chemical properties of acids and bases
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A chemical property is any of a material's properties that becomes evident during a chemical reaction; that is, any quality that can be established only by changing a substance's chemical identity. Simply speaking, chemical properties cannot be determined just by viewing or touching the substance; the substance's internal structure must be affected for its chemical properties to be investigated.
Chemical properties can be contrasted with physical properties, which can be discerned without changing the substance's structure. However, for many properties within the scope of physical chemistry, and other disciplines at the border of chemistry and physics, the distinction may be a matter of researcher's perspective. Material properties, both physical and chemical, can be viewed as supervenient; i.e., secondary to the underlying reality. Several layers of superveniency are possible.
Chemical properties can be used for building chemical classifications.
Examples of chemical properties
- Reactivity against other chemical substances
- Heat of combustion
- Enthalpy of formation
- Chemical stability in a given environment
- Preferred oxidation state(s)
- Coordination number
- Capability to undergo a certain set of transformations, for example molecular dissociation, chemical combination, redox reactions under certain physical conditions in the presence of another chemical substance
- Preferred types of chemical bonds to form, for example metallic, ionic, covalent
For example hydrogen has the potential to ignite and explode given the right conditions. This is a chemical property.
Metals in general do they have chemical properties of reaction with an acid. Zinc reacts with hydrochloric acid to produce hydrogen gas. This is a chemical property.
acids and bases two related classes of chemicals; the members of each class have a number of common properties when dissolved in a solvent, usually water. Properties Acids in water solutions exhibit the following common properties: they taste sour; turn litmus paper red; and react with certain metals, such as zinc, to yield hydrogen gas. Bases in water solutions exhibit these common properties: they taste bitter; turn litmus paper blue; and feel slippery. When a water solution of acid is mixed with a water solution of base, water and a salt are formed; this process, called neutralization , is complete only if the resulting solution has neither acidic nor basic properties. Classification Acids and bases can be classified as organic or inorganic. Some of the more common organic acids are: citric acid , carbonic acid , hydrogen cyanide , salicylic acid, lactic acid , and tartaric acid . Some examples of organic bases are: pyridine and ethylamine. Some of the common inorganic acids are: hydrogen sulfide , phosphoric acid , hydrogen chloride , and sulfuric acid . Some common inorganic bases are: sodium hydroxide , sodium carbonate , sodium bicarbonate , calcium hydroxide , and calcium carbonate . Acids, such as hydrochloric acid, and bases, such as potassium hydroxide, that have a great tendency to dissociate in water are completely ionized in solution; they are called strong acids or strong bases. Acids, such as acetic acid, and bases, such as ammonia, that are reluctant to dissociate in water are only partially ionized in solution; they are called weak acids or weak bases. Strong acids in solution produce a high concentration of hydrogen ions, and strong bases in solution produce a high concentration of hydroxide ions and a correspondingly low concentration of hydrogen ions. The hydrogen ion concentration is often expressed in terms of its negative logarithm, or p H (see separate article). Strong acids and strong bases make very good electrolytes (see electrolysis ), i.e., their solutions readily conduct electricity. Weak acids and weak bases make poor electrolytes. See buffer ; catalyst ; indicators, acid-base ; titration . Acid-Base Theories There are three theories that identify a singular characteristic which defines an acid and a base: the Arrhenius theory, for which the Swedish chemist Svante Arrhenius was awarded the 1903 Nobel Prize in chemistry; the BrÃ¶nsted-Lowry, or proton donor, theory, advanced in 1923; and the Lewis, or electron-pair, theory, which was also presented in 1923. Each of the three theories has its own advantages and disadvantages; each is useful under certain conditions. The Arrhenius Theory When an acid or base dissolves in water, a certain percentage of the acid or base particles will break up, or dissociate (see dissociation ), into oppositely charged ions. The Arrhenius theory defines an acid as a compound that can dissociate in water to yield hydrogen ions, H + , and a base as a compound that can dissociate in water to yield hydroxide ions, OH - Â . For example, hydrochloric acid, HCl, dissociates in water to yield the required hydrogen ions, H + , and also chloride ions, Cl - Â . The base sodium hydroxide, NaOH, dissociates in water to yield the required hydroxide ions, OH - , and also sodium ions, Na + . The BrÃ¶nsted-Lowry Theory Some substances act as acids or bases when they are dissolved in solvents other than water, such as liquid ammonia. The BrÃ¶nsted-Lowry theory, named for the Danish chemist Johannes BrÃ¶nsted and the British chemist Thomas Lowry, provides a more general definition of acids and bases that can be used to deal both with solutions that contain no water and solutions that contain water. It defines an acid as a proton donor and a base as a proton acceptor. In the BrÃ¶nsted-Lowry theory, water, H 2 O, can be considered an acid or a base since it can lose a proton to form a hydroxide ion, OH - , or accept a proton to form a hydronium ion, H 3 O + (see amphoterism ). When an acid loses a proton, the remaining species can be a proton acceptor and is called the conjugate base of the acid. Similarly when a base accepts a proton, the resulting species can be a proton donor and is called the conjugate acid of that base. For example, when a water molecule loses a proton to form a hydroxide ion, the hydroxide ion can be considered the conjugate base of the acid, water. When a water molecule accepts a proton to form a hydronium ion, the hydronium ion can be considered the conjugate acid of the base, water. The Lewis Theory Another theory that provides a very broad definition of acids and bases has been put forth by the American chemist Gilbert Lewis. The Lewis theory defines an acid as a compound that can accept a pair of electrons and a base as a compound that can donate a pair of electrons. Boron trifluoride, BF 3 , can be considered a Lewis acid and ethyl alcohol can be considered a Lewis base.
