Cadmium bromide is a cream-coloured crystalline ionic cadmium salt of hydrobromic acid, soluble in water. It is toxic, along with other cadmium compounds, it is used in the manufacturing of photographic film and lithography. Cadmium bromide is prepared by heating cadmium with bromine vapor; the compound can be prepared by the treatment of dry cadmium acetate with glacial acetic acid and acetyl bromide. Alternatively, it can be obtained by dissolving cadmium or cadmium oxide in hydrobromic acid and evaporating the solution to dryness under helium in an inert atmosphere
Yttrium is a chemical element with symbol Y and atomic number 39. It is a silvery-metallic transition metal chemically similar to the lanthanides and has been classified as a "rare-earth element". Yttrium is always found in combination with lanthanide elements in rare-earth minerals, is never found in nature as a free element. 89Y is the only stable isotope, the only isotope found in the Earth's crust. In 1787, Carl Axel Arrhenius found a new mineral near Ytterby in Sweden and named it ytterbite, after the village. Johan Gadolin discovered yttrium's oxide in Arrhenius' sample in 1789, Anders Gustaf Ekeberg named the new oxide yttria. Elemental yttrium was first isolated in 1828 by Friedrich Wöhler; the most important uses of yttrium are LEDs and phosphors the red phosphors in television set cathode ray tube displays. Yttrium is used in the production of electrodes, electronic filters, superconductors, various medical applications, tracing various materials to enhance their properties. Yttrium has no known biological role.
Exposure to yttrium compounds can cause lung disease in humans. Yttrium is a soft, silver-metallic and crystalline transition metal in group 3; as expected by periodic trends, it is less electronegative than its predecessor in the group and less electronegative than the next member of period 5, zirconium. Yttrium is the first d-block element in the fifth period; the pure element is stable in air in bulk form, due to passivation of a protective oxide film that forms on the surface. This film can reach a thickness of 10 µm; when finely divided, yttrium is unstable in air. Yttrium nitride is formed; the similarities of yttrium to the lanthanides are so strong that the element has been grouped with them as a rare-earth element, is always found in nature together with them in rare-earth minerals. Chemically, yttrium resembles those elements more than its neighbor in the periodic table, if physical properties were plotted against atomic number, it would have an apparent number of 64.5 to 67.5, placing it between the lanthanides gadolinium and erbium.
It also falls in the same range for reaction order, resembling terbium and dysprosium in its chemical reactivity. Yttrium is so close in size to the so-called'yttrium group' of heavy lanthanide ions that in solution, it behaves as if it were one of them. Though the lanthanides are one row farther down the periodic table than yttrium, the similarity in atomic radius may be attributed to the lanthanide contraction. One of the few notable differences between the chemistry of yttrium and that of the lanthanides is that yttrium is exclusively trivalent, whereas about half the lanthanides can have valences other than three; as a trivalent transition metal, yttrium forms various inorganic compounds in the oxidation state of +3, by giving up all three of its valence electrons. A good example is yttrium oxide known as yttria, a six-coordinate white solid. Yttrium forms a water-insoluble fluoride and oxalate, but its bromide, iodide and sulfate are all soluble in water; the Y3+ ion is colorless in solution because of the absence of electrons in the d and f electron shells.
Water reacts with yttrium and its compounds to form Y2O3. Concentrated nitric and hydrofluoric acids do not attack yttrium, but other strong acids do. With halogens, yttrium forms trihalides such as yttrium fluoride, yttrium chloride, yttrium bromide at temperatures above 200 °C. Carbon, selenium and sulfur all form binary compounds with yttrium at elevated temperatures. Organoyttrium chemistry is the study of compounds containing carbon–yttrium bonds. A few of these are known to have yttrium in the oxidation state 0; some trimerization reactions were generated with organoyttrium compounds as catalysts. These syntheses use YCl3 as a starting material, obtained from Y2O3 and concentrated hydrochloric acid and ammonium chloride. Hapticity is a term to describe the coordination of a group of contiguous atoms of a ligand bound to the central atom. Yttrium complexes were the first examples of complexes where carboranyl ligands were bound to a d0-metal center through a η7-hapticity. Vaporization of the graphite intercalation compounds graphite–Y or graphite–Y2O3 leads to the formation of endohedral fullerenes such as Y@C82.
