D-Block and Transition Elements

D-Block Elements

• These elements are located in the middle of the periodic table.
• They are positioned between groups 2 and 13.
• The top row of the d-block elements is particularly important.

Transition Elements

• A transition element is defined as a d-block element that can form at least one stable ion with a partially filled d subshell.
• D subshells can hold up to 10 electrons.
• For period 4 D block elements, only 8 are transition elements.
• Scandium and zinc are not transition elements because they do not form stable ions with partially filled d subshells.

Electron Configurations of Transition Elements

• After argon, electrons fill the 3d and 4s orbitals.
• When writing electron configurations using argon as a shorthand, the configuration represents 1s2 2s2 2p6 3s2 3p6.
• Electrons fill singly first in orbitals due to repulsion.
• The configuration for titanium is two electrons in the 3d orbital and two in the 4s.
• Chromium has five electrons in the 3d and only one in the 4s orbital because one electron from the 4s orbital jumps down to fill up the 3d orbitals, achieving a more stable half-filled subshell configuration.
• Manganese has five electrons in the 3d and two in the 4s orbital.
• Iron starts pairing up electrons in the 3d orbitals.
• Copper behaves similarly to chromium; an electron from the 4s orbital moves into the 3d orbital, creating a full 3d subshell.

Scandium and Zinc

• Scandium forms only one stable ion, Sc3+, which has an empty d subshell, thus not a transition element.
• Zinc forms only one stable ion, Zn2+, which has a full d subshell, and is therefore not a transition element.
• When transition metals lose electrons to form ions, they lose electrons from the 4s orbital first, then from the 3d.

Properties of Transition Metals

• Transition metals exhibit variable oxidation states, form colored ions in solution, and are good catalysts.
• Variable oxidation states are due to the close energy levels of the 4s and 3d orbitals, allowing electrons to transition readily between them.

Common Ions of Transition Metals

• Vanadium(II) or V2+ forms violet-colored ions.
• Vanadium(III) or V3+ forms green-colored ions.
• Vanadium oxide (VO2) is blue, and vanadium dioxide (VO2) is yellow
• Chromium(III) or Cr3+ is green when surrounded by six water molecules and normally substituted, so that it looks green.
• Dichromate is orange.
• Manganese(II) or Mn2+ is pale pink.
• Manganate(VI) or MnO42- is green.
• Permanganate(VII) or MnO4- is purple.
• Iron(II) or Fe2+ is pale green.
• Iron(III) or Fe3+ is yellow-brown.
• Cobalt(II) or Co2+ is pink.
• Nickel(II) or Ni2+ is green.
• Copper(II) or Cu2+ is blue.
• Titanium(III) or Ti3+ is purple, and titanium(II) or Ti2+ is violet.

Transition Metals as Catalysts

• Transition metals' ability to have variable oxidation states makes them good catalysts.
• They can absorb and desorb substances on their surfaces, which lowers the activation energy of reactions.
• Using transition metals as catalysts can lower the temperature required for a reaction, which is beneficial for industrial processes and the environment.
• Copper exposure can damage the liver, indicated by a ring of copper in the eye.
• Long-term exposure to manganese can cause psychiatric issues and physical tremors.

Examples of Reactions Using Transition Metals as Catalysts

• MnO2 catalyzes the decomposition of hydrogen peroxide into water and oxygen; manganese has an oxidation state of +4.
• Iron catalyzes the Haber process for making ammonia; iron has an oxidation state of 0.
• Copper sulfate catalyzes which catalyzes zinc and acid reactions and zinc has an oxidation state of +2

Complex Ion Solution Colors

• Transition metal ions in water form complexes with water molecules, resulting in various colors.
• Transition metal ions react with water, forming complexes where six water molecules surround the metal ion, simplified as [M(H2O)6]n+.
• The color changes are expected in various chemical reactions.

Transition Metal Ions in Solution

• Copper (Cu2+) forms six ligands around it in water, written as Cu(H2O)62+.
• Copper solutions are blue.
• Iron(II) or Fe2+ is pale green.
• Iron(III) or Fe3+ is yellow tending to brown, similar to rust.
• Manganese(II) or Mn2+ is pale pink.
• Chromium or Cr is green.

Reactions with Hydroxide (OH-)

• Adding a small amount of hydroxide to copper results in a partial reaction, forming Cu(OH)2(H2O)4, a pale blue precipitate.
• Fe(OH)2(H2O)4 is a dirty green precipitate formed by adding hydroxide to iron(II).
• Fe 3 forms 3 o h and then h 2 o 3 orange precipitate
• Manganese forms pink precipitate
• Greeny grey precipitate forms chromium
• 2+ ions react with hydroxide to form neutral compounds with the formula OH2(H2O)4.
• These neutral complexes precipitate out of the solution.
• 3+ ions react to form three OH ligands instead of two, resulting in a neutral complex and precipitation.
• Copper, iron(II), iron(III), and manganese do not change with excess sodium hydroxide, remaining insoluble.
• Chromium reacts with excess hydroxide to form Cr(OH)6, a dark green solution, as all water ligands disappear
• This is a charged substance and forms a solution.

