydrochloric acid. After the reaction was complete, the excess
hydrochloric acid required 3.50 cm3 of 0.200 mol dm–3 NaOH to
reach the equivalence point in a titration. Calculate the
percentage of NaHCO3 in the tablet.
Let’s Practice
Lesson 10
Challenge Questions
Concentration Ratios
Concentration Description
Percent by mass
Mass of solute
x 100
Mass of solution
Percent by volume
Volume of solute
x 100
Volume of solution
Parts per million
Grams of solute
x 1,000,000
Grams of solution
Moles of solute
Liters of solution
Moles of solute
Kilograms of solvent
Challenge Questions
72 The fertilizer tri-ammonium phosphate is made from ‘phosphate rock’ by:
• 1 reacting the phosphate rock with sulfuric acid, H SO , to produce phosphoric acid, H PO ;
• 2 reacting the phosphoric acid with ammonia, NH , to give tri-ammonium phosphate, (NH ) PO .
4 3
If the phosphate rock contains 90% by mass Ca (PO ) from which the overall yield of tri-ammonium
phosphate is 95%, calculate the mass of phosphate rock required to make 1000 tonnes of triammonium phosphate.
73 The combustion of both ammonia, NH , and hydrazine, N H , in oxygen gives nitrogen and water
only. When a mixture of ammonia and hydrazine is burned in pure oxygen, the volumetric N : H O
ratio in the product gas is 0.40. Calculate the % by mass of ammonia in the original mixture. What
assumptions are being made here?
4 2
74 Sulfuric acid, H SO , is produced from sulfur in a three-step process:
• 1 S(s) + O (g) → SO (g)
• 2 2SO (g) + O (g) → 2SO (g)
• 3 SO (g) + H O(l) → H SO (l)
Assuming 100% conversion and yield for each step, what is the minimum mass of sulfur in kg needed
to produce 980 tonnes of H SO ?
Challenge Questions
75 The concentration of hydrogen peroxide, H O , in excess aqueous sulfuric acid, H SO , can be
determined by redox titration using potassium permanganate, KMnO as follows: 2KMnO (aq) +
5H O (l) +3H SO (aq) → 2MnSO (aq) + K SO (aq) + 8H O(l) + 5O (g)
A 10.00 cm sample of H O solution requires 18.00 cm of a 0.05 mol dm solution of KMnO to
reach the equivalence point in a titration. Calculate the concentration of H O in the solution.
76 Mixtures of sodium carbonate, Na CO , and sodium hydrogencarbonate, NaHCO , in
aqueous solution are determined by titration with hydrochloric acid, HCl, in a two-step
• 1 Titrate to the phenolphthalein end-point: Na CO (s) + HCl(aq) → NaHCO (aq) + NaCl(aq)
• 2 Continue titration to the methyl orange end-point: NaHCO + HCl → NaCl + H O + CO
For an X cm sample of a sodium carbonate / sodium hydrogencarbonate mixture titrated with
Y mol dm HCl, the respective end-points are Step 1 = P cm HCl, Step 2 = Q cm HCl. Derive
relationships between X, Y, P, and Q to obtain the concentrations of sodium carbonate and
sodium hydrogencarbonate in the original mixture.
77 A sealed vessel with fixed total internal volume 2.00 m contains 0.720 kg pentane, C H , and
oxygen only. The pentane is ignited and undergoes 100% conversion to carbon dioxide and water.
Subsequently the temperature and pressure in the vessel are respectively 740 K, 400 kPa.
Calculate the initial amount and mass in kg of oxygen in the vessel.
• 72 1217 tonne
Lesson 11
Review Day
Review Questions
• See handout
Boyles Law and Breathing
How does Boyle’s Law explain breathing?
• As your diaphragm moves downward, the lungs expand.
• As your diaphragm moves upward, the lungs shrink in volume.
Boyles Law and Breathing
• Why does air enter the lungs when you inhale, and exit the lungs
when you exhale?
• When a person is ill with emphysema, parts of the lungs lose their
elasticity and become enlarged. Why does this affect their
• Why are beginning scuba divers taught to never hold their breath
when they are under water?
Temperature vs. Pressure
• Volume is held constant
• Why does pressure
increase as temperature is
• Why is a higher temperature
required to bake at higher
Temperature vs. Pressure
Temperature vs. Volume
• Volume is proportional to absolute
Volume vs. Temperature