Acids & Bases
Acids:
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Tastes sour, (but remember, in the lab, you test, not taste).
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Produces painful sensation on the skin.
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Reacts with certain metals (magnesium, zinc, and iron) to produce hydrogen gas.
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Reacts with limestone and baking soda to produce carbon dioxide.
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Reacts with litmus paper and turns it red.
Bases:
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Tastes bitter (again, in the lab, you test, not taste).
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Slippery on skin.
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Reacts with oils and greases.
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Reacts with acids to produce a salt and water.
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Reacts with litmus paper and turns it blue.
The Arrhenius Theory,
Must have Water:
In this theory, an acid is a substance that yields H+ (hydrogen) ions when dissolved in water, and a base is a substance that yields OH- (hydroxide) ions when dissolved in water.
Ex. of acid:
HCl(aq) --> H+(aq) + Cl-(aq)
Ex. of base:
NaOH(aq) --> Na+ (aq) + OH-(aq)
Arrhenius also classified the reaction between an acid and a base as a neutraliztion reaction, because if you mix an acidic solution with a basic solution, you end up with a neutral solution composed of water and a salt.
Distinguishing between Strong and Weak Acids and Bases:
Acid-base strength is not the same as concentration.
Strength refers to the amount of ionization or breaking apart that a particular acid or base undergoes.
Concentration refers to the amount of acid or base that you initially have.
You can have a concentrated solution of a weak acid, or a dilute solution of a strong acid, or a concentrated solution of a strong acid, or...well, im sure you get the idea.
Ionizing completely; Strong Acids:
Acids that ionize completely are considered strong.
Ex: If you dissolve hydrogen chloride gas in water, the HCl reacts with the water molecules and donates a proton to them:
HCl(g) + H2O(L) --> Cl-(aq) +H3O+(aq)
The H3O+ ion is called the hydronium ion. The reactants keep creating the product until they're all used up. In this case, all the HCl ionizes to H3O+ and Cl-; no more HCl is present. The water in this case acts as a base, accpeting the proton from the hydrogen chloride.
Because strong acids ionize completely, calculating the concentration of the hydronium ion and chloride ion is easy if you know the initial concentration of the strong acid.
For example, suppose that you bubble 0.1 moles of HCl gas into a liter of water. You can say that the initial concnetration of HCl is 0.1 M (0.1mol/L). M stand molarity, and mol/L stands for moles of solute per liter.
Because the HCl completely ionizes, you see from the balanced equation that for every HCl that ionizes, you get one hydronium ion and one chloride ion. So the concentration of ions in that 0.1 M HCl solution is
[H3O+] = 0.1 M and [Cl-] = 0.1 M.
Monoprotic acids only donate one proton and diprotic acids donate two.
Falling to Pieces: Strong Bases:
A strong base dissociates (breaks apart) completely in water. You normally only see one strong base, the hydroxide ion, OH-. Calculating the hydroxide ion concentration is really straightfoward.
Example:
You have 1.5 M (1.5mol/L) NaOH solution. The sodium hydroxide completely dissolves into ions. NaOH --> Na+(aq) + OH- (aq).
If you start with 1.5mol/L NaOH, then you have the same concentration of ions: [Na+] = 1.5 M and [OH-] = 1.5 M
Finding Equilibrium with Water; Weak Bases:
Weak basses react withwater to establish an equilibrium system. Aammonia is a typical weak base. It reacts with water to form the ammonium ion and hydroxide ion:
NH3(g) + H2O(L) <=> NH4+ +OH-
Like a weak acid, a weak base is only partially ionized. The modified equilibrium constant expression for weak bases is Kb. You use it exactly the saem way you use the Ka
(see "Ionizing Partway Weak Acids") except you solve for the [OH-]
Ionizing Partway: Weak Acids:
Acids that only partially ionize are called weak acids. One example is acetic acid (CH3COOH). If you dissolve acetic acid in water, it reacts with the water molecules, donating a proton and forming hydronium ions. It also establishes an equilibrium in which you have a significant amount of unionized acetic acid. ( In reactions thhat go to completion, the reactants are completely used up creating the products. But in equilibrium systems, two exactly opposite chemical reactions -- one on each side of the reaction arrow -- are occuring at the same place, at the same time, with the same speed of reaction.)
The acetic acid reaction with water looks like this:
CH3COOH(L) + H2O(L) <=> CH3COO-(aq) + H3O+ (aq)
The acetic acid that you added to water is only partially ionized, so it's a weak acid. In the case of acetic acid, about 5 percent ionized, and 95 percent remains in the molecular form. The amount of hydronium ion that you get in solutions of acids that don't ionize completely is much less than it is with a strong acid.
Calculating the hydronium ion concentration in weak acid solutions isn't as straightfoward as it is in strong solutions, because not all the weak acid that dissolves initially has ionized. In order to calculate the hydronium ion concentration, you must use the equilibrium conctant expression for the weak acid. This expression is called the Ka -- acid ionization constant.
The Ka for acetic acid is 1.8 x 10^-5. The Ka expression for the acetic acid ionization is
Ka= 1.8x10^-5= [H3O+][CH3COO-]
[CH3COOH]
You would solve for x, which is the [H3O+].
Remember, one way to distinguish between a weak acid and a strong is if the acid ionization constant is Ka value. If it has a Ka value, then it is a weak acid.
Competing for Protons;
Bronsted-Lowry acid-base reactions:
With the Arrhenius theory, acid-base reactionsa re neutralization reactions. With the Bronsted-Lowry theory, acid-base reactionsa are a competition for a proton.
For example, the reaction of ammonia with water:
NH3(g) + H2O(L) <=> NH4+ (aq) + OH-(aq)
Ammonia is a base (it accepts the proton), and water is an acid (it donates the proton) in the foward (left to right) reaction. But in the reverse reaction (right to left), the ammonium ion is an acid and the hydroxide ion is a base.
If water is a stronger acid than the ammonium ion, then a relatively large concentration of ammonium and hydroxide ions are at equilibrium. If, however, the ammonium ion is a stronger acid, much more ammonia than ammonium ion is present at equilibrum.
Bronsted and Lowry said that an acid reacts with a base to form conjugate acid-base pairs, which differ by a single H+. For example, NH3 is a base and NH4+ is an acid in the reaction between ammonia and water, and OH- is its conjugate base. In this reaction, the hydroxide ion is a strong base and ammonia is a weak base, so the equilibrium is shifted to the left -- not much hydroxide is present at equilibrium.
Click Learn More to learn more about Identifying Acids and Bases with Indicators and the Exponential Nature of pH Values in terms of Concentrations!