# Acid–base chemistry

Acid/Base chemistry began with the Arrhenius model of acids and bases. This model states molecules in water that release hydrogen ions (H$^+$) are acids, while molecules in water that release hydroxide (OH$^-$) are bases. This is not complete. The current common definition of an acid and a base is based upon how the substance releases or attracts hydrogen ions (H$^+$). Acids release H$^+$ ions that can turn neutral molecules into positively charged ions, while bases can attract H$^+$ ions from neutral molecules to produce negatively charged ions. This definition allows for bases such as ammonia which does not contain a hydroxide ion.
In water solutions, acids affect water molecules, producing hydronium (H$_3$O$^+$) and bases also affect water molecules, producing hydroxide (OH$^-$) ions. The relative strength of acids and bases is measured by their respective ion concentrations once dissolved. The product of the hydronium ion (H$_3$O$^+$) concentration in water times the hydroxide ion concentration equals 1*10 to the -14th power. When water is not the solvent, the product of the concentrations of the positive and negative ions produced by acids and bases also has a constant value, but the value is different for each solvent.
In water solutions, the pH is equal to the negative log of the hydronium ion (H$_3$O$^+$) concentration. The pOH is equal to the negative log of the hydroxide ion (OH$^-$) concentration. For typical solutions, pH varies between 0 and 14, with 7 (the pH of water) as neutral. Each "step" below 7 is ten times more acidic (since it was derived from a power of ten). For concentrated solutions (> 1 M) of strong acids or bases, solutions can have pH below 0 or above 14 respectively. However, considering the water system at 25 °C, the pH plus pOH of a substance is equal to 14.
The notion of pH in fact depends on the considered solvent and can be extended to other non aqueous polar solvent systems with different autoprotolysis constant. In this case, the sum of pH and pOH is equal to the value of -log of this constant.
For acids and bases that fit the Arrhenius model, a reaction between them always produce water and a salt. Reactions between acids and metal always produce a salt and hydrogen gas.
An electrolyte is a substance, that forms solutions capable to carry electric current.

## Operational and Conceptual definitions

* Operational definition deals with observable things.
* Conceptual definition deals with interpretation.

### Operational definition of bases

The fact that water dissociates to form H$^+$ and OH$^-$ ions in a reversible reaction is the basis for an operational definition of acids and bases that is more powerful than the definitions proposed by Arrhenius. In an operational sense, an acid is any substance that increases the concentration of the H$^+$ ion when it dissolves in water. A base is any substance that increases the concentration of the OH$^-$ ion when it dissolves in water.
These definitions tie the theory of acids and bases to a simple laboratory test for acids and bases. To decide whether a compound is an acid or a base we dissolve it in water and test the solution to see whether the H+ or OH- ion concentration has increased.