Electron spin resonance studies indicated the formation of 3 − ion pairs. The carbides Y3C, Y2C, YC2 can be hydrolyzed to form hydrocarbons. Yttrium in the Solar System was created through stellar nucleosynthesis by the s-process, but by the r-process; the r-process consists of rapid neutron capture of lighter elements during supernova explosions. The s-process is a slow neutron capture of lighter elements inside pulsating red giant stars. Yttrium isotopes are among the most common products of the nuclear fission of uranium in nuclear explosions and nuclear reactors. In the context of nuclear waste management, the most important isotopes of yttrium
In chemistry, a salt is an ionic compound that can be formed by the neutralization reaction of an acid and a base. Salts are composed of related numbers of cations and anions so that the product is electrically neutral; these component ions can be inorganic, such as organic, such as acetate. Salts can be classified in a variety of ways. Salts that produce hydroxide ions when dissolved in water are called alkali salts. Salts that produce acidic solutions are acidic salts. Neutral salts are those salts that are neither basic. Zwitterions contain an anionic and a cationic centres in the same molecule, but are not considered to be salts. Examples of zwitterions include amino acids, many metabolites and proteins. Solid salts tend to be transparent. In many cases, the apparent opacity or transparency are only related to the difference in size of the individual monocrystals. Since light reflects from the grain boundaries, larger crystals tend to be transparent, while the polycrystalline aggregates look like white powders.
Salts exist in many different colors, which arise either from the cations. For example: sodium chromate is yellow by virtue of the chromate ion potassium dichromate is orange by virtue of the dichromate ion cobalt nitrate is red owing to the chromophore of hydrated cobalt. copper sulfate is blue because of the copper chromophore potassium permanganate has the violet color of permanganate anion. Nickel chloride is green of sodium chloride, magnesium sulfate heptahydrate are colorless or white because the constituent cations and anions do not absorb in the visible part of the spectrumFew minerals are salts because they would be solubilized by water. Inorganic pigments tend not to be salts, because insolubility is required for fastness; some organic dyes are salts, but they are insoluble in water. Different salts can elicit all five basic tastes, e.g. salty, sour and umami or savory. Salts of strong acids and strong bases are non-volatile and odorless, whereas salts of either weak acids or weak bases may smell like the conjugate acid or the conjugate base of the component ions.
That slow, partial decomposition is accelerated by the presence of water, since hydrolysis is the other half of the reversible reaction equation of formation of weak salts. Many ionic compounds exhibit significant solubility in water or other polar solvents. Unlike molecular compounds, salts dissociate in solution into cationic components; the lattice energy, the cohesive forces between these ions within a solid, determines the solubility. The solubility is dependent on how well each ion interacts with the solvent, so certain patterns become apparent. For example, salts of sodium and ammonium are soluble in water. Notable exceptions include potassium cobaltinitrite. Most nitrates and many sulfates are water-soluble. Exceptions include barium sulfate, calcium sulfate, lead sulfate, where the 2+/2− pairing leads to high lattice energies. For similar reasons, most alkali metal carbonates are not soluble in water; some soluble carbonate salts are: potassium carbonate and ammonium carbonate. Salts are characteristically insulators.
