Pearson Baccalaureate Chemistry Higher Level 2nd Edition Print and Online Edition for the IB Diploma
IB Chemistry Notes from a student who scored a 7 in IB HL Chemistry and 44/45 overall. Overall score of 88/100 in HL Chemistry. Option D (Medicinal Chemistry taken).
Band 7 SL Chemistry Notes from a 44/45 Novemeber 2022 Graduate
Topic 1 and 11 Notes
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Door: alinaplcalinac • 2 maanden geleden
very good notes, highly recommended
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Voorbeeld van de inhoud
Topic 1: Stoichiometric relationships
1.1: Introduction to the particulate nature of matter and chemical change
Types of Matter
- Elements: Single substances composed of the same type of atoms and are always homogeneous.
- Compounds: single substances containing elements combined in fixed ratios that have different physical and
chemical properties to their component elements which are always homogeneous.
- Mixtures: a group of substances that are not chemically bonded together and thus retain their
physical/chemical properties which can be both homogeneous when in the form of an aqueous solution (solute
is dissolved in solvent) or heterogenous.
Matter Uniformity
- Homogeneous: a substance is uniform throughout
- Heterogeneous: A substance isn’t uniform throughout and can be separated into different components.
All matter can also be found in 3 states of matter that change depending on their temperature (degree of kinetic
energy that a substance possesses) or pressure.
States of Matter
Solid Liquid Gas
- Densely packed - Somewhat spread out - Very spread out
- Strong intermolecular - Weaker intermolecular - Weak intermolecular
forces/bonds; Vibrate in forces/bonds; particles move forces/bonds; particles move
fixed positions past/around each other randomly
- Fixed shape - Takes shape of container - No fixed shape
- Fixed volume - Fixed volume - No fixed volume
Changes of State
, Chemical Formulae
Example chemical formula:
2Na(s) + 2H2O(l) -> 2NaOH(aq) + H2(g)
The reactants are the substances that are being reacted to create a product. Reactants are found on the left while
products are found on the right of the formula.
States of matter can be found in the chemical formula through the following subscripts:
- (s) meaning solids
- (l) meaning liquids
- (g) meaning gases
- (aq) meaning aqueous solutions (solvent is water)
When balancing equations, both sides’ individual elements must be equal in number. This occurs due to how
matter cannot be created nor destroyed, only rearranged. This can be done by changing the coefficients of the
formula.
1.2: The mole concept
The Mole
The mole is a fixed number of particles (atoms/molecules/subatomic particles) and refers to the amount, ‘n’ of that
substance, acting much like a dozen (12). Number of moles is measured in the unit mol.
One mole = 6.02 x 1023 (also known as Avogadro’s number/constant). This number is the number of atoms in 12.01
grams of C-12. This number can be found in DB2.
Molar mass is the mass in grams of a pure substance required to have 1 mol of said substance’s particles. Molar mass
(M) has the units g mol-1.
Types of Molecular Formulas
- Molecular formula: the actual numbers of atoms present in a single molecule
- Empirical formula: the simplest ratio of atoms present in a single molecule.
, Calculating Number of Moles from Mass
Calculating Number of Particles from Number of Moles
Percentage (%) Composition
Percentage composition is the mass of each individual element make-up of a substance in percent form. To convert
from percentage composition to empirical/molecular formula:
1. Assume full mass of substance equals 100g, thus each 1% = 1g.
𝑚
2. Using 𝑛 = 𝑀
, find the number of moles.
3. Using ratio from n values, determine the empirical formula
4. If given full mass of substance, divide this by empirical formula mass then multiply empirical formula by that
amount to get the molecular formula.
Waters of Crystallisation (Prescribed Practical)
Waters of Crystallisation: Water molecules (H2O) chemically bonded or filling the space between the lattice
structures of crystals. Examples include:
- Copper sulfate pentahydrate (CuSO4 · 5H2O)
- Cobalt (II) Chloride Hexahydrate (CoCl2 · 6H2O)
Anhydrous: contains no water molecules.
Heated to constant mass: heating a substance until two consecutive masses are the same (usually to get rid of water
molecules)
1.3: Reacting masses and volumes
Limiting and Excess Reagents
Limiting Reagent (reacting substance) is the substance that is completely used up in a chemical reaction.
Contrastingly, the excess reagent is the substance that is not completely used up in a chemical reaction.
This can be found by calculating the number of moles and dividing it by the coefficient of each reactant. The lower
n(mol) number dictates which reactant is the limiting reagent. Any further calculations concerning the products’
quantities should be based on the limiting reagent.
MAKE SURE TO ALWAYS FIND THE LIMITING REAGENT IF YOU ARE GIVEN MASSES OF REACTANTS!
, Yields
Theoretical yield: the maximum amount of product formed according to the balanced chemical equation.
Experimental yield: the actual amount of product formed when the experiment is conducted
Using these two yields, we can calculate the percentage (%) yield of an experiment:
Solution Stoichiometry (Moles in Solutions)
Gas Stoichiometry
Avogadro’s Law: At the same temperature and pressure, equal volumes of different gases will contain the same
number of particles, therefore, at a particular temperature and pressure, 1 mol of any gas will occupy the same
volume.
Molar Volume: One mole of any gas at STP occupies 22.7dm3. This is found in DB2.
This equation can only be used at STP (standard temperature and pressure (100kPa, 273K)
Combined Gas Law
For a fixed amount of gas, when 1 of these variables changes (pressure, volume, or temperature) the other variables
will change due to these being either directly or indirectly proportional to each other.
Ideal Gas Laws
Ideal gases deviate from real gases because
- Gas particles are assumed to have zero volume (at high pressure, due to decreased empty space, this
assumption is less valid)
- It is assumed there are no intermolecular attractions between particles (at low temp and high pressure,
attractions are most common)
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