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Answers:ACIDS- Substances that produce H+ ions in aqueous solution Physical Properties: 1. Acids have a sour taste eg.lemon juice,vinegar 2.they have a pH of less than 7 3.turns blue litmus paper red in color Chemical Properties: 1.Acids react with bases to form salt+water 2.Acids are corrosive in nature 3. Acids release H+ ions in water(aqueuos solution) BASES-Substances that produce OH- ions in aqueous solution Physical Properties: 1.Bases have a bitter taste 2. they have a pH of more than 7 3.turn red litmus paper blue Chemical Properties: 1.they are corrosive as well 2.reacts with acids to form salt+water 3. Bases release OH- ions in a water solution
Answers:From my favorite science website: Chem4kids Acid: A solution that has an excess of H+ ions. It comes from the Latin word acidus that means "sharp" or "sour". Base: A solution that has an excess of OH- ions. Another word for base is alkali. Aqueous: A solution that is mainly water. Think about the word aquarium. AQUA means water. Strong Acid: An acid that has a very low pH (0-4). Strong Base: A base that has a very high pH (10-14). Weak Acid: An acid that only partially ionizes in an aqueous solution. That means not every molecule breaks apart. They usually have a pH close to 7 (3-6). Weak Base: A base that only partially ionizes in an aqueous solution. That means not every molecule breaks apart. They usually have a pH close to 7 (8-10). Neutral: A solution that has a pH of 7. It is neither acidic nor basic. Strong Electrolyte A strong electrolyte is compound that ionizes one hundred percent in solution. Strong acids, bases, and salts are all strong electrolytes. Electrolyte Question: Proof Information Below http://library.thinkquest.org/3659/acidbase/electrolytes.html Certain substances that are called electrolytes produce ions when they dissolve in solution. Because these ions are free to move in solution, the solution conducts electricity. Ions can be produced in solution in either of two ways. Electrolytes can be either ionic compounds (i.e. sodium hydroxide, potassium nitrate) that dissolve in water, giving solutions of ions, or they may be covalent compounds that react with water and form ions in solution as a result. When an ionic substance such as NaCl dissolves in H2O, the water then separates the ions present in the NaCl crystal lattice. This process, known as dissociation, is shown below: Na+Cl-(s) --> Na+(aq) + Cl-(aq) When a polar covalent substance such as HCl dissolves in water, ions are created by the interaction between HCl and H2O molecules. This process, known as ionization is shown below: HCl(g) + H 2O(l) --> H3O+(aq) + Cl-(aq) Lastly, examples of acids and bases (Wikipedia) Perchloric acid HClO4 Hydroiodic acid HI Hydrobromic acid HBr Hydrochloric acid HCl Sulfuric acid H2SO4 (Ka1/first dissociation only) Nitric acid HNO3 Hydronium ion H3O+ or H+. For purposes of simplicity, H3O+ is often replaced in a chemical equation with H+. However, it should be noted that a bare proton simply does not exist in water but instead is bound to one of the lone pairs of electrons on the H2O molecule. This creates the hydronium ion and gives its single O atom a formal charge of +1. Some chemists include chloric acid (HClO3), bromic acid (HBrO3), perbromic acid (HBrO4), iodic acid (HIO3), and periodic acid (HIO4) as strong acids, although these are not universally accepted. Potassium hydroxide (KOH) Barium hydroxide (Ba(OH)2) Caesium hydroxide (CsOH) Sodium hydroxide (NaOH) Strontium hydroxide (Sr(OH)2) Calcium hydroxide (Ca(OH)2) Lithium hydroxide (LiOH) Rubidium hydroxide (RbOH)
Answers:Some people think acidic, means corrosive. This is not entirely true. Yes, the more acidic a solution is, the more corrosive it can be. But, bases can be corrosive to. The acidity or basicity of a solution is determined by the pH scale, or also, the pOH scale. Both scales span from 0-14. On the pH scale, the left side (0-6ish) of the scale is acidic, while the left side(8-14) is basic. We will get more into the terms of basic and acidic later, but for now, lets look at strong and weak acids. Some people think acids are dangerous corrosive liquids that have the ability of melting the flesh from your hand. This is not true, weak acids are used almost every day, in stuff you eat and use. For example; Orange