Concentration terms and units - Concentration, % , ppm, Density, Solutions, Solubility, Mixtures, Water Solutions, Aqueous - Engineering - http://maxpages.com/jpchemical/Basic_Chemistry

Concentration: Definition The concentration of a chemical substance expresses the amount of a substance present in a mixture. There are many different ways to express concentration. Concentration measures the amount of one substance (the solute) contained within a unit measure of a mixture (the solution). Chemists use the term solute to describe the substance of interest and the term solvent to describe the material in which the solute is dissolved. For example, in a can of soda pop (a solution of sugar in carbonated water), there are approximately twelve tablespoons of sugar (the solute) dissolved in the carbonated water (the solvent). In general, the component that is present in the greatest amount is termed the solvent. For example, a concentration of 2.7 millilitres per cubic metre would mean that 2.7 mL of the solute are contained in every cubic metre of the solution. Note that, in the calculator, the fluid ounces [fl.oz] are (UK) or (US) to match the gallon. Additional Info It can also be expressed as mass per unit mass. (mg/kg etc.) Another common form for concentration is to give it: in parts per million [ppm]; parts per thousand [‰]; or parts per hundred or 'per cent' [%]. There are many ways to express concentrations. Some of the more common concentration units are: Mass per unit volume. Some MSDS's use milligrams per milliliter (mg/mL) or milligrams per cubic centimeter (mg/cm3). Note that 1 mL = 1 cm3 and that cm3 is sometimes denoted as a "cc" (see volume units and mass units). Mass per unit volume is handy when discussing how soluble a material is in water or a particular solvent. For example, "the solubility of substance X is 3 grams per liter". Percent by Mass. Also called weight percent or percent by weight, this is simply the mass of the solute divided by the total mass of the solution and multiplied by 100%: The mass of the solution is equal to the mass of the solute plus the mass of the solvent. For example, a solution consisting of 30 grams of sodium chloride and 70 grams of water would be 30% sodium chloride by mass: [(30 g NaCl)/(30 g NaCl + 70 g water)] * 100% = 30%. To avoid confusion whether a solution is percent by weight or percent by volume, the symbol "w/w" (for weight to weight) is often used after the concentration such as "10% potassium iodide solution in water (w/w)". Percent by Volume. Also called volume percent or percent by volume, this is typically only used for mixtures of liquids. Percent by volume is simply the volume of the solute divided by the sum of the volumes of the other components multiplied by 100%: If we mix 30 mL of ethanol and 70 mL of water, the percent ethanol by volume will be 30% BUT the total volume of the solution will NOT be 100 mL (although it will be close). That's because ethanol and water molecules interact differently with each other than they do with themselves. To avoid confusion whether we have a percent by weight or percent by volume solution, we could label this as "30% ethanol in water (v/v)" where v/v stands for "volume to volume". Molarity. Molarity is the number of moles of solute dissolved in one liter of solution. For example, if we have 90 grams of glucose (molar mass = 180 grams per mole) this is (90 g)/(180 g/mol) = 0.50 moles of glucose. If we place this in a flask and add water until the total volume = 1 liter we would have a 0.5 molar solution. Molarity is usually denoted with an italicized capital M, i.e. a 0.50 M solution. Recognize that molarity is moles of solute per liter of solution, not per liter of solvent!! Also recognize that molarity changes slightly with temperature because the volume of a solution changes with temperature. Molality. Molality is the number of moles of solute dissolved in one kilogram of solvent. Notice the two key differences between molarity and molality. Molality uses mass rather than volume and uses solvent instead of solution. Unlike molarity, molality is independent of temperature because mass does not change with temperature. If we were to place 90 grams of glucose (0.50 moles) in a flask and then add one kilogram of water we would have a 0.50 molal solution. Molality is usually denoted with a small italicized m, i.e. a 0.50 m solution. Note: m also has other possible meaninsg on MSDS's, so look at the context carefully. Parts per million (PPM). Parts per million works like percent by mass, but is more convenient when there is only a small amount of solute present. PPM is defined as the mass of the component in solution divided by the total mass of the solution multiplied by 106 (one million): A solution with a concentration of 1 ppm has 1 gram of substance for every million grams of solution. Because the density of water is 1 g per mL and we are adding such a tiny amount of solute, the density of a solution at such a low concentration is approximately 1 g per mL. Therefore, in general, one ppm implies one milligram of solute per liter of solution. Finally, recognize that one percent = 10,000 ppm. Therefore, something that has a concentration of 300 ppm could also be said to have a concentration of (300 ppm)/(10,000 ppm/percent) = 0.03% percent by mass. Parts per billion (PPB). This works like above, but we multiply by one billion (109; caution: the word billion has different meanings in different countries). A solution with 1 ppb of solute has 1 microgram (10-6 g) of material per liter. Parts per trillion (PPT). This works like parts per million and parts per billion except that we multiply by one trillion (1012). There are few, if any, solutes which are harmful at concentrations as low as 1 ppt. Note: "ppt" is sometimes used as laboratory shorthand for precipitate, which is entirely unrelated. Mixture: Definition According to the OSHA Hazard Communication Standard, 29 CFR 1910.1200, a mixture is "any combination of two or more chemicals if the combination is not, in