Transition Metals Flashcards
oxidation states and successive ionisation energies of metals in groups 1, 2 and 13
The metals in groups 1, 2 and 13 form ions with only one charge/ oxidiation state. The successive ionization energies to remove their valence electrons are relatively low, but after that, the successive ionization energies increase drastically. This makes it energetically unfavorable for those metals to lose any additional electrons from their inner shells.
oxidation states and successive ionisation energies of transition metals
With transition metals, however, their valence s-electrons and inner d-electrons are very close in energy due to the convergence of electron energy levels. This results in successive ionization energies that gradually increase through the removal of both s & d electrons. Therefore, it is possible for not just the outer s-electrons to be removed, but also many, if not all, of the inner d-electrons to be removed. This can help explain how multiple stable oxidation states are able to be formed.
transition metal
A transition metal is an element with an incomplete d-sublevel (1-9 d electrons) in one or more of its oxidation states (except zinc).
an element that forms at least one ion with a partially filled d-subshell. these elements are typically found in the d-block of the periodic table (group 3-12) and exhibit properties sucuh as variable oxidation states, the ability to form complex ions, coloured copounds, and catalytic activity.
physical properties of transition metals
1) delocalised d electrons: Although not valence electrons, the d electrons of transitional elements become delocalized along with the valence s electrons. This is because the s and d levels are so close in energy.
2) high melting points 3) electrical conductivity: The extra delocalized eletrons increase the strength of the metallic bonds, leading to higher melting points and greater electrical conductivity than other metals.
3) easily form alloys: Because of their similar sizes, transition metals can easily form alloys with one another, as their atoms can readily replace one another in the lattice arrangement.
catalyst
A catalyst speeds up a chemical reaction by providing an alternate pathway for a chemical reaction with lower activation energy. It is not consumed (used up) by the reaction.
catalytic behaviour of transition metals
only required to focus on their variable oxidation states
In their solid form, transition metals and their compounds have common lattice arrangements of their atoms, which can have porous surfaces. This allows reactant particles to bind to their surfaces.
The fact that transition metals have many stable oxidation states allows them to gain and lose electrons easily with the reactant particles, which can often help break reactant bonds more easily (with less energy). This speeds up the rate of the reaction.
The products, once formed, leave the surface of the transition metal catalyst, or can be easily separated by some form of filtration.
why can transition metals act as catalysts?
because they can form variable oxidaton states.
example with iron ( which can be electron donors and recievers)
1) Fe^+2 (aq) + X (aq) –> Fe^3+ + X^- (aq)
2) Fe^3+ (aq) + Y (aq) –> Fe^2+ + Y^+ (aq)
heterogeneous catalysts
Heterogeneous Catalysts are in a different state of matter than the reactants they are helping to react.
homogeneous catalysts
Homogeneous catalysts are in the same state of matter as the reactants. Homogenous catalysts are commonly aqueous along with aqueous reactants.
what creates a magnetic field?
Spinning electric charges create magnetic fields. Each electron, therefore, creates its own very tiny magnetic field. With most substances, the electrons are arranged randomly and these magnetic fields cancel out, so there is no net magnetic field.
If subjected to an external magnetic field, the electrons within an element will align with, or against, the applied magnetic field. Depending on the electron configuration of an element, it will respond differently to these applied magnetic fields.
Distinctions between different types of magnetism are not
required for the IB Chemistry Exam. Remembering that magnetism is related to unpaired electrons is probably enough.
magnetism of transition metals
Transition metals have at least one unpaired electron in their ground state configuration (lowest energy electron configuration of an atom). This means that all transition metals are paramagnetic (having at least one unpaired electron causes them to be attracted to an external magnetic field). However, only a small number of transition metals are ferromagnetic, which is the type of magnestism we mean when we think of the term “magnetic.”
Distinctions between different types of magnetism are not
required for the IB Chemistry Exam. Remembering that magnetism is related to unpaired electrons is probably enough.
which transition metals are ferromagentic?
Ferromagnetism is only displayed by iron, cobalt and nickel (amongst 3d transition metals). Under an applied magnetic field, the electrons align with the magnetic field and the elements show attraction in the direction of the magnetic field. When the magnetic field is removed, the electrons remain aligned,
causing the elements to maintain its own magnetic field. This means the metal will now act as a permanent magnet.
Distinctions between different types of magnetism are not required for the IB Chemistry Exam. Remembering that magnetism is related to unpaired electrons is probably enough.
ligands
packet: species with at least 1 nonbonding electron pair that form coordinate covalent bonds with a metal ion.
chatgpt: A molecule or ion that donates a pair of electrons to a central metal atom or ion to form a coordinate covalent bond in a complex.
examples of common ligands:
I- < Br- < S2- < Cl- < F- < OH- < SCN- < NH3 < CN- ≈ CO
asked to draw lewis structures
It is important to know whether or not the ligand has a negative charge or not. Usually they do, but some important and common exceptions are water (H2O) and ammonia (NH3). The list of common ligands is shown above and in section 15 of the data booklet.
complex / complex ion
formation of complexes: The high (positive) charge density of the metal ions is so great, it attracts the electron pairs of nearby molecules and ions. This results in the formation of a coordinate covalent bond between the metal ion and the ion/molecule (now called a ligand).
