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20-September-2008 09:55:59 - Base chemistry Acids and bases: Acid dissociation constant Acid-base extraction Acid-base reaction Acid-base physiology Acid-base homeostasis Dissociation constant Acidity function Buffer solutions pH Proton affinity Self-ionization of water Acids: Lewis acids Mineral acids Organic acids Strong acids Superacids Weak acids Bases: Lewis bases Organic bases Strong bases Superbases Non-nucleophilic bases Weak bases In chemistry, a base is most commonly thought of as an aqueous substance that can accept protons. This refers to the Brønsted-Lowry theory of acids and bases. Alternate definitions of bases include electron pair donors Lewis, as sources of hydroxide anions Arrhenius and can be commonly thought of as any chemical compound that, when dissolved in water, gives a solution with a pH higher than 7.0. Examples of simple bases are sodium hydroxide and ammonia. Bases can be thought of as the chemical opposite of acids. A reaction between an acid and base is called neutralization. Bases and acids are seen as opposites because the effect of an acid is to increase the hydronium ion H3O+ concentration in water, whereas bases reduce this concentration. Bases react with acids to produce water and salts or their solutions. Contents 1 Definitions 2 Properties 3 Bases and pH 3.1 Common bases 4 Neutralization of acids 5 Alkalinity of non-hydroxides 6 Strong bases 7 Bases as heterogeneous catalysts 8 See also 9 External links 10 References Definitions Main article: acid-base reaction theories A strong base is a base which hydrolyzes completely, raising the pH of the solution towards 14. Strong bases, like strong acids, attack living tissue and cause serious burns. They react differently to skin than acids do, so while strong acids are corrosive, we say that strong bases are caustic. Superbases are a class of especially basic compounds and non-nucleophilic bases are a special class of strong bases with poor nucleophilicity. Bases may also be weak bases such as ammonia, which is used for cleaning. Arrhenius bases are water-soluble and these solutions always have a pH greater than 7. An alkali is a special example of a base, where in an aqueous environment, hydroxide ions also viewed as OH- are donated. There are other more generalized and advanced definitions of acids and bases. The notion of a base as a concept in chemistry was first introduced by the French chemist Guillaume François Rouelle in 1754. He noted that acids which in those days were mostly volatile liquids like acetic acid turned into solid salts only when combined with specific substances. These substances form a concrete base for the salt 1 and hence the name. Properties Some general properties of bases include: Bitter taste opposed to sour taste of acids and sweetness of aldehydes and ketones Slimy or soapy feel on fingers, due to saponification of the lipids in human skin Concentrated or strong bases are caustic corrosive on organic matter and react violently with acidic substances Aqueous solutions or molten bases dissociate in ions and conduct electricity Reactions with indicators: bases turn red litmus paper blue and phenolphthalein red Bases and pH The pH of impure water is a measure of its acidity. In pure water, about one in ten million molecules dissociate into hydronium ions H3O+ and hydroxide ions OH-, according to the following equation: 2H2Ol → H3O+aq + OH-aq The concentration, measured in molarity M or moles per dm³, of the ions is indicated as H3O+ and OH-; their product is the dissociation constant of water with and has the value 10-7 M. The pH is defined as -log H3O+; thus, pure water has a pH of 7. These numbers are correct at 23 °C and slightly different at other temperatures. A base accepts removes hydronium ions H3O+ from the solution, or donates hydroxide ions OH- to the solution. Both actions will lower the concentration of hydronium ions, and thus raise pH. By contrast, an acid donates H3O+ ions to the solution or accepts OH-, thus lowering pH. For example, if 1 mole of sodium hydroxide 40 g is dissolved in water to make 1 litre of solution, the concentration of hydroxide ions becomes OH- = 1 mol/L. Therefore H+ = 10-14 mol/L, and pH = -log 10-14 = 14. Note that in this calculation, it is assumed that the activity is equivalent to the concentration, which is not realistic at concentrations over 0.1 mol dm-3. The base dissociation constant or Kb is a measure of basicity. pKb is the negative log of Kb and related to the pKa by the simple relationship pKa + pKb = 14. Alkalinity is a measure of the ability of a solution to neutralize acids to the equivalence points of carbonates or bicarbonates. Common bases Baking Soda Ammonia Neutralization of acids When dissolved in water, the strong base sodium hydroxide decomposes into hydroxide and sodium ions: NaOH → Na+ + OH- and similarly, in water hydrogen chloride forms hydronium and