Chapter 22 Notes PDF

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Combination of textbook notes for Chapter 22, as well as content from the lectures. Includes important visuals to understand the topic at hand....

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Chapter 22 Notes: Transition Metals 22.1 - Transition Metals in Biology (Complex Ions) Interactions between transition metals and accompanying nonmetal ions and molecules influence the solubility of the metals, which is a key factor toward making them chemically reactive and biologically available to plants and animals. ● Interactions also influence other properties, including wavelengths of visible light the metals absorb and therefore the colors of their compounds and solutions. Acid-Base Properties: ● The Arrhenius definition for acids and bases only refers to compounds dissolved in WATER. ○ Arrhenius acid - A molecule that when dissolved in water will donate an H+ in solution (proton donor). ○ Arrhenius base - A molecule that when dissolved in water will break down to yield an OH- or hydroxide in solution. ● The Bronsted-Lowry definition for acids and bases is similar to the Arrhenius definition, but the solution does not have to be water. ○ Bronsted-Lowry acid - A compound that breaks down to give an H+ in solution. ○ Bronsted-Lowry base - Any atom or ion capable of accepting or bonding to a free proton in solution. To understand the interactions between transition metal ions and the other ions and molecules that surround them in solids and solutions, we need to review Lewis acids and Lewis bases used in Chapter 16. ● A Lewis base is a substance that donates a lone pair of electrons in a chemical reaction. ● A Lewis acid is a substance that accepts a lone pair of electrons in a chemical reaction. Anything with a lone pair of electrons is a BASE and anything that can accept a lone pair of electrons is an ACID.

In Chapter 10 we described that ions dissolved in

water are hydrated - they are surrounded by water molecules oriented with their positive dipoles directed toward anions and their negative dipoles directed toward cations. ● When these ion-dipole interactions lead to the sharing of lone-pair electrons with empty valence-shell orbitals on the cations, they meet our definition of covalent bonds, and in this case, are called coordinate covalent bonds. Coordinate Bonds - A covalent bond in which one of the atoms contributes both of the electrons in the shared pair. Once formed, a coordinate bond is the same as any other covalent bond. ● Such bonds form when either a molecule or an anion donates a lone pair of electrons to an empty valence-shell orbital of an atom, cation, or molecule. ● When a cation forms coordinate bonds to one or more molecular or ionic electron-pair donors, the resulting structure is called a complex ion. ● The electron-pair donors in complex ions are called ligands.

→ BF3 (a 6 e- complex) bonds with a NH3 (with a lone pair). → The bond is a coordinate covalent bond.

Molecules or anions that function as Lewis bases and form coordinate bonds with metal cations are called ligands. The resulting species, which are composed of central metal ions and the surrounding ligands, are called complex ions, or simply complexes. ● Direct bonding to a central atom means that the ligands in a complex occupy the inner coordination sphere of the cation. ● The capacity to form two kinds of bonds means that transition metal ions have two kinds of bonding capacity, or valence.

○ First kind involves ionic bonds and is based on the number of electrons a metal atom loses when it forms an ion. This valence is equivalent to its oxidation number. ○ The second kind of valence is based on the capacity of the metal ions to form coordinate bonds. This property corresponds to an ion’s coordination number. Coordination Number - The number of sites occupied by ligands around a metal ion in a complex. Coordination Compound - A compound made up of at least one complex ion. Counterion - An ion whose charge balances the charge of a complex ion in a coordination compound. One can use Lewis structures to show how bonds are formed or broken in reactions. We can identify the Lewis acid and the Lewis base and see how electrons rearrange to give the product.

Examples of ligands with lone pairs: → NH3 has one lone pair → H2O has two lone pairs → Cl- has three lone pairs → OH- has four lone pairs These four atoms are all similar because they are monodiante.

→ CN- has lone pairs on each end, making it capable of holding two metal ions, one at each end. Notice the lone pairs that can coordinate to a metal ion. Complex ions are formed when metal ions coordinate (covalently) to a molecule or ion with lone

pairs of electrons. Example: This is a TETRAHEDRAL COMPLEX.

For this complex, the NH3 ligands are the Lewis base. They are considered to be in the ‘inner-sphere’ and donates 2 electrons. The overall charge on the complex is 2+. In the solution, there must be two counter ions, which reside in the ‘outer-sphere’ and are there to balance the charge. Anions to balance the charge are spectator ions. They do not participate directly in the reaction, but rather exist in the same oxidation state on both the reactant and product side of the chemical equation. When we write out the formula of the product, spectator ions are written outside the square brackets to emphasize their location.

