Inorganic Chemistry Final Exam – Flashcards
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Any of a class of compounds whose aqueous solutions react with and dissolve certain metals to form salts and react with bases to form salts.
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Acid
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An acid with one or more molecules of water removed; for example, SO₃ is the acid anhydride of H₂SO₄, sulfuric acid. Nonmetal oxides are ______.
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Acid anhydride / Acidic oxide
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The energy, in excess over the ground state, that must be added to an atomic or molecular system to allow a particular process to take place.
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Activation energy
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An element or other substance that has two or more different forms of structures that are most frequently stable in different temperature ranges, such as different crystalline forms of carbon as charcoal or diamond. Ex: O₂ and O₃
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Allotrope
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Any of a large number of substances having metallic properties and consisting of two or more elements; with few exceptions, the components are usually metallic elements.
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Alloy
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A reagent or substance that can have two or more attacking sites
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Ambident
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Pertaining to a solid that is noncrystalline, having neither definite form nor structure.
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Amorphous
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Having both acidic and basic characteristics.
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Amphiprotic / Amphoteric
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An ion that is negatively charged.
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Anion
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A solvent that does not yield or accept a proton.
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Aprotic solvent
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Any chemical species, ionic or molecular, capable of accepting or receiving a proton (hydrogen ion) from another substance; the other substance acts as an acid in giving of the proton.
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Base
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A metallic oxide that is a base or that forms a hydroxide when combined with water.
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Basic anhydride / Basic oxide
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Improving the chemical or physical properties of an ore so that the metal can be recovered at a profit. Also known as mineral dressing.
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Beneficiation
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The heat of formulation of a molecule from its constituent atoms. [Note: Usually tabulated as positive numbers, which represent the energy required to dissociate (break apart) a gaseous diatomic molecule into its constituent gaseous atoms, or a corresponding fraction for a larger molecule composed of two elements.]
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Bond energy
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A sequence of chemical and physical processes by means of which the cohesive energy of an ionic crystal can be deduced from experimental quantities; it leads from an initial state in which a crystal is at zero pressure and 0°K to a final state that is an infinitely dilute gas of its constituent ions, also at zero pressure and 0°K. ∆H = (atom. NRG) + (ion. NRG) + (atom. NRG) + (EA) - (U lattice NRG)
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Born-Haber cycle
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A ligand in which an atom or molecular species that is able to exist independently is simultaneously bonded to two or more atoms.
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Bridging ligand
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Substance that alters the velocity of a chemical reaction and may be recovered essentially unaltered in form and amount at the end of the reaction.
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Catalyst
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The property of an element to link itself to form molecules, as with carbon.
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Catenation
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One of the elements that form Group 16/VIA of the periodic table; included are oxygen, sulfur, selenium, tellurium, and polonium.
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Chalcogen
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Referring to a crystal structure in which the lattice points are centers of spheres of equal radius arranged so that the volume of the interstices between the spheres is minimal.
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Close packed
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A wire, cable, or other body or medium that is suitable for carrying electric current.
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Conductor
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A chemical bond between two atoms in which a shared pair of electrons forms the bond, the pair having been supplied by one of the two atoms.
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Coordinate covalent bond
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Gradual destruction of a metal or alloy due to chemical processes such as oxidation or the action of a chemical agent.
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Corrosion
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A bond in which each atom of a bound pair contributes one electron to form a pair of electrons.
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Covalent bond
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The effective radius of an atom in a covalent bond.
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Covalent radius
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Any departure from crystal symmetry caused by free surfaces, disorder, impurities, vacancies, and interstitials, dislocations, lattice vibrations, and grain boundaries. Also known as lattice defect.
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Crystal defect
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The absorption of atmospheric water vapor by a crystalline solid until the crystal eventually dissolves into a saturated solution.
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Deliquescence
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A soluble or insoluble chemical substance that has such a great affinity for water that it will abstract water from a great many fluid materials.
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Desiccant
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A chemical reaction in which an atom, radical, or molecule displaces and sets free an element of a compound.
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Displacement reaction
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The changing of a substance, usually by simultaneous oxidation and reduction, into two or more dissimilar substances. 1 atom both oxidizes and reduces into 2 separate atoms.
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Disproportionation
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A method by which chemical reactions are carried out by passage of electric current through a solution of an electrolyte or through a molten salt.
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Electrolysis
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1. The difference in electric potential that exists between two dissimilar electrodes immersed in the same electrolyte or otherwise connected by ionic conductors. 2. The resultant of the relative electrode potential of the two dissimilar electrodes at which electrochemical reactions occur.
