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Inorganic Chemistry Final Exam

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Acid
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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 anhydride / Acidic oxide
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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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Activation energy
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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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Allotrope
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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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Alloy
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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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Ambident
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A reagent or substance that can have two or more attacking sites
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Amorphous
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Pertaining to a solid that is noncrystalline, having neither definite form nor structure.
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Amphiprotic / Amphoteric
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Having both acidic and basic characteristics.
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Anion
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An ion that is negatively charged.
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Aprotic solvent
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A solvent that does not yield or accept a proton.
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Base
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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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Basic anhydride / Basic oxide
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A metallic oxide that is a base or that forms a hydroxide when combined with water.
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Beneficiation
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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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Bond energy
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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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Born-Haber cycle
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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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Bridging ligand
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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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Catalyst
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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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Catenation
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The property of an element to link itself to form molecules, as with carbon.
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Chalcogen
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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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Close packed
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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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Conductor
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A wire, cable, or other body or medium that is suitable for carrying electric current.
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Coordinate covalent bond
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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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Corrosion
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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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Covalent bond
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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 radius
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The effective radius of an atom in a covalent bond.
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Crystal defect
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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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Deliquescence
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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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Desiccant
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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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Displacement reaction
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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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Disproportionation
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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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Electrolysis
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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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Electromotive force (emf)
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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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Electron affinity / Electron gain enthalpy
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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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Electronegative
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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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Enthalpy (∆H)
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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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Entropy (∆ S)
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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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Equilibrium constant
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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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Explosion
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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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Fischer-Tropsch process
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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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Gibbs free energy (∆ G)
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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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Glass
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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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Group
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A family of elements wither similar chemical properties.
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Halogen
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Any of the elements of Group 17, consisting of flourine, chlorine, bromine, iodine, and astatine.
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Hexagonal close packing
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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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Homogeneous catalysis
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Catalysts occurring within a single phase, usually a gas or liquid
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Hydration energy / Heat of hydration
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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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Hydrolysis
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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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Hydrothermal synthesis
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Mineral synthesis in the presence of heated water.
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Intercalation
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A layer located between layers of different character.
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Ionic bonding
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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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Isoelectronic
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Pertaining to atoms having the same number of electrons outside the nucleus of atoms.
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Isomorphous / Isomorphic minerals
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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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Lattice energy
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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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Layer structure
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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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Madelung constant
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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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Metallic bond
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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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Molecular orbital
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A wave function describing an electron in a molecule.
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Native
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Pertaining to an element found in nature in a nongaseous state.
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Oxidation-reduction reaction
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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 state/number
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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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Oxidizing agent
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Compound that gives up oxygen easily, removes hydrogen from another compound, or attracts negative electrons.
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Period
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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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Periodic table
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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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Pi bonding
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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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Polar covalent bond
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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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Polymorphism
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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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Radiochemistry
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That area of chemistry concerned with the study of radioactive substances.
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Radius ratio
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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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Redox system
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A chemical system in which reduction and oxidation reactions occur.
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Reducing agent
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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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Reduction
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Chemical reaction in which an element gains an electron (has a decrease in positive valence).
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Refractory
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A material of high melting point.
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Resonance structure
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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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Sedimentation
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The act or process of accumulating sediment in layers.
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Semiconductor
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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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Sigma bond
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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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Sintering
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Forming a coherent bonded mass by heating metal powders without melting; used mostly in powder metallurgy.
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Slag
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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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Solvolysis
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A reaction in which a solvent reacts with the solute to form a new substance.
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Spectrum
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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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Spin-orbit coupling
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The interaction between a particle’s spin (ms) and its orbital angular momentum (ml).
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Steel
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An iron-based alloy, malleable under proper conditions, containing up to about 2% carbon.