Molten salts or solutions of salts conduct electricity. For this reason, liquified salts and solutions containing dissolved salts are called electrolytes. Salts characteristically have high melting points. For example, sodium chloride melts at 801 °C; some salts with low lattice energies are liquid near room temperature. These include molten salts, which are mixtures of salts, ionic liquids, which contain organic cations; these liquids exhibit unusual properties as solvents. The name of a salt starts with the name of the cation followed by the name of the anion. Salts are referred to only by the name of the cation or by the name of the anion. Common salt-forming cations include: Ammonium NH+4 Calcium Ca2+ Iron Fe2+ and Fe3+ Magnesium Mg2+ Potassium K+ Pyridinium C5H5NH+ Quaternary ammonium NR+4, R being an alkyl group or an aryl group Sodium Na+ Copper Cu2+Common salt-forming anions include: Acetate CH3COO− Carbonate CO2−3 Chloride Cl− Citrate HOC2 Cyanide C≡N− Fluoride F− Nitrate NO−3 Nitrite NO−2 Oxide O2− Phosphate PO3−4 Sulfate SO2−4 Salts with varying number of hydrogen atoms, with respect to the parent acid, replaced by cations can be referred to as monobasic, dibasic or tribasic salts: Sodium phosphate monobasic Sodium phosphate dibasic Sodium phosphate tribasic Salts are formed by a chemical reaction between: A base and an acid, e.g. NH3 + HCl → NH4Cl A metal and an acid, e.g. Mg + H2SO4 → MgSO4 + H2 A metal and a non-metal, e.g. Ca + Cl2 → CaCl2 A base and an a
Cadmium oxide is an inorganic compound with the formula CdO. It is one of the main precursors to other cadmium compounds, it crystallizes in a cubic rocksalt lattice like sodium chloride, with octahedral cation and anion centers. It occurs as the rare mineral monteponite. Cadmium oxide can be found as brown or red crystals. Cadmium oxide is an n-type semiconductor with a band gap of 2.18 eV at room temperature. Since cadmium compounds are found in association with zinc ores, cadmium oxide is a common by-product of zinc refining, it is produced by burning elemental cadmium in air. Pyrolysis of other cadmium compounds, such as the nitrate or the carbonate affords this oxide; when pure, it is red, but CdO is unusual in being available in many differing colours due to its tendency to form defect structures resulting from anion vacancies. Cadmium oxide is prepared commercially by oxidizing cadmium vapor in air. Cadmium oxide is used in cadmium plating baths, electrodes for storage batteries, cadmium salts, ceramic glazes and nematocide.
Major uses for cadmium oxide are as an ingredient for electroplating baths, in pigments. CdO is used as a transparent conductive material, prepared as a transparent conducting film as early as 1907 by Karl Baedeker. Cadmium oxide in the form of thin films has been used in applications such as photodiodes, photovoltaic cells, transparent electrodes, liquid crystal displays, IR detectors, anti reflection coatings. CdO microparticles undergo bandgap excitation when exposed to UV-A light and is selective in phenol photodegradation. Most commercial electroplating of cadmium is done by electrodeposition from cyanide baths; these cyanide baths consist of cadmium oxide and sodium cyanide in water, which form cadmium cyanide and sodium hydroxide. A typical formula is 75 g/L sodium cyanide; the cadmium concentration may vary by as much as 50%. Brighteners are added to the bath and the plating is done at room temperature with high purity cadmium anodes. CdO is a basic oxide and is thus attacked by aqueous acids to give solutions of 2+.
Upon treatment with strong alkaline solutions, 2− forms. A thin coat of cadmium oxide forms on the surface of cadmium in moist air at room temperature. Cadmium will oxidize at room temperatures to form CdO. Cadmium vapor and steam will form hydrogen in a reversible reaction. Cadmium oxide information at Webelements
Cadmium is a chemical element with symbol Cd and atomic number 48. This soft, bluish-white metal is chemically similar to the two other stable metals in group 12, zinc and mercury. Like zinc, it demonstrates oxidation state +2 in most of its compounds, like mercury, it has a lower melting point than the transition metals in groups 3 through 11. Cadmium and its congeners in group 12 are not considered transition metals, in that they do not have filled d or f electron shells in the elemental or common oxidation states; the average concentration of cadmium in Earth's crust is between 0.5 parts per million. It was discovered in 1817 by Stromeyer and Hermann, both in Germany, as an impurity in zinc carbonate. Cadmium is a byproduct of zinc production. Cadmium was used for a long time as a corrosion-resistant plating on steel, cadmium compounds are used as red and yellow pigments, to color glass, to stabilize plastic. Cadmium use is decreasing because it is toxic and nickel-cadmium batteries have been replaced with nickel-metal hydride and lithium-ion batteries.