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Electromotive force (emf)
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The work needed in removing an electron from a negative ion, thus restoring the neutrality of an atom or molecule. It increases right and up the periodic table. The atom with the lowest ______ is the central atom.
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Electron affinity / Electron gain enthalpy
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Pertaining to an atom or group of atoms that has a relatively great tendency to attract to itself electrons [in a molecule].
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Electronegative
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The sum of the internal energy of a system plus the product of the system's volume multiplied by the pressure exerted on the system by its surroundings. Also known as heat content, sensible heat, and total heat. Spontaneous when exothermic/negative.
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Enthalpy (∆H)
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1. Measure of the disorder of a system, equal to the Boltzman constant times the natural logarithm of the number of microscopic states corresponding to the thermodynamic state of the system. 2. Function of the state of thermodynamic system whose changed in any differential reversible process is equal to the heat absorbed by the system from its surroundings divided by the absolute temperature of the system. Spontaneous when more disordered/positive.
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Entropy (∆ S)
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A constant at a given temperature such that when a reversible chemical reaction cC + dD ⇌ gG + hH has reached an equilibrium, the value of this constant K° is equal to (G)^g (H)^h / (C)^c (D)^d
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Equilibrium constant
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A chemical reaction or change of state that is effected in an exceedingly short space of time with the generation of a high temperature and generally a large quantity of gas.
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Explosion
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A catalytic process to synthesize hydrocarbons and their oxygen derivatives by the controlled reaction of hydrogen and carbon monoxide. n CO + (2n+1) H₂ → CnH₂n+₂ + n H₂O
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Fischer-Tropsch process
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The thermodynamic function G = H - TS, where H is enthalpy, T absolute temperature, and S entropy. Spontaneous when negative.
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Gibbs free energy (∆ G)
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A hard, amorphous, inorganic, usually transparent, brittle substance made by fusing silicates, sometimes borates and phosphates, with certain basic oxides and then rapidly cooling to prevent crystallization.
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Glass
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A family of elements wither similar chemical properties.
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Group
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Any of the elements of Group 17, consisting of flourine, chlorine, bromine, iodine, and astatine.
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Halogen
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Close-packed crystal structure characterized by the regular alteration of two layers; the atoms in each layer lie at the vertices of a series of equilateral triangles, and the atoms in one layer lie directly above the centers of the triangles in neighboring layers. ABA packing
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Hexagonal close packing
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Catalysts occurring within a single phase, usually a gas or liquid
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Homogeneous catalysis
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The increase in enthalpy accompanying the formation of 1 mole of a hydrate from the anhydrous form of the compound and from water at constant pressure.
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Hydration energy / Heat of hydration
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In aqueous solutions of electrolytes, the reactions of cations with water to produce a weak base or of anions to produce a weak acid.
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Hydrolysis
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Mineral synthesis in the presence of heated water.
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Hydrothermal synthesis
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A layer located between layers of different character.
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Intercalation
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A type of chemical bonding in which one or more electrons are transferred completely from one atom to another, thus converting neutral atoms into electrically charged ions; these ions are approximately spherical and attract one another because of their opposite charge.
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Ionic bonding
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Pertaining to atoms having the same number of electrons outside the nucleus of atoms.
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Isoelectronic
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Any two or more crystalline mineral compounds having different chemical compositions but identical structure, such as the garnet series or the feldspar group.
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Isomorphous / Isomorphic minerals
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1. The energy required to separate ions in an ionic crystal an infinite distance from each other. 2. The energy released when ions come together from infinite separation to form an ionic crystal.
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Lattice energy
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A crystalline structure found in substances such as graphites and clays in which the atoms are largely concentrated in a set of parallel planes, with the regions between the planes comparatively vacant.
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Layer structure
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A dimensionless constant that determines the electrostatic energy of a three-dimensional periodic crystal lattice consisting of a large number of positive and negative point charges when the number and magnitude of the charges and the nearest-neighbor distance between them is specified.
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Madelung constant
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The type of chemical bond that is present in all metals, and may be thought of as resulting from a sea of valence electrons that are free to move throughout the metal lattice.
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Metallic bond
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A wave function describing an electron in a molecule.
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Molecular orbital
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Pertaining to an element found in nature in a nongaseous state.
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Native
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An oxidizing chemical charge, where an element's positive valence is increased (electron loss), accompanied by a simultaneous reduction of an associated element (electron gain).