juice contains Citric acid which, with .1M (amount of moles dissolved in 1L of water), has a pH of about 3, which says its in the medium range of acids. There are a bunch of other acids that are used on a daily basis, you can find it in; soda pop, dyes, and vinegar (and many more). A strong acid, that you unknowningly use every day, is Hydrochloric acid. Hydrochloric acid, with 12M(12 moles of HCl/1L water) has a pH of about .5. Very acidic. Hydrochloric acid is found in your stomach, to dissolve your food. Strong acids dissociates completely in water. Did you know water is an acid? Actually, its and acid, and a base. One important piece of information would be that Acids are willing Hydrogen Donors. This means they will willingly give up an H atom (all acids MUST have an H atom; HCl, CH3COOH, HF). Bases willingly accept Hydrogen atoms. With this information, you can determine that Bases and Acids react to neutralize eachother. All bases have OH, so, we can also determine that this chemical equation is true: HCl + KOH -> Cl+K+H20. So a byproduct of acid+base reaction is water. This is just a small bit of information on acids and bases. There is much more out there, and much that we don't know about acids and bases. Keep searching and futher your understanding THE BALANCE EQUATION OF ACIDS AND BASES: Generally, when a reaction occurs between an acid and a base, water and a salt will be produced. Keep that in mind, you'll know then that H2O is almost guaranteed to be a product. The other product will be a salt made up of the anion in the Acid and the cation in the Base. a. HCl + KOH ---> H20 + KCl (see how water is one product and the other product is just what's left over) b. H2SO4 + Mg(OH)2 ---> 2H2O + MgSO4. (Remember the 2 after the OH because OH has a 1- charge and Mg has a 2+ charge. And since you have 2 Hydroxides and 2 Hydrogens as reactants, you'll have 2 waters as products when you balance.) c. H2SO3 + 2NaOH ---> 2H2O + Na2SO3 (Just remember "a light is on in the house" [corny, I know] so since you have sulfurOUS acid, it's a sulfITE compound, and sulfite is SO3 with a 2+ charge. Then just be careful with the charges and balancing) d. HCl + NaOH ---> H2O + NaCl (Shouldn't be too hard). Hope this helps you!
Answers:Bond polarities and the electronegativities of nearby atoms are all important in determining these properties. When something has an -OH group attached to it, if the bond between the O and H is more easily split, it will be an acid, releasing H+ into solution. If the bond between the -OH unit and the rest of the molecule is more easily split, it will be a base, releasing OH- into solution. In the case of NaOH, the electronegativity difference between Na and O is very large, so much so that the bond is ionic. When placed in water, the surrounding water molecules can much more easily pull the Na+ and OH- ions apart, as opposed to separating the O and H, which are covalently bonded. Even though the O-H bond is polar, the nearly complete charges at either end of the Na-O bond are much more readily grabbed by water than the mere partial charges at either end of the O-H bond. In the case of H3PO4, the electronegativity difference between P and O is much smaller, making it a polar covalent bond. In this case, the high electronegativities of the O atoms take center stage. All of them, especially the one that is double-bonded, tend to draw electrons toward themselves. This doesn't do much to the central P atom because it's having electrons only slightly withdrawn from all directions, so there is little net effect on it. The H atoms, however, are having electrons pulled away from them from one net direction by multiple competing O atoms. This weakens the O-H bond, polarizing it more than it already is, and allowing surrounding water molecules to remove just H+. When in doubt about a chemical property, always look to the electrons. When highly electronegative atoms such as O are present, they have a tendency to steal electrons from close neighbors, which affects charges, bond polarities, and bond strengths. In the case of NaOH, the lone O atom steals more electron density from Na than H, separating Na+ from OH-. In the case of H3PO4, the four O atoms working together steal more electron density from the respective H atoms than the P atom, separating H+ from the rest of the molecule.