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Aqueous Reactions Study Guide (page 2)

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Updated on Sep 24, 2011

Acid-Base Reactions

The most general definition of an acid and a base is dependant on water. Svante Arrhenius recognized that certain molecules dissolve in water and produce protons (H+) and hydroxide ions (OH). These substances also react with each other to produce water:

H+ + OH → H2O

Arrhenius defined acids as substances that donate protons in water, and bases as substances that donate hydroxide ions in water. However, not all acids and bases contain components that can donate protons or hydroxide ions. Johannes Brønsted and Thomas Lowry defined acids and bases on the proton; acids donate protons and bases accept protons. The Brønsted-Lowry theory of acid-base chemistry is arguably the most widely used for aqueous reactions.

A third definition is based on the valance electron structures developed by Gilbert N. Lewis and does not involve the components of water. The Lewis definition states that substances that can accept electrons in an aqueous solution are acids, and substances that can donate electrons in an aqueous solution are bases.

Table 6.2 The Three Definitions of Acid-Base

Writing equations of acid-base reactions follows similar rules as the precipitation reactions: Strong electrolytes are soluble and Table 6.1 can identify solids. What acids are strong electrolytes? All strong acids are strong electrolytes, and the remaining weak acids are weak electrolytes. The strong acids are listed in Table 6.3 and can be separated into their ions. However, the remaining weak acids, although soluble, cannot be separated into their ions. These aqueous weak acids are still labeled (aq) but must be treated as a molecule in the ionic and net ionic equations.

Table 6.1 Solubility Rules for Aqueous Solutions

Table 6.1 Solubility Rules for Aqueous Solutions

Table 6.3 The Six Strong Acids

Example 1:

Write the net ionic equation for the reaction of sodium hydroxide and hydrochloric acid.

  • Write the molecular equation: NaOH(aq) + HCl(aq) → NaCl(aq) + HOH(l) (or H2O)
  • Write the ionic equation: Na+(aq) + OH(aq) + H+(aq) + Cl(aq) → Na+(aq) + Cl(aq) + H2O(l)
  • Write the net ionic equation: OH(aq) + H+(aq) → H2O(l)

Example 2:

Write the net ionic equation for the reaction of sodium hydroxide and phosphoric acid.

  • Write the molecular equation: 3NaOH(aq) + H3PO4(aq) → Na3PO4(aq) + 3HOH(l) (or H2O)
  • Write the ionic equation: 3Na+(aq) + 3OH(aq) + H3PO4(aq) → 3Na+(aq) + PO43–(aq) + 3H2O(l)
  • Write the net ionic equation: 3OH(aq) + H3PO4(aq) → PO43–(aq) + 3H2O(l)

Notes

The water molecule is liquid (not aqueous) and not separated in the ionic equation. HCl is a strong acid and separated into its ions.

Phosphoric acid is a weak acid (i.e., not a strong acid) and not separated in the ionic or net ionic equation.

Oxidation-Reduction Reactions

The oxidation state (or oxidation number) for an atom is the number of charges carried by an ion or that an atom would have in a (neutral) molecule if electrons were transferred completely. Oxidation numbers enable the identification of oxidized (an increase in the oxidation number) and reduced (a reduction in the oxidation number) elements.

The rules for assigning oxidation state the following:

  • The oxidation state of an atom in an element at its standard state is always zero.
The atoms in H2(g), O2(g), O3(g), Fe(s), and S8(s) all have an oxidation number of zero.
  • The oxidation state of a monatomic (one atom) ion is its charge.
NaCl: Na is +1 and Cl is –1.
ZnS: Zn is + 2 and S is –2.
  • Oxygen has a –2 oxidation state in compounds, except hydrogen peroxide (H2O2), where O is –1.
SO2; O is –2. Consequently, S is + 4.
N2O5; O is –2. Consequently, N is +5.
H2O2; O is –1 and H is +1.
  • Hydrogen has a +1 oxidation state when it is bonded covalently to nonmetals.
H2S; H is +1. Consequently, S is –2.
CH4; H is +1. Consequently, C is –4.
  • Hydrogen has a –1 oxidation state when it is bonded ionically to metals.
NaH; H is –1 and Na is +1.
CaH2; H is –1 and Ca is +2.
  • Fluorine always has a –1 oxidation state in compounds.
  • The sum of the oxidation states must equal zero for compounds and the net charge for polyatomic ions.
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