lecture 8 study guide

Acids and bases

Acids—when added to water lower the pH, increase concentration of protons (H+ ion).
Bases—readily accept protons, so much so that they will rip them from other molecules. The hydroxide ion (OH-) is the simplest base.

Stronger acids completely disassociate in water. Examples include hydrochloric acid (HCl), sulfuric acid (H2SO4) , nitric acid (HNO3).

—>practice drawing a strong acid outside of water and disassociated in water. Hint: review the slide of NaCl in water.

In general, strong acids and bases are small molecules while weaker acids and bases are larger molecules.
Carboxylic acids (organic acids) are weaker acids, they are larger molecules. A common example is acetic acid/ethanoic acid in vinegar.

Conjugate acids and bases

When an acid disassociates in water, the basic ion leftover is known as its conjugate base. Similarly, when a base dissociates in water, the acidic ion that is left over is known as its conjugate acid.

Strong acids have weak conjugate bases. Strong bases have weak conjugate acids.

Hydrochloric acid (HCl) is a strong acid. Cl- is a weak base. Sodium hydroxide (NaOH) is a strong base. Na+ is a weak acid. Weak acids have conjugate bases that are not as weak.

—>Acetic acid (vinegar) is a weak acid. Is acetate, its conjugate base, strong or weak? Why?

Chemical equilibriums

Weak acids and bases do not completely disassociate in water. They exist in a dynamic chemical equilibrium. Here is an example equation showing acetic acid and water on the left side, and acetate and H3^O+ (an acidic proton and a water molecule) on the right side.

20051103062911!Acetic_acid_deprotonation.png

Increasing the concentration of the molecules on one side of the equilibrium will drive the reaction in the other direction.

—>illustrate this on the image above

Carbonic Acid

Carbonic acid (H2CO3) is formed when carbon dioxide is dissolved in water. Carbonic acid exists in an equilibrium with bicarbonate (HCO3-) and carbonate (CO32-). Carbonic acid is the most acidic chemical in this system, bicarbonate is intermediate, carbonate is the most basic.

Carbonic-acid-2D.png

http://upload.wikimedia.org/wikipedia/commons/d/d5/Carbonic-acid-2D.png

—>Draw the equilibrium between the chemicals in the carbonic acid complex. Carbonic acid has been provided above as a starting point.

What does this graph mean?

4.1-2.png

As the concentration of CO2 in the ocean increases, the water becomes more acidic
As the acidity increases, the concentration of the bicarbonate ion, and the carbonate ion decrease greatly
—Bottom line: increasing CO2 in water increases acidity and decreases the concentration of carbonate

Why is this important from an ecological perspective?

Lots of ocean critters (corals, bivalves and others) use calcium carbonate (CaCO3) to make their shells and. Increasing the acidity of the ocean will decrease the amount of calcium carbonate available to these critters, making it so they cannot build their shells. Sufficiently acidic conditions will draw the carbonate from their shells, essentially dissolving their shells. This can disrupt many important and highly productive marine ecosystems.

—>given, our knowledge of equilibriums, how might this happen?

There is evidence that this is happening now. i.e. coral death. Hood and Kolbert articles.

Acid rain, NOx and SO2 emissions

Acid rain is caused by NOx and SO2 emissions. NOx is a bigger player now due to clean air act. NOx combines with water in the atmosphere to form nitric acid (HNO3). SO2 combines with water in the atmosphere to form sulfuric acid (H2SO4). Acid rain that is caused by nitric acid leads to an increase in available nitrogen in an ecosystem.

Acidity of soil impacts cation exchange with plants. Acid rain can wash essential nutrients from soil but also contributes to the addition of nutrients by weathering rocks.

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