→ In this diagram, the RED atoms are the ligands. → When writing the name of the complex ions, the METAL always goes inside the bracket. → Whatever atom is written inside the bracket is present in the ‘inner-sphere’ and those written outside the bracket are present in the ‘outer-sphere.’

The identity and location of the ligand will dramatically change the color and behavior of the atom.

→ If a ligand can easily lose a water molecule, it is an ACID. → Metal ions are only positively charged, and pull electron density, making it easier for a ligand to lose an acid.

In this diagram, a hydrated Fe3+ cation draws electron density away from the water molecule of its inner coordination sphere, which makes it possible for one or more of these molecules to donate H+ ion to water a molecule outside the sphere. The hydrated Fe3+ ion has one fewer H2O ligands with an electrostatic attraction between the complex ion and the OH- ion.

Simply, if a ligand can lose H+ (aq), usually in water, it is an acid. → 2+ charges are relatively weaker than 3+ charges, as we would expect given the greater electronwithdrawing power of the more highly charged central ions. → In the diagram, one of the six water molecules of hydration surrounding the Fe3+ ion is converted into a hydroxide ion as a result of the acid ionization reaction. This reduces the charge of the complex ion from 3+ to 2+. → If the pH of a solution is raised by adding a small quantity of a strong base, the ions will undergo additional acid ionization reactions. Complex Ion Formation: The equilibrium constant Kf is called a formation constant - it describes the formation of a metal complex from a free metal ion and its ligands. ● Formation constants can be used to calculate the concentration of free, uncomplexed

metal ions, in equilibrium with a given concentration of a ligand. ● Kf is usually quite large, therefore equilibrium concentrations of uncomplexed metal ions are usually very small. One approach to calculating the concentration of an uncomplexed metal ion is to couple two reactions - Ksp and Kf. Example: AgCl (s) ← → Ag+ (aq) + Cl- (aq)

Ksp = 1.8E-10 (very small number)

AgCl (s) is soluble, but if NH3 (aq) is dissolved in the solution, more solid dissolves Ag + (aq) + 2NH3 (aq) ← → Ag(NH3)2 2+ (aq) Kf = 1.7E7 (very large number) In this reaction, Ag and NH3 turn into a complex. Find Keq for the coupled reaction by multiplying (Ksp)(Kf). Example:

→ Weak ligands will have small Kf values. → The table is ordered from strongest to weakest complex ion, based on the Kf values.

22.2 - Naming Complex Ions and Coordination Compounds The names of complex ions and coordination compounds tell us the identity and oxidation state of the central ion, the names and numbers of ligands, the charge in the case of complex ions, and the identity of counterions. Complex Ions with a Positive Charge: 1. Start with the identities of the ligand(s). Names of common ligands appear in the table below. If there is more than one kind of ligand, list the names alphabetically. 2. Use the usual prefixes in front of the game written in the first step to indicate the number of each type of ligand (di, tri, etc.). 3. Write the name of the metal ion with a Roman numeral indicating its oxidation state.

Examples: Ni(H2O)6 2+ = Hexaaquanickel(II) Co(NH3)6 3+ = Hexaamminecobalt(III) [Cu(NH3)4(H2O)2]2+ = Tetraamminediaquacopper(II)

Complex Ions with a Negative Charge: 1. Follow the steps for naming positively charged complexes. 2. Add -ate to the name of the central metal ion to indicate that the complex ion carries a negative charge (just as we use -ate to end the names of oxoanions). ● For some metals, the base name changes too. ● The two most common examples are iron, which becomes ferrate, and copper, which becomes cuprate.

→ Notice that the overall complex is NEGATIVE!

Coordination Compounds: 1. If the counterion of the complex is a cation, the cation’s name goes first, followed by the name of the anionic complex ion. 2. If the counterion of the complex is an anion, the name of the cationic complex ion goes first, followed by the name of the anion. A key to naming coordination compounds is to recognize from their formulas that they are coordination compounds! To help with this, look for formulas that have the atomic symbols of a metallic element and one or more ligands, all in brackets. → ‘Inner-sphere’ ligands are in the bracket, and ‘outer-sphere’ ligands are outside the bracket.

