Week 7 Transition Metals Workbook - Monash Chem PDF

9/2/24, 12:51 PM Week 7: Introduction to Transition Metals - workbook | MonashELMS1 Week 7: Introduction to Transition Metals - workbook Site: Monash Moodle1 Printed by...

9/2/24, 12:51 PM Week 7: Introduction to Transition Metals - workbook | MonashELMS1 Week 7: Introduction to Transition Metals - workbook Site: Monash Moodle1 Printed by: Kaltham Alzaabi Unit: CHM1022 - Chemistry II - S2 2024 Date: Monday, 2 September 2024, 12:51 PM Book: Week 7: Introduction to Transition Metals - workbook https://learning.monash.edu/mod/book/tool/print/index.php?id=2780818 1/24 9/2/24, 12:51 PM Week 7: Introduction to Transition Metals - workbook | MonashELMS1 Table of contents 1. Pre-workshop material 1.1. Periodic Trends of Transition Metals 1.2. Atomic Orbitals of Transition Metals 1.3. Revision of basic Redox 1.4. Oxidation states of Transition Metals 2. Summary 3. Preparation quiz 4. Online lectures 5. Workshops https://learning.monash.edu/mod/book/tool/print/index.php?id=2780818 2/24 9/2/24, 12:51 PM Week 7: Introduction to Transition Metals - workbook | MonashELMS1 1. Pre-workshop material Transition d-block elements The elements from group 3 to group 12 are collectively known as the transition metals or d-block elements (Figure 1). The properties of these elements often follow a set trend within the d-block and within each group, such as the ground state electronic configuration. They exist in low (1+, 2+) and high (3+ to 7+ or higher) oxidation states. The low oxidation states are more reminiscent of related elements in the s and p-blocks. However, the high oxidation states are much more interesting due to the coordination chemistry. Coordination chemistry is the study of compounds formed between metal ions and other neutral or negatively charged molecules, which will be further discussed in later weeks. Figure 1: d-block elements and the transition metal series Our focus will be on the first-row transition metals, Scandium to Zinc (Figure 2). Unlike main group elements, all transition elements are metals. Transition metals often form coloured compounds due to their electronic configuration and can be paramagnetic (attracted to a magnet). This magnetism is due to the presence of unpaired electrons and is the reason why metals are attracted to external magnetic sources. The d-electrons of these transition metal causes the bright colours often observed. A change in oxidation state can also influence the colour produced by the metal. Please read section 13.1 Metals in the periodic table of Chemistry, Blackman et al. (4th Ed.)(page 749) Figure 2: d-block elements and the transition metal series https://learning.monash.edu/mod/book/tool/print/index.php?id=2780818 3/24 9/2/24, 12:51 PM Week 7: Introduction to Transition Metals - workbook | MonashELMS1 1.1. Periodic Trends of Transition Metals Periodic trends Across the periodic table and down each column there are trends often observed in terms of electronic configuration, atomic radii, ionisation energy and so on. The video below will go through some of these major trends and how it is arranged in the periodic table before we go into more detail. The Periodic Table: Atomic Radius, Ionization Energy, and … https://www.youtube.com/watch?v=hePb00CqvP0 Some of the trends mentioned are atomic radii, ionisation energy, atomic size and electron affinity. It was seen from the video that as we move across the table we have a decrease in the atomic radii, but as we go down a group that radii increases etc. These effects can be rationalised by the change in the effective nuclear charge of each element as we go across and down the table. The value is known as Zeff and it represents to net positive charge of the nucleus. These trends have been summarised in Figure 3. Please read Section 4.5 and 4.8 of Chemistry, Blackman et al. (4th ed.)(page 194-199)(page 213-214) Figure 3: Trends in the periodic table Shielding and approximation of the effective nuclear charge (Zeff) Below we will be using the term shielding to describe trends of the transition metals in the periodic table. Shielding is the cancelling out of a portion of the attraction of the nucleus to the electron, due to the repulsion of these electrons to each other. In other words, the shielding effect is related to the effective nuclear charge (Zeff) because Zeff is defined as the nuclear charge that valance electrons experience i.e. the valence electrons are pulled towards the nucleus by Zeff. https://learning.monash.edu/mod/book/tool/print/index.php?id=2780818 4/24 9/2/24, 12:51 PM Week 7: Introduction to