Ch 14.1 Properties of Acids and Bases

Acids  Are sour to taste  React with bases to produce salts and water.  React with metals and release H 2 gas  Turn litmus paper red  Conduct an electric current.  Examples: citrus fruit, tomatoes, HCl, and stomach acid

Acid Nomenclature  Binary Acid: an acid that contains only 2 different elements: Hydrogen &one of the more electronegative elements.  Naming: begins with hydro- followed by the root of the 2nd element & ends in -ic.  Oxyacid: an acid that is a compound of hydrogen, oxygen, and third element, usually a nonmetal.

Common Industrial Acids  Sulfuric Acid  Nitric Acid  Phosphoric Acid  Hydrochloric Acid  Acetic Acid

Bases  Are bitter to taste  Feel slippery  React to neutralize acids, forming salt and water  Turn litmus paper blue  Conduct an electric current.  Examples: Baking Soda, Soaps, Ammonia, and NaOH

Arrhenius Acid  Produces H + in an aqueous solution. Once present, the H + combines with water to form H 3 O + (called hydronium).  The dissociation of HCl looked like: HCl  H + + Cl -  But it was actually: HCl + H 2 O  Cl - + H 3 O +

Arrhenius Base  Produces OH - in solution.  Examples: NaOH  Na + + OH - NH 3 + H 2 O  NH OH -

Dissociation  The strength of an acid or base depends on the amount of dissociation that occurs. This depends on the polarity of the bond and the ease at which the bond can break.  Organic Acids containing carboxyl/acid groups COOH, like vinegar, are generally weak.

Ch 14.2 Acid-Base Theories

Bronsted-Lowry Acids & Bases  Do not require aqueous solutions.  Acids are defined as proton (H + ) donors and bases are proton (H + ) acceptors.  Examples: HCl + NH 3  NH Cl - A B H 2 O + NH 3  NH OH - A B

Monoprotic and Polyprotic Acids  Monoprotic: An acid that can donate only one proton (H + ) per molecule. HCl + H 2 O  H 3 O + + Cl -  Polyprotic: An acid that can donate more than one proton (H + ) per molecule. Diprotic: Can donate 2 protons. H 2 SO 4 Triprotic: Can donate 3 protons. H 3 PO 4

Lewis Acids and Bases  Lewis of electron dot structure fame  Where both Arrhenius and Bronsted- Lowry acids contain or produce H +, Lewis acids don’t have to.  An acid is defined as an atom, ion, or molecule that accepts an electron pair to form a covalent compound.  A base donates an electron pair.

Lewis Acids and Bases  They can exist in solid, liquid, or gas phase.  Examples H + + NH 3  NH 4 + H 2 O + HCl  Cl - + H 3 O + BF 3 + F -  BF 4 -

Ch 14.3 Acid-Base Reactions

Conjugate Acids and Bases  When a Bronsted-Lowry acid gives up a proton, it becomes a B-L Base because it can then accept a proton. Reactions usually involve two acid- base pairs. HCl + NH 3  NH Cl - A B CACB H 2 O + NH 3  NH OH - A B CACB

 In general, strong acids have weak conjugate bases and weak acids have strong conjugate bases.  Proton-transfer reactions favor the production of the weaker acid and weaker base.

Amphoteric Substances  Substances which can be an acid or a base.  Whether it acts as one or the other depends on the strength of what it reacts with.  Example: H 2 O: H 2 O + NH 3  NH OH - H 2 SO 4 + H 2 O  H 3 O + + HSO 4 -

Neutralization Reactions  A base neutralizes an acid forming salt and water. But things are not quite so simple. The return of the SPECTATOR ION! HCl + NaOH  NaCl + H 2 O  Base Dissociation NaOH  Na + + OH -  Acid Dissociation HCl +H 2 O  Cl - + H 3 O +  Then write Overall and then Net Ionic.

Acid Rain  Rain that is very acidic  Industrial processes produce gases such as NO, NO 2, CO 2, SO 2, and SO 3. These can dissolve in water producing acidic solutions that fall to the ground in the form of rain. This Acid Rain can erode statues and affect ecosystems.