"Kid and other Academy students become full members of the Legion, as replacements for Legion members believed deceased. Chemical King is a mutant with the power to act as a human catalyst. He can speed up or slow down any chemical reaction; for example, he can cause iron to rapidly rust, or even burst into flames from rapid oxidation. His powers can affect the chemical reactions in the human body, rendering an opponent unconscious from shock, but take several seconds and can only focus on a single opponent at a time. Late in his career, Chemical King displayed the ability"
"Catalysis Catalysis is the process of increasing the rate of a chemical reaction by adding a substance known as a catalyst (), which is not consumed in the catalyzed reaction and can continue to act repeatedly. Because of this, only very small amounts of catalyst are required to alter the reaction rate in principle. In general, chemical reactions occur faster in the presence of a catalyst because the catalyst provides an alternative reaction mechanism with a lower activation energy than the non-catalyzed mechanism. In catalyzed mechanisms, the catalyst usually reacts to form a temporary intermediate, which then regenerates the original"
"Accelerant Accelerants are substances that can bond, mix or disturb another substance and cause an increase in the speed of a natural, or artificial chemical process. Accelerants play a major role in chemistry—most chemical reactions can be hastened with an accelerant. Accelerants alter a chemical bond, speed up a chemical process, or bring organisms back to homeostasis. Accelerants are not necessarily catalysts as they may be consumed by the process. In fire protection, the term accelerant is used very broadly to include any substance or mixture that ""accelerates"" the development of fire to commit arson. Chemists would distinguish an accelerant"
"on temperature as observed by the van 't Hoff equation. Adding a catalyst will affect both the forward reaction and the reverse reaction in the same way and will not have an effect on the equilibrium constant. The catalyst will speed up both reactions thereby increasing the speed at which equilibrium is reached. Although the macroscopic equilibrium concentrations are constant in time, reactions do occur at the molecular level. For example, in the case of acetic acid dissolved in water and forming acetate and hydronium ions, a proton may hop from one molecule of acetic acid on to a water"
"however, are the enzymes. Virtually every reaction in a living cell requires an enzyme to lower the activation energy of the reaction. These molecules recognize specific reactant molecules called ""substrates""; they then catalyze the reaction between them. By lowering the activation energy, the enzyme speeds up that reaction by a rate of 10 or more; a reaction that would normally take over 3,000 years to complete spontaneously might take less than a second with an enzyme. The enzyme itself is not used up in the process, and is free to catalyze the same reaction with a new set of substrates."
"and not what actually does happen, or how fast, away from equilibrium. Which reactions do occur and how fast is the subject of chemical kinetics, another branch of physical chemistry. A key idea in chemical kinetics is that for reactants to react and form products, most chemical species must go through transition states which are higher in energy than either the reactants or the products and serve as a barrier to reaction. In general, the higher the barrier, the slower the reaction. A second is that most chemical reactions occur as a sequence of elementary reactions, each with its own"
"speeds the setting time and thus cure time starts earlier. This allows concrete to be placed in winter with reduced risk of frost damage. Concrete is damaged if it does not reach a strength of before freezing. Typical chemicals used for acceleration today are calcium nitrate (Ca(NO)), calcium formate (Ca(HCOO)) and sodium nitrate (NaNO). Accelerant Accelerants are substances that can bond, mix or disturb another substance and cause an increase in the speed of a natural, or artificial chemical process. Accelerants play a major role in chemistry—most chemical reactions can be hastened with an accelerant. Accelerants alter a chemical bond,"
"to changes in viscosity. This also implies that at the onset of autoacceleration the overall rate of reaction increases relative to the rate of unautoaccelerated reaction given by the overall rate of reaction equation for free-radical polymerization: Approximately, as the termination decreases by a factor of 4, the overall rate of reaction will double. The decrease of termination reactions also allows radical chains to add monomer for longer time periods, raising the mass-average molecular mass dramatically. However, the number-average molecular mass only increases slightly, leading to broadening of the molecular mass distribution (high dispersity, very polydispersed product). Autoacceleration Autoacceleration (gel"
"methionine adenosyltransferase. This continuous regeneration means that small amounts of coenzymes can be used very intensively. For example, the human body turns over its own weight in ATP each day. As with all catalysts, enzymes do not alter the position of the chemical equilibrium of the reaction. In the presence of an enzyme, the reaction runs in the same direction as it would without the enzyme, just more quickly. For example, carbonic anhydrase catalyzes its reaction in either direction depending on the concentration of its reactants: The rate of a reaction is dependent on the activation energy needed to form"