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Oxidation-reduction reaction
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The number of electrons to be added to (or subtracted from) an atom in a combined state to convert it to elemental form.
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Oxidation state/number
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Compound that gives up oxygen easily, removes hydrogen from another compound, or attracts negative electrons.
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Oxidizing agent
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A family of elements with consecutive atomic numbers in the periodic table and with closely related properties.
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Period
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A table of the elements, written in sequence in the order of atomic number and arranged in horizontal rows (periods) and vertical columns (groups) to illustrate the occurrence of similarities in the properties of the elements as a periodic function of the sequence.
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Periodic table
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Covalent bonding in which the greatest overlap between atomic orbitals is along a plane perpendicular to the line joining the nuclei of the two atoms.
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Pi bonding
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A bond in which a pair of electrons is shared in common between two atoms, but the pair is held more closely by one of the atoms.
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Polar covalent bond
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The property of a chemical substance of crystallizing into two or more forms having different structures, such as diamond and graphite.
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Polymorphism
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That area of chemistry concerned with the study of radioactive substances.
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Radiochemistry
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The ratio of the radius of a cation to the radius of an anion [cat/an]; relative ionic radii are pertinent to crystal lattice structure, particularly the determination of coordination number. This can be used to rank compounds in order of melting points.
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Radius ratio
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A chemical system in which reduction and oxidation reactions occur.
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Redox system
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1. A material that adds hydrogen to an element or compound. 2. A material that adds an electron to an element or compound, that is, decreases the positiveness of its valence.
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Reducing agent
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Chemical reaction in which an element gains an electron (has a decrease in positive valence).
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Reduction
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A material of high melting point.
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Refractory
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Any of two or more structures of the same compound that have identical geometry but different arrangements of their paired electrons; none of the structures has physical reality or adequately accounts for the properties of the compound, which exists as an intermediate form.
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Resonance structure
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The act or process of accumulating sediment in layers.
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Sedimentation
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A solid crystalline material whose electrical conductivity is intermediate between that of a conductor and an insulator, ranging from about 10⁵ mho to 10⁻⁷ mho per meter, and is usually strongly temperature-dependent.
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Semiconductor
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The chemical bond resulting from the formation of a molecular orbital by the end-on overlap of atomic orbitals.
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Sigma bond
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Forming a coherent bonded mass by heating metal powders without melting; used mostly in powder metallurgy.
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Sintering
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A nonmetallic product resulting from the interaction of flux and impurities in the smelting and refining of metals.
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Slag
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A reaction in which a solvent reacts with the solute to form a new substance.
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Solvolysis
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A display or plot of intensity of radiation (particles, photons, or acoustic radiation) as a function of mass, momentum, wavelength, frequency, or some related quantity.
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Spectrum
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The interaction between a particle's spin (ms) and its orbital angular momentum (ml).
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Spin-orbit coupling
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An iron-based alloy, malleable under proper conditions, containing up to about 2% carbon.
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Steel
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The process by which solids are transformed directly to the vapor state, or vice versa, without passing through the liquid phase.
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Sublimation
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A mixture of gases prepared as feedstock for a chemical reaction, for example, carbon monoxide and hydrogen to make hydrocarbons or organic chemicals, or hydrogen and nitrogen to make ammonia.
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Synthesis gas
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A positive number that characterizes the combining power of an element for other elements, as measured by the number of bonds to other atoms that one atom of the given elemental forms upon chemical combination; hydrogen is assigned ______ 1, and the ______ is the number of hydrogen atoms, or the equivalent, with which an atom of the given element combines.
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Valence
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An electron that belongs to the outermost shell of an atom. [Note: This definition is not satisfactory for d- or f-block elements.]
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Valence electron
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A mixture of carbon monoxide and methane produced by passing steam through deep beds of incandescent coal; used for industrial heating and as a gas engine fuel.
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Water gas
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A function of the coordinates of the particles of a system and of time that is a solution of the Schrodinger equation and that determines the average result of every conceivable experiment on the system.
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Wave function / Schrodinger wave function
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Physical disintegration and chemical decomposition of earthly and rocky materials on exposure to atmospheric agents, producing an in-place mantle of waste.
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Weathering
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The number of spots around a central atom that can either bond or hold an unshared pair of electrons
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Maximum coordination number
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Group 1 metals
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Alkali metals
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Group 2 metals
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Alkaline earth metals
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An external magnetic field is attracted to an unpaired electron, making the molecule heavier. Degree dependent on ligands.