22.3 - Polydentate Ligands and Chelation Many ligands can donate only one pair of electrons to a single metal ion. Even atoms with more than one lone pair usually donate only one pair at a time to a given metal ion because the other lone pair or pairs are oriented away from the metal ion. ● Monodentate Ligands - A species that forms only a single coordinate bond to a metal ion in a complex. Certain molecules larger than ammonia and water may be able to donate more than one lone pair of electrons and therefore form more than one coordinate bond to a central metal ion. ● Polydentate Ligands - A species that can form more than one coordinate bond per molecule. ● More specifically, these ligands are called bidentate, tridentate, and so on.

→ The bidentate ligand pictured in (a) has two N atoms that can each donate a pair of electrons to empty orbitals of adjacent octahedral bonding sites on the same Ni2+ (aq) ion, represented by the gold sphere. This displaces two molecules of water. → Pictured in (b), is three molecules occupying all six octahedral coordination sites of a Ni2+ ion.

Larger molecules may have even more atoms per molecule that can bond to a single metal ion. The interaction of a metal ion with a ligand having multiple donor atoms is called chelation. Chelation means ‘claw’ and indicates multiple coordination to the central metal ion. Polydentate ligands that take part in these interactions are called chelating agents. We know that under standard conditions, the change in free energy (Delta G) is related to changes in enthalpy and entropy that accompany the reaction. To understand why there is such a large increase in entropy, consider that there are 4 moles of reactants but 7 moles of products in an equation. Nearly doubling the number of moles of products over reactants translates into a large gain in entropy. Entropy gains drive many complexation reactions that involve polydentate ligands.

Chelate Effect - The greater affinity of metal ions for polydentate ligands than for monodentate ligands. Usually, polydentate ligands will DISPLACE monodentate ligands due to increased entropy. ● It is important to understand that some ligands react ‘more strongly’ with metal ions. ● For example, NiCl2 6H2O dissolves in water, but changes color when NH3 is added. ● Many chelating agents have more than one kind of electron-pair-donating group. ● Sequestering agents are chelating agents that bind metal ions so tightly that they are ‘sequestered’ and prevented from reacting with other substances. 22.4 - Crystal Field Theory Many metal complexes are brightly colored in solution. For these systems, colors are due to electron transitions in the d orbitals. Remember the periodic table specifies how many electrons are “valence” electrons. In the “d” row, there are 5 d orbitals that can hold up to 10 electrons total. Four of the orbitals have similar shapes, but POINT in different directions. → A Cr3+ ion has the electron configuration [Ar]3d^3 in the gas phase. When a Cr3+ ion is in the gas phase, all of the orbitals in a given subshell have the same energy. → When a Cr3+ ion is in an aqueous solution and surrounded by an octahedral array of water molecules, the energies of its 3d orbitals are no longer all the same.

The 3d(xy), 3d(yz), and 3d(xz) orbitals experience some increase in energy. They are located between axes and will behave similarly. The energies of the 3d(x^2-y^2) and 3d(x^2) orbitals increase even more because the lobes of their orbitals point directly toward the H2O molecules’ oxygen atoms at the corners of the octahedron formed by the ligands and are repelled by the electrons on those O atoms. These lobes are along the axes, as represented in the diagram above. The first row’s lobes do not point directly toward the corners of the octahedron, so the electron repulsion they experience is weaker.

Crystal Field Splitting - The separation of a set of d orbitals into subsets with different energies as a result of interactions between electrons in those orbitals and lone pairs of electrons in ligands. Crystal Field Splitting Energy (Delta) - The difference in energy between subsets of d orbitals split by interactions in a crystal field.

→ A Cr3+ ion in an octahedral field can ABSORB a photon of light that has energy (hv) equal to Deltao. → This energy raises a 3d electron from (a) one of the lower-energy d-orbitals to (b) one of the higher-energy d orbitals. Think of it like this: when you ABSORB a photon of light that has energy, you INCREASE in energy level. The wavelength of the absorbed photon is related to the energy difference between the two groups of orbitals - in other words, to

the crystal field splitting energy. The equation is: Energy and wavelength of a photon are inversely proportional to each other. Therefore, the larger the crystal field splitting in a complex ion, the shorter the wavelength of the photons the ion absorbs. ● Larger Deltao = stronger ligand = high energy absorbed = short wavelength ● Smaller Deltao = weaker ligand = low energy absorbed = large wavelength The size of the energy gap between split d orbitals corresponds to radiation in the visible region of the electromagnetic spectrum. ● The colors of solutions of metal complexes depend on the strengths of metal-ligand interactions that affect Deltao. ● The color we perceive for any object is not the color it absorbs but rather the color that it transmits. ● Red and green are complementary colors; therefore if a solution absorbs green light it will appear red to us. Square Planar Crystal Field Splitting: ● The d orbitals of a transition metal ion in a square planar field are split into several energy levels depending on the relative orientations of the metal orbitals and ligand electrons at the four corners of the square. ● The d(x^2-y^2) orbital has the highest energy because its lobes are directed right at the four corners of the square plane.

