Week 7: Introduction to Transition Metals PDF

Week 7: Introduction to Transition metals Weekly objectives 1. Recall the first row transition metals and identify some uses 2. Describe the periodic trends: ionisation energy, electron affinity, atomic radius and effective nuclear charge 3. Determine the various transition metal oxida...

Week 7: Introduction to Transition metals Weekly objectives 1. Recall the first row transition metals and identify some uses 2. Describe the periodic trends: ionisation energy, electron affinity, atomic radius and effective nuclear charge 3. Determine the various transition metal oxidation states 4. Describe an atomic orbital and discuss their shapes and how they relate to quantum numbers with the emphasis on d orbitals 5. Apply the Aufbau principle, Hund’s rule and the Pauli exclusion principle to write full and condensed electronic configurations 6. Discuss the use of the four quantum numbers Weekly objectives (Part 1) 1. Recall the first row transition metals and identify some uses 2. Describe the periodic trends: ionisation energy, electron affinity, atomic radius and effective nuclear charge 3. Determine the various transition metal oxidation states 4. Describe an atomic orbital and discuss their shapes and how they relate to quantum numbers with the emphasis on d orbitals 5. Apply the Aufbau principle, Hund’s rule and the Pauli exclusion principle to write full and condensed electronic configurations 6. Discuss the use of the four quantum numbers Examples of Transition metals and their uses https://www.compoundchem.com/category/iypt-elements/ Examples of Transition metals and their uses https://www.compoundchem.com/category/iypt-elements/ Activity 1 What are some of the general properties of the first row d-block elements? Recap from pre-workshop material Periodic trends Activity 2 Why does atomic size stay relatively constant for the transition metals across a row? Recap from pre-workshop material Effective nuclear charge (Zeff) approximate values for Zeff for boron (B) and nickel (Ni) B: 1s2 2s2 2p1 Zeff(B) = 5 – 2 = 3 Ni:1s2 2s2 2p63s23p64s23d8 Zeff(Ni) = 28 – 18 = 10 Recap from pre-workshop material Electron shielding The shielding ability of electrons is dependent on the shape of the orbitals s p d f s p d f decreasing shielding effect Activity 3: Differences in Atomic Radii 1st Row 2nd Row 3rd Row Create a concise explanation for why there is a negligible difference in atomic radii between the second and third row transition elements of any group. Activity 3: Discussion and feedback Differences in Atomic Radii Atomic radii Poor shielding of nuclear charge by the f-orbitals: - Shielding by different orbitals: s > p > d > f - The third row of the transition metals (starting with Hf) are the first to have filled f-subshells (i.e. orbitals). - This phenomenon is called lanthanoid contraction. Recap from pre-workshop material ACTIVITY 4. Demonstration Demo - We will be mixing a solution of VO3 with different amounts of Zn. While this is occurring, in your group consider and discuss the context 15 mins. question: “Why does the colour change?” Demo ACTIVITY 4. Demonstration Demo Questions Q1: What is the nature of the reactions being shown? Q2: What is the role of Zn in the reaction? Demo Q3: What is the oxidation state of V in VO₃⁻ ? Q4: What are the most probable oxidation states of V found in the blue, green, and violet solutions? Weekly objectives 1. Recall the first row transition metals and identify some uses 2. Describe the periodic trends: ionisation energy, electron affinity, atomic radius and effective nuclear charge 3. Determine the various transition metal oxidation states 4. Describe an atomic orbital and discuss their shapes and how they relate to quantum numbers with the emphasis on d orbitals 5. Apply the Aufbau principle, Hund’s rule and the Pauli exclusion principle to write full and condensed electronic configurations 6. Discuss the use of the four quantum numbers Weekly objectives (Part 2) 1. Recall the first row transition metals and identify some uses 2. Describe the periodic trends: ionisation energy, electron affinity, atomic radius and effective nuclear charge 3. Determine the various transition metal oxidation states 4. Describe an atomic orbital and discuss their shapes and how they relate to quantum numbers with the emphasis on d orbitals 5. Apply the Aufbau principle, Hund’s rule and the Pauli exclusion principle to write full and condensed electronic configurations 6. Discuss the use of the four quantum numbers Recap from the pre-workshop material You were asked to compare the dxy, dyz and dxz orbitals. Similarity: same shape Difference: spatial orientation i.e. each of these three orbitals is pointing in a different direction ACTIVITY 5: Explaining Orbitals PART 1: orbital shapes Below is the shape of dx2- y2. Create a description of the similarities/difference of dx2- y2 compared to dxy, dyz and dxz orbitals. dxy dyz dxz Answer: Same shape but different spatial orientation dx 2 - y 2 ACTIVITY 5: Explaining Orbitals PART 2: orbital shapes Below is the shape of dz2. Why is the shape of