### Conceptual definition of acids and bases.

Arrhenius theory
Svante Arrhenius concocted the first successful concept of acids and bases. He did this by defining acids and bases according to the effect these substances have on water. The Arrhenius concept of acids and bases is as follows: an acid is a substance that when dissolved in water increases the concentration of the hydrogen ion, H+. A base is a substance that when dissolved in water increases the concentration of the hydroxide ion, OH-. The hydrogen ion, is not just a bare proton, it is a proton bonded to a water molecule, H$_2$O. This results in a hydronium ion, H$_3$O$^+$. In Arrhenius's theory, something that is a strong acid is a substance that completely ionizes in aqueous solution to give a hydronium ion, H$_3$O$^+$, and an anion. An anion is a negatively charged ion. An example of a strong acid is perchloric acid:
What is going on above is that we have perchloric acid in an aqueous solution. This perchloric acid ionizes entirely and results in a hydronium ion and a perchlorate anion. Some other examples of strong acids would be: HI, HBr, HCl, HNO$_3$, and H$_2$SO$_4$.
Now on to bases...A strong base is something that completely ionizes in aqueous solution to give a hydroxide ion and a cation. A cation is a positively charged ion. Strong bases are most of the hydroxides of Group IA elements and Group IIA elements including LiOH, NaOH, KOH, Ca(OH)$_2$, Sr(OH)$_2$ and Ba(OH)$_2$.
Many of the acids and bases that we encounter in our everyday lives are not strong acids, they are considered weak. Weak acids and bases do not completely ionize in solution, but exist in equilibrium.
Bronsted-Lowery Theory
The Bronsted-Lowery concept of acids and bases is that acid-base reactions can be seen as proton-transfer reactions. This results in acids and bases being able to be defined in terms of this proton (H+) transfer. According to the Bronsted-Lowery concept, acids donate a proton in a proton-transfer reactions. Bases accept the proton in a proton-transfer equation. As an example, lets look at the reaction of hydrochloric acid with ammonia. When we write it as an ionic equation we get:
which reduces to:
because there are two Cl$_{-(aq)}$ (one each side). We now have the net ionic equation after we cancel out the "spectator ions"(Cl$^-$).
What happens in this reaction in aqueous solution is a proton transfer. According to the Bronsted-Lowery concept, acids donate a proton in a proton-transfer reactions. Bases accept the proton in a proton-transfer equation. As an example, let's look at the reaction of hydrochloric acid with ammonia shown above. What happens in this reaction in aqueous solution is that a proton is transferred from H$_3$O$^+$ to NH$_3$. This results in H3O+ losing a (H+), resulting in H2O. The NH3 gains the transferred proton, resulting in NH$^{4+}$. We call H$_3$O$^+$ the proton donor, or acid. We call NH$_3$ the proton acceptor, or base.
The Bronsted-Lowery concept defines something as either an acid or base depending on its function in the acid-base (proton transfer) reaction. Some things can act as either an acid or a base. These are called amphoteric species, they can either lose or gain a proton, depending on the other reactant. An example of an amphoteric species would be HCO$^{3-}$. In the presence of OH$^-$, it acts as an acid. In the presence of HF it acts as a base. Water is also amphoteric, as are most anions with ionizable hydrogens and certain solvents. Water as an amphoteric species is very important to the acid-base reactions.
In the Bronsted-Lowery concept we have found that:
1. A base is a species that accepts protons, while an acid is a species that donates protons.
2. Acids and bases can be ions as well as molecular substances.
3. Some species can act as either acids or bases, depending on what the other reactant is.

## Acid - Base reactions

* Neutralization reaction occurs, when equal quantities of acids and bases occurs. For example: H$^+$ ions react with OH$^-$ ions, to make the H$_2$O molecule.
* Acid - base titration is the process of adding known molarity of acid or base to unknown molarity, until the neutralization will occur. The acid or base of known molarity called standard solution. The endpoint of the reaction can be determined
by acid - base indicators.

### Comolarity calculation

The unknown molarity can be calculated from used volumes and molarity of known standard solution. To do this, we have to balance the equation, and determine the ratio between the Acid and Base.
* Moles of acids = molarity × liters of acid.
* Moles of base = molarity × liters of base.
** If the ratio of this two equations is 1, then
** Liters of acids × Moles of acid = Liters of base × Moles of base.
** The unknown molarity can be calculated, from this relationship.
1) 23.45 mL of 0.275 M sodium hydroxide was used to titrate against mL of acetic acid. What was the concentration in M of acetic acid?
2) 17.05 mL of 0.247 M barium hydroxide was used to titrate against 10 mL of nitric acid. What was the concentration in M of nitric acid?
3) 35.79 mL of 0.275 M sodium hydroxide was used to titrate against 15 mL of sulfuric acid. What was the concentration in M of sulfuric acid?
4) 24.92 mL of 0.00199 M silver nitrate was used to titrate against 5 mL of sodium chloride solution. What was the concentration of NaCl?
(a) 1.29 M (b) 0.842 M (c) 0.328 M (d) 9.92 M

### Normality of the solution

Normality of the solution is the number of moles of (H$^+$) or (OH$^-$) ions per solution. In general volumes and normality of reacting acids and bases are in inverse proportion.
* Volume of Acids = Normality of bases.
* Volume of solution = Normality of acids.

### Conjugate acid - base pairs

have chemical formula that differ by one H$^+$ (they differ by both one H atom and by a +1 charge). They typically appear in a chemical equation for an acid-base reaction, where one is a reactant and the other is a product. Stronger acids have weaker conjugate bases, while stronger bases have weaker conjugate acids. The strongest acids have conjugate bases that are so weak as to be non-basic. The strongest bases have conjugate acids that are so weak as to be non-acidic.