Transition Metals - workbook | MonashELMS1 In a very simplistic approach the Zeff value could be approximated by Zeff = Z – #CE Z = total number of protons in the nucleus (NOTE: you do not need to know this formula for the exam. It is only to reinforce the main point of this discussion) #CE = the number of the Core Electrons For example: what is the approximate value of Zeff for boron (B), carbon (C), vanadium (V) and nickle (Ni)? Therefore, the Zeff value for carbon is larger than for boron which means that the valence electrons for carbon will be pulled towards the nucleus a bit stronger than in the case of boron which means that carbons has a smaller atomic radius, a larger value for the first ionisation energy as well as a larger electronegativity value than boron. The same idea applies for vanadium vs nickel. It is also worth mentioning that the shielding of core electrons also depends on the nature of the orbital i.e. in general, electrons in the s orbital shield better than the electrons in the p orbital and so on. In other words, the core electrons found in the d and f orbitals do not shield the valence electrons from the influence of the nuclear charge to the same extent as the core electrons found in the s and p orbitals. s > p > d > f (relative shielding effect of electrons found in different orbitals) Atomic radius of transition metals For the d-block elements in particular we can see that filled inner d-orbitals shield the outer s-electrons from the increasing Zeff (Figure 4). Therefore the atomic radii across the period of transition metals only decrease slightly. The atomic radii of the d-block elements also decrease down each row. Figure 4: Atomic radii in the periodic table. Note that for transition metals between Sc to Zn, the atomic radii only decreases slightly compared to the main group elements. Please read section 4.8 atomic radii of Chemistry, Blackman et al. (4th ed.) (page 213) Ionisation energy Ionisation energy (IE) is the energy required for the complete removal of 1 mol of electrons from 1 mol of atoms or ions in the gaseous state. It can be summarised by the following: IE1 = first ionisation energy: removes an outermost electron from the gaseous atom: atom(g) → ion+(g) + e- ∆E = IE1 > 0 IE2 = second ionisation energy: removes a second electron from the gaseous ion: ion+(g) → ion+2(g) + e- ∆E = IE2 > IE1 Atoms with a low IE tend to form cations during reactions, whereas the opposite can be said for elements with a high ionisation energy, where they form anions (excluding the noble gases). The ionisation energy tends to correlate with the atomic size of the elements. For example, as the size of the atom decreases, the energy it takes to remove electrons increases. https://learning.monash.edu/mod/book/tool/print/index.php?id=2780818 5/24 9/2/24, 12:51 PM Week 7: Introduction to Transition Metals - workbook | MonashELMS1 This is due to the smaller size of the atom contracting the electrons closer to its positively charged nucleus, thus the electrostatic attraction requires more energy to overcome it. Following the trends in Figure 5, it can be seen that IE generally decreases down a group, while increasing across the period. This however is only a general trend. There are instances where this is not the case. In Figure 5, it can be seen that the first ionisation energy does not increase from Beryllium (Be) to boron (B) to Nitrogen (N) to Oxygen (O). This is due to the difference in the electronic configuration of these elements and the stability of half and fully-filled sub-shells. Figure 5: First ionisation energy of all the elements Please read section 4.8 ionisation energy of Chemistry Blackman et al. (4th ed) (Page 215) For example, any s-subshell (or s-orbital) can be empty (i.e. no electrons present in this orbital), half-filled (i.e. when one electron is present) or fully-filled (i.e. when two electrons are present). For a p-subshell there are three p-orbitals. The half-filled electronic configuration is achieved when three electrons are filled in this subshell according to the Hund’s rule i.e one electron is placed in each orbital with the electron spins aligned. The fully-filled scenario is possible only with six electrons in the p-subshell. All other possibilities (i.e. when we have 1, 2, 4 or 5 electrons is the p-subshell) are considered as partially filled and not energetically favourable. https://learning.monash.edu/mod/book/tool/print/index.php?id=2780818 6/24 9/2/24, 12:51 PM Week 7: Introduction to Transition Metals - workbook | MonashELMS1 Therefore, achieving either half- of fully-filled electronic