whole or in part, the result of a chemical reaction." A chemist defines a mixture as a combination of two or more substances in which each substance retains its own chemical identity and properties. In general, mixtures have no fixed composition. The amount of one or more components (substances) can usually vary over a wide range. Additional Info Mixtures generally fall into one of two categories: A homogeneous mixture has uniform chemical composition, appearance and properties throughout. A simple example is air, which is a homogeneous mixture of gases consisting primarily of nitrogen and oxygen. Another example would be dissolving a small amount of sugar in a glass of water. After stirring, every section of the aqueous solution would be identical in composition, appearance and physical properties (such as boiling point, for example). A heterogenous mixture has different compositions, appearance and properties at various points in the mixture. A simple example is a chocolate chip cookie; at some points we encounter a chocolate chip and at others we encounter cookie dough. Another example is a mixture of ice and water; at one point we might encounter a solid, but at another we might have a liquid. An important property of mixtures is that they can (usually) be separated into their individual components without requiring any chemical reactions. For example, we can evaporate our sugar-water solution to obtain pure water or pure sugar. Likewise, we could put our chocolate chip cookie into water to dissolve away the dough and collect pure chocolate chips. A microscopic view When chemicals are combined, they can either form a mixture or they can chemically react to form new chemical species. For example, consider atoms A (blue) and B (red). If they do not react chemically, we obtain a mixture: If each atom of A reacts with exactly one atom of B, we can form a new molecule AB that contains one atom of A chemically bonded to one atom of B. This is a chemical reaction because the new material will have different physical and chemical properties than either A or B. Solutions: Solubility What is a solution? Solutions are homogeneous mixtures of two or more pure substances. A homogeneous mixture is a physical combination of two or more pure substances whose distribution throughout the mixture is uniform. What this means is that if we were to make a solution and take only a portion of the solution at random called an aliquot, the proportion of each pure substance in the aliquot would be the same as the proportion of that pure substance in the whole solution. We call these proportions to the whole solution the concentration. Most aqueous solutions involving liquid or solid solutes will have endothermic Heats of Solution. However, a few are exceptions to this statement. Other solutions involving gaseous solutes in water will release thermal energy during the solution formation process. These solutions are said to have an exothermic Heat of Solution. Solute(g) + Solvent -----> Solution + Thermal Energy Solutions can be unsaturated, saturated, or supersaturated. Unsaturated solutions are those that are below the solubility limits of the solute in that solvent. Saturated solutions are those that are at the solubility limits. Supersaturated solutions are those solutions that are above the solubility limits. Solubility of Solutions The solubility of a particular solute in a solvent is the maximum amount of solute that will dissolve in a specified amount of solution or solvent. It represents the saturated level of the solution where no more solute will dissolve within the solution. This saturated condition creates a dynamic physical equilibrium between the solute and solvent and the solution: solute + solvent = solution Such dynamic equilibrium involve two processes, a forward process and a reverse process. When the rate of the forward process is equal to the rate of the reverse process, then the system is said to be in a dynamic equilibrium. Dynamic Equilibria will involve two opposing processes occuring simultaneously. We can prove this by taking a salt crystal and chipping off one end. Then we suspend the deformed cystal in a saturated solution of the salt and allow the saturated solution to be in contact with the deformed cystal for several weeks. When we come back we will have discovered that the deformed crystal will have been reformed. This could only happen if the solution process and its reverse process, the dissolution process, were occuring simultaneously. Aqueous : Definition Aqueous refers to a solution in water. A more exact definition is a solution in which the solute (the substance dissolved) initially is a liquid or a solid and the solvent is water. Additional Info Aqueous solutions fall into three general categories based on how well they conduct electricity: Strong electrolytes when dissolved in water dissociate completely into ions and conduct electricity. For example, sodium chloride, NaCl, dissociates into Na+ and Cl- ions in water. Other examples of strong electrolytes are nitric acid (HNO3) and sodium hydroxide (NaOH). Weak electrolytes when dissolved in water do not dissociate to any large extent and therefore do not conduct electricity very well. Examples include ammonia (NH3) and acetic acid (CH3COOH). Non-electrolytes do not dissociate to ions in water and do not conduct electricity. Examples include sugar (sucrose = C12H22O11), ethanol (CH3CH2OH) and methanol (CH3OH). MSDS Relevance The term aqueous may appear on an MSDS in terms of a special instruction (such as first aid procedures), a particular chemical property, or a chemical incompatibility. Aqueous solutions are not usually flammable, but may be able to carry toxic materials into your body through skin contact or ingestion. Be careful with terminology. A solution of ammonia gas (NH3) in water is often called ammonium hydroxide, NH4OH, ammonia water, or simply ammonia. Do not confuse this aqueous solution with ammonia gas (anhydrous ammonia)!

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