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Paramagnetism
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An external magnetic field is repelled by paired electrons
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Diamagnetism
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The charge of an atom in a compound. = (# of valence e⁻) - (½ # of bond e⁻) - (# of lone e⁻)
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Formal charge
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Increases left and down the table
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Periodic trend of radii size
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The net positive charge that an electron experiences. Increases right across the table. Z* = Z - ∑S (screening constants from Slater's Rules)
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Effective nuclear charge (Z*)
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Equation used to calculate heat of hydration. deltaH = 60900Z^2 / (r + 50) kJ/mol
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Latimer equation
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M(H₂O)₆ⁿ⁺
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Hydrated ion
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[M(H₂O)₅OH]ⁿ⁺
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Hydroxy cation
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M(OH)₂
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Metal hydroxide
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MO₂ⁿ⁻
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Oxoanion
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MO₂ⁿ⁺
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Oxocation
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The difference in energy levels of the d-orbitals creates color when electrons in the lower levels use energy in light to advance to a higher level. Levels are separated by 10Dq.
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Crystal Field Theory
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dz² dx²-y² eg set dyz dxz dxy t₂g set
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Octahedral field
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dxz dyz dxy t₂ set dz² dx²-y² eg set
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Tetrahedral field
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Determines the acidity of a metal cation: 0.00 - 0.01 = nonacidic 0.02 - 0.04 = feebly acidic 0.05 - 0.10 = weakly acidic 0.11 - 0.15 = moderately acidic 0.16 - 0.22 = strongly acidic
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Z²/r
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Determines the basicity of an oxoanion Excess oxo groups = nonbasic Complete cancellation = feebly basic -½ ≤ x ≤ -1 charge = moderately basic x ≥ -1 charge = strongly basic
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2:1 cancellation
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Groups 7, 8, 17, 18 All others per- -ate ox # = group # - -ate ox # = group # - 2 ox # = group # -ite ox # = group # - 4 ox # = group # - 2 hypo- -ite ox # = group # - 6 ox # = group # - 4
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Oxoanion nomenclature
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3 oxo groups = very strong 2 oxo groups = strong 1 oxo group = moderate 0 oxo groups = weakly -ate → -ic -ite → -ous
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Oxo acids, strength and nomenclature
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A solid with a well-defined form/shape/structure
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Crystalline solid
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Regular, repeating 3-D array of atoms, molecules, or ions
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Crystalline lattice
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The smallest part of a lattice from which the entire crystal can be repeated by repeating the ______ in 3-D
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Unit cell
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cubic tetragonal orthorhombic monoclinic hexagonal rhomohedral triclinic
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7 Bravais crystal systems
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[volume (unit cell) - volume (atoms) / volume (unit cell)] x 100%
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% empty space
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Acidic + nonbasic Nonacidic + basic
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Soluble salts
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Acidic + basic Nonacidic + nonbasic
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Insoluble salts
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Equation used to determine energy of attraction
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Born-Lande Equation
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Equation used to determine the lattice energy
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Kapustinski Equation
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The most efficient closest packing; least empty space. Occurs in face-centered cubes. CBA packing
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Cubic closest packing
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vacancies of cations = vacancies of anions
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Schottky Defect
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ions shift to interstitial vacancy
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Frenkel Defect
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The ability of a solvent to be the strongest acid or base in solution. Example: in H₂O, the strongest acid is H₃O⁺ and the strongest base is ⁻OH.
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Leveling effect
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CaO + H₂O ⇌ Ca²⁺ + 2 ⁻OH
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Slaking process
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A substance that pulls H₂O out of whatever it contacts
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Drying agent
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Electron pair acceptor
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Lewis acid
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Proton donor
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Bronsted-Lowry acid
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Oxide ion acceptor
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Lux-Flood acid
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Electron pair donor
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Lewis base
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Proton acceptor
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Bronsted-Lowry base
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Oxide ion donor
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Lux-Flood base
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Hydrogen ion producer
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Arrhenius acid
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Hydroxide ion producer
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Arrhenius base
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1 atom oxidizes and 1 atom reduces to form one atom.
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Comproportionation
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NH₃ + ⁵/₄ O₂ → NO + ³/₂ H₂O NO + ½ O₂ → NO₂ NO₂ + H₂O → HNO₃
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Ostwald Process
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A solution that dissolves metals that HCl cannot 1 HNO₃ : 3 HCl
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Aqua regia (Kingly water)
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V₂O₅ H₂SO₄ H₂O ↓ ↓ ↓ SO₂ + ½ O₂ → SO₃ → H₂S₂O₇ → H₂SO₄
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Production of sulfuric acid
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Series of reactions used to discover the presence of arsenic. Zn + 2H⁺ → Zn²⁺ + H₂ H₃AsO₄ + H₂ → AsH₃ + 4 H₂O 2 AsH₃ → 2 As(s) + 3 H₂
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Marsh Test
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The mixture that forms concrete when ground with a gypsum.