### Ionization constant of acids

Ionization constant (K$_a$) is used to characterize the strength of acids. It is calculated like the constant of the following reaction.
In language, the equilibrium expression reads; "The dissociation constant of an acid is equal to the concentration of hydrogen ions times the concentration of the conjugate base of the acid divided by the concentration of un-ionized acid.
Similarly, the equilibrium expression for a weak base reads: "The dissociation constant of a base equals concentration of hydroxide ions times concentration of conjugate acid divided by the concentration of un-ionized base.
Where K$_w$ is the equilibrium constant for water (unitless)
[H$^+$] is the molar concentration of hydrogen
[OH$^-$] is the molar concentration of hydroxide
K$_a$ x K$_b$ = K$_w$ = 1.00x10$^{-14}$

### Hydrogen ion concentration.

pH is defined as the negative value of the logarithm of the hydrogen ion concentration. In acidic solutions H$^+$ concentration is larger than 10$^{-7}$. For example it can be 10$^{-3}$, which means pH = 3. Solutions with pH values less than 7 are acidic; solutions with pH values greater than 7 are basic (alkaline). The range of pH is typically from 0 (strongly acidic) to 14 (strongly basic); more extreme values sometimes occur.
The acidity or basicity of a substance is defined most typically by the pH value, defined as below:
At equilibrium, the concentration of H+ is 10$^{-7}$, so we can calculate the pH of water at equilibrium as:
Solutions with a pH of seven (7) are said to be neutral, while those with pH values below seven (7) are defined as acidic and those above pH of seven (7) as being basic.
pOH gives us another way to measure the acidity of a solution. It is just the opposite of pH. A high pOH means the solution is acidic while a low pOH means the solution is basic.
Strong Acids: These acids completely ionize in solution so they are always represented in chemical equations in their ionized form. There are only seven (7) strong acids: HCl, HBr, HI, H$_2$SO$_4$, HNO$_3$, HClO$_3$, HClO$_4$ To calculate a pH value, it is easiest to follow the standard "Start, Change, Equilibrium" process. Example Problem: Determine the pH of a 0.25 M solution of HBr.
Weak Acids: These are the most common type of acids. They follow the equation:
The equilibrium constant for the dissociation of an acid is known as Ka. The larger the value of Ka, the stronger the acid.
Example Problem: Determine the pH of .30 M acetic acid (HC$_2$H$_3$O$_2$) with the K$_a$ of 1.8x10$^{-5}$.
Strong Bases: Like strong acids, these bases completely ionize in solution and are always represented in their ionized form in chemical equations. There are only eight (8) strong bases: LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH)$_2$, Sr(OH)$_2$, Ba(OH)$_2$
Example Problem: Determine the pH of a 0.010 M solution of Ba(OH)$_2$.
Weak Bases: They follow the equation:
Weak Base + H$_2$O <---> conjugate acid + OH$^-$
example: NH$_3$ + H$_2$O <---> NH$^{4+}$ + OH$^+$
K$_b$ is the base-dissociation constant:
K$^a$ x K$_b$ = K$_w$ = 1.00x10$^{-14}$
To calculate the pH of a weak base, we must follow a very similar "Start, Change, Equilibrium" process as we did with the weak acid, however we must add a few steps.
Example Problem: Determine the pH of 0.15 M ammonia (NH$_3$) with a K$_b$=1.8x10$^{-5}$.
When dealing with weak acids and weak bases, you also might have to deal with the "common ion effect". This is when you add a salt to a weak acid or base that contains one of the ions present in the acid or base. To be able to use the same process to solve for pH when this occurs, all you need to change are your "start" numbers. Add the molarity of the ion, which comes from the salt, and then solve the Ka or Kb equation as you did earlier.
Example Problem: Find the pH of a solution formed by dissolving 0.100 mol of HC$_2$H$_3$O$_2$ with a K$_a$ of 1.8x10$^{-8}$ and 0.200 mol of NaC$_2$H$_3$O$_2$ in a total volume of 1.00 L.

## Salts

Salt is one of the products of a neutralization. It is a compound containing at least one positive and one negative ion. All salts are strong electrolytes completely ionized by water. Some salts form acidic or basic solution. When it occurs this process is called hydrolysis.