TOPIC 1: QUANTITATIVE CHEMISTRY
IB Chemistry SL
STATES OF MATTER (P. 2-4)
Determined by the kinetic energy of the
particles
Kinetic Molecular theory states…
All matter consists of particles in motion
As the temperature increases, the
movement of the particles increase
STATES OF MATTER
Three common states of matter exist
Solid –particles closely packed in fixed positions; forces
between particles strong enough to restrict movement;
have fixed shape
Liquid –particles closely packed but not in fixed
position; forces between particles weak enough to allow
movement but fix the volume; no fixed shape
Gas –forces between particles negligible; particles free
to move around in space provided; no fixed shape or
volume
KINETIC MOLECULAR THEORY
When temperature increases enough for the particles to have
sufficient energy to overcome the intermolecular forces, a change
of state occurs
THINK ABOUT IT (THEN SHARE!)
1. Explain why the temperature of a boiling liquid does not
increase despite energy being constantly applied.
2. Deduce which would be more painful, scalding your skin
with water vapour or boiling water.
3. Explain why you might feel cold and shiver when you get
out of the water at the beach on a very hot, windy day.
Matter: any substance that
occupies space and has mass
Mixture: a combination of two or
more pure substances that retain
their individual properties
Homogeneous
solution: has
both uniform
composition and
properties
throughout, eg
salt water, metal
alloys
Heterogeneous
solution: has nonuniform
composition and
varying properties,
eg salad dressing,
paint, garden soil
p. 4-6
Pure substance: has a definite
and constant composition
Element: made
up of atoms that
each have the
same atomic
number, eg Pb,
Hg, Br
Compound: made up
of a combination of
atoms or ions in a fixed
ratio and having
different properties
from the constituent
elements, eg H2O, CO2,
NaCl
THE SEPARATION OF MIXTURES PRE-LAB
Complete the pre-lab portion of the investigation.
Materials:
Filter Funnel
Hot plates
Filter paper
Electronic balance
Test tubes
Distillation
apparatus
Beakers
Bunsen burners
Conical flasks
Ring Stands
Sugar
Salt
Coffee
Shells
Sand
Pepper
•This experiment will be
conducted in groups of
two.
•Please turn in one lab per
group.
•We will meet in the
Chemistry Lab tomorrow!
What else would you like? – Different apparatus? Alternative soluble
and insoluble solids?
CHEMICAL EQUATIONS (P. 7-10)
Shows the changes during a reaction as
reactants are rearranged into products
Provides the ratio of reactants (what starts the
reaction) and products (what is made)
Provides information on the state of the
substances: (s), (g), (l), (aq)
The ratio, in the form of coefficients, illustrates
the conservation of matter–matter is neither
created nor destroyed, simply rearranged
Generally•Balance what occurs only once (or smallest amount)
on each side first.
•Balance H and O last.
•Worry about charges when we get to the chapter
on Redox.
TRY IT!
Zinc metal reacts with hydrochloric acid to form the salt zinc
chloride. Hydrogen gas is evolved.
Hydrogen gas and oxygen react together to form water.
At a high temperature, calcium carbonate decomposes into calcium
oxide and carbon dioxide.
TYPES OF REACTIONS
Combination (synthesis): combination of two or more substances
to produce a single product
Ex: C(s) + O2(g)  CO2(g)
Decomposition: a single reactant being broken down into two or
more products
Ex: CaCO3(s)  CaO(s) + CO2(g)
TYPES OF REACTIONS
Single replacement: one element replaces another in a compound
Ex: Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g)
Double replacement: occurs between ions in solution to form
insoluble substancesand weak or non-electrolytes (AKA metathesis)
Ex: HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
Combustion: a reaction between a hydrocarbon and oxygen
 Ex: 2C4H10 + 13 O2  8CO2 + 10H2O
HOMEWORK
Read and take notes on the section titled “The Atom
Economy” on page 11 in your textbook.
Complete the activity at the end of the section (a-c) in your
notebook. (I will check tomorrow when you come in.)