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Clinker
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Ni(impure) + 4 CO → Ni(CO)₄ → Ni(pure) + 4 CO
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Mond Process
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Compounds that look like oxo salts but are actually oxides formed from multiple metal cations.
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Mixed-metal oxides
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Mixed-metal oxides that are in cubic closest packing of the formula AB₂O₄.
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Spinels
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A is in the tetrahedral holes. B is in the octahedral holes. A[B]₂O₄
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Normal spinel
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½ B is in the tetrahedral holes. A and the other ½ B are in the octahedral holes. B[AB]O₄
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Inverse spinel
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Used in determining the energy of electrons in the tetrahedral vs octahedral holes when determining whether a spinel is normal or inverse. octahedral field = (# of e⁻ in t₂g set)(-4Dq) + (# of e⁻ in eg set)(6Dq) tetrahedral field = (# of e⁻ in e set)(-6Dq) + (# of e⁻ in t₂ set)(4Dq)
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Crystal field stabilization energy (CFSE)
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Used in determining which metal ion in a spinel is more stable in the octahedral holes. The most negative value is more stable in the octahedral holes. = CFSE(octa) - CFSE(tetra)
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Octahedral site stabilization energy (OSSE)
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The basis for determining the standard reduction potentials of molecules. E = 0.00V A container with a 1M solution of H⁺ ions and a electrode of platinum black hanging from platinum wire with H₂ gas pumped into the solution.
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Standard Hydrogen Electrode (SHE)
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A cell with one electrode the SHE and the other a metal, connected by a salt bridge. Used to determine the standard reduction potential of that metal.
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Daniell Cell
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The connector in cells that keeps electrical neutrality by sending negative ions opposite the movement of electrons.
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Salt bridge
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∆G° = -nFE° n → moles of e⁻ F → Faraday's constant E° → standard reduction potential
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Gibbs free energy of a cell
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The ability of a molecule to act as an oxidizing agent / be reduced. A large positive reactant = good O.A. A large negative product = good R.A.
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Standard reduction potential (SRP)
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Equation used to determine the cell potential under non-standard conditions. E = E° - RTlnQ / nF Q → [products]^coeff / [reactants]^coeff *DO NOT include solids
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Nernst Equation
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Increases right across the table. Most stable (most negative) in the 3rd period.
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Periodic trends in SRP
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A series of oxides, ions, etc. of a metal that gives the difference in SRP between each subsequent molecule. Can be used to determine the SRP of two molecules that are non-adjacent by adding each SRP multiplied by the number of electrons transferred and dividing the sum by the total # of electrons transferred.
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Latimer Diagrams
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A diagram of one metal's oxides, ions, etc. in a chart of SRP(y) vs pH(x). The molecules at the top are good O.A. The molecules within the dotted lines are stable in aqueous solution. The molecules at the bottom are good R.A. E° = E at pH of 0
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Pourbaix Diagrams
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Metals with low Xp (<1.4) and E° ≤ -1.6. These include Groups 1, 2 and f-block of Group 3. They are good R.A. and cannot exist in aqueous solution.
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Very electropositive metals
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Metals with 1.4 ≤ Xp ≤ 1.9 and -1.6 < E° < 0. These include the d-block of Period 4 and p-block of Periods 4, 5. They do not react with H₂O but do react with acids.
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Electropositive metals
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Metals with 1.9 < Xp 0. These include the d-block of Periods 5,6 and p-block of Period 6. They are O.A. and are not oxidized by acids.
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Electronegative metals
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Cations of electronegative metals oxidize (very)electropositive metals to produce cations of more active metals.
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Electronegativity displacement reaction
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Ore of aluminum. Al₂O₃∙2H₂O
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Bauxite
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Used in production of molten aluminum. Add cryolite(Na₃AlF₆) to corundum(Al₂O₃) to reduce its melting point. Electrons travel from a graphite anode to a graphite cathode that lines the container holding the molten aluminum solution. The electrons reduce the Al³⁺ to molten aluminum.