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Sublimation
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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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Synthesis gas
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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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Valence
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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 electron
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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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Water gas
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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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Wave function / Schrodinger wave function
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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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Weathering
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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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Maximum coordination number
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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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Alkali metals
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Group 1 metals
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Alkaline earth metals
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Group 2 metals
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Paramagnetism
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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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Diamagnetism
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An external magnetic field is repelled by paired electrons
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Formal charge
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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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Periodic trend of radii size
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Increases left and down the table
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Effective nuclear charge (Z*)
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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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Latimer equation
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Equation used to calculate heat of hydration. deltaH = 60900Z^2 / (r + 50) kJ/mol
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Hydrated ion
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M(H₂O)₆ⁿ⁺
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Hydroxy cation
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[M(H₂O)₅OH]ⁿ⁺
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Metal hydroxide
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M(OH)₂
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Oxoanion
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MO₂ⁿ⁻
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Oxocation
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MO₂ⁿ⁺
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Crystal Field Theory
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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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Octahedral field
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dz² dx²-y² eg set dyz dxz dxy t₂g set
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Tetrahedral field
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dxz dyz dxy t₂ set dz² dx²-y² eg set
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Z²/r
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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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2:1 cancellation
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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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Oxoanion nomenclature
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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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Oxo acids, strength and 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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Crystalline solid
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A solid with a well-defined form/shape/structure
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Crystalline lattice
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Regular, repeating 3-D array of atoms, molecules, or ions
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Unit cell
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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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7 Bravais crystal systems
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cubic tetragonal orthorhombic monoclinic hexagonal rhomohedral triclinic
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% empty space
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[volume (unit cell) – volume (atoms) / volume (unit cell)] x 100%
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Soluble salts
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Acidic + nonbasic Nonacidic + basic
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Insoluble salts
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Acidic + basic Nonacidic + nonbasic
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Born-Lande Equation
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Equation used to determine energy of attraction
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Kapustinski Equation
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Equation used to determine the lattice energy
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Cubic closest packing
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The most efficient closest packing; least empty space. Occurs in face-centered cubes. CBA packing
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Schottky Defect
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vacancies of cations = vacancies of anions
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Frenkel Defect
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ions shift to interstitial vacancy
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Leveling effect
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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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Slaking process
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CaO + H₂O ⇌ Ca²⁺ + 2 ⁻OH
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Drying agent
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A substance that pulls H₂O out of whatever it contacts
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Lewis acid
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Electron pair acceptor
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Bronsted-Lowry acid
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Proton donor
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Lux-Flood acid
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Oxide ion acceptor
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Lewis base
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Electron pair donor
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Bronsted-Lowry base
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Proton acceptor
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Lux-Flood base
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Oxide ion donor
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Arrhenius acid
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Hydrogen ion producer
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Arrhenius base
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Hydroxide ion producer
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Comproportionation
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1 atom oxidizes and 1 atom reduces to form one atom.
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Ostwald Process
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NH₃ + ⁵/₄ O₂ → NO + ³/₂ H₂O NO + ½ O₂ → NO₂ NO₂ + H₂O → HNO₃
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Aqua regia (Kingly water)
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A solution that dissolves metals that HCl cannot 1 HNO₃ : 3 HCl
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Production of sulfuric acid
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V₂O₅ H₂SO₄ H₂O ↓ ↓ ↓ SO₂ + ½ O₂ → SO₃ → H₂S₂O₇ → H₂SO₄
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Marsh Test
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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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Clinker
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The mixture that forms concrete when ground with a gypsum.
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Mond Process
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Ni(impure) + 4 CO → Ni(CO)₄ → Ni(pure) + 4 CO
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Mixed-metal oxides
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Compounds that look like oxo salts but are actually oxides formed from multiple metal cations.
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Spinels
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Mixed-metal oxides that are in cubic closest packing of the formula AB₂O₄.