The large structure formed between the covalently bonded transition metal ion and ligands is referred to as a complex. If they have an overall charge, they are called complex ions.
coordinate (covalent) bonds
bonds where the shared electron pair has been donated by one of the two bonding atoms (as opposed to each contributing one electron to the shared pair).
chemical formulas of complexes
When written as chemical formulas, complex ions are usually surrounded by brackets [ ]. The overall charge is written outside of the brackets.
[Co(NH3)6]^3+
[CuCl4]^2-
Fe(H2O)3(CN)3
complex ion charge
determined by the combination of the oxidation state of the metal and the charge and number of the ligands.
In complexes, although the electron pairs are shared between the ligand and the metal, the electrons are considered to belong more to the ligand than to the metal. The metal therefore ends up with a positive oxidation state, which can be determined by considering the overall charge of the complex, as well as the number and charges of the ligands on the complex.
Example:
[Co(H2O)6]^2+
Overall Charge = 2+
Charge on Ligand (H2O) = 0
Number of Ligands = 6
Total Charge from the ligands = 0
Oxidation State of the Metal = +2
Overall Charge = 2-
Charge on Ligand (Cl^-) = 1-
Number of Ligands = 4
Total Charge from the ligands = 4-
Oxidation State of the Metal = +2
coordination number
The number of coordinate covalent bonds to ligands around a given transition metal ion.
Usually, this can be determined by looking at the chemical formula. With a few exceptions, the number of ligands is equal to the coordination number.
geometry
the 3D shape of the complex
Coordination number and 4 possible geometries
- linear: 2 e domains and coordination number of 2
- square planar: 4 e domains and coordination number of 4
- tetrahedral: 4 e domains and coordination number of 4
- octahedral: 6 e domains and coordination number of 6
Complexes with 4 ligands can form both tetrahedral and square planar geometries. It is not important to be able to predict which one of these geometries is more likely in this topic, and so predicting tetrahedral is probably the simpler and safer approach.
why do things appear to have colour?
The molecules in the substance absorb the wavelengths of visible light that are complementary to that color.
The other wavelengths of visible light transmit through and reflect off the substance, combining to create the appearance of a particular color.
(electrons)
absorption of visible light
Electrons absorb photons of visible light, which causes them to transition from a lower energy level to a higher energy level.
The difference in energy between the where the electrons begin & end determines which wavelengths of visible light are absorbed.
what can it be used to calculate?
the wavelength of max absorbance
In practice, a complex usually absorbs a wide range of wavelengths around the complementary color. Measuring the wavelength of maximum absorbance can be used to calculate the Energy difference between the 2 different sets of d-orbitals.
which transition metals in compounds do not have a color and why?
compounds of Zn^2+ ions, Cu^+ ions and Sc^3+, do not have a color. they all have complete d-sublevels, where an incomplete d-sublevel is necesarry for colour (absorbing visible light).
How does the geometry of an octahedral complex ion align with the d-orbitals of the metal ion?
explained with [Fe(H2O)6]^3+ in packet
the orientation of 2 of the d-orbitals is directly aligned with 2 ligands in octahedral complexes. This means that electrons in the 2 d orbitals with the same orientation experience more repulsion than electrons in the other 3 d orbitals. Although the 5 d orbitals originally had the exact same energy, in the presence of ligands they split into 2 different energy sublevels. 3 d orbitals are lower in energy and 2 are higher in energy in the presence.
The splitting of the d sublevel provides just the right amount of energy difference for electrons to absorb visible light.
*Note, this can only happen if the d-sublevel is incomplete, as this means there is at least 1 empty space in the higher half of the d-sublevel for an electron to be promoted into.
a markscheme explanation for Cu^2+ complexes (+ other colourless complexes) producing color (worth up to 4 marks)
1) The d-orbitals split into two energy levels in the presence of ligands.
2) Electrons absorb energy from visible light to move to the higher energy d-orbital (d-d transition).
3) The d-sublevel is incomplete, allowing these transitions to occur.
4) The color observed is the complementary color to the wavelength of light absorbed.
the relationship of concentration and light absorbance
for solution with a colour, the Beer-Lambert Law states that for relatively low concentrations, the relationship between concentration and light absorption is linear.
using a colorimeter or spectrometer, absorbance data can be recorded for standard solutions of known concentration, to create a calibration curve.
the trendline can be used to determine the concentration of a solution with an unknown concentration, based on its absorbance.
Metallic Bond
Electrostatic attraction between a lattice of positive ions and sea of delocalized electrons
What would changing the ligand do?
Changing the ligand (in other words the coordination number/geometry) changes the energy difference between the d-orbitals
