Chemistry Module 1 Common Mistakes

As the first module of Stage 6 Chemistry, it is crucial for you to set up a solid foundation for your Module 1 Chemistry exam in year 11. Module 1: Properties and Structure of Matter covers the fundamental concepts in Chemistry which will carry over well into year 12 and the HSC!

Whilst the majority of Module 1 Chemistry may not seem too conceptually difficult, there is much room for error and silly mistakes. These silly mistakes are going to cost you a lot of marks in the long-run, so it's important that you have a firm grip on them!

IUPAC Nomenclature of Inorganic Compounds

Not specifying the valency of transition metals when naming ionic compounds. (eg. FeCl₂ should be named iron(II) chloride, not iron chloride)
Not truncating the prefix in the nomenclature of covalent compounds where necessary. (eg. N₂O₄ is dinitrogen tetroxide, not dinitrogen tetraoxide)
Using common names instead of systematic names for polyatomic ions. (eg. The systematic name for CH₃COO^- is ethanoate, not acetate)
Providing explicit numbers in chemical formula where they are not necessary. (eg. The +1 charge on a sodium ion is expressed as a superscript with a '+', not a '+1' or a '1+')
Not balancing the charge of ions in ionic compounds correctly. (eg. Iron(II) phosphate has the formula, Fe₃(PO₄)₂, as iron(II) ions have a +2 charge, whereas phosphate ions have a 3- charge)

Mixtures & Separation Techniques

Misclassifying heterogeneous mixtures as homogeneous ones. (eg. Blood is considered heterogeneous in spite of its apparent homogeneous appearance)
Not following the correct HSC conventions for drawing diagrams. (eg. Drawing the filter funnel without a retort stand / clamp in the illustration of the filtration apparatus)
Not rounding the answer the percentage composition calculation questions to the correct number of significant figures. (eg. Rounding 3.47 g / 100.0 g to 2 significant figures, instead of the correct rounding to 3 significant figures)
Drawing flowcharts for separating mixtures with arrows moving upwards. In flowcharts, arrows always head downwards (or sideways, if practically necessary).
Not labelling flowcharts sufficiently. Flowcharts in the context of separating mixtures should explicitly specify the separation technique being used, the substances involved, and the observations which lead to different pathways down the flowchart.

Atomic Models

Describing electrons as "moving" between different shells (for Bohr's model) or "swirling" inside an electron orbital (for Schrödinger's model). The energy level of electrons are discrete, and electrons do not physically "move" between different shells. Similarly, electron orbitals in Schrödinger's model are projections of probability curves - not an actual cloud in which electrons zap about.
Placing commas or spaces within electron configurations. (eg. The correct electron configuration* for sodium is '1s²2s²2p⁶3s¹', without any spaces or commas)
Understanding the third shell to have a maximum capacity of 8 electrons. In practice, the third shell can hold up to 18 electrons. However, the lower energy of the 3s and 3p subshells compared to the 4s subshell is what results in a stable third shell with 8 electrons.
Applying the Aufbau Principle or Hund's Rule in excited species. These principles only apply to ground state atoms and ions, and can be disregarded for excited chemical species.
Writing the atomic number over the mass number in the symbolic representation of atoms. The mass number is placed above the atomic number in such symbolic representations in practice.

Evidence of Atomic Models

Quoting simplified flame colours for metal cations tested via flame test. (eg. Describing the flame colour for calcium as simply 'red', rather than 'brick-red')
Remembering the absorption and emission spectra the wrong way around. The hydrogen absorption spectrum consists of the electromagnetic spectrum with blank spectral lines, whereas the hydrogen emission spectrum consists of a black band with solid spectral lines.
Forgetting to mention the process of atomisation which takes place during flame test. When heated, metal cations are first atomised into neutral atoms before the electrons in the metal atom are excited to higher energy levels.
Interpreting light as a continuous source of energy. (eg. When an electron collides with light of 100 J of energy, it is unable to absorb 80 J of said energy and release the rest)
Not understanding that flame test and atomic spectra are evidence against Rutherford's model of the atom. The presence of discrete spectral lines on atomic spectra and the production of specific, consistent colours in flame test indicates that the energy level of electrons must be discrete (as compared to the continuous energy levels in Rutherford's model).