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Normal spinel
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A is in the tetrahedral holes. B is in the octahedral holes. A[B]₂O₄
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Inverse 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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Crystal field stabilization energy (CFSE)
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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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Octahedral site stabilization energy (OSSE)
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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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Standard Hydrogen Electrode (SHE)
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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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Daniell Cell
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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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Salt bridge
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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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Gibbs free energy of a cell
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∆G° = -nFE° n → moles of e⁻ F → Faraday’s constant E° → standard reduction potential
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Standard reduction potential (SRP)
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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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Nernst Equation
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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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Periodic trends in SRP
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Increases right across the table. Most stable (most negative) in the 3rd period.
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Latimer Diagrams
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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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Pourbaix 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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Very electropositive metals
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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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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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Electronegative 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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Electronegativity displacement reaction
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Cations of electronegative metals oxidize (very)electropositive metals to produce cations of more active metals.
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Bauxite
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Ore of aluminum. Al₂O₃∙2H₂O
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Hall-Heroult Process
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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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Downs Cell
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Used in sodium production. (Look up chart in book)
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Pig iron
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Impure iron mixture with a lot of carbon
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Basic oxygen furnace (BOF)
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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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Alnico V
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Magnetized metal (steel) containing Al, Ni, Co, Fe, and Cu
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Inconel
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Heat-resistant steel containing Ni, Cr, Fe, and Cu
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Pearlite
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Very strong form of iron (in Pourbaix diagram)
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Iron corrosion with excess H₂O, little O₂
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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 and 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 little 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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Galvanized
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Protect underlying surface from corrosion by adding a zinc coating
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Sacrificial electrode
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Adding a metal that corrodes easier than iron to prevent corrosion
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Magnesium production
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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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Pseudohalides
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Compounds similar halides that are good R.A. Examples: CN⁻, SCN⁻, S₂O₃²⁻, NCO⁻, N₃⁻, H⁻
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Unstable, explosive bond
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O-O, N-N, O-Cl, O-N, N-Cl, N-Br, N-I
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Chemical conversion
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A reaction where a salt produces a different salt. Ex: Cu₂S → Cu₂O
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Hydrometallurgy
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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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Very electronegative nonmetals
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Nonmetals with Xp > 2.8 and high E°. These include F₂, Cl₂, Br₂, O₂, and N₂.
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Electronegative nonmetals
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Nonmetals with 1.9 < Xp < 2.8.
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Monoatomic nonmetal anions
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Nonmetal anions with the lowest possible oxidation number. These include S²⁻, O²⁻, P, Cl.
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Heat of vaporization
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Enthalpy necessary to convert liquid to vapor at constant T and P.
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Heat of atomization
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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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Periodic trends of atomization energy
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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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Ranking of radii size
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Van der Waals > anionic > metallic > covalent > cationic
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Mullike electronegativity
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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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Superfluid
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A liquid that behaves with zero viscosity, finding the lowest possible spot.
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Ligand
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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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Coordination compound
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Compound that contains 1 or more complex ion
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Donor atom
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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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Inner coordination sphere
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The metal and the ligands directly bonded to it.
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Most common coordination number of coordination compounds
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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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Complex charge
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The sum of charges on the metal and ligands.
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Neutral charge of coordination compound
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The sum of charges on the metal, ligands, and counterbalancing ions.
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Classification of ligands
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Monodentate = 1 donor atom Bidentate = 2 Tetradentate = 4 Hexadentate = 6 Polydentate = 2 or more
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Chelating ligands
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Ligands that bond to one metal ion in multiple places.
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Monodentate ligand examples
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H₂O, CN⁻, NH₃, NO₂⁻, SCN⁻, OH⁻, X⁻, CO, O²⁻
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Ambidentate ligands
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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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Common geometries of complexes
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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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Coordination compound nomenclature
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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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Isomers
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Compounds that have the same composition but a different arrangement of atoms.
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Structural isomers
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Isomers that have different bonds
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Coordination-sphere 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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Linkage 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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Stereoisomers
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Isomers that have the same bonds, but different spatial arrangements.
