1 / 15

ACIDS and bases

ACIDS and bases. Theories, definitions, neutralization, and pH/ pOH. Dissociation of Water. Hydroxide ion (OH-) – water that loses an H+ Hydronium ion (H 3 O+) – water that gains an H+ Pure water is neutral… for every OH- there is an H+/H 3 O+. Arrhenius Acids and Bases.

zeki
Download Presentation

ACIDS and bases

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. ACIDS and bases Theories, definitions, neutralization, and pH/pOH

  2. Dissociation of Water • Hydroxide ion (OH-) – water that loses an H+ • Hydronium ion (H3O+) – water that gains an H+ • Pure water is neutral… for every OH- there is an H+/H3O+

  3. Arrhenius Acids and Bases • Arrhenius acid – a compound that ionizes to produce H+ • Ex. HCl, H2CO3 • Note: Only H’s in very polar bonds will ionize • Arrhenius bases – compound that ionizes to produce OH- • Ex. NaOH

  4. ______- protic Acids • Monoprotic- acids that donate one H+ • HCl • Diprotic- donates 2 H+ • H2SO4 • Triprotic- donates 3 H+ • H3PO4

  5. Bronsted-Lowry Acids & Bases • Arrhenius doesn’t explain NH3— no hydroxide group! • Bronsted-Lowry Theory: • Acid – proton donor • Base – proton acceptor • Water = amphoteric – acts as both an acid & a base.

  6. Conjugate acid-base Pairs • Conjugate acid-base pair – two substances related by gain or loss of a proton • Conjugate acid –formed when a base gains a proton • Conjugate base –formed when an acid has lost a proton

  7. Practice • For each below, label the acid, base, conjugate acid and conjugate base. • NH3 + H2O  NH4+ + OH- • OH- + H3O+ H2O + H2O • HPO4-2 + H2O PO4-3 + H3O+

  8. Strong vs. Weak • STRONG acids ionize completely in water • High [H+] • more polar = easily ionized • Ex. HCl • WEAK acids do not ionize completely • low [H+] • Ex. Acetic Acid (HC2H3O2), organic acids • STRONGbases completely dissociate (ionize) • High [OH-] • Ex. NaOH • WEAKbases don’t completely dissociate • low [OH-] • sometimes not very soluble in water • Ex. NH3, covalently bonded bases

  9. STRONG ACIDS: • HNO3 nitric acid • HClhydrochloric acid • H2SO4 sulfuric acid • HClO4perchloric acid • HBrhydrobromicacid • HI hydroiodic acid • STRONG BASES: • RULE:hydroxides with alkali or alkaline earth metals (except Be) • Ex:NaOH, Ca(OH)2, KOH

  10. Neutralization • Neutralization Reactions: • Strong acid and strong base react in water to produce a salt and water – pH of ~7 • Ex. • HCl + NaOH • H2SO4 + KOH • These are a common way to prepare a pure salt. • How could you prepare low-sodium table salt (KCl)? • Balancing: • H+ and OH- react in a one to one ratio (H2O). BUT, not all compounds ionize just once. (ex: H2SO4 , Ca(OH)2 ) • Must use appropriate mole ratio.

  11. Neutralization Practice • How many moles of sodium hydroxide are required to neutralize 0.60 mol of nitric acid? • How many moles of potassium hydroxide are needed to neutralize 2.63 mol of phosphoric acid? • What is the molarity of phosphoric acid if 18.0 mL of the solution is neutralized by 35.5 mL of 0.15 M NaOH?

  12. Titrations • Titrations: • The molarity of an unknown acid/base can be determined by the carefully measured addition of a base/acid of known molarity. • Standard solution – the solution of known molarity • Equivalence point (S) – the point at which the two solutions are in chemically equivalent amounts (ie. They neutralize) • Indicator- Solution that changes color at a certain pH • End point – the point at which the indicator changes color • Which indicator used determines if the end point and equivalence point are the same.

  13. pH & pOH Calculations • Calculating pH: • pH = -(log [H+]) • ranges from 0-14. • 7 = neutral • pH < 7 = acid (high [H+]) • pH > 7 = base (low [H+]) • Calculating pOH • pOH = -(log [OH-]) • ranges from 0-14. • 7 = neutral • pOH < 7 = base (high [OH-]) • pOH > 7 = acid (low [OH-]) • pOH + pH = 14

  14. Practice • Determine the pH for the following solutions: • [H+] = 1 x 10-6M • [H+] = 0.0001 M • [OH-] = 3 x 10-2M • [OH-] = 4.3 x 10-11 M

  15. Calculating Molarity • Calculating Concentration(M) from pH: • [ ]=molarity/concentration • [H+] = 10-pH • [OH-] = 10-pOH • Practice: • Calculate the [H+] for the following solutions: • pH = 3.0 • pOH = 6.0 • pH = 7.4 • pOH = 1.34

More Related