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Hall-Heroult Process
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Used in sodium production. (Look up chart in book)
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Downs Cell
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Impure iron mixture with a lot of carbon
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Pig iron
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Used to remove P and S impurities from molten Fe. 4 P + 5 O₂ → 2 P₂O₅ S + O₂ → SO₂ P₂O₅ + 3 CaO → Ca₃(PO₄)₂ (part of slag)
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Basic oxygen furnace (BOF)
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Magnetized metal (steel) containing Al, Ni, Co, Fe, and Cu
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Alnico V
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Heat-resistant steel containing Ni, Cr, Fe, and Cu
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Inconel
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Very strong form of iron (in Pourbaix diagram)
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Pearlite
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Cathode = 2 H₂O + 2 e⁻ → H₂ + 2 ⁻OH Anode = Fe → Fe²⁺ + 2 e⁻ Fe²⁺ → Fe(OH)₂
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Iron corrosion with excess H₂O, little O₂
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Cathode = O₂ + 2 H₂O + 4 e⁻ → 4 ⁻OH Anode = Fe → Fe²⁺ + 2 e⁻ 4 Fe(OH)₂ + O₂ → 2 Fe₂O₃∙H₂O (red-brown rust) + 2 H₂O
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Iron corrosion with excess H₂O and O₂
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6 Fe(OH)₂ + O₂ → 2 Fe₃O₄∙H₂O (green hydrated magnetite) + 4 H₂O Fe₃O₄∙H₂O → H₂O + Fe₃O₄ (black magnetite)
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Iron corrosion with little H₂O and O₂
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Protect underlying surface from corrosion by adding a zinc coating
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Galvanized
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Adding a metal that corrodes easier than iron to prevent corrosion
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Sacrificial electrode
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MgCl₂ + Ca(OH)₂ (sea shells) → Mg(OH)₂ + CaCl Mg(OH)₂ → MgCl₂ + 2 H₂O MgCl₂ → Mg + Cl₂ ↑ e⁻
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Magnesium production
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Compounds similar halides that are good R.A. Examples: CN⁻, SCN⁻, S₂O₃²⁻, NCO⁻, N₃⁻, H⁻
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Pseudohalides
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O-O, N-N, O-Cl, O-N, N-Cl, N-Br, N-I
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Unstable, explosive bond
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A reaction where a salt produces a different salt. Ex: Cu₂S → Cu₂O
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Chemical conversion
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Beneficiation and chemical conversion is performed at the same time. Ex: 2 Cu₂S + SO₂ + 4 H⁺ → 4 Cu²⁺ + 2 SO₄²⁻ + 2 H₂O
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Hydrometallurgy
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Nonmetals with Xp > 2.8 and high E°. These include F₂, Cl₂, Br₂, O₂, and N₂.
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Very electronegative nonmetals
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Nonmetals with 1.9 < Xp < 2.8.
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Electronegative nonmetals
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Nonmetal anions with the lowest possible oxidation number. These include S²⁻, O²⁻, P, Cl.
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Monoatomic nonmetal anions
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Enthalpy necessary to convert liquid to vapor at constant T and P.
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Heat of vaporization
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Enthalpy necessary to convert element in usual form at given T and P into individual gaseous atoms.
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Heat of atomization
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Highest in middle of period (more e⁻ needed to accept/donate). Highest at top of p-block (more allotropes and e⁻ are closer to nucleus/effective nuclear charge). Highest at bottom of d-block (delta bonding).
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Periodic trends of atomization energy
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Van der Waals > anionic > metallic > covalent > cationic
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Ranking of radii size
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Electronegativity that shows the relationship between electron affinity and ionization energy (average of the two). Xm = EA + IE / 2
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Mullike electronegativity
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A liquid that behaves with zero viscosity, finding the lowest possible spot.
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Superfluid
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Generally an anion or neutral molecule that contains at least one pair of electrons to be donated to the metal center. Donates both electrons to form the bond.
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Ligand
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Compound that contains 1 or more complex ion
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Coordination compound
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The atom(s) of the ligand that possesses the pair of electrons that are donated to the metals to form a bond.
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Donor atom
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The metal and the ligands directly bonded to it.
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Inner coordination sphere
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4 and 6, determined by the ligands. Larger ligands and those that transfer substantial negative charge to the metal favor lower coordination numbers.
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Most common coordination number of coordination compounds
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The sum of charges on the metal and ligands.
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Complex charge
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The sum of charges on the metal, ligands, and counterbalancing ions.
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Neutral charge of coordination compound
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Monodentate = 1 donor atom Bidentate = 2 Tetradentate = 4 Hexadentate = 6 Polydentate = 2 or more
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Classification of ligands
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Ligands that bond to one metal ion in multiple places.