CHEMICAL EQUATIONS
Write the reaction equation for the combustion of methane.
Balance the reaction.
Nitrogen (N2) and oxygen (O2) react in the cylinders of car engines to form
nitrogen monoxide. Give a balanced equation for this reaction.
CHEMICAL EQUATIONS
After nitrogen monoxide escapes into the environment, it reacts with oxygen to
produce nitrogen dioxide. Give a balanced equation for this reaction.
Nitrogen dioxide can further react with oxygen and water to produce nitric
acid, one of the ingredients in acid rain. Give a balanced equation for this
reaction.
MEASUREMENT AND UNITS - SYSTEME
INTERNATIONAL (SI)
Mass – kg
 1000 g = 1 kg
 1000 mg = 1 g
Time – s
 60 sec = 1 min
 60 min = 1 hr
Volume –m3(cubic metre)
 1 L = 1 dm3= 1000 mL
Pressure –Pa (pascal)
 1 atmosphere = 1.01 x 105Pa =
101 kPa
Temperature – K
 Based on thermal motion
 Absolute zero (all thermal motion ceases)
 K = ˚C + 273
1.1.1 APPLY THE MOLE CONCEPT TO SUBSTANCES
AMOUNTS OF SUBSTANCES
All matter composed of varying types of substances
 Atoms – building blocks of matter; extremely small; exist as varieties in
universe (elements)
 Compounds – substances composed of atoms joined together in wholenumber ratios
Because of size, atoms and compounds cannot be counted directly
 Counted based on mass and using mole concept
1.1.1 APPLY THE MOLE CONCEPT TO SUBSTANCES
THE MOLE
•Latin for heap; defined as
number of carbon atoms in a
12.01 g sample
•1 mole = average atomic
mass
• 1 mole = 6.02 x 1023
particles (Avogadro’s number)
•The mole is just a way to
define a quantity
•Similar to the word ‘dozen’
1.1.1 APPLY THE MOLE CONCEPT TO SUBSTANCES
THE MOLE
1 mole of an element is equal to relative atomic
mass (Ar) of that element –this mass found on
the periodic table
 Molar Mass = Expressed in units g mol-1
1 mole of a compound/molecule is equal to the
relative mass (Mr) of that compound/molecule
 Mr= sum of all Ar of the elements making up
compound
 Subscripts in compound indicate amount of atoms
 Ar/Mr and mass measured allow for determination
of number of moles present
MOLAR MASS
Molar mass is equal to mass (g) of 1 mol of the substance
Determine molar mass by summing masses of all component atoms
Find the molar mass of ammonia.
Find the molar mass of water.
What is the mass of 1 mole of methane?
What is the mass of 1 mole of calcium carbonate?
1.1.2 DETERMINE THE NUMBER OF PARTICLES AND THE AMOUNT OF SUBSTANCE (IN MOLES)
MOLE CONVERSIONS
Number of moles (n) = mass (m)/ molar mass (M)
Number of particles (N) = number of moles(n) x
Avogadros’s constant(L)
•Calculate the number of moles in 4.00 g of NaOH
•What is the mass of 2.50 moles of water?
•How many molecules are in 2.50 moles of water?
CLARIFICATI0N OF UNITS
Relative molecular mass (Mr) of CO2 = 44
Molar Mass of CO2 = 44 g/mol
Molecular Mass of CO2 = 44 a.m.u
PRACTICE
1. What is the mass of 2 moles of Lithium?
How many atoms are in 2 moles of Li atoms?
2. What is the mass of 4.5 moles of CaCO3?
MORE PRACTICE
1. 6g of hydrogen reacts with fluorine to produce hydrogen
fluoride. What mass of HF fluoride is produced?
2. 0.1 moles of nitrogen reacts with hydrogen to make ammonia.
What mass of ammonia is produced?