Reactions with Excess Ammonia (NH3)

• With copper, a dark blue solution results from partial ligand substitution, forming Cu(NH3)4(H2O)2 with ammonia and water ligands
• This is charged and in solution so there's no precipitate
• Iron(II), iron(III), and manganese do not change, remaining insoluble in excess ammonia.
• Chromium undergoes full ligand substitution, forming Cr(NH3)6, a purple solution.
• This is because all ligands, OH3 and H2O, disappear and 6 ammonia ligands surround it

Equations for Reactions

• Copper reacts with hydroxide: Cu(H2O)62+ + 2OH- → Cu(OH)2(H2O)4 + 2H2O.
• Copper reacts with ammonia: Cu(H2O)62+ + 2NH3 → Cu(OH)2(H2O)4 + 2NH4+.
• Adding excess ammonia to copper forms Cu(NH3)4(H2O)22+.
• Iron(II) reactions are similar to copper(II), with all being insoluble.
• Iron(III) reacts with 3OH to form a neutral substance.
• Manganese reactions are similar to iron and copper, with all being insoluble.
• Chromium reacts similarly to iron(III), but with excess hydroxide, Cr(OH)63- forms.
• With excess ammonia, Cr(NH3)6 forms

Complex Ions

• Complex ions feature a central metal ion surrounded by ligands, bonded by dative covalent or coordinate bonds.
• A ligand has a lone pair of electrons donated into the metal ion.
• The overall charge of the complex sits outside the square brackets.
• Complex ions can have different shapes.

Ligands

• A ligand is an ion, atom, or molecule with at least one lone pair of electrons.
• Ligands can be monodentate, bidentate, or polydentate.
• Monodentate ligands have one lone pair: water, ammonia, cyanide, and chloride.
• Bidentate ligands have two pairs of electrons, like ethanedioate and ethane-1,2-diamine, so you need to remember the structures of these ligands
• Polydentate ligands have multiple lone pairs, such as EDTA4- with six lone pairs forming six coordinate bonds.

Shapes of Complexes

• The shape depends on ligand size and coordination number.
• The coordination number is the number of coordinate bonds, not the number of ligands.
• Small ligands, like water and ammonia, can fit six around a central metal ion, forming an octahedral shape.
• Larger ligands, like chloride, can only fit four around the metal ion.
• The size impacts the number of ligands and the complex type.
• Complexes with six coordinate bonds form an octahedral shape, with 90-degree bond angles.
• Complexes with four coordinate bonds form tetrahedral or square planar shapes.

Tetrahedral and Square Planar Complexes

• Chlorine is a bigger ligand so we can only fit four of them around a central metal line
• Copper with four chlorine ligands has a tetrahedral shape (CuCl42-), with 109.5-degree bond angles.
• Most square planar shapes are default with just a handful needing to be known.
• An important example is cisplatin, an anti-cancer drug, with two ammonia and two chlorine ligands, forming a square planar complex.
• Square planar bond angles are 90 degrees.

Overall Charge of Complexes

• The overall charge is the total oxidation state of the metal minus the total oxidation state of the ligands.
• For CuCl42-, each Cl- ligand is -1, totaling -4.
• Thus, the oxidation state of copper is +2.

Optical Isomerism

• Complex ions show optical isomerism, a type of stereoisomerism.
• Optical isomers are non-superimposable mirror images and rotate plane-polarized light.
• Octahedral complexes with three bidentate ligands show optical isomers.

Cis-Trans Isomerism

• Complex ions also demonstrate cis-trans isomerism, a type of stereoisomerism.
• Octahedral complexes can have 4 ligands of the same type and 2 of a different type
• So octahedral complexes with four identical ligands and two different ligands display cis-trans isomerism.
• If the two different ligands are opposite each other, it's a trans isomer.
• If the two different ligands are adjacent, it is a cis isomer.
• Square planar complexes with two identical and two different ligands display cis-trans isomerism.

Cisplatin

• Cisplatin is an anti-cancer drug with a square planar complex containing a platinum metal center.
• Has two ammonia ligands and two chloride ligands
• It binds to DNA in cancer cells, preventing reproduction and causing cell death.
• The chloride ions in the complex are easily displaced and bond to nitrogen atoms in cancerous cell DNA
• It also affects healthy cells, suppressing the immune system, increasing infection risk, and causing side effects like hair loss.
• However, the benefits of destroying cancer cells outweigh the risks from the side effects

Color Change and Ligand Exchange

• A color change occurs when ligands in a complex exchange or substitute with each other
• Ligand substitution reactions occur when ligands of a similar size are exchanged with each other.
• In a cobalt reaction with excess ammonia, there is a change from pink to straw colored.

