2014-21-09 tabs

Le Châtelier's Principle. Le Châtelier's principle states that if a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium shifts to counteract the change to reestablish an equilibrium.

If a chemical reaction is at equilibrium and experiences a change in pressure, temperature, or concentration of products or reactants, the equilibrium shifts in the opposite direction to offset the change. This page covers changes to the position of equilibrium due to such changes and discusses briefly why catalysts have no effect on the equilibrium position. Introduction An action that changes the temperature, pressure, or concentrations of reactants in a system at equilibrium stimulates a response that partially offsets the change while a new equilibrium condition is established (2).

Hence, Le Châtelier's principle states that any change to a system at equilibrium will adjust to compensate for that change. Concentration Changes. Gibbs–Helmholtz equation. The Gibbs–Helmholtz equation is a thermodynamic equation useful for calculating changes in the Gibbs energy of a system as a function of temperature.

It is named after Josiah Willard Gibbs and Hermann von Helmholtz. Equation Nernst equation. In electrochemistry, the Nernst equation is an equation that relates the reduction potential of a half-cell (or the total voltage (electromotive force) of the full cell) at any point in time to the standard electrode potential, temperature, activity, and reaction quotient of the underlying reactions and species used.

When the reaction quotient is equal to the equilibrium constant of the reaction for a given temperature, i.e. when the concentration of species are at their equilibrium values, the Nernst equation gives the equilibrium voltage of the half-cell (or the full cell), which is zero; at equilibrium, Q=K, ΔG=0, and therefore, E=0. It is named after the German physical chemist who first formulated it, Walther Nernst.[1][2] Watch California Dry Up Right Before Your Eyes In 6 Jaw-Dropping GIFs. “This is a big deal,” California Governor Jerry Brown said at a ceremony Tuesday as he signed into law a trio of bills regulating, for the first time, the state’s groundwater use.

As of Thursday, almost 60 percent of the state is facing “exceptional drought,” the most severe level of dryness measured by the U.S. Drought Monitor. But if you’re not living in a community dependent on bottled water rations, farming land that’s projected to lose \$800 million in crop revenue or watching raging wildfires ravage your drought-parched town, the historic California drought may still feel like little more than a headline. To fully grasp how desperate California is for relief, we’ve created six before-and-after GIFs that will show you how badly the drought has dehydrated the state in just the last three years.

The Green Bridge passes over full water levels near Bidwell Marina on July 20, 2011, in Oroville, California, and much lower levels on Aug. 19, 2014. Also on HuffPost: Getty Images. 'We're on edge of large war in MidEast' - Daniel Ellsberg. Lecture-8.pdf. Phase rule. Gibbs' phase rule[1][2] was proposed by Josiah Willard Gibbs in his landmark paper titled On the Equilibrium of Heterogeneous Substances, published from 1875 to 1878.

The rule is the equality The number of degrees of freedom is the number of independent intensive variables, i. e. the largest number of properties such as temperature or pressure that can be varied simultaneously and arbitrarily without affecting one another. An example of one-component system is a system involving one pure chemical, while two-component systems, such as mixtures of water and ethanol, have two chemically independent components, and so on. Typical phases are solids, liquids and gases.

Foundations The basis for the rule (Atkins and de Paula,[2] justification 6.1) is that equilibrium between phases places a constraint on the intensive variables. To be more specific, the composition of each phase is determined by C – 1 intensive variables (such as mole fractions) in each phase. Two-component systems Phase diagram. Overview Common components of a phase diagram are lines of equilibrium or phase boundaries, which refer to lines that mark conditions under which multiple phases can coexist at equilibrium.

Phase transitions occur along lines of equilibrium. Triple points are points on phase diagrams where lines of equilibrium intersect. Triple points mark conditions at which three different phases can coexist. Miscibility. By contrast, substances are said to be immiscible if a significant proportion does not form a solution.

