Acid and Bases

Anion

Negatively charged ion

Cation

Positively charged ion

Acid

Contains a Hydrogen atom and dissolves to form (H+)

Base 

Contains an hydroxide and dissolves in water to form (OH-)

Fact: Arrhenius correcty predicts behavior of many acids and bases but is sometimes inaccurate

-H⁺ does not actually exist in water! Quickly forms hydronium ion

• H⁺ + H₂O → H₃O

Compounds contain no hydroxide anions (OH^-), yet still acts like bases

• Ammonium: NH₃

• Sodium Carbonate: NaCO₃

Johannes Bronsted and Thomas Lowry definition of Acid and Bases

• A Bronsted-Lowrey acid is a proton donor

• A Bronsted-Lowret base is a proton acceptor

Ex: HCl (g) + H₂O → H₃O⁺ (aq) + Cl^- (aq)

the HCL donated its hydrogen to H₂O that resulted in H₃O that is now postively charged: Acid

Cl^- resulted in negative charge: Base

in an acid and base reaction the product should always be neutral 

A Bronsted-Lowry acid must contain a hydrogen atom 

Common anion pairs: HF, HSO₃^-

Organic: CH₃COOH 

Bronsted-Lowry Acid is often written as 

HA

naming an acid that ends with

-ide 

add prefix hydro,and change suffix to ic acid

Cl^- chloride →

HCl hydrochloric acid

naming acids 

-ate 

change suffix to -ic acid 

SO₄^2- sulfate →

H₂SO⁴ sulfuric acid

naming an acid

-ite

chage suffix to ous acid

SO₃^2- sulfite → 

H₂SO₃ sulfurous acid

A Bronsted-Lowry acid may contain one or more protons

• monoprotic: HCl

• diprotic: H₂SO₄

• triprotic: H₃PO₄

A base must contain a lone pair of electrons

ex: NH₃ + H₂O (l) → NH₄⁺ + OH^-

: and often written as B:

NH₃ + H₂O (l) → NH₄⁺ + OH^-

• NH_3: a weak base, and is a lone pair 

• H₂O: acts as an acid

• NH₄⁺: conjugate acid

• OH^-: conjugate base

Common Bronsted-Lowry Bases 

• NaOH

• H₂O

• Mg(OH)₂

• NH₃

• KOH

• Ca(OH)₂