Arrhenius Theory, Acid-Base Naming, and Neutralization Reactions

Acids
- Taste: Acids have a characteristic sour taste.
- Chemical Cause: The sour taste and acidic properties are due to the presence of $H^{+}$ in water.
- Chemical Formulas: Acids are often identified by a hydrogen symbol written at the front of the chemical formula (e.g., $HCl$ , $HNO_{3}$ , $H_{2}SO_{4}$ ).
- Acidic vs. Non-Acidic Hydrogens: In formulas like acetic acid, the acidic hydrogen is written at the front to signify it will leave the compound, while non-acidic hydrogens are written later in the formula.
Bases
- Tactile Properties: Bases feel soapy or slippery to the touch.
- Common Uses: Most cleaning products are bases.
- Chemical Behavior: Bases increase the concentration of $OH^{-}$ in water.
- Formula Identification: While many bases have $OH^{-}$ in their formula (like $NaOH$ , $KOH$ , $LiOH$ , $Sr(OH)_{2}$ ), not all do.
- Ammonia ( $NH_{3}$ ): Often written as $NH_{4}OH$ to signify its nature as a base.
- Ionic Requirement: Not all compounds containing $-OH$ groups are basic. To be a base, the compound must be ionic so that the $OH^{-}$ ion can dissociate/leave the compound.

Historical Context: Proposed by Arrhenius in 1887.
Acid Definition: Arrhenius realized that all acids dissociate in water to produce $H^{+}$ .
Hydronium Ion Formation: The produced $H^{+}$ reacts with water ( $H_{2}O$ ) to create the hydronium ion ( $H_{3}O^{+}$ ).
Quantitative Dissociation Example: If you add $1.0\,M$ $HCl$ to water, you yield $1.0\,M$ $H_{3}O^{+}$ because $HCl$ is a strong acid.
Strong Bases: These are strong electrolytes that dissociate $100\%$ . They are measured by the concentration of $OH^{-}$ they produce in solution.

An acid is considered strong if it dissociates $100\%$ in solution. Any acid not on this list is automatically considered a weak acid.
The List of Strong Acids:
1. Hydrochloric Acid ( $HCl$ ): Dissociates into $H^{+}$ and $Cl^{-}$ . It is commercially available as "mutacritic acid."
2. Hydrobromic Acid ( $HBr$ ): Dissociates into $H^{+}$ and $Br^{-}$ .
3. Hydroiodic Acid ( $HI$ ): Dissociates into $H^{+}$ and $I^{-}$ .
4. Nitric Acid ( $HNO_{3}$ ): Dissociates into $H^{+}$ and $NO_{3}^{-}$ .
5. Sulfuric Acid ( $H_{2}SO_{4}$ ): Contains two hydrogens at the front, though both may not dissociate equally (the second hydrogen's behavior is discussed later).
6. Chloric Acid ( $HClO_{3}$ ): Contains the chlorate ion.
7. Perchloric Acid ( $HClO_{4}$ ): Contains the perchlorate ion.
Reactivity and Safety:
- At the same concentration, all seven acids provide the same amount of $H^{+}$ .
- Chloric and perchloric acids are extremely oxidizing and potentially explosive.
- The first five acids ( $HCl$ , $HBr$ , $HI$ , $HNO_{3}$ , $H_{2}SO_{4}$ ) are more commonly used in General or Inorganic Chemistry because they are cheaper, less reactive (aside from their acidity), and less explosive.

General Logic: Acid naming is based on the negative ion (anion) left behind when the $H^{+}$ ion is removed.
The "-ate" Rule: If the anion ends in "-ate," the acid name ends in "-ic acid."
- Example: $HNO_{3}$ contains the nitrate ion ( $NO_{3}^{-}$ ) $\rightarrow$ Nitric acid.
- Example: $OCN^{-}$ (cyanate ion) $\rightarrow$ Cyanic acid.
- Example: Carbonate ion ( $CO_{3}^{2-}$ ) $\rightarrow$ Carbonic acid.
The "-ite" Rule: If the anion ends in "-ite," the acid name ends in "-ous acid."
- Example: $H_{2}SO_{3}$ containing the sulfite ion ( $SO_{3}^{2-}$ ) $\rightarrow$ Sulfurous acid.
The "-ide" Rule (Hydro- acids): This applies to acids containing no oxygen. The name uses the prefix "hydro-" and the suffix "-ic acid."
- Example: $HCl$ (chloride ion) $\rightarrow$ Hydrochloric acid.
- Example: $HI$ (iodide ion) $\rightarrow$ Hydroiodic acid.
- Example: Cyanide ( $CN^{-}$ ) $\rightarrow$ Hydrocyanic acid.
Special Pronunciation Rules (Sulfur and Phosphorus):
- For Sulfur acids, add the syllable "-ur-". Instead of "sulphic acid," it is sulfuric acid. Instead of "sulfous acid," it is sulfurous acid.
- For Phosphorus acids, add the syllable "-ur-". Instead of "phosphous acid," it is phosphorous acid. Instead of "phosphic acid," it is phosphoric acid.

Strong bases are ionic salts paired with Group 1 or Group 2 metal ions.
Group 1 Hydroxides: Formed with metals like Sodium ( $Na$ ), Lithium ( $Li$ ), Potassium ( $K$ ), Rubidium ( $Rb$ ), and Cesium ( $Cs$ ).
- Commonality: Sodium hydroxide ( $NaOH$ ) is the cheapest and most common. Potassium hydroxide ( $KOH$ ) is also frequently used.
- Naming: Simply use the ionic name: Sodium hydroxide, Lithium hydroxide, etc.
Group 2 Hydroxides: Specifically Calcium ( $Ca$ ), Strontium ( $Sr$ ), and Barium ( $Ba$ ).
- Stoichiometry: These metals have a $+2$ charge, meaning there are two hydroxides per metal atom (e.g., $Ca(OH)_{2}$ ). Mole for mole, these provide twice as much base as Group 1 bases.
- Solubility Limitations: These bases have very low solubility compared to Group 1 bases (which can easily reach solution concentrations of $5$ to $10\,M$ ).
- Magnesium Hydroxide ( $Mg(OH)_{2}$ ): Technically should be a strong base, but its solubility is too low to be tested.

Basic Reaction Formula: Acid + Base $\rightarrow$ Salt + Water.
Net-lonic Perspective: The fundamental chemical change in these reactions is the combination of hydrogen ions and hydroxide ions.
- Equation: $H^{+} + OH^{-} \rightarrow H_{2}O$
Case Study: $NaOH$ and $HCl$
1. $NaOH$ and $HCl$ are both strong electrolytes.
2. In solution, they exist as $Na^{+}$ , $OH^{-}$ , $H^{+}$ , and $Cl^{-}$ .
3. The $H^{+}$ and $OH^{-}$ react completely to form $H_{2}O$ .
4. The $Na^{+}$ and $Cl^{-}$ ions do not change; they start as ions in solution and end as ions in solution (spectator ions). Since there is no real change, they are crossed out of the net-ionic equation.
Stoichiometric Examples:
- One acidic hydrogen + one $OH^{-}$ creates one water molecules.
- Two acidic hydrogens + two $OH^{-}$ creates two water molecules.
- Phosphoric acid ( $H_{3}PO_{4}$ ) + $3\,NaOH$ yields three water molecules and a salt ( $Na_{3}PO_{4}$ ).