Scottish Advanced Higher Chemistry Syllabus

Chemistry syllabus — full chapter and topic list

Expand any chapter to see its topics and sub-topics. This is the whole examinable outline for Chemistry in Scottish Advanced Higher, not a summary of it.

1. 1

   Inorganic and Physical Chemistry

   4 topics

   • Electromagnetic radiation and atomic spectra
     • Emission and absorption spectra
     • Quantum numbers and orbitals
     • Electronic configurations and the aufbau principle
   • Transition metals
     • Oxidation states and d-block trends
     • Ligands and complex ions
     • Colour and the spectrochemical series
     • Catalysis
   • Chemical equilibrium
     • Equilibrium constants
     • Acid and base dissociation constants
     • Buffer solutions
   • Thermodynamics
     • Enthalpy, entropy and Gibbs free energy
     • Spontaneity of reactions
2. 2

   Organic Chemistry and Instrumental Analysis

   4 topics

   • Molecular structure and bonding
     • Hybridisation and shapes of molecules
     • Stereochemistry and chirality
   • Synthesis and reaction mechanisms
     • Nucleophilic substitution
     • Electrophilic addition and addition-elimination
     • Oxidation and reduction of organic compounds
   • Spectroscopy
     • Infrared spectroscopy
     • Nuclear magnetic resonance spectroscopy
     • Mass spectrometry
   • Chromatography
     • Thin-layer and column chromatography
     • Gas and high performance liquid chromatography
3. 3

   Pharmaceutical and Medicinal Chemistry

   3 topics

   • Medicines and drug action
     • Agonists and antagonists
     • Receptors and enzymes
   • Structure-activity relationships
     • Functional groups in drug molecules
     • Drug design principles
   • Natural product and synthetic chemistry
4. 4

   Researching Chemistry and Practical Techniques

   3 topics

   • Common chemical apparatus and techniques
     • Volumetric analysis and titrations
     • Distillation, reflux and recrystallisation
   • Stoichiometric calculations
     • Percentage yield and atom economy
     • Gravimetric and back titration calculations
   • Practical investigation
     • Planning and methodology
     • Analysis, uncertainty and evaluation

Chemistry flashcards for Scottish Advanced Higher

21 of 54 cards from the Chemistry deck — real questions with worked answers.

1. State the relationship between the energy of a photon and the frequency of electromagnetic radiation.

   $E = h\nu$, where $E$ is photon energy (J), $h$ is Planck's constant ($6.63 \times 10^{-34}\ \text{J s}$) and $\nu$ is frequency (Hz).

2. Write the equation linking photon energy to wavelength, combining the wave and quantum relationships.

   $E = \frac{hc}{\lambda}$, where $c$ is the speed of light ($3.00 \times 10^{8}\ \text{m s}^{-1}$) and $\lambda$ is wavelength.

3. How is the wavenumber of radiation defined and what is its commonly used unit?

   Wavenumber is the reciprocal of wavelength, $\bar{\nu} = \frac{1}{\lambda}$, usually expressed in $\text{cm}^{-1}$.

4. In an atomic emission spectrum, what produces a spectral line and how is its energy determined?

   A line is produced when an electron falls from a higher to a lower energy level; the photon energy equals the difference between the two levels, $\Delta E = E_{\text{high}} - E_{\text{low}} = h\nu$.

5. Distinguish between an atomic emission spectrum and an atomic absorption spectrum.

   An emission spectrum shows bright coloured lines on a dark background (light emitted as electrons fall); an absorption spectrum shows dark lines on a continuous spectrum (specific wavelengths absorbed as electrons are promoted).

6. Give the order of subshell filling for the first four principal energy levels using the aufbau principle.

   $1s\ 2s\ 2p\ 3s\ 3p\ 4s\ 3d\ 4p$, filling lowest-energy subshells first.

7. State the four quantum numbers and what each describes.

   Principal $n$ (shell/size), angular momentum $l$ (subshell shape), magnetic $m_l$ (orbital orientation) and spin $m_s$ ($+\tfrac{1}{2}$ or $-\tfrac{1}{2}$).

8. State Hund's rule for filling degenerate orbitals.

   Electrons occupy degenerate (equal-energy) orbitals singly with parallel spins before any orbital is doubly occupied, minimising electron-electron repulsion.

9. Define a transition metal in terms of electron configuration.

   A metal that forms at least one stable ion with a partially filled set of d orbitals (a partially filled d subshell).

10. Why are Sc and Zn often excluded from the strict definition of transition metals?

    $\ce{Sc^3+}$ has an empty d subshell ($d^{0}$) and $\ce{Zn^2+}$ has a full d subshell ($d^{10}$); neither forms an ion with a partially filled d subshell.

11. Define a ligand and a complex (coordination compound).

    A ligand is a molecule or ion with a lone pair that donates electrons to a central metal ion; a complex is a central metal ion surrounded by dative-bonded ligands.

12. State the rule used to determine the oxidation state of a transition metal in a complex ion.

    Oxidation state of metal $=$ overall charge of complex $-$ sum of ligand charges (neutral ligands contribute $0$).

13. What is the coordination number of a complex, and what shapes correspond to coordination numbers 4 and 6?

    The coordination number is the number of dative bonds from ligands to the metal. Four typically gives tetrahedral or square planar; six gives octahedral.

14. Explain, using d-orbital splitting, why many transition metal complexes are coloured.

    Ligands split the d orbitals into two energy levels. Visible light of energy matching the gap is absorbed to promote a d electron ($d$–$d$ transition); the complementary colour is transmitted, with $\Delta E = h\nu$ for the absorbed light.

15. How does the spectrochemical series relate ligand identity to the colour observed?

    It ranks ligands by the size of d-orbital splitting they cause; stronger-field ligands give a larger $\Delta E$, so higher-energy (shorter-wavelength) light is absorbed and the complementary colour changes.

16. Write the expression for the equilibrium constant $K_c$ for the reaction $\ce{aA + bB <=> cC + dD}$.

    $K_c = \dfrac{[\ce{C}]^{c}[\ce{D}]^{d}}{[\ce{A}]^{a}[\ce{B}]^{b}}$, using equilibrium concentrations.

17. State Le Chatelier's principle.

    If a system at equilibrium is subjected to a change in concentration, pressure or temperature, the position of equilibrium shifts to counteract (partially oppose) that change.

18. How does increasing temperature affect the equilibrium position and $K$ for an exothermic forward reaction?

    The equilibrium shifts to the left (the endothermic reverse direction) and the value of $K$ decreases.

19. What effect does a catalyst have on the position of equilibrium and on $K$?

    None on the position of equilibrium or on $K$; a catalyst speeds up the forward and reverse reactions equally, so equilibrium is reached faster.

20. Define the ionic product of water $K_w$ and give its value at $25\ ^\circ\text{C}$.

    $K_w = [\ce{H+}][\ce{OH-}] = 1.0 \times 10^{-14}\ \text{mol}^2\,\text{dm}^{-6}$ at $25\ ^\circ\text{C}$.

21. Write the definitions of pH and pOH and state their sum at $25\ ^\circ\text{C}$.

    $\text{pH} = -\log_{10}[\ce{H+}]$ and $\text{pOH} = -\log_{10}[\ce{OH-}]$, with $\text{pH} + \text{pOH} = 14$ at $25\ ^\circ\text{C}$.

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