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Chelating ligands
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H₂O, CN⁻, NH₃, NO₂⁻, SCN⁻, OH⁻, X⁻, CO, O²⁻
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Monodentate ligand examples
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Ligands that may bind through one of two donor atoms, but not both at the same time. Ex: thiocyanato-S or thiocyanato-N-, nitrito-N or nitrito-O
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Ambidentate ligands
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C# 2 = linear C# 4 = tetrahedral (most common) square planar (metals with d⁸ e⁻) C# 6 = octahedral
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Common geometries of complexes
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Cation named before anion. Ligands named before metal ion in alphabetical order. Anion endings: -ide → -o -ite → -ito -ate → -ato Neutral ligands stay same. Multiple of same ligand prefixes: di-, tri-, tetra-, penta-, hexa- 4 or more syllables prefixes (ligand in parantheses): bis-, tris-, tetrakis-, pentakis-, hexakis- Metal anion has -ate ending and uses Latin name.
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Coordination compound nomenclature
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Compounds that have the same composition but a different arrangement of atoms.
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Isomers
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Isomers that have different bonds
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Structural isomers
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Isomers that differ in a ligand bonded to the metal in the complex, as opposed to being outside the coordination-sphere. Ex: [Co(NH₃)₅Cl]Br vs [Co(NH₃)₅Br]Cl
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Coordination-sphere isomers
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Isomers that differ in the atom of a ligand bonded to the metal in the complex. Ex: ambidentate ligands
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Linkage isomers
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Isomers that have the same bonds, but different spatial arrangements.
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Stereoisomers
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Isomers that differ in the spatial arrangements of the ligands. They have different chemical/physical properties. Ex: cis vs trans isomers
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Geometric isomers
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Isomers that are nonsuperimposable mirror images (chiral = no plane of symmetry). They have similar chemical/physical properties. Differ in interactions with plane polarized light.
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Optical isomers / Enantiomers
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-∆H = EaEb + CaCb
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Drago-Wayland Equation
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The susceptibility to undergo electrostatic interaction (ionic, dipole-dipole), relative to iodine, E = 1
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E parameter of DW equation
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The susceptibility to form covalent bonds, relative to iodine, C = 1
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C parameter of DW equation
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Nonaqueous solvent for nonpolar molecules. Example: NH₃
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Basic protic solvent
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Example: H₂SO₄
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Acidic protic solvent
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Example: SbF₅
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Superacidic protic solvent
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Nonaqueous solvents that don't interact much. Examples: CCl₄, C₆H₁₂, benzene, tholuene
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Nonpolar, nonionizing, weakly solvating aprotic solvent
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Examples: CH₃CN, DMSO, SO₂
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Highly polar, nonionizing, strongly solvating aprotic solvent
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Example: BrF₃
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Highly polar, autoionizing aprotic solvent
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Ranking of the stability of metal cations right across the periodic table
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Irving-Williams Series
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Alkali, alkaline earth, and lighter, highly charged metal ions. Ex: Ti⁴⁺, Fe³⁺, Co³⁺, Al³⁺
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Chatt, Arhland, Davis Class (a) metals
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Heavier transition metals in low oxidation states. Ex: Cu⁺, Pd²⁺, Ag⁺, Pl²⁺, Au⁺, Hg²⁺, Tl⁺, Pb²⁺
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Chatt, Arhland, Davis Class (b) metals
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Class (a) metals. They tend to prefer F>Cl>Br>I O>>S>Se>Te N>>P>As>Sb
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Pearson "Hard" metals
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Class (b) metals. They tend to prefer F<Cl<Br<I O<Se≈Te N<
As>Sb
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Pearson "Soft" metals
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0.7 < Xp < 1.6 small ionic radii high ionic charge
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Hard acids
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1.9 < Xp < 2.5 large ionic radii low ionic charge
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Soft acids
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2.1 < Xp < 3.0 large ionic radii
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Soft bases
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3.4 < Xp < 4.0 small ionic radii
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Hard bases
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∆H [E-X] = ½ [∆H(E-E) + ∆H(X-X) + 96.5(Xp(X) - Xp(E))²]
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Enthalpy with Xp
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∆H = 96.5 ( [ (Xp(HA) - Xp(SB))² + (Xp(SA) - Xp(HB))² ] - [ (Xp(HA) - Xp(HB))² + (Xp(SA) - Xp(SB))² ] )
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Enthalpy with HSAB
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Once there's a ligand already attached to a metal, the metal will prefer additional matching ligands. Metals with hard ligands prefer more hard ligands and vice versa.