configuration with respect to any subshell is favourable and if an atom needs to lose an electron (i.e. through ionisation) the energy required for this process will not be substantial. However, if an element already has one of these configurations the energy to lose one of the electrons will be greater. For example, beryllium’s electronic configuration (considering only the valence shell) is 2s2 which means that its s-orbital is fully filled and the first ionization energy of beryllium is higher than of boron’s because the boron’s configuration is 2s2 2p1 (the p-orbital is partially filled) which means that if boron were to lose the electron from the p-orbital (i.e. to ionise this atom) it will result in an electronic configuration of 2s2 which is fully-filled (i.e. the electronic configuration of boron cation (B+) is the same as of beryllium (Be) atom). Thus removing an electron to achieve fully-filled electronic configuration (i.e. ionising boron atom) requires less energy than for an electron from a fully-filled electronic configuration (i.e. ionising beryllium atom) was to be removed. Another example is the difference in the first ionisation energies of nitrogen and oxygen. The nitrogen’s electronic configuration (considering only the valence shell) is 2s2 2p3 which means the s-orbital is fully filled while the p-orbital is half-filled. This is a stable electronic configuration and more energy is needed to remove i.e. to ionise this atom. On the other hand, the oxygen’s electronic configuration is 2s2 2p4 which means that its p- orbital is partially filled and when oxygen is ionised to form oxygen cation (O+) it will form it would then form the nitrogen’s electronic configuration (which is favourable) and less energy is required. The same idea applies for the kinks observed in the first ionisation energy in Figure 5 as we progress from Mg to Al and P and S. It can be seen from the Figure 6 below that once the valence electrons from an atom has been removed the ionisation energy increases drastically. https://learning.monash.edu/mod/book/tool/print/index.php?id=2780818 7/24 9/2/24, 12:51 PM Week 7: Introduction to Transition Metals - workbook | MonashELMS1 Figure 6: Changes in ionisation energy, general trends Electron Affinity Electron affinity can be described as the energy change accompanying the addition of 1 mol of electrons to 1 mol atoms or ions in a gaseous state. It can be generalised in the equation below. The value of EA1 is in most cases negative atom(g) + e- → ion-(g) ∆U = EA1 When it comes to the value of EA2 it is always positive. Electronegativity Simply put, electronegativity is a measure of an atoms tendency to attract a bonding pair of electrons. This is measured commonly with the Pauling scale, which defines fluorine (F) as the most electronegative element with an assigned value of 4.0, and ranges down to the alkali metals caesium (Cs) and francium (Fr) which are the least electronegative elements with values of 0.7. Similar to the ionisation energy, the electronegativity of transition metals increases across the row but then decrease down a group. The change in electronegativity for the transition elements is relatively small across the transition series (Figure 7). It can also be noted that not all metals will follow the trend exactly, this can be seen in figure 7, where metals such as Rhenium (Re) actually increase in electronegativity down the group. Figure 7: Electronegativity across the transition metal series https://learning.monash.edu/mod/book/tool/print/index.php?id=2780818 8/24 9/2/24, 12:51 PM Week 7: Introduction to Transition Metals - workbook | MonashELMS1 1.2. Atomic Orbitals of Transition Metals Background: Quantum numbers, Atomic orbitals and electronic configurations The basics of atomic orbitals, quantum numbers and electronic configurations is summarised in the video below, further explanations are also provided. Quantum Numbers, Atomic Orbitals, and Electron Configur… Configur… https://www.youtube.com/embed/Aoi4j8es4gQ Extra videos from CHM1011 to help Week 1 from CHM1011: Understanding Quantum Numbers Week 1: Understanding Quantum Numbers Week 1: Understanding Quantum Numbers play_arrow Powered by Panopto get_app closed_caption keyboard_arrow_up Panopto Video: Understanding Quantum Numbers Note: Recommend Chrome or Firefox for playback on a MAC Atomic Orbitals-Pt1 Atomic Orbitals-Pt1 Atomic Orbitals-Pt1 play_arrow Powered by Panopto get_app closed_caption keyboard_arrow_up Panopto Video: Atomic Orbitals Pt. 1 Note: Recommend Chrome or Firefox for playback on a MAC Atomic Orbitals-Pt2 