Question 34 — The Chemistry of Art (25 marks) Answer parts (a) and (b) of the question on pages 2–4 of the Section II Writing Booklet - HSC - SSCE Chemistry - Question 34 - 2015 - Paper 1

The coloration of solutions is largely due to the presence of d-electrons and how they interact with light.

• For Cr^2+(aq): The electronic configuration is ext{[Ar] } 3d^4. In this state, there are unpaired d-electrons, which can absorb certain wavelengths of light, causing electronic transitions between energy levels, leading to the observed color.

• For Zn^2+(aq): The electronic configuration is ext{[Ar] } 3d^{10}. In this case, all d-orbitals are fully filled; thus, there are no available d-electrons for transition. This lack of electronic transitions means that Zn^2+(aq) does not absorb visible light and consequently appears colourless.

To investigate the flame colors of metal ions, the following safe procedure can be utilized:

1. Equipment Needed: Prepare a Bunsen burner, a clean nichrome wire, and samples of metal salt solutions such as sodium nitrate (NaNO3), potassium nitrate (KNO3), and copper(II) nitrate (Cu(NO3)2).
2. Safety Measures: Ensure to wear safety goggles and lab coats to protect against splashes.
3. Procedure:
   • Dip the nichrome wire into hydrochloric acid (HCl) to clean it.
   • Then, dip the cleaned wire into the metal salt solution, ensuring no contaminations.
   • Place the wire in the hottest part of the Bunsen flame and observe the color produced.
   • Record the flame colors for each metal ion solution tested.
4. Specific Metal Ions Observed: Sodium ions will produce a yellow flame, potassium ions will give a lilac flame, and copper(II) ions produce a green flame.

Certain metal ions can be identified by flame tests due to their ability to emit specific colors when heated. This property depends on the electronic configuration of the metal ions.

Some metals, like aluminum (Al), do not produce a distinctive flame color due to the lack of accessible d-electrons that can undergo transitions. As a result, aluminum cannot be reliably identified using this method. In contrast, metals with partially filled d-orbitals can provide distinct colors that allow for identification.

First ionisation energy refers to the energy required to remove the outermost electron from a neutral atom. This energy varies across the periodic table, providing evidence for the existence of sub-shells:

1. As we move across a period, the ionisation energy generally increases, indicating that electrons are held more tightly due to increasing nuclear charge.
2. The data for elements in the third row of the Periodic Table shows larger jumps in ionisation energy between groups. For instance, there is a noticeable increase between magnesium (Mg) and aluminum (Al), suggesting that the third electron in aluminum is being removed from a higher energy p-orbital rather than a filled s-orbital.
3. The presence of such gaps indicates that electrons in different sub-shells are at varying energy levels, thus confirming the sub-shell structure in atomic configurations.

Electronegativity is a property that reflects an atom's ability to attract electrons, which is influenced by its electronic arrangement:

1. The graph shows a trend where electronegativity generally increases across a period and decreases down a group. This trend is closely related to the arrangement of electrons.
2. In a period, as the atomic number increases, electrons are added to the same energy level while protons are added to the nucleus, resulting in greater effective nuclear charge, thus enhancing electronegativity.
3. Conversely, down a group, the number of electron shells increases, leading to increased distance between the nucleus and the valence electrons, which diminishes the effective nuclear charge felt by these electrons, thereby reducing electronegativity.
4. These trends in the graph reinforce the concept that the arrangement of electrons and their energy levels directly impact an element's electronegativity.

The Bohr model, proposed by Niels Bohr, was pivotal in shaping our understanding of atomic structure:

1. Strengths: It introduced the concept that electrons orbit the nucleus at fixed distances, or energy levels, which are quantized. This explained the stability of atoms and their discrete emission spectra.
2. The model also explained how electrons can transition between orbits, absorbing and emitting specific energies corresponding to visible wavelengths of light. This provided a theoretical framework to understand spectral lines observed in hydrogen and other elements.
3. Weaknesses: However, it has limitations, particularly in explaining multi-electron systems and the finer details of atomic transitions, which necessitated advancements in quantum mechanics.
4. Overall, while the Bohr model simplified atomic structure significantly, its foundational ideas continue to influence modern atomic theory, showcasing the interplay between energy levels and emission spectra.