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Principle of Symbiosis
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When a metal ion receives another ligand, the Xp should increase, resisting hard bases. However, some Xp increase when a hard base is added. Pearson's concept of HSAB wins because the metal prefers another hard base when it already has hard bases attached to it.
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Pauling-Pearson Paradox
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The process of using HSAB to precipitate individual metal ions out of a solution at different times so that the metal ions can be separated from eachother.
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Qualscheme
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Metals that are hard acids. Oxides, silicates, sulfates, carbonates
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Lithophiles
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Metals that are soft acids. Sulfides, tellurides, arsenides
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Chalcophiles
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Examples: N₂, noble gases
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Atmophiles
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Metals that occur in elemental form. Au, Pt, Ir, Co, Fe, Mn, Mo, Ni, Os, Pd, Rh, Re, Ru
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Siderophiles
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Elements important for life. All within the first 35 elements except Mo and I
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Essential elements
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MLCT = Metal to ligand charge transfer LMCT = Ligand to metal charge transfer
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Charge-transfer transition
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Halides formed with highly electronegative elements and high oxidation state metals
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Molecular halides
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Molecular, T-metal, ionic
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Classes of halides
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Ionic, covalent, metallic, diamond-like [BN, (SN)x]
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Classes of nitrides
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S²⁻ and S₄²⁻
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Anions of sulfides
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(total # e⁻ of metal) + (# of e⁻ donated by ligands) = 36, 54, or 86 → extra stable (filled subshells, similar to noble gases)
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Effective Atomic Number (EAN) Rule
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(# of valence e⁻ of metal) + (# of e⁻ donated) = 18 → extra stable
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18 Electron Rule
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# of carbons equidistant from metal center
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η = hapticity
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A reaction that increases the oxidation state and coordination number of a metal center.
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Oxidation addition reaction
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A carbonyl molecule donates e⁻ through sigma bonds but accepts e⁻ from pi bonds
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Carbonyl stretching
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Hydrogen ions occupy holes in metal lattices.
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Metallic hydrides
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Compounds that do not obey stoichiometry. Ex: TiH₁.₇
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Interstitional compounds
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Hydrides formed with two different metals. Ex: FeTiHx and NiMH
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Mixed-metal hydrides
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Hydrides formed from Groups 1, 2 metals (except Be). They have the NaCl lattice and react with protic solvents. Ex: NaH, LiH
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Saline hydrides
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Hydrides formed with p-block elements (except Al, Bi). Ex: BH₃, GaH₃, NH₃, CH₄, boranes
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Molecular hydrides
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Carbides formed with C₂²⁻. Ex: CaC₂, [CuC₂, Ag₂C₂, AuC₂ are explosive]
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Dicarbides (acetylides)
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Interstitional carbides formed with strong M-C bonds. Ex: Fe₃C
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Metallic carbides
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Carbides with diamond-like structures. Ex: SiC
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Polymeric carbides
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Carbides formed with C⁴⁻
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Methides
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E = -2π²m(Z*)²e⁴ / n²h² m → mass of e⁻ Z* → effective nuclear charge e → e⁻ charge n → principle quantum number h → Plank's constant
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Energy of an electron equation
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1.) Group e⁻ together as (1s)(2s, 2p)(3s, 3p)(3d)(4s, 4p)(4d)(4f)... 2.) For e⁻ in s- and p-block, value of "s" = ∑ of following: a.) Nothing from e⁻ in groups to right of e⁻ under consideration b.) 0.35 from each e⁻ in group under consideration (0.30 for 1s) c.) 0.85 from each e⁻ with quantum number of n-1 d.) 1.00 from all remaining e⁻ 3.) For e⁻ in d- and f-block, for same rules as 2, excluding 2.c)
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Slater's Rules
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r = a₀n² / (Z*)² a₀ = 4πEh² / me²
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Radius equation with Z*
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Electrons in the 1s and other lower orbitals must move much faster than electrons in higher subshells in order to stay in orbit. As the speed increases, so does the mass, contracting the radius. This does not affect electrons in farther subshells, causing the radii to remain steady in later periods.
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Relativistic contraction
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Numbers of protons and neutrons where the element is stable. The element is magic if it has one of the numbers, double magic if it has both. All the numbers are even.
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Magic numbers
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2, 8, 20, 28, 50, 82, 114
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Magic numbers of protons
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2, 8, 20, 28, 50, 82, 126, 184
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Magic numbers of neutrons