this orbital different from the other four d orbitals? dz 2 dxy dyz dxz dx 2 - y 2 Recap from pre-workshop material Pauli Exclusion Principle: any orbital can contain a maximum of two electrons but they must be in opposing spins Hund’s Rule: electrons will fill the lowest unoccupied orbital singularly before pairing up. Aufbau Principle: electrons fill orbitals from the lowest available energy level As a result we would fill in the 4s orbital before putting electrons in the 3d level. BUT… Recap from pre-workshop material IMPORTANT: In reality, the 3d orbitals are actually lower in energy than 4s for Sc through Zn (not for K and Ca). The 3d-4s energy gap is relatively small for Sc but it progressively increases along the series! Why is then electronic configuration (gas phase) for Sc [Ar]4s23d1 rather than [Ar]3d3? - If we start with Sc3+ (gas phase) no issue there as its electronic configuration is [Ar] Do we then place an electron in the 3d or 4s orbital to make Sc2+? - We place it in the 3d level resulting in [Ar]3d1 How about the next electron to make Sc+ (gas phase)? - We would think that it would go to the 3d level but having two electrons in the 3d level would create significant electron-electron repulsion (which increases the overall energy of the system i.e. destabilizes the system) resulting in this electron being then placed in the 4s orbital (remember the energy gap between 3d and 4s is quite small) leading to [Ar]3d14s1 as the lowest energy state!! Recap from pre-workshop material IMPORTANT: In reality, the 3d orbitals are actually lower in energy than 4s for Sc through Zn (not for K and Ca). The 3d-4s energy gap is relatively small for Sc but it progressively increases along the series! How to make neutral Sc(gas phase)? - To make a neutral Sc atom we would need add one more electron and placing it to either a 3d or 4s orbital would create electron-electron repulsion so we would expect this electron to go to the lower energy level i.e. a 3d orbital. However, it turns out it is not that simple because electron-electron repulsion is “stronger” in the 3d orbitals (which are compactly arranged round the nucleus) than in the 4s orbital. As a result, [Ar]3d14s2 represents the lowest energy “solution” for Sc. Recap from pre-workshop material If we do the same exercise with vanadium (5 valence electrons), we will notice that the first time we place electrons in the 4s orbitals is in the last step i.e. when we want to make a neutral V atom from V+ simply due to the increasing 3d-4s energy gap: V5+: [Ar] V4+: [Ar]3d1 V3+: [Ar]3d2 V2+: [Ar]3d3 V+: [Ar]3d4 V: [Ar]3d34s2 Electron configuration of transition metals in gas phase A few important points: - For transition metals the nd orbitals are lower in energy than the corresponding (n+1)s orbital. (n = 3, 4 or 5) - It is then expected that the nd orbitals will be populated with electrons before the (n+1)s orbital. - However, due to several factors (the size of the energy gap between nd and (n+1)s orbitals and the magnitude of the electron-electron repulsion in these two respective set of orbitals) electrons are also added to the (n+1)s orbital in order to achieve the lowest energy state. - In other words, the Aufbau Principle is not always obeyed when it comes to the electron configuration of the transition metals. A few important points: - For the first row (i.e. 3d block) transition metals it appears that the 4s orbital is filled in first followed by the 3d orbitals (except for Cr and Cu). However, the electron configuration of the second row (i.e. 4d block) metals is more suggestive that the 4d is filled first before 5s. - Most importantly, the gas-phase electronic configurations of transition metals of the 3d block will be the only examinable material because it is very difficult to generally predict electronic configurations of transition metal ions in gas phase! Recap from pre-workshop material naked/free atom “dressed” atom in gas phase i.e. it has other atoms/ligands around it Fe(0): [Ar]4s23d6 Fe(0): [Ar]3d8 NOTE 1: the energy gap between 3d and 4s in a complex is significantly larger than in gas phase resulting in all valence electrons being allocated to the d level. NOTE 2: In transition metal chemistry we always talk about electronic configuration of an atom/ion as a part of a complex i.e. we are only concerned about the d orbitals and the electrons they hold. Activity 6 Give the condensed electron configurations for the following ions when (i) in gas phase (ii) as a part of a coordination complex: a) Sc2+ b) Fe0 c) V5+ d) Co2+ Activity 7 Using your knowledge of charges on common cations and anions, calculate the oxidation state of the transition metal in each of the following compounds: a) KMnO4 b) (NH4)2WO4 c) Na3VO4 d) Au2O3 Activity 8 (extension) There are numerous compounds/complexes that contain an oxygen atom (i.e. as part of a ligand) in between two transition metals. For the scenarios a, b and c given below, determine the oxidation state for each metal that are caused by the sole presence of oxygen atom(s) (the presence of any other atoms/ligands on these metals is not important).