One of its few new uses is cadmium telluride solar panels. Although cadmium has no known biological function in higher organisms, a cadmium-dependent carbonic anhydrase has been found in marine diatoms. Cadmium is a soft, ductile, bluish-white divalent metal, it forms complex compounds. Unlike most other metals, cadmium is resistant to corrosion and is used as a protective plate on other metals; as a bulk metal, cadmium is not flammable. Although cadmium has an oxidation state of +2, it exists in the +1 state. Cadmium and its congeners are not always considered transition metals, in that they do not have filled d or f electron shells in the elemental or common oxidation states. Cadmium burns in air to form brown amorphous cadmium oxide. Hydrochloric acid, sulfuric acid, nitric acid dissolve cadmium by forming cadmium chloride, cadmium sulfate, or cadmium nitrate; the oxidation state +1 can be produced by dissolving cadmium in a mixture of cadmium chloride and aluminium chloride, forming the Cd22+ cation, similar to the Hg22+ cation in mercury chloride.
Cd + CdCl2 + 2 AlCl3 → Cd22The structures of many cadmium complexes with nucleobases, amino acids, vitamins have been determined. Occurring cadmium is composed of 8 isotopes. Two of them are radioactive, three are expected to decay but have not done so under laboratory conditions; the two natural radioactive isotopes are 116Cd. The other three are 106Cd, 108Cd, 114Cd. At least three isotopes – 110Cd, 111Cd, 112Cd – are stable. Among the isotopes that do not occur the most long-lived are 109Cd with a half-life of 462.6 days, 115Cd with a half-life of 53.46 hours. All of the remaining radioactive isotopes have half-lives of less than 2.5 hours, the majority have half-lives of less than 5 minutes. Cadmium has 8 known meta states, with the most stable being 113mCd, 115mCd, 117mCd; the known isotopes of cadmium range in atomic mass from 94.950 u to 131.946 u. For isotopes lighter than 112 u, the primary decay mode is electron capture and the dominant decay product is element 47. Heavier isotopes decay through beta emission producing element 49.
One isotope of cadmium, 113Cd, absorbs neutrons with high selectivity: With high probability, neutrons with energy below the cadmium cut-off will be absorbed. The cadmium cut-off is about 0.5 eV, neutrons below that level are deemed slow neutrons, distinct from intermediate and fast neutrons. Cadmium is created via the s-process in low- to medium-mass stars with masses of 0.6 to 10 solar masses, over thousands of years. In that process, a silver atom captures a neutron and undergoes beta decay. Cadmium was discovered in 1817 by Friedrich Stromeyer and Karl Samuel Leberecht Hermann, both in Germany, as an impurity in zinc carbonate. Stromeyer found the new element as an impurity in zinc carbonate, for 100 years, Germany remained the only important producer of the metal; the metal was named after the Latin word for calamine. Stromeyer noted that some impure samples of calamine changed color when heated but pure calamine did not, he was persistent in studying these results and isolated cadmium metal by roasting and reducing the sulfide.