Ligands of Different Sizes

• Replacing it with ligands of a different size can change the shape and the coordination number
• Chloride is a larger ligand than water, so only four chloride ligands can fit around the central metal ion.
• This forces a different shape.
• Cobalt reacting with chloride ions in tetrachloro complex changes from pink (octahedral) to blue (tetrahedral).
• Copper reacting with chloride ions also changes from blue (octahedral) to yellow (tetrahedral).

Factors Affecting Color of Complex

• color is affected by the change in oxidation state
• color is affected by the coordination number
• color is affected by the change of ligand
• A change in ligand or shape can affect the size of Delta e (energy change).

Hemoglobin

• Heme is a large, multi-dentate ligand in hemoglobin.
• Hemoglobin is a protein that transports oxygen in the body.
• The structure is octahedral.
• Four nitrogens come from the heme ligand, with one coordinate bond from a globin protein, and the final from oxygen or water.
• In the lungs, oxygen substitutes the water ligand to form oxyhemoglobin.
• Oxyhemoglobin gives up oxygen to needed places and water takes its place, returning to the lungs for re-oxygenation.

Carbon Monoxide Poisoning

• Carbon monoxide replaces the water ligand in hemoglobin, bonding strongly.
• Carbon monoxide is not easily replaced by oxygen or water.
• A lower oxygen level results.
• This leads to oxygen starvation, causing unconsciousness and death.

Metal Aqua Complexes

• Metal aqua complexes form when a transition metal compound is added to water.
• Generally, six water molecules form coordinate bonds with the metal ions.
• The lone pair on oxygen allows bonds to form.

Redox Reactions

• Transition metals participate in redox reactions with variable oxidation states, gaining or losing electrons.
• Complete equations from half equations can be derived
• Balance half equations
• Write you fill ionic question
• Comment on what has been oxidized and/or reduced and how
• Comment on the color changes.

Reactions involving Iron(II) and Iron(III) are interconverted

• Use manganate half equations to write an ionic equation
• Manganese ions have been reduced, while the iron ions have been oxidized.
• The color change has been from green of Fe2+ to yellow of Fe3+.

Reactions involving Iodide Ions

• Use iodine half equations to write an ionic equation
• The iron ions are reduced, while the iodine ions are oxidized.
• Results in a color change from yellow to green.

Interconversion between Chromium(III) and Dichromate

• Use hydrogen peroxide in equations.
• Oxygen has been reduced, while the chromium has been oxidized.
• Results in a color change from dark green of Cr3+ to chromate yellow of CrO42-.
• Addition of dilute sulfuric acid to chromate produces orange dichromate solution (Cr2O7).

Dichromate can be converted to Chromium(III) using Zinc

• Chromium is reduced, while zinc is oxidized.
• This causes a color change from orange to green, under acidic conditions (h2SO4).

Reactions with Copper

• Copper(II) is reduced while iodide ions are oxidized.
• Results in a color change from pale blue of Cu2+ to off-white precipitate of copper iodide (CuI).
• Copper(I) can be oxidized or reduced
• Copper(I) ions can simultaneously be reduced and oxidized in disproportionation reaction.
• Reaction with Cu+ is unstable.
• Products are Cu2+-blue precipitate
• Products are Cu Solid

Test Tube Reactions with Sodium Hydroxide

• Transition metals are identified. Identify reaction is dependent on the color of precipitate when sodium hydroxide is added
• Copper(II): Pale blue precipitate.
• Iron(II): Green precipitate.
• Iron(III): Orange/yellow precipitate.
• Manganese(II): Pink precipitate.
• Chromium(III): Dirty green precipitate.

Tests for Halide Ions with Silver Nitrate

• Halide ions can be checked for using silver, and confirmed with ammonia solution
• Add nitric acid then add silver nitrate solution, the precipitate will then tell us what halide ion is in the test tube from the precipitate color.
• Chloride: White precipitate (silver chloride). Dissolves in dilute ammonia and it has a reaction forming silver ions reacts with chloride
• Bromide: Cream precipitate (silver bromide). Soluble in concentrated ammonia only.
• Iodide: Yellow precipitate (silver iodide). Insoluble in both dilute and concentrated ammonia.
• NOTE Fluid ions do not form a precipitate silver fluoride forms instead and it is soluble

Tests for Carbonates and Sulfates

• Hydrochloric acid and Barium chloride are used
• Carbonates in solution test is carbonate adding an acid to the acid to form/producing carbon dioxide and you'll see fizzing
• Bubble carbon dioxide through lime water it turns the water cloudy
• Sulfates tests need you need to add an acid to remove any other impurities
• Sulfates also need barium chloride added too resulting in a white precipitate that forms from precipitating barium sulfate

Tests for Ammonium Compounds and Hydroxides

• Add sodium hydroxide with gentle heat to ammonium compounds.
• Ammonia gas is produced, recognized by pungent smell, turns damp red litmus paper blue

If a litmus paper is turned blue from a red litmus dump then it means that the substance given has alkaline presence.