Tetrahedral Crystal Field Splitting: ● In a tetrahedral complex ion, the d orbitals of the metal ion are split by a tetrahedral crystal field. ● The lobes of the higher-energy orbitals - d(xy), d(yz) and d(xz) - are closer to the ligands at the four corners of the tetrahedron than the lobes of the lower-energy orbitals.

We

use the parameter field strength to

describe the relative magnitude of the split in the energies of the d orbitals in metal ions, ranking ligands in what is called a spectrochemical series. Spectrochemical Series - A list of ligands rank-ordered by their ability to split the energies of the d orbitals of transition metal ions. ● As field strength of the ligand increases from the bottom to top of the table pictured below, the crystal field splitting energy increases. ● High-field-strength ligands form complexes that absorb short-wavelength, highenergy light. ● Low-field-strength ligands absorb long-wavelength, low-energy light. 22.5 - Magnetism and Spin States Crystal field splitting influences the magnetic properties of transition metal ions because these properties depend on the number of unpaired electrons in the valence shell d orbitals. The more unpaired electrons, the more paramagnetic the ion. ● If Deltao is small, then Hund’s rule holds and a complex is called HIGH-SPIN. ○ The spin of all five electrons is in the same direction, resulting in the maximum magnetic field produced by the spins. ○ Metals and their ions with their d electrons evenly distributed across all the d-orbitals in the valence shell represent a high-spin state. ● If Deltao is large, than electrons fill all low energy orbitals first and is called LOW-SPIN. ○ Has only one electron unpaired and has a lower magnetic field. ○ Strong repulsions and large values of crystal field splitting energy can lead to electron pairing in low-energy orbitals and an electron configuration called a lowspin state. ● Somewhere in the middle we switch from low spin to high spin. Remember that high-spin or low-spin depends on whether less energy is needed to promote an electron to a higher-energy orbital or to overcome the repulsion experienced by two electrons sharing the same lower-energy orbital. Several factors affect the size of Deltao. ● Different field strengths of different ligands. ● Oxidation state of the metal ion - higher oxidation state (ionic charge), the stronger the attraction of the electron pairs on the ligands for the ion, and greater attraction leads to

more ligand-d orbital interactions and therefore a larger Deltao. 22.6 - Isomerism in Coordination Compounds Some arrangements of ligands can have geometric isomers for example, if a compound is square planar, there could be two ways to draw it and arrange the ligands about the central metal ion. Coordination compounds like these two that have the same composition and the same connections between parts but differ in the 3D arrangement of those parts are stereoisomers. ● The two members of each pair of ligands can be (a) on the same side of the square, or (b) at opposite corners. ○ (a) CIS = next to ○ (b) TRANS = opposite/across from ● The relative positions of ligands will make a BIG difference. Different arrangements change a lot of the physical properties. Can the three common ligand arrangements have geometric isomers? 1. Square Planar - YES 2. Octahedral - YES 3. Tetrahedral - NO

This image illustrates another kind of stereoisomerism that is possible in complex ions and coordination compounds. The complex ion here is chiral, which means that its mirror image is not superimposable on the original complex. If you rotate the image 180 degrees about its vertical axis to look as much like the original as possible, the top ligand is located behind the plane of the page in the original, but in front of the plane of the page in the rotated mirror image. Thus the mirror images are not superimposable. In other words - there is no way to rotate the mirror image so that its atoms align exactly with those in the original. These stereoisomers are called enantiomers. 22.7 - Coordination Compounds in Biochemistry Many proteins and enzymes contain metal atoms at the active site. An enzyme does chemistry, so the metal is often involved. Chlorophyll a is one of the main light harvesting proteins in plants. It absorbs everything except green. All molecules of chlorophyll contain ring-shaped tetradentate ligands call...