# Acids and Bases

Acids and bases are compounds that one or more "loose" hydrogens (H) or hydroxides (OH) attached to them. These hydrogens or hydroxides do not have a very strong bond to the compound, and they will break free very easily. It is the hydrogens and hydroxides, that give acids and bases their properties.

# How Our View of Acids and Bases Changed

People have noticed differences between acids and bases for ages. Differences between the two had been noticed since the middle ages. Discoveries about the pH scale and acid-base reactions began in the early 1900s.
Most of the early theories about acids and bases weren’t right but at least they were trying. In the mid-seventeen hundreds, a guy named Antoine Lavoisier came up with a theory. He concluded that oxygen and water made things acidic. Sulfuric acid was H$_2$SO$_4$ which could be made of H$_2$O and SO$_3$. So He thought it oxygen made it acidic. This wasn’t the case. Some solutions with oxygen in them, like CaO and are basic. There are also acids without oxygen in them.
A guy named Sir Humphrey Davy concluded that Hydrogen was responsible for acidity. Another guy by the name of Justus von Liebig agreed with Humphrey and elaborated the theory. He hypothesized that hydrogen can form ionic bonds were the hydrogen act like the metal. He decided that acids were the salt of hydrogen. This was a good theory, but it didn't explain why loads of compounds with hydrogen in them are neutral, like water or basic, like NH$_3$.
At last in 1887, Svante Arrhenius came up with a pretty good explanation. He concluded that acids were ionized in an aqueous solution. In the solution, they formed bond with the hydrogen ions. Bases did the same thing, except they formed ionic bonds with hydroxide. This might have been a little off, but it did explain somethings, like water being neutral. Because H$_2$O is a combination of Hydrogen (H$^+$) and hydroxide (OH$^-$), it is neutral.

## Revising Arrhenius' Definition

Arrhenius was right about acids dropping hydrogen ions and bases picking up hydrogen ions but that theory wasn't complete. When acids drop hydrogen, it doesn't stay a lone hydrogen for very long, if there is any water nearby, the positive hydrogen ion will be attracted to the slightly negative side of the oxygen atom and forms H$_3$O$^+$. This is not a covalent bond. The hydrogen is just attracted to the slightly negative oxygen.

## Hydrolysis of Amphoteric Ions

Amphoteric ions are things that alone are not acidic or basic but, it will become so if you just add water. Take NaHCO$_3$ for example. Alone, it doesn't look too acidic but, if you mix it in water, the sodium will disconnect. It's still close to the rest of the ions, still technically bonded to the HCO$_3$. It's kind of like at a dance, if your not holding hands with your dance partner, your just close by. Then, something happens in the solution. The water wants to bond with the HCO$_3$ so they just leave the Na there. It would be like if one of the dancers found someone else to dance with and just left. Really rude. Anyway, the HCO$_3$ and the H$_2$O bond, forming OH$^-$ and H$_2$CO. The OH$^-$ is of course, a base. Another example is NaHSO$_4$. In water the Na is again, dissociated with the ion and H$_2$O comes in. They form HSO$_4$ and an extra hydrogen (acid). This hydrogen finds some water and makes H$_3$O.

### Hydrolysis of Metal and Nonmetal Oxides

These are a lot like Amphoteric ions except there simple metal and nonmetals instead of polyatomic ions. There is just a simple rule for these: If a metal ion is mixed with water, it makes a basic solution. If the compound is nonmetal, it forms an acid.

### Hydrolysis of Metal Ions

When a metal ion is mixed into a water solution, bonds with the water to form complex ions with the water. These will drop hydrogen in the water to make an acid.

## Predicting the Results of Hydrogen

There is a way to find out the concentration of an acid (usually hydrogen). See, in a reaction not all the chemicals react. If you mixed Na and Cl, you'd get a lot of NaCl and also some lone Na and Cl. This equation shows us what the concentration is: K$_a$× K$_b$=K$_w$. Say you have a reaction like HF + H$_2$O → H$_3$O + F. K is HF × H$_2$O/H$_3$O × F. Since H$_3$O is an acid, K is replaced by K$_a$ which is basically the same thing, it just specifies that we're finding the acid concentration. If we use K$_b$ we are looking for a base. K$_w$ is water. So say we have the reaction, CO$_{3(aq)}$ + H$_2$O$_{(aq)}$ → OH$_{(aq)}$ + HCO$_{3(aq)}$. If we know that there is a 0.100 mol/L concentration of CO$_3$ and the pH is 11.66. Since we are working with bases, we will use K$_b$. To find the concentration, we multiply each side of the equation and divide the sides. We divide the side of the reaction with the base by the side without the base. It looks like this:
    K$_b$ = [OH] × [HCO$_3$]/ [CO$_3$]
Notice, we don't put water in the equation. This is because everything is dissolved in water and the CO$_3$ concentration is measured in moles per liter of water. Now, let's put some numbers in the equation! We know the pH is 11.66, this means the pOH is 14 - 11.66 or 2.34 or 10$^{-2.34}$. If we turn 10$^{-2.34}$ into a proper exponent, we will know the concentration of OH. If you put 10$^{-2.34}$ in your calculator, it will come out as the proper exponent: 4.57×10$^{-3}$. This is the concentration of OH and since there is an equal concentration between OH and HCO$_3$, the number is also the concentration of HCO$_3$. Now we know the three numbers needed!
    K$_b$ = (4.57 × 10$^{-3}$)(4.57 × 10$^{-3}$)/ 0.100
    K$_b$ = 2.0 × 10$^{-4}$

