Acid and Bases

• Acidic in taste
• Acids disassociate one proton (H⁺) at a time
• Types of acids: Monoprotic¹⁾, Diprotic²⁾, and Polyprotic³⁾
• Bronsted-Lowry Acids are important and they donate H⁺ in water

• Bitter/poisonous
• Slippery when wet
• Most cleaners are basic == alkaline
• Bronsted-Lowry Bases are similar to their acid counterparts, they accept an H⁺

• Strong acids/bases completely ionize or disassociate in water, weak ones do not
• Less than 1% of a weak acid/base will dissociate

Note: Weak acids are extremely dangerous, HF on your skin will put amputation in your near future.

• HCl
• HBr
• HI
• HlO₄
• H₂SO₄
• HNO₃

All group I + II Hydroxides are strong bases⁴⁾

Also called Amphoteric, these are ions that behave as either acids or bases. Water is an example of one.

pH is a logarithmic scale to measure the acidity of a solution. Acids are between 0-7, bases are 7-14, and an exact value of 7 represents a neutral substance.

These equations can be used to calculate the pH of a substance. pOH is useful for determining the basicness of a solution.

• $\text{pH} = -\log{H^+}$
• $\text{pOH} = -\log{OH^-}$

Taking the inverse log⁵⁾ can return the amount of H⁺ and OH^- as indicated by the formulas above.

• $H^+ = 10^{-\text{pH}}$
• $OH^- = 10^{-\text{pOH}}$

Since the pH scale goes between 0 and 14, $\text{pH} + \text{pOH} = 14$. Also, using the rules in the section above, it can also be determined that $H^+ \times OH^- = 1 \times 10^{-14}$. This all assumes that the temperature is 25 °C. You can still use this if the temperature isn't stated.