The potential for cadmium yellow as pigment was recognized in the 1840s, but the lack of cadmium limited this application. Though cadmium and its compounds are toxic in certain forms and concentrations, the British Pharmaceutical Codex from 1907 states that cadmium iodide was used as a medication to treat "enlarged joints, scrofulous glands, chilblains". In 1907, the International Astronomical Union defined the international ångström in terms of a red cadmium spectral line. This
The boiling point of a substance is the temperature at which the vapor pressure of a liquid equals the pressure surrounding the liquid and the liquid changes into a vapor. The boiling point of a liquid varies depending upon the surrounding environmental pressure. A liquid in a partial vacuum has a lower boiling point than when that liquid is at atmospheric pressure. A liquid at high pressure has a higher boiling point than when that liquid is at atmospheric pressure. For example, water at 93.4 °C at 1,905 metres altitude. For a given pressure, different liquids will boil at different temperatures; the normal boiling point of a liquid is the special case in which the vapor pressure of the liquid equals the defined atmospheric pressure at sea level, 1 atmosphere. At that temperature, the vapor pressure of the liquid becomes sufficient to overcome atmospheric pressure and allow bubbles of vapor to form inside the bulk of the liquid; the standard boiling point has been defined by IUPAC since 1982 as the temperature at which boiling occurs under a pressure of 1 bar.
The heat of vaporization is the energy required to transform a given quantity of a substance from a liquid into a gas at a given pressure. Liquids may change to a vapor at temperatures below their boiling points through the process of evaporation. Evaporation is a surface phenomenon in which molecules located near the liquid's edge, not contained by enough liquid pressure on that side, escape into the surroundings as vapor. On the other hand, boiling is a process in which molecules anywhere in the liquid escape, resulting in the formation of vapor bubbles within the liquid. A saturated liquid contains as much thermal energy. Saturation temperature means boiling point; the saturation temperature is the temperature for a corresponding saturation pressure at which a liquid boils into its vapor phase. The liquid can be said to be saturated with thermal energy. Any addition of thermal energy results in a phase transition. If the pressure in a system remains constant, a vapor at saturation temperature will begin to condense into its liquid phase as thermal energy is removed.
A liquid at saturation temperature and pressure will boil into its vapor phase as additional thermal energy is applied. The boiling point corresponds to the temperature at which the vapor pressure of the liquid equals the surrounding environmental pressure. Thus, the boiling point is dependent on the pressure. Boiling points may be published with respect to the NIST, USA standard pressure of 101.325 kPa, or the IUPAC standard pressure of 100.000 kPa. At higher elevations, where the atmospheric pressure is much lower, the boiling point is lower; the boiling point increases with increased pressure up to the critical point, where the gas and liquid properties become identical. The boiling point cannot be increased beyond the critical point; the boiling point decreases with decreasing pressure until the triple point is reached. The boiling point cannot be reduced below the triple point. If the heat of vaporization and the vapor pressure of a liquid at a certain temperature are known, the boiling point can be calculated by using the Clausius–Clapeyron equation, thus: T B = − 1, where: T B is the boiling point at the pressure of interest, R is the ideal gas constant, P is the vapour pressure of the liquid at the pressure of interest, P 0 is some pressure where the corresponding T 0 is known, Δ H vap is the heat of vaporization of the liquid, T 0 is the boiling temperature, ln is the natural logarithm.
Saturation pressure is the pressure for a corresponding saturation temperature at which a liquid boils into its vapor phase. Saturation pressure and saturation temperature have a direct relationship: as saturation pressure is increased, so is saturation temperature. If the temperature in a system remains constant, vapor at saturation pressure and temperature will begin to condense into its liquid phase as the system pressure is increased. A liquid at saturation pressure and temperature will tend to flash into its vapor phase as system pressure is decreased. There are two conventions regarding the standard boiling point of water: The normal boiling point is 99.97 °C at a pressure of 1 atm. The IUPAC recommended standard boiling point of water at a standard pressure of 100 kPa is 99.61 °C. For comparison, on top of Mount Everest, at 8,848 m elevation, the pressure is about 34 kPa and the boiling point of water is 71 °C; the Celsius temperature scale was defined until 1954 by two points: 0 °C being defined by the wate
Gibbs free energy
In thermodynamics, the Gibbs free energy is a thermodynamic potential that can be used to calculate the maximum of reversible work that may be performed by a thermodynamic system at a constant temperature and pressure. The Gibbs free energy is the maximum amount of non-expansion work that can be extracted from a thermodynamically closed system; when a system transforms reversibly from an initial state to a final state, the decrease in Gibbs free energy equals the work done by the system to its surroundings, minus the work of the pressure forces. The Gibbs energy is the thermodynamic potential, minimized when a system reaches chemical equilibrium at constant pressure and temperature, its derivative with respect to the reaction coordinate of the system vanishes at the equilibrium point. As such, a reduction in G is a necessary condition for the spontaneity of processes at constant pressure and temperature; the Gibbs free energy called available energy, was developed in the 1870s by the American scientist Josiah Willard Gibbs.