EMPIRICAL AND MOLECULAR FORMULAS
•Empirical formula –lowest whole-number ratio; does not
necessarily show the actual number of atoms involved
•Molecular formula – shows the number of each atom in a
molecule
CALCULATING EMPIRICAL FORMULAS
If experimental data is given
1. Determine mass of elements involved (if percent composition
data is given, assume a 100g sample)
2. Convert mass to moles ( moles = mass/molar mass)
3. Divide all molar amounts by the smallest molar amount; the
numbers found are the amount of each atom in the empirical
formula
If necessary force the ratio into whole numbers (multiply
everything by 2, 3, etc.)
EMPIRICAL FORMULAS PRACTICE
Smog is common in cities throughout the world. One component of smog is
PAN (peroxyacylnitrate) which consists of 20.2 % C, 11.4 % N, 65.9 % O,
and 2.50 % H by mass. Determine the empirical formula of PAN, showing
your work.
MOLECULAR FORMULA
The actual number of atoms in the molecule
1. Determine the empirical formula
2. Divide the Mr of the molecule (given) by the Mr of the empirical;
the number is the factor that all the subscripts of the empirical must
be multiplied by
MOLECULAR FORMULA PRACTICE
Upon analysis, a sample of an acid with a molar mass of 194.13 g mol-1
was found to contain 0.25 g of hydrogen, 8.0 g of sulfur, and 16.0 g of
oxygen. Determine the empirical and molecular formula.
Which is both an empirical and a molecular formula?
A. C5H12
B. C5H10
C. C4H8
D. C4H10
(Total 1 mark)
MASS RELATIONSHIPS IN CHEMICAL REACTIONS
(STOICHIOMETRY) (P. 20-22)
Balanced chemical equations allow for the
determination of theoretical yields of
products and required amounts of
reactants needed
Convert known masses to moles (if
needed)
Use the molar ratio (from the balanced
chemical equation) to convert to the
unknown
The combustion of hydrocarbon fuels is an environmental concern as
it add to the carbon dioxide levels in the atmosphere. Calculate
the mass of CO2 produced when 100 g of propane is burned
according to the equation:
C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(l)
LIMITING REAGENTS
If amounts of all reactants are given, determine the limiting
reagent (produces the lowest amount of product) and the reagent
in excess
Calculate theoretical yield for one product starting from each
reactant; limiting reagent will produce the lowest amount.
If 14 g of each reactant is present, which is in excess? How many
moles of product is made?
N2 + 3H2  2NH3
ASSIGNMENT (DUE TODAY)
In the textbook:
Complete the Quick Questions (#1-3) on page 22.
THEORETICAL AND EXPERIMENTAL YIELD (P. 23-24)
Comparison of the experimental yield to the theoretical yield
%𝑌𝑖𝑒𝑙𝑑 =
𝐸𝑥𝑝𝑒𝑟𝑖𝑚𝑒𝑛𝑡𝑎𝑙 𝑌𝑖𝑒𝑙𝑑
𝑇ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙 𝑌𝑖𝑒𝑙𝑑
× 100
Why is the actual yield of a reaction almost always smaller than
the theoretical yield?
Calcium carbonate decomposes on heating as shown below.
CaCO3  CaO + CO2
When 50g of calcium carbonate are decomposed, 7g of calcium
oxide are formed. What is the percentage yield of calcium oxide?
Sodium thiosulfate may be produced by boiling solid sulfur and an
excess of Na2SO3, determine the theoretical yield of the product.
If only 50.0 g of product is collected, determine the percent yield.
Your starting amount of S8 is 15.50 g.
S8(s) + 8Na2SO3(aq)  8Na2S2O3(aq)
GASEOUS VOLUMES (P. 24-30)
Molar ratio from a balanced chemical
equation provides information on
volume of gases that are
used/produced; coupling this with the
fact that 1 mole of any gas is always
the same volume at a given
temperature pressure allows for easy
calculations when gases are reacting.
STANDARD TEMPERATURE AND PRESSURE
•Standard conditions of temperature and
pressure (STP) are 273 K (0°C) and
101.3 kPa (1 atm) pressure.
•The molar gas volume is 22.4 dm3 (L).
•Room Temperature and Pressure (RTP) =
298 K (25°C) and 1 atm.
•Molar volume at RTP is 24 dm3.
