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Noble gases

Noble gases (inert gases; rare gases; group 18 elements) is a group of monatomic gaseous elements forming group 18 (formerly group 0) of the periodic table: helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe) and radon (Rn). The electron configuration of helium is Is². The configurations of the others terminate in ns²np⁶ and all inner shells are fully occupied. The elements thus represent the termination of a period and have closed-shell configuration and associated high ionization energies (He 2370 to Rn 1040 kj ^mol-1) and lack of chemical reactivity. Being monatomic the noble gases are spherically symmetrical and have very weak interatomic interactions and consequent low enthalpies of vaporization. The behaviour of the lighter members approaches that of an ideal gas at normal temperatures; with the heavier members increasing polarizability and dispersion forces lead to easier liquefaction under pressure. Four types of compound have been described for the noble gases but of these only one can be correctly described as compounds in the normal sense. One type consists of such species as HHe⁺. He₂⁺, Ar₂⁺, HeLi⁺, which form under highly energetic conditions, such as those in arcs and sparks. They are short-lived and only detected spectroscopically. A second group of materials described as inert-gas-metal compounds do not have defined compositions and are simply noble gases adsorbed onto the surface of dispersed metal. The third type, previously described as 'hydrates' are in fact clathrate compounds with the noble gas molecule trapped in a water lattice. True compounds of the noble gases were first described in 1962 and several fluorides, oxyfluorides, fluoroplatinates, and fluoroantimonates of 'xenon are known. A few krypton fluorides and a radon fluoride are also known although the short half-life of radon and its intense alpha activity restrict the availability of information. Apart from argon, the noble gases are present in the atmosphere at only trace levels. Helium may be found along with natural gas (up to 7%), arising from the radioactive decay of heavier elements (via alpha particles).

Equilibrium and equilibrium constant

Equilibrium is astate in which a system has its energy distributed in the statistically most probable manner: a state of a system in which forces, influences, reactions and etc. balance each other out so that there is no net change. A body is said to be in thermal equilibrium if no net heat exchange is taking place within it or between it and its surroundings. A system is in chemical equilibrium when a reaction and its reverse are proceeding at equal rates (see also equilibrium constant). These are examples of dynamic equilibrium,in which activity in one sense or direction is in aggregate balanced by comparable reverse activity.

Equilibrium constant for a reversible reaction of the type is the following :

xA + yB « zC + wD.

Chemical equilibrium occurs when the rate of the forward reaction equals the rate of the back reaction, so that the concentrations of products and reactants reach steady-state values. It can be shown that at equilibrium the ratio of concentration

is a constant for a given reaction and fixed temperature, called the equilibrium constant K_C (where the c indicates concentrations have been used). Note that, by convention, the products on the right-hand side of the reaction are used on the top line of the expression for equilibrium constant. This form of the equilibrium constant was originally introduced in 1863 by C. M. Guldberg and P. Waage using the law of mass action. They derived the expression by taking the rate of the forward reaction:

k_f [A]^x [B]^y

and that of the back reaction

k_b [C]^z [D]^w.

Since the two rates are equal at equilibrium, the equilibrium constant K_C is the ratio of the rate constants k_f /k_b. The principle that the expression is a constant is known as the equilibrium law or law of chemical equilibrium.

The equilibrium constant shows the position of equilibrium. A low value of K_C indicates that [C] and [D] are small compared to [A] and [B]; i.e. that the back reaction predominates. It also indicates how the equilibrium shifts if concentration changes. For example, if [A] is increased (by adding A) the equilibrium shifts towards the right so that [C] and [D] increase, and K_C remains constant.

For gas reactions, partial pressures are used rather than concentrations. The symbol K_P is then used. Thus, in the example above

K_P = P_C^z P_D^w / P_A^x P_B^y

It can be shown that, for a given reaction K_P =K_C(RT)^Dⁿ where Dv is the difference in stoichiometric coefficients for the reaction (i.e. z + w - x - y). Note that the units of K_P and K_C depend on the numbers of molecules appearing in the stoichiometric equation. The value of the equilibrium constant depends on the temperature. If the forward reaction is exothermic, the equilibrium constant decreases as the temperature rises: if endothermic it increases.

The expression for the equilibrium constant can also be obtained by thermodynamics: it can be shown that the standard equilibrium constant K^O is given by expression (-DG ^O/ RT), where DG^O - is the standard Gibbs free energy change for the complete reaction.

<== попередня сторінка	\|	наступна сторінка ==>
Ammonium hydrogen carbonate	\|	Lesson 13

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