Otherwise, the substances are considered miscible. For example, butanone is significantly soluble in water, but these two solvents are not miscible because they are not soluble in all proportions. Organic compounds In organic compounds, the weight percent of hydrocarbon chain often determines the compound's miscibility with water. For example, among the alcohols, ethanol has two carbon atoms and is miscible with water, whereas 1-octanol with eight carbons is not.  There also exist metals that are immiscible in the liquid state.

Effect of entropy Substances with extremely low configurational entropy, especially polymers, are likely to be immiscible in one another even in the liquid state. Determination Euler characteristic. In mathematics, and more specifically in algebraic topology and polyhedral combinatorics, the Euler characteristic (or Euler–Poincaré characteristic) is a topological invariant, a number that describes a topological space's shape or structure regardless of the way it is bent.

It is commonly denoted by Polyhedra The Euler characteristic was classically defined for the surfaces of polyhedra, according to the formula This result is known as Euler's polyhedron formula or theorem. Vapor–liquid equilibrium. Vapor–liquid equilibrium (VLE) is a condition where a liquid and its vapor (gas phase) are in equilibrium with each other, a condition or state where the rate of evaporation (liquid changing to vapor) equals the rate of condensation (vapor changing to liquid) on a molecular level such that there is no net (overall) vapor–liquid interconversion.

A substance at vapor–liquid equilibrium is generally referred to as a saturated fluid. For a pure chemical substance this implies that it is at its boiling point.[1] The notion of "saturated fluid" includes saturated liquid (about to vaporize), saturated liquid–vapor mixture, and saturated vapor (about to condense). Although in theory equilibrium takes forever to reach, such an equilibrium is practically reached in a relatively closed location if a liquid and its vapor are allowed to stand in contact with each other long enough with no interference or only gradual interference from the outside.

VLE data introduction ; and where and x1 + x2 = 1. Relative volatility. Relative volatilities are not used in separation or absorption processes that involve components reacting with each other (for example, the absorption of gaseous carbon dioxide in aqueous solutions of sodium hydroxide).

Definition For a liquid mixture of two components (called a binary mixture) at a given temperature and pressure, the relative volatility is defined as When their liquid concentrations are equal, more volatile components have higher vapor pressures than less volatile components. Clausius–Clapeyron relation. Latent heat. Latent heat is the energy released or absorbed by a body or a thermodynamic system during a constant-temperature process. A typical example is a change of state of matter, meaning a phase transition such as the melting of ice or the boiling of water.[1][2] The term was introduced around 1762 by Scottish chemist Joseph Black.

It is derived from the Latin latere (to lie hidden). Black used the term in the context of calorimetry when referring to the heat transferred that caused a change of volume while the thermodynamic system was held at constant temperature. In contrast to latent heat, an energy is called a sensible energy or heat when it causes processes that do result in a change of the temperature of the system.

Usage Specific volume. Van 't Hoff equation. Van 't Hoff factor. The van 't Hoff factor (named after J. H. van 't Hoff) is a measure of the effect of a solute upon colligative properties such as osmotic pressure, relative lowering in vapor pressure, elevation of boiling point and freezing point depression. The van 't Hoff factor is the ratio between the actual concentration of particles produced when the substance is dissolved, and the concentration of a substance as calculated from its mass.

For most non-electrolytes dissolved in water, the van' t Hoff factor is essentially 1. Gibbs–Helmholtz equation. Enthalpy. Enthalpy is a defined thermodynamic potential, designated by the letter "H", that consists of the internal energy of the system (U) plus the product of pressure (P) and volume (V) of the system:[1] Since enthalpy, H, consists of internal energy, U, plus the product of pressure (P) and the volume (V) of the system, which are all functions of the state of the thermodynamic system, enthalpy is a state function. The unit of measurement for enthalpy in the International System of Units (SI) is the joule, but other historical, conventional units are still in use, such as the British thermal unit and the calorie.

The enthalpy is the preferred expression of system energy changes in many chemical, biological, and physical measurements, because it simplifies certain descriptions of energy transfer. Enthalpy change accounts for energy transferred to the environment at constant pressure through expansion or heating.

The total enthalpy, H, of a system cannot be measured directly. Origins where.