In 1873, Gibbs described this "available energy" as the greatest amount of mechanical work which can be obtained from a given quantity of a certain substance in a given initial state, without increasing its total volume or allowing heat to pass to or from external bodies, except such as at the close of the processes are left in their initial condition. The initial state of the body, according to Gibbs, is supposed to be such that "the body can be made to pass from it to states of dissipated energy by reversible processes". In his 1876 magnum opus On the Equilibrium of Heterogeneous Substances, a graphical analysis of multi-phase chemical systems, he engaged his thoughts on chemical free energy in full. According to the second law of thermodynamics, for systems reacting at STP, there is a general natural tendency to achieve a minimum of the Gibbs free energy. A quantitative measure of the favorability of a given reaction at constant temperature and pressure is the change ΔG in Gibbs free energy, caused by the reaction.
As a necessary condition for the reaction to occur at constant temperature and pressure, ΔG must be smaller than the non-PV work, equal to zero. ΔG equals the maximum amount of non-PV work that can be performed as a result of the chemical reaction for the case of reversible process. If the analysis indicated a positive ΔG for the reaction energy — in the form of electrical or other non-PV work — would have to be added to the reacting system for ΔG to be smaller than the non-PV work and make it possible for the reaction to occur. We can think of ∆G as the amount of "free" or "useful" energy available to do work; the equation can be seen from the perspective of the system taken together with its surroundings. First, assume that the given reaction at constant temperature and pressure is the only one, occurring; the entropy released or absorbed by the system equals the entropy that the environment must absorb or release, respectively. The reaction will only be allowed if the total entropy change of the universe is positive.
This is reflected in a negative ΔG, the reaction is called exergonic. If we couple reactions an otherwise endergonic chemical reaction can be made to happen; the input of heat into an inherently endergonic reaction, such as the elimination of cyclohexanol to cyclohexene, can be seen as coupling an unfavourable reaction to a favourable one such that the total entropy change of the universe is greater than or equal to zero, making the total Gibbs free energy difference of the coupled reactions negative. In traditional use, the term "free" was included in "Gibbs free energy" to mean "available in the form of useful work"; the characterization becomes more precise if we add the qualification that it is the energy available for non-volume work.. However, an increasing number of books and journal articles do not include the attachment "free", referring to G as "Gibbs energy"; this is the result of a 1988 IUPAC meeting to set unified terminologies for the international scientific community, in which the adjective "free" was banished.
This standard, has not yet been universally adopted. The quantity called "free energy" is a more advanced and accurate replacement for the outdated term affinity, used by chemists in the earlier years of physical chemistry to describe the force that caused chemical reactions. In 1873, Willard Gibbs published A Method of Geometrical Representation of the Thermodynamic Properties of Substances by Means of Surfaces, in which he sketched the principles of his new equation, able to predict or estimate the tendencies of various natural processes to ensue when bodies or systems are brought into contact. By studying the interactions of homogeneous substances in contact, i.e. bodies composed of part solid, part liquid, part vapor, by using a three-dimensional volume-entropy-internal energy graph, Gibbs was able to determine three states of equilibrium, i.e. "necessarily stable", "neutral", "unstable", whether or not changes woul