"the rate of a chemical reaction, thermodynamics determines the extent to which reactions occur. In a reversible reaction, chemical equilibrium is reached when the rates of the forward and reverse reactions are equal [the principle of dynamic equilibrium ] and the concentrations of the reactants and Products no longer change. This is demonstrated by, for example, the Haber–Bosch process for combining nitrogen and hydrogen to produce ammonia. Chemical clock reactions such as the Belousov–Zhabotinsky reaction demonstrate that component concentrations can oscillate for a long time before finally attaining the equilibrium. In general terms, the free energy change (ΔG) of a"
"found that typical electron-withdrawing or electron-donating substituents had little effect on the rate, but ester substituents such as acetate or pivaloate seemed to cause noticeable rate acceleration. This observation is difficult to rationalize given that the X group is far from the reactive center during the course of the reaction and electronics do not seem to be the cause. In general, despite the relative ease of electronic tuning with organic catalysts, chemists have not yet reached a useful understanding of these modifications. To date, there have been few examples of hydrogen-bond catalysis in the synthesis of natural products despite the"
"the temperature dependence of reaction. At maximum concentration (typically only found in laboratory settings), mixing the chemicals results in a furious reaction, producing large amounts of light for only a few seconds. The same effect can be achieved by adding copious amounts of sodium salicyate or other bases. Heating a glow stick also causes the reaction to proceed faster and the glow stick to glow more brightly for a brief period. Cooling a glow stick slows the reaction a small amount and causes it to last longer, but the light is dimmer. This can be demonstrated by refrigerating or freezing"
"the reactants. Reaction can occur only at their area of contact; in the case of a liquid and a gas, at the surface of the liquid. Vigorous shaking and stirring may be needed to bring the reaction to completion. This means that the more finely divided a solid or liquid reactant the greater its surface area per unit volume and the more contact it with the other reactant, thus the faster the reaction. To make an analogy, for example, when one starts a fire, one uses wood chips and small branches — one does not start with large logs right"
"the sense that the reaction products are more stable than the starting material, though the uncatalysed reaction is slow. In fact, the decomposition of hydrogen peroxide is so slow that hydrogen peroxide solutions are commercially available. This reaction is strongly affected by catalysts such as manganese dioxide, or the enzyme peroxidase in organisms. Upon the addition of a small amount of manganese dioxide, the hydrogen peroxide reacts rapidly. This effect is readily seen by the effervescence of oxygen. The manganese dioxide is not consumed in the reaction, and thus may be recovered unchanged, and re-used indefinitely. Accordingly, manganese dioxide ""catalyses"""
"catalyst, which only facilitates the reaction, so that its total concentration, free plus combined, formula_25 is a constant. This conservation law can also be observed by adding the first and third equations above. In their original analysis, Michaelis and Menten assumed that the substrate is in instantaneous chemical equilibrium with the complex, which implies From the enzyme conservation law, we obtain Combining the two expressions above, gives us Upon simplification, we get where formula_30 is the dissociation constant for the enzyme-substrate complex. Hence the velocity formula_1 of the reaction – the rate at which P is formed – is where"
"the side with the fewer moles of gas. The reaction yield stabilizes at equilibrium, but can be increased by removing the product from the reaction mixture or changed by increasing the temperature or pressure. A change in the concentrations of the reactants does not affect the equilibrium constant, but does affect the equilibrium position. Chemical reactions are determined by the laws of thermodynamics. Reactions can proceed by themselves if they are exergonic, that is if they release energy. The associated free energy of the reaction is composed of two different thermodynamic quantities, enthalpy and entropy: Reactions can be exothermic, where"
"reaction by either stabilizing the transition state structure or destabilizing a key reaction intermediate, and as only a small amount of catalyst is required it can provide economic access to otherwise expensive or difficult to synthesize organic molecules. Catalysts may also influence a reaction rate by changing the mechanism of the reaction. Although a rate law provides the stoichiometry of the transition state structure, it does not provide any information about breaking or forming bonds. The substitution of an isotope near a reactive position often leads to a change in the rate of a reaction. Isotopic substitution changes the potential"