chloride ions: HCl + H2O → H3O+ + Cl- When the two solutions are mixed, the H3O+ and OH- ions combine to form water molecules: H3O+ + OH- → 2 H2O If equal quantities of NaOH and HCl are dissolved, the base and the acid exactly neutralize, leaving only NaCl, effectively table salt, in solution. Weak bases, such as soda or egg white, should be used to neutralize any acid spills. Neutralizing acid spills with strong bases, such as sodium hydroxide or potassium hydroxide can cause a violent exothermic reaction, and the base itself can cause just as much damage as the original acid spill. Alkalinity of non-hydroxides Bases are generally compounds that can neutralize an amount of acids. Both sodium carbonate and ammonia are bases, although neither of these substances contains OH- groups. Both compounds accept H+ when dissolved in water: Na2CO3 + H2O → 2 Na+ + HCO3- + OH- NH3 + H2O → NH4+ + OH- From this, a pH, or acidity, can be calculated for aqueous solutions of bases. Bases also directly act as electron-pair donors themselves: CO32- + H+ → HCO3- NH3 + H+ → NH4+ Carbon can act as a base as well as nitrogen and oxygen. This occurs typically in compounds such as butyl lithium, alkoxides, and metal amides such as sodium amide. Bases of carbon, nitrogen and oxygen without resonance stabilization are usually very strong, or superbases, which cannot exist in a water solution due to the acidity of water. Resonance stabilization, however, enables weaker bases such as carboxylates; for example, sodium acetate is a weak base. Strong bases A strong base is a basic chemical compound that is able to deprotonate very weak acids in an acid-base reaction. Compounds with a pKa of more than about 13 are called strong bases. Common examples of strong bases are the hydroxides of alkali metals and alkaline earth metals like NaOH and CaOH2. Very strong bases are even able to deprotonate very weakly acidic C-H groups in the absence of water. Hydroxide compounds in order of strongest to weakest:citation needed Potassium hydroxide KOH Barium hydroxide BaOH2 Caesium hydroxide CsOH Sodium hydroxide NaOH Strontium hydroxide SrOH2 Calcium hydroxide CaOH2 Lithium hydroxide LiOH Rubidium hydroxide RbOH The cations of these strong bases appear in the 1st and 2nd groups of the periodic table alkali and earth alkali metals. Group 1 salts of carbanions, amides, and hydrides tend to be even stronger bases due the conjugate acids, which are stable hydrocarbons, amines, and water. Usually these bases are created by adding pure alkali metals such as sodium into the conjugate acid. They are called superbases and it is not possible to keep them in water solution, due to the fact they are stronger bases than the hydroxide ion and as such it will deprotonate the conjugate acid water. For example the ethoxide ion conjugate base of ethanol in the presence of water will undergo this reaction. CH3CH2O- + H2O → CH3CH2OH + OH- Butyl lithium n-BuLi Lithium diisopropylamide LDA C6H14LiN Sodium amide NaNH2 Sodium hydride NaH Bases as heterogeneous catalysts Basic substances can be used as insoluble heterogeneous catalysts for chemical reactions. Examples are metal oxides such as magnesium oxide, calcium oxide, and barium oxide as well as potassium fluoride on alumina and some zeolites. A great deal of transition metals make good catalysts, many of which form basic substances. Basic catalysts have been used for hydrogenations, the migration of double bonds, in the Meerwein-Ponndorf-Verley reduction, the Michael reaction, and many other reactions. See also Acid-base reactions Acids External links Acid-Base equilibrium diagrams, pH calculation and titration curves simulation and analysis - freeware Link References ^ The Origin of the Term Base William B. Jensen Journal of Chemical Education 1130 Vol. 83 No. 8 August 2006 Retrieved from http://en..org/wiki/Base_chemistry Categories: Chemical compounds | BasesHidden categories: All articles with statements | Articles with statements since December 2007 Views Article Discussion this page History Personal tools Log in / create account Navigation Main page Contents Featured content Current events Random article Search Go Search Interaction Community portal Recent changes Contact Donate to Help Toolbox What links here Related changes Upload file Special pages Printable version Permanent link Cite this page Languages اردو Afrikaans العربية Bosanski Català Česky Dansk Deutsch Eesti Ελληνικά Español Esperanto Euskara Français í•œêµì–´ हिनà¥?दी Hrvatski Ido Bahasa Indonesia Ã?slenska Italiano עברית Latina LatvieÅ¡u Lietuvių Magyar МакедонÑ?ки മലയാളം Nederlands 日本語 ‪Norsk bokmÃ¥l‬ ‪Norsk nynorsk‬ Novial Occitan Plattdüütsch Polski Português Română Runa Simi РуÑ?Ñ?кий Simple English SlovenÅ¡Ä?ina СрпÑ?ки / Srpski Srpskohrvatski / СрпÑ?кохрватÑ?ки Suomi Svenska ไทย Tiếng Việt Türkçe УкраїнÑ?ька ä¸æ–‡ This page was last modified on 17 August 2008, at 11:3
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