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Geometric isomers
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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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Optical isomers / Enantiomers
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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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Drago-Wayland Equation
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-∆H = EaEb + CaCb
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E parameter of DW equation
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The susceptibility to undergo electrostatic interaction (ionic, dipole-dipole), relative to iodine, E = 1
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C parameter of DW equation
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The susceptibility to form covalent bonds, relative to iodine, C = 1
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Basic protic solvent
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Nonaqueous solvent for nonpolar molecules. Example: NH₃
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Acidic protic solvent
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Example: H₂SO₄
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Superacidic protic solvent
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Example: SbF₅
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Nonpolar, nonionizing, weakly solvating aprotic solvent
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Nonaqueous solvents that don’t interact much. Examples: CCl₄, C₆H₁₂, benzene, tholuene
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Highly polar, nonionizing, strongly solvating aprotic solvent
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Examples: CH₃CN, DMSO, SO₂
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Highly polar, autoionizing aprotic solvent
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Example: BrF₃
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Irving-Williams Series
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Ranking of the stability of metal cations right across the periodic table
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Chatt, Arhland, Davis Class (a) metals
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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 (b) 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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Pearson “Hard” 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 “Soft” 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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Hard acids
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0.7 < Xp < 1.6 small ionic radii high ionic charge
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Soft acids
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1.9 < Xp < 2.5 large ionic radii low ionic charge
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Soft bases
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2.1 < Xp < 3.0 large ionic radii
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Hard bases
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3.4 < Xp < 4.0 small ionic radii
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Enthalpy with Xp
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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 HSAB
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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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Principle of Symbiosis
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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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Pauling-Pearson Paradox
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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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Qualscheme
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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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Lithophiles
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Metals that are hard acids. Oxides, silicates, sulfates, carbonates
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Chalcophiles
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Metals that are soft acids. Sulfides, tellurides, arsenides
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Atmophiles
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Examples: N₂, noble gases
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Siderophiles
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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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Essential elements
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Elements important for life. All within the first 35 elements except Mo and I
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Charge-transfer transition
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MLCT = Metal to ligand charge transfer LMCT = Ligand to metal charge transfer
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Molecular halides
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Halides formed with highly electronegative elements and high oxidation state metals
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Classes of halides
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Molecular, T-metal, ionic
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Classes of nitrides
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Ionic, covalent, metallic, diamond-like [BN, (SN)x]
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Anions of sulfides
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S²⁻ and S₄²⁻
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Effective Atomic Number (EAN) Rule
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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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18 Electron Rule
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(# of valence e⁻ of metal) + (# of e⁻ donated) = 18 → extra stable
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η = hapticity
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# of carbons equidistant from metal center
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Oxidation addition reaction
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A reaction that increases the oxidation state and coordination number of a metal center.
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Carbonyl stretching
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A carbonyl molecule donates e⁻ through sigma bonds but accepts e⁻ from pi bonds
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Metallic hydrides
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Hydrogen ions occupy holes in metal lattices.
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Interstitional compounds
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Compounds that do not obey stoichiometry. Ex: TiH₁.₇
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Mixed-metal hydrides
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Hydrides formed with two different metals. Ex: FeTiHx and NiMH
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Saline 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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Molecular hydrides
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Hydrides formed with p-block elements (except Al, Bi). Ex: BH₃, GaH₃, NH₃, CH₄, boranes
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Dicarbides (acetylides)
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Carbides formed with C₂²⁻. Ex: CaC₂, [CuC₂, Ag₂C₂, AuC₂ are explosive]
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Metallic carbides
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Interstitional carbides formed with strong M-C bonds. Ex: Fe₃C
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Polymeric carbides
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Carbides with diamond-like structures. Ex: SiC
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Methides
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Carbides formed with C⁴⁻
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Energy of an electron equation
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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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Slater’s Rules
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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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Radius equation with Z*
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r = a₀n² / (Z*)² a₀ = 4πEh² / me²
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Relativistic contraction
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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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Magic numbers
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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 of protons
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2, 8, 20, 28, 50, 82, 114
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Magic numbers of neutrons
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2, 8, 20, 28, 50, 82, 126, 184