https://learning.monash.edu/mod/book/tool/print/index.php?id=2780818 9/24 9/2/24, 12:51 PM Week 7: Introduction to Transition Metals - workbook | MonashELMS1 Atomic Orbitals-Pt2 Atomic Orbitals-Pt2 play_arrow Powered by Panopto get_app closed_caption keyboard_arrow_up Panopto Video: Atomic Orbitals Pt. 2 Note: Recommend Chrome or Firefox for playback on a MAC Atomic orbitals and transition metals d-orbitals are of particular interest for the transition metals. They are five-fold degenerate (all equal energy) and like p-orbitals are directional (Figure 8). However, unlike the p-orbitals they have two nodal planes (which is when the probability of finding an electron is zero) and the sign change does not occur through the centre (a sign change for example is when the two p atomic orbitals are in phase with each other but the two lobes have opposite signs). Please read Section 4.5 of Chemistry Blackman et al, (4th ed.) Figure 8: The five equivalent 3d-orbitals, reproduced from: Chemistry: The Central Science by Brown, LeMay, Busten, Murphy, and Woodward Comparison of d orbitals Compare the 3dxy, 3dyz and 3dxz orbitals. Find one similarity and one difference between the 3 orbitals (hint: consider the axes and shapes of each!). The animations below may help you visualise the different shapes and axes. https://learning.monash.edu/mod/book/tool/print/index.php?id=2780818 10/24 9/2/24, 12:51 PM Week 7: Introduction to Transition Metals - workbook | MonashELMS1 Please read The Pauli exclusion principle (page 188-189) and Section 4.6 and 4.7 the Aufbau principle and order of orbital filling, including Hund's rule, Blackman et al, (4th ed.) Orbital filling and electronic configuration Hund’s rule in relation to electronic configurations state that when you are filling a degenerate set of orbitals in the ground state the pairing of electrons cannot begin until each orbital in the set contains one electron. Electrons in a degenerate set will singly occupy these orbitals in parallel spins (Figure 9) Figure 9: Hund’s rule Pauli Exclusion Principle The Pauli exclusion principle states that no two electrons in an atom can be defined by the same four quantum numbers n, I, ml, and ms. What this tells us is that every electrons in an atom is uniquely defined by its set of four quantum numbers. The spins of those electrons in each orbital therefore cannot be the same (Figure 10). Figure 10: Visual representation of the Pauli Exclusion Principle The Aufbau principle The Aufbau principle states that in the ground state of an atom or ion, electrons fill atomic orbitals of the lowest available energy levels before occupying higher levels. For example, the 1s shell is filled before the 2s subshell is occupied (Figure 11) https://learning.monash.edu/mod/book/tool/print/index.php?id=2780818 11/24 9/2/24, 12:51 PM Week 7: Introduction to Transition Metals - workbook | MonashELMS1 Figure 11: Visual representation of the Aufbau principle Hund’s rule, Pauli exclusion principle and Aufbau principle Three Rules-pt1 Three Rules-pt1 play_arrow Powered by Panopto get_app closed_caption keyboard_arrow_up Panopto Video: Three Rules Pt. 1 Note: Recommend Chrome or Firefox for playback on a MAC Three Rules-pt2 Three Rules-pt2 play_arrow Powered by Panopto get_app closed_caption keyboard_arrow_up When filling orbitals it is easiest to envision this as a graph, starting from the orbital with the lowest sum (Figure 12) https://learning.monash.edu/mod/book/tool/print/index.php?id=2780818 12/24 9/2/24, 12:51 PM Week 7: Introduction to Transition Metals - workbook | MonashELMS1 Figure 12: Orbital filling This sum is derived from the principle quantum number (n) and angular quantum number (l) and dictates the order of electron addition in orbital filling (n + l). However, there are instances where the sum may be equal, for example, both 2p and 3s have the same n + l value of 3, therefore the one with the lowest value of n, in this case 2p, will be filled first. Generally, when filling electronic configurations, it is only the outer-shell orbital that are taken into account, the remaining sub-shells can be noted by the electronic configuration of a noble gas. This is as the noble gases are inert due to fully filled sub-shells.For example, the full electronic configuration of nickel (Ni) is 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d8 but its condensed electronic configuration is [Ar] 4s2, 3d8 as the electronic configuration of argon (Ar) is equal to: 1s2, 2s2, 2p6, 3s2, 3p6. A general rule for the filling of these sub-shell is [noble gas] ns2 (n -1)dx where n ranges from 4 – 7 and x ranges from 1 to 10. IMPORTANT: Even though in Figure 12 it is depicted that, for example, 4s is filled before 3d, for Sc through Zn i.e. the first row of transition metals (not for K and Ca) the 3d orbitals are actually lower in energy than 4s. The same observation applies for the second and third rows of transition metals i.e. 4d and 5d are lower in energ compared to 5s and 6s, respectively. 