Isotopes & Radioactivity

Leaving nuclear equations unbalanced. The mass number on each side of the equation should add up, as should the atomic / charge number, as per the Law of Conservation of Mass and Charge.
Providing a unit for relative atomic mass. Relative atomic mass is a unitless ratio which relies on the mass of natural isotopes in atomic mass units (amu).
Using the mass number instead of the atomic mass for isotopes when calculating relative atomic mass. (eg. Mg-24 has a mass number of 24, but an atomic mass of 23.99 amu)
Failing to distinguish between the two different causes for radioactivity. Radioactivity may be caused by an excessively high atomic number, which results in alpha decay; or by the presence of an unstable neutron to proton ratio, which results in beta decay.
Failing to memorise examples of natural and human-made radioisotopes, as outlined in the Syllabus.

Periodic Trends

Memorising the periodic trends prescribed by the Syllabus. The Syllabus specifies that the examinable content is not limited to the prescribed periodic trends. Students must seek to understand the underlying reasoning behind each trend, and may be assessed on other trends, such as ionic radius, second ionisation energy, and polarisability.
Applying the periodic trends onto transition metals. Transition metals generally do not follow the same periodic trends as the main group elements. (eg. Zinc has a higher atomic radius than nickel, even though it's further across on the Periodic Table)
Applying the periodic trend for electronegativity onto the noble gases. Electronegativity is defined as the tendency of an atom to attract a shared pair of electrons towards itself. As noble gases do not share electrons, they do not tend to possess electronegativity values (in particular, the small noble gases).
Using a single periodic trend to probe atoms' reactivity with water. Metals and non-metals react differently with water (the former loses electrons, whilst the latter gains them), so they possess different periodic trends when it comes to reactivity. (eg. Metals increase in reactivity going down the Periodic Table, whereas non-metals decrease in their reactivity in the same direction)
Applying the same periodic trends for ions. In general, ions do NOT follow the same periodic trend as the corresponding atoms. (eg. Sodium ion has a very low first ionisation energy, whereas sodium atom has a very high first ionisation energy)

Lewis Dot Diagrams & VSEPR Theory

Using lines or crosses instead of dots in Lewis Dot Diagrams. Whilst it is common for online sources or even textbooks to use lines and crosses, this is generally frowned upon at most schools.
Treating double bonds and triple bonds as separate electron pairs. The VSEPR Theory considers double bonds and triple bonds as one whole 'electron group'. (eg. Carbon dioxide, which has two double bonds on the central carbon atom, forms a linear structure)
Not denoting the charge on ions in Lewis Dot Diagrams. (eg. The +1 charge of ammonium ions should be denoted with square brackets and a superscript, even though it consists of covalent bonding internally)
Applying the Octet Rule blindly onto all atoms. (eg. Boron is stable with 6 valence electrons, whereas phosphorus is stable with 10 valence electrons)
Drawing the ionic Lewis Dot Diagram for covalently bonded substances. In general, ionic compounds are only formed by a combination of metals and non-metals. (eg. Magnesium chloride is ionic, whereas nitrogen dioxide is covalent in its bonding)

Forces & Bonding

Referring to dispersion force as a force between two momentary dipoles. In practice, dispersion force refers to the electrostatic attraction between the partially charged end of a molecule's momentary dipole and that of another molecule's consequently induced dipole.
Describing a hydrogen bond (or any other intermolecular force) as a chemical bond. A hydrogen bond isn't really a chemical bond (poor choice of name, I know). Rather, it's simply a particularly strong form of a dipole-dipole force.
Describing electricity strictly as a flow of electrons. In practice, electricity can be conducted by any charged particles which are free to move about (called mobile charge carriers).
Memorising the properties of allotropes of carbon. The Syllabus never actually prescribes carbon as an example to study on the topic of allotropes. Exams frequently assess students on properties of other allotropes, such as phosphorus or oxygen.
Not knowing the prescribed examples of covalent networks. The Syllabus prescribes diamond and silicon dioxide as examples of covalent networks to study. The actual method of classifying covalently bonded substances as networks or molecules is more complex, and taught only as extension material even at Cognito Tuition.

And that's about it - 5 of the most common mistakes from each major section of Module 1!

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