"contributes only a very small proportion to the increase in rate of reaction. Much more important is the fact that the proportion of reactant molecules with sufficient energy to react (energy greater than activation energy: ""E"" > ""E"") is significantly higher and is explained in detail by the Maxwell–Boltzmann distribution of molecular energies. The 'rule of thumb' that the rate of chemical reactions doubles for every 10 °C temperature rise is a common misconception. This may have been generalized from the special case of biological systems, where the ""α"" (temperature coefficient) is often between 1.5 and 2.5. A reaction's kinetics"
"redistribution of substances in the human body. Faster reactions are usually studied with ultrafast laser spectroscopy where utilization of femtosecond lasers allows short-lived transition states to be monitored at time scaled down to a few femtoseconds. Chemical reaction A chemical reaction is a process that leads to the chemical transformation of one set of chemical substances to another. Classically, chemical reactions encompass changes that only involve the positions of electrons in the forming and breaking of chemical bonds between atoms, with no change to the nuclei (no change to the elements present), and can often be described by a chemical"
"Autoacceleration Autoacceleration (gel effect, Trommsdorff–Norrish effect) is a dangerous reaction behavior that can occur in free-radical polymerization systems. It is due to the localized increases in viscosity of the polymerizing system that slow termination reactions. The removal of reaction obstacles therefore causes a rapid increase in the overall rate of reaction, leading to possible reaction runaway and altering the characteristics of the polymers produced. Autoacceleration of the overall rate of a free-radical polymerization system has been noted in many bulk polymerization systems. The polymerization of methyl methacrylate, for example, deviates strongly from classical mechanism behavior around 20% conversion; in this"
"the substrate concentration up to a certain level called V; at V, increase in substrate concentration does not cause any increase in reaction rate as there no more enzyme (E) available for reacting with substrate (S). Here, the rate of reaction becomes dependent on the ES complex and the reaction becomes a unimolecular reaction with an order of zero. Though the enzymatic mechanism for the unimolecular reaction <chem> ES ->[k_{cat}] E + P </chem> can be quite complex, there is typically one rate-determining enzymatic step that allows this reaction to be modelled as a single catalytic step with an apparent"
"Chemical equilibrium In a chemical reaction, chemical equilibrium is the state in which both reactants and products are present in concentrations which have no further tendency to change with time, so that there is no observable change in the properties of the system. Usually, this state results when the forward reaction proceeds at the same rate as the reverse reaction. The reaction rates of the forward and backward reactions are generally not zero, but equal. Thus, there are no net changes in the concentrations of the reactant(s) and product(s). Such a state is known as dynamic equilibrium. The concept of"
"temperature increase would be about ), and in any case the rule is only a rough approximation and cannot always be relied on. The same is true, up to a point, of the chemical reactions of living things. They are usually catalyzed by enzymes which change reaction rates, but with no variation in catalytic action, the rule of thumb is still mostly applicable. In the case of bacteria and fungi, the reactions needed to feed and reproduce speed up at higher temperatures, up to the point that the proteins and other compounds in their cells themselves begin to break down,"
"(see also thermodynamics). The second law of thermodynamics describes why a catalyst does not change the chemical equilibrium of a reaction. Suppose there was such a catalyst that shifted an equilibrium. Introducing the catalyst to the system would result in a reaction to move to the new equilibrium, producing energy. Production of energy is a necessary result since reactions are spontaneous only if Gibbs free energy is produced, and if there is no energy barrier, there is no need for a catalyst. Then, removing the catalyst would also result in reaction, producing energy; i.e. the addition and its reverse process,"
"by heterocycles under prolonged reaction at elevated temperature. To avoid this, chemists often use pressure reactors to accelerate reactions at high temperature and pressure. Q-Tube and microwave synthesizer are available safe pressure reactors. Coupling reactions include (not exhaustive): In one study, an unusual coupling reaction was described in which an organomolybdenum compound, [Mo(CCH)(OAc)(HO)](CFSO) not only sat on a shelf for 30 years without any sign of degradation but also decomposed in water to generate 2-butyne, which is the coupling adduct of its two ethylidyne ligands. This, according to the researchers, opens another way for aqueous organometallic chemistry. One method for"
"intermediate time, ""t"", in another (the parent reaction). One would expect the reaction progress, described by the rate vs. substrate concentration plots detailed above, to map directly onto each other from that interception point onward. This will hold true, however, only if the rate of the reaction is not altered by changes to the active substrate/catalyst concentration (such as by catalyst activation, catalyst deactivation, or product inhibition) before that interception. A perfect overlay of multiple experiments with the same-excess but different initial substrate loadings suggests that no changes in the active substrate/catalyst concentration occur over the course of the reaction."
