In VCE Chemistry Unit 2, you will be learning all about acid-base chemistry. Acids and bases are required for many different aspects of life, such as enzyme function, effective cleaning products and even baking a cake!
In this article, we will go through the different applications of acid-base chemistry that you are expected to know, and the theory behind it.
Brief Overview of Acids and Bases
Before we jump into the applications, let’s do a quick review on what acids and bases are.
According to the Bronsted-Lowry Theory on Acids and Bases:
- Acids are proton (H+) donors.
- Bases are proton (H+) acceptors.
In an acid-base reaction, an acid donates a proton to a base. The acid then turns into a base, while the base turns into an acid.
H2SO4 + Cl- --> HCl + HSO4-
Acids are usually corrosive and sour tasting, while bases are more slippery and bitter tasting.
Want to learn more? You can download some detailed notes on acids and bases here.
Perhaps you’ve heard of, or even taken the medicines Gaviscon or Mylanta before. These medicines are known as antacids, which are designed to relieve heartburn, indigestion, or an upset stomach. These symptoms are usually caused by an excess of hydrochloric acid in the stomach.
An antacid is usually a metal carbonate (XCO32-) - it reacts with the excess acid in the stomach to neutralise it. Remember, the products of the reaction with an acid and metal carbonate are a salt, water, and carbon dioxide.
Acid + Metal Carbonate --> Salt + Water + Carbon Dioxide
Some common types of antacids are calcium carbonate, sodium hydrogen carbonate and magnesium carbonate.
Carbon Dioxide in the Environment
We are burning more fossil fuels than we ever have before. We have become so dependent on them as an energy source to produce heat, electricity, and power. When fossil fuels are burnt, an extreme amount of carbon dioxide (and other toxins) are released into the atmosphere.
In 1950, the world emitted 6 billion tonnes of CO2 into the atmosphere. Shockingly, in the present, we emit 34 billion tonnes of CO2 a year! This increase in CO2 levels has major consequences for the environment.
Carbon dioxide is a gas, but it can dissolve in water and become aqueous. Most of the time, it remains in the aqueous form, but it sometimes reacts with the water to form carbonic acid (a weak acid).
Carbonic acid is diprotic (as seen by the two hydrogen atoms), which means that it can ionise in two steps.
Step 1 forms hydrogen carbonate ions:
H2CO3(aq)+ H2O(l) ⇌H3O+(aq)+HCO3-(aq)
Step 2 forms carbonate ions:
HCO3-(aq) + H2O(l) ⇌ H3O+(aq)+CO3-2(aq)
As you can see, the more carbon dioxide reacts with water, the more hydronium ions are produced. More hydronium ions produced means an increase in acidity!
Since our atmosphere contains a small level of CO2, uncontaminated rain is naturally a bit acidic – it has a pH of about 5.6. As mentioned above, this is due to the formation of carbonate ions. The carbon dioxide in the air does not contribute to acid rain, it impacts the natural acidity of rain.
However, due to human activities, the acidity of rain can decrease to a pH of 4.2-4.4. This is what is known as acid rain. As we burn fossil fuels through various practices, sulphur and nitrogen oxides are released into the atmosphere. These pollutants react with water, oxygen, and other gases to form acidic products.
Sulphuric Acid Formation: 2SO2 (g) + O2 (g) + 2H2O (l) → 2H2SO4 (aq)
Nitric Acid Formation: 4NO2 (g) + O2 (g) + 2H2O (l) → 4HNO3 (aq)
Some consequences of acid rain include:
- Corrosion of steel structures
- Weathering of stone
- Extracts aluminium from soil – harmful to plants and animals
- Removes minerals and nutrients from soil.
- Respiratory issues in humans and animals
As the levels of CO2 rises in the atmosphere, so does the dissolved content of CO2 in the oceans. As you’d expect, this causes the pH of the ocean to decrease over time, as the weak carbonic acid is produced. Prior to the Industrial Revolution, the ocean’s pH was 8.2, but now, it is about 8.14. Although this may not seem like a big difference, since pH is measured in logarithmic scale, this is about a 30% increase in the H+ concentration!
This has consequences for the shell growth of marine invertebrates (shellfish, starfish crabs, etc). These creatures have a protective shell which is mostly composed of calcium carbonate (CaCO3). To maintain this shell, the animals take up the calcium (Ca+2) and carbonate ions (CO32-) that are naturally present in the ocean. This process is called calcification.
Ca+2+ CO32-⇌ CaCO3
As we saw above, CO32- can react with hydronium ions in a reversible reaction:
H3O+(aq)+CO3-2(aq) ⇌HCO3-(aq) + H2O(l)
The more hydronium ions that are present in the ocean, the more this reaction will occur. This means that concentration of CO3-2 ions in the ocean will decreases, which means that the organisms will not be able to build their protective layers. This process is called decalcification, and it can have severe impacts on the health of our oceans and marine life.
That’s it for the applications that are specifically mentioned in the study design – however, but as mentioned in the introduction, acids and bases truly are everywhere! If you’d like to learn more, see if you can do some research into the acids and bases that you can find around your own home.