MOLAR VOLUME
Moles(n) = Volume (V) / molar volume (Vmolar)
1. Calculate the moles of chlorine in 44.8 cm3 of the gas at STP
2. Calculate the volume occupied by 4.40 g of carbon dioxide at
STP.
3. What volume of hydrogen gas is produced when 0.056 g of
lithium reacts with water? Assume the volume is measured at STP.
IDEAL GAS LAW
Gas volumes are not only affected by the number
of moles present but also by pressure and
temperature
This leads to the development of the Ideal Gas Law
– how gases behave in ideal situations
PV = nRT
Gases only deviate from this law at high pressures
and low temperature/volumes due to increasing
IMFs
P = pressure (kPa)
V = volume (dm3 or L)
n = moles (mol)
R = 8.31J K-1 mol -1
T = temperature (K) - MUST BE
IN KELVIN!!! (if given data in °C,
convert by adding 273)
NOTES - BOOK PAGES 24-30
Kinetic theory of gases
Definitions:
Ideal gas
Real gas
STP
Molar volume
Formulas and graphical
relationships (if possible):
Ideal gas law
Boyle’s law
Units for variables
Charles’ law
Gay-Lussac’s law
Combined gas law
Equation
Gas constant
Worked example
SOLVE THESE AT THE END OF YOUR NOTES
(SHOW YOUR WORK)
1. Calculate the molar mass of a gas if it is found that 15.0g of the
gas occupies 3.00 x 104 mL at 581°C and 50.0 kPa. The actual
molar mass is found to be different by 10%, account for this.
2. Gas in a 2.0 dm3 box is at a pressure of 100kPa and 200 K
temperature. If the gas is heated to 300K, what is the new
pressure?
3. A gas has it’s pressure doubled and volume tripled. What is the
effect on it’s temperature (in K)?
SOLVE PROBLEMS USING THE IDEAL GAS EQUATION PV=NRT
Calculate the molar mass of a gas if it is found that 15.0g of the gas occupies
3.00 x 104 mL at 581°C and 50.0 kPa. The actual molar mass is found to be
different by 10%, account for this.

PV = nRT
mass
Molar mass
PV = mass x R x T
Molar mass
15.0 x 8.31 x 854
50.0 x 30.0L
Mass x R x T= Mmass
PV
=Mmass
Mmass = 71.1 g/mol
(3 SF)
How do you
account for the
10% difference?
SOLVE PROBLEMS BETWEEN T,P, AND V FOR A
FIXED MASS OF AN IDEAL GAS
P
V
1
1
__________
T1
=
P
V
2
2
__________
T2
Gas in a 2.0 dm3 box is at a pressure of 100kPa and 200 K temperature. If the gas
is heated to 300K, what is the new pressure?
SOLUTIONS (P. 31-33)
•Solute – the less abundant
component in a solution
•Solvent – the more abundant
component in a solution
SOLVENT
•Solution – mixtures of two
components
SOLUTION
SOLUTE
CONCENTRATION OF SOLUTIONS
Concentration is the composition of a solution. As more and more solute
dissolves in the solvent, the solution becomes more and more concentrated.
When the solvent cannot dissolve any more solute, it is saturated.
Ex: [HCl] = 1 mol dm-3 or 1M
𝑚𝑜𝑙𝑒𝑠 𝑠𝑜𝑙𝑢𝑡𝑒 (𝑚𝑜𝑙)
𝐶𝑜𝑛𝑐𝑒𝑛𝑡𝑟𝑎𝑡𝑖𝑜𝑛 𝑀 =
𝑣𝑜𝑙𝑢𝑚𝑒 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 (𝑑𝑚3 )
Units: mol/dm3 or mol dm-3
Determine the concentration (in mol dm-3) of the following solutions:
A) 250 cm3 containing 45.0 g of MgSO4
B) 0.25 dm3 containing 0.210 g of NaHCO3
Determine the mass of solute required to make the following
solutions:
A) 250.0 mL of 0.1 M AgNO3
B) 0.500 L of 0.40 M Mg(NO3)2