"is positive, the reaction is nonspontaneous. Chemical reactions can be “coupled” together if they share intermediates. In this case, the overall Gibbs Free Energy change is simply the sum of the ∆""G"" values for each reaction. Therefore, an unfavorable reaction (positive ∆""G"") can be driven by a second, highly favorable reaction (negative ∆""G"" where the magnitude of ∆""G"" > magnitude of ∆""G""). For example, the reaction of glucose with fructose to form sucrose has a ∆""G"" value of +5.5 kcal/mole. Therefore, this reaction will not occur spontaneously. The breakdown of ATP to form ADP and inorganic phosphate has a ∆""G"""
"formula_33 is the maximum reaction velocity. An alternative analysis of the system was undertaken by British botanist G. E. Briggs and British geneticist J. B. S. Haldane in 1925. They assumed that the concentration of the intermediate complex does not change on the time-scale of product formation – known as the quasi-steady-state assumption or pseudo-steady-state-hypothesis. Mathematically, this assumption means formula_34. Combining this relationship with the enzyme conservation law, formula_35, gives us Moving [ES] to same side then factoring it out, the concentration of the complex is where is known as the Michaelis constant, where formula_9, formula_10, and formula_41 are, respectively,"
"a fireplace in the presence of oxygen, but it does not when it is stored at room temperature. The reaction is spontaneous at low and high temperatures but at room temperature its rate is so slow that it is negligible. The increase in temperature, as created by a match, allows the reaction to start and then it heats itself, because it is exothermic. That is valid for many other fuels, such as methane, butane, and hydrogen. Reaction rates can be independent of temperature (""non-Arrhenius"") or decrease with increasing temperature (""anti-Arrhenius""). Reactions without an activation barrier (e.g., some radical reactions), tend"
"an ordinary combustion reaction. Unless the reaction occurs very rapidly, the thermally expanding gases will be moderately dissipated in the medium, with no large differential in pressure and there will be no explosion. Consider a wood fire. As the fire burns, there certainly is the evolution of heat and the formation of gases, but neither is liberated rapidly enough to build up a sudden substantial pressure differential and then cause an explosion. This can be likened to the difference between the energy discharge of a battery, which is slow, and that of a flash capacitor like that in a camera"
"Determination of the rate law was historically accomplished by monitoring the concentration of a reactant during a reaction through gravimetric analysis, but today it is almost exclusively done through fast and unambiguous spectroscopic techniques. In most cases, the determination of rate equations is simplified by adding a large excess (“flooding”) all but one of the reactants. The study of catalysis and catalytic reactions is very important to the field of physical organic chemistry. A catalyst participates in the chemical reaction but is not consumed in the process. A catalyst lowers the activation energy barrier (ΔG), increasing the rate of a"
"competing reactions, parallel reactions, consecutive reactions, chain reactions, reversible reactions, etc. [10] If one reaction occurs much faster than the other one (k > k), it (k) will be called the main reaction, the other one (k) side reaction. If both reactions roughly of same speed (k ≅ k) is spoken of parallel reactions. If the reactions <chem>{A} + B ->[{k_1}] P1</chem> and <chem>{A} + C ->[{k_2}] P2</chem> are irreversibly (without reverse reaction), then the ratio of P and P corresponds to the relative reactivity of B and C compared with A: Side reaction A side reaction is a chemical"
"than the chemical reaction time and the system explodes with high temperature since there is not enough time for conduction to remove the heat. On the other hand, when formula_27 is very low, heat conduction time is much faster than the chemical reaction time, such that all the heat produced by the chemical reaction is immediately conducted to the wall, thus there is no explosion, it goes to an almost steady state, Amable Liñán coined this mode as slowly reacting mode. At a critical Damköhler number formula_47 the system goes from slowly reacting mode to explosive mode. Therefore, formula_48, the"