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]. If we are to make Sc2+ we will place an electron in the 3d orbital resulting in [Ar]3d1. How about the next electron to make Sc+ (gas phase)? We would think that it would again 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!! Laslty, 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. 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+ and this is 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 If we look at the electronic configuration of the first row transition metals (Figure 13) it does appear that the 4s orbiral is filled before the 3d level. https://learning.monash.edu/mod/book/tool/print/index.php?id=2780818 13/24 9/2/24, 12:51 PM Week 7: Introduction to Transition Metals - workbook | MonashELMS1 Figure 13: Electronic configuration of the first row. However this is deceiving because the electronic configuration of the secon row (Figure 14) clearly shows that the s orbital is not filled before the d level. Figure 14: Electronic configuration of the second row. Another misconception is that the electronic configurations (Figure 13) of Cr and Cu of [Ar]3d54s1 and [Ar]3d104s1, respectively, are simply due to the half/fully filled concept. However that is not true as this is not observed for tungsten which is in the third row (Figure 15). Figure 15: Electronic configuration of the third row. 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. https://learning.monash.edu/mod/book/tool/print/index.php?id=2780818 14/24 9/2/24, 12:51 PM Week 7: Introduction to Transition Metals - workbook | MonashELMS1 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. 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 only the first row transition metals in their neutral state will be examinable as it is generally very difficult to predict electronic configurations of transition metal ions in gas phase! Electronic configuration of transition metals when they form complexes The gap between, for example, the 3d and 4s levels significantly widens when a transition metal forms a complex i.e. when it contains other atoms/molecules bound to it. As a result, all valence electrons exclusively fill the 3d orbitals. Thus: If you have a "naked" Fe(0) atom in the gas phase its electronic configuration would be [Ar]4s23d6 (Figure 13). On the other hand, if have an Fe(0) ion in a complex (e.g. Fe(CO)5) its electronic configuration would be [Ar]3d10 The same idea applies for Zn ions i.e. "naked" Zn2+: [Ar]4s23d8 ; Zn2+ in a complex: [Ar]3d10 Conclusion: when we form transition metal complexes (which will be the main theme of this part), the electrons found in the 4s orbital technically "drop" to the 3d level and we only talk about the d electrons. NOTE: this also applies to the 2nd and the 3rd row transition metals. https://learning.monash.edu/mod/book/tool/print/index.php?id=2780818 15/24 9/2/24, 12:51 PM Week 7: Introduction to Transition Metals - workbook | MonashELMS1 1.3. Revision of basic Redox Oxidation Please read Section 12.1 redox reactions and oxidation numbers of Chemistry, Blackman et al. (4th ed.) Oxidation was originally considered addition of oxygen atom or a loss of hydrogen atom to a molecule. However, oxidation is the LOSS of electrons. e.g. Na(s) → Na+(aq) + e- Cu(s) → Cu2+(aq) + 2 e- Oxidation half reactions have electrons as a product Remember OIL “Oxidation Is Loss” Sodium in Water Sodium in Water https://youtu.be/ODf_sPexS2Q Reduction The opposite half reaction to oxidation is reduction, the GAIN of electrons. e.g. Na+(aq) + e- → Na(s) Cu2+(aq) + 2 e- → Cu(s) Reduction half reactions have electrons as a reactant Remember RIG “Reduction Is Gain” Redox Reaction Redox reaction requires both OXIDATION and REDUCTION. Electrons from the oxidation reaction are used for reduction reaction https://learning.monash.edu/mod/book/tool/print/index.php?id=2780818 16/24 9/2/24, 12:51 PM Week 7: Introduction to Transition Metals - workbook | MonashELMS1 i.e. redox reactions involve electron transfer e.g. Na(s) → Na+(aq) + e- (oxidation) Cu2+(aq) + 2 e- → Cu(s) (reduction) Redox Terminology Oxidising Agent/Oxidant A substance that causes another to be oxidised the oxidant itself undergoes reduction Reducing