# Measuring Acids and Bases
## The pH Scale

Acids and bases are measured on something called the pH scale. It is a scale going from 0 to 14, where 1 is the most acidic and 14 is the most basic. 7 is in the middle of the scale, and so it is neutral.

## Hydrogen Ion Concentration

The pH scale measures the relative concentration of hydrogen in a compound. Starting at seven, each level of the scale going up or down represents 10 times more concentration of hydrogen. Therefore, a pH level of 6 is ten times more acidic than one of 7, and a pH level of 5 is one hundred times more acidic than seven. Alternately, a pH level of 8 is ten times more basic than one of seven, and so on. This is what is known as a logarithm. In this case, it is a negative logarithm, because acids have a higher concentration of hydrogen than bases. This principle can be applied to help you know the pH from a hydrogen ion concentration. If you are given an Hydrogen concentration of 0.000006 M. You can find it's pH by converting it to scientific notation. The decimal place has to move back six times, so the scientific notation would be 6.0 x 10^-6. This means you multiply 6.0 by negative 10 six times, and therefore if it were a pH it would be one of 6. Reference: The virtual chembook (to be added as actual reference)

# Naming Acids

Depending on what cation the hydrogen is attached to, acids will have different names. Simple acids, known as binary acids, have only one cation and one hydrogen. These acids are named by taking the prefix hydro, then putting the first syllable of the cation, then the suffix -ic. For example, HCl, which is hydrogen and chlorine, is hydrochloric acid. More complex acids are acids that have oxygen in the compound. There is a simple rule for these acids. Take nitrate (NO$_3$) for an example. Any polyatomic ion having the suffix -ate uses the suffix -ic as an acid. So HNO$_3$ will be nitric acid. When you have a polyatomic ion with one more oxygen than the -ate ion, then your acid will have the prefix per and the suffix -ic. With one less oxygen than the -ate ion, it will have the suffix -ous. With two less than the -ate ion, the prefix will be hypo and the suffix will be -ous. Example: (some of these acids do not exist and are only used for example purposes)

# The Acid-base Concepts of Bronsted and Lowry

Both Bronsted and Lowry came up with a theory to explain how acids and bases work. Acids are compounds that will drop hydrogen ions. These loose hydrogens will proceed to bond onto other compounds. Bases are the opposite; they pick up hydrogen ions. Stronger acids drop hydrogens easier than weak acids. Strong bases take hydrogen ions more than weak bases. The really strong bases will take hydrogens from compounds that don't want to give up their hydrogens.

## Conjugate Acids and Bases

After a strong acid gives up its hydrogen, it becomes a weak base. Redundantly, a strong base becomes a weak acid after it gets a hydrogen. There is a simple explanation for this. An acid has a hydrogen and has a bond with it. The bind may be weak, but it is a bond none the less. After the acid get rid of the hydrogen, it doesn't really want to get another one, but it does have room for the bond. It works just opposite to that with bases.

## Acid-base Indicators

When a substance changes colour when it is mixed with an acid or base, it is called an acid-base indicator. Litmus paper is the most common type. When you put a piece of litmus paper in an acid, it turns red. When you put it in a base, it turns blue. Remember: Acid Red, Base Blue.

## Buffers

A buffer is a solution that will keep something at about the same pH. For an acid, it is a weak base and for a base, it is a weak acid. When you mix the buffer to an acid so the solution is about half acid and half buffer, it is acting as a buffer. During this time, the acid is giving its hydrogen ions to the base buffer. However, when all the hydrogen ions are taken by the buffer, the solution reaches the buffer capacity. Then, if you add any more of the buffer, the pH will just go up to the pH of the buffer.

# References

* 1 2 3 4 5 6 7 8 9 10 11 12 Jenkins, F., van Kessel, H., Tompkins, D., Lantz, O., (1996). Nelson Chemistry, British Columbia, Nelson Canada.
* ↑ Author Unknown. (n.d). "Naming Acids". <http://www.sciencegeek.net/Chemistry/chempdfs/NamingAcids.pdf>