Agent/Reductant A substance that causes another to be reduced the reductant itself undergoes oxidation For example: Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq) Copper is going from Cu2+ to Cu, thus Cu2+ is being reduced to Cu. Cu2+(aq) + 2 e- → Cu(s) [Reduction - Gain of Electrons] Zinc is going from Zn to Zn2+, thus Zn is being oxidised to Zn2+. Zn(s) → Zn2+(aq) + 2 e- [Oxidation - Loss of Electrons] The oxidant is Cu2+ (as it causes Zn to undergo oxidation) The reductant is Zn (as it causes Cu2+ to undergo reduction) Oxidation Numbers “Hypothetical charge that an individual atom/ion in a molecule would possess if the shared electrons in a covalent bond were assigned to most electronegative element in the bond” Oxidation number (Ox#) has two parts: 1. Sign (positive/negative) 2. Value Notation: The sign is before the value for oxidation numbers Basic Rules: Ox# on neutral atom or molecule with one element is 0 e.g. Pb = 0, Ar = 0, H2 each H = 0, S8 each S = 0 Ox# on single atom ions is equal to the charge on ion e.g. Cl- = -1, Pb2+ = +2, S2- = -2 Ox# for each element adds up to total charge on multiple element molecules https://learning.monash.edu/mod/book/tool/print/index.php?id=2780818 17/24 9/2/24, 12:51 PM Week 7: Introduction to Transition Metals - workbook | MonashELMS1 e.g. H2O charge is 0, thus 2H and 1O ox# must equal 0 NO3- charge is 1-, thus N and 3O ox# must equal -1 Ox# for F in molecules is -1 e.g. [AlF6]3- each F = -1, ClF4 each F = -1 Ox# for H in most compounds is +1 (excludes metal hydrides, e.g. NaH, where it is -1) e.g. H2O each H = +1, C6H12O6 each H = +1 Ox# for O in most compounds is -2 e.g. H2O the O = -2, NO3- each O = -2 Ox# for group 1 ions are +1, group 2 ions are +2 e.g. Na+ = +1, Mg2+ = +2 Oxidation numbers can determine which element is undergoing oxidation and which one is undergoing reduction. 1. Increase in oxidation number = oxidation (e- lost) e.g. Cu(s) → Cu2+(aq) + 2 e- Cu = 0, Cu2+ = +2 2. Decrease in oxidation number = reduction (e- gained) e.g. Na+(aq) + e- → Na(s) Na+ = +1, Na = 0 https://learning.monash.edu/mod/book/tool/print/index.php?id=2780818 18/24 9/2/24, 12:51 PM Week 7: Introduction to Transition Metals - workbook | MonashELMS1 1.4. Oxidation states of Transition Metals Please read Section 13.2 oxidation states of transition metals of Chemistry, Blackman et al. (4th ed.) Oxidation Multiple oxidation states are a common feature of the transition metal series. As with most other metals, transition metals will prefer to lose electrons and form positively charged ions. However, in some instances a transitional metal could have a negative oxidation state (e.g. the formal oxidation state of Fe in K2[Fe(CO)4] is -2). The oxidation state of +2 is very common for transition elements because ns2 electrons are more easily lost than those in the inner sub-shells. The formation of +3 oxidation states is also very common but other oxidation states are also readily accessible with manganese having the highest stable state of +7 (Figure 14). When going along the row up to Mn, it can be seen that the highest stable oxidation state, matches the group number (Figure 14). Higher oxidation states (e.g. +8, etc.) are very uncommon due to the increased ionisation energy. It should be also noted that Zn has really only one accessible oxidation state +2 which is primarily due to Zn2+ having a d10 electronic configuration i.e. a fully-filled d sub-shell. Figure 14: Common oxidation states of first row transition metals. Most stable oxidation states are highlighted. One of the most significant differences therefore between the d-block elements and those from the s and p-block is this wide range of stable oxidation states. Most often elements from the s-block will only have one accessible stable oxidation state, and those form the p-block one or two. Elements from the d-block have a wide range of accessible and stable oxidation states that often dictate a wide variety of chemical properties and chemical reactivity (Figure 14). https://learning.monash.edu/mod/book/tool/print/index.php?id=2780818 19/24 9/2/24, 12:51 PM Week 7: Introduction to Transition Metals - workbook | MonashELMS1 2. Summary You have learnt the basics of transition metal chemistry and the periodic trends observed as we go across the group. We’ve also discovered that these trends are not final and there are exceptions to the rules. We’ve also looked at the more stable oxidation states of the first-row transition metals and learned how to write the electronic configurations. https://learning.monash.edu/mod/book/tool/print/index.php?id=2780818 20/24