Approximately 3.4 billion years ago, the first life form sprung up in a hydrothermal vent deep in the ocean near present-day Quebec. As hundreds of millions of years passed, life evolved from little specs to multicelled organisms like plants and fish. Eventually, species evolved to walk on land, and even more millions of years later, the first human was born!
What do all these life forms have in common? They are all made of carbon. In this article, we will be walking through the different properties of carbon and see why it is one of, if not the most important element on the periodic table.
- This article is about the properties of carbon.
- First, we will look at why carbon is so important.
- Next, we will look into the physical properties of carbon.
- Then, we will look at its atomic structure, so we can understand its chemical properties in the next section.
- Lastly, we will talk about carbon's many uses.
Purpose of carbon
Due to carbon's properties, it is a building block of life. All living things are made up of carbon. The whole field of organic chemistry is dedicated to carbon, that's one important element!
Now that I have sufficiently hyped up carbon, let's look at its properties, shall we?
Physical properties of carbon
Carbon has a melting point of 3,652 °C and a boiling point of 4,827 °C, so it is almost always present as a solid. It is soft and either dull grey or black. Some of its physical properties are dependent on the allotrope.
An allotrope is a version of an element with a different physical structure. The bonding of the atoms within the molecule differs, leading to different physical properties.
Two of the most common carbon allotropes are graphite and diamond.
Graphite is soft, light, and a good conductor of heat and electricity. It has a density of 2.26 g/cm3. Diamond, on the other hand, is hard, strong, and a poor conductor of electricity (but is a good conductor of heat). Its density is 3.51 g/cm3.
Not only are they different physically, but chemically as well. Graphite can react violently with strong oxidizing agents like fluorine, and graphite dust and air are explosive when ignited. Diamond doesn't react at all at room temperature, but can catch fire around 690 °C if put in an environment of pure oxygen.
These differences are dependent on how the carbon atoms are arranged. Below is a comparison between the two structures:
Fig. 1: Comparison of the structures of diamond and graphite. Wikimedia commons.
Diamond's atoms are arranged in a hexagonal lattice and have a crystalline structure. The atoms in graphite are also arranged hexagonally but are in flat sheets, which is why it's much weaker.
The different allotropes make carbon a jack-of-all-trades. While these two are the most common, there are many more. Here is a table with some more carbon allotropes:
| Name | Notable properties |
| Lonsdaleite (hexagonal diamond) | Is an "in-between" of graphite and carbon. It has graphite's hexagonal lattice, but its density is much closer to that of diamond. |
| Diamene | A 2D form of diamond |
| Amorphous carbon | Doesn't have a crystalline structure |
| Fullerene (buckyballs) | Are curved molecules that can be shaped like a hollow sphere, ellipsoid, or tube. |
| Glassy carbon | Commonly used as an electrode material |
| Carbon nanofoam | Low-density cluster of carbon atoms |
Several of these allotropes (like diamine), only exist in certain conditions like high-pressure environments.
Atomic structure of carbon
To help us understand the chemical properties of carbon, we should first look at its atomic structure.
Carbon's entry on the periodic table tells us the basics about its structure:
Fig. 2: Carbon's entry on the periodic table
Looking at the atomic number, we know that carbon has 6 protons. Carbon has an atomic weight of 12.01 amu (atomic mass units), and its most common isotope has 6 neutrons.
Isotopes are different forms of one element that have varying numbers of neutrons.
For example, carbon-14 is an isotope of carbon that has 8 neutrons. It is also the isotope used in carbon dating.
Here is carbon's atomic structure:
Fig. 3: Neutral carbon's atomic structure. Wikimedia commons.
Carbon has two core electrons and four valence electrons.
Valence electrons are the electrons in the outermost ring/the highest energy level. These electrons are used in bonding.
Carbon's four valence electrons are a main reason why it is so important, which we will get to in the next section.
Chemical properties of carbon
Since carbon has four valence electrons, it can form 4 covalent bonds.
In a covalent bond, the electrons are shared between the two elements.
Carbon is incredibly versatile. It can perform 4 types of reactions:
1. Combustion: Carbon combines with oxygen to produce heat energy
$$CH_4+2O_2 \rightarrow CO_2 + 2H_2O$$
2. Oxidation: Oxygen is added to carbon
$$CH_3CH_2OH + O \rightarrow CH_3CHO + H_2O$$
3. Addition: An element/compound is added to an element/compound
$$HCCH + HI \rightarrow H_2CCHI$$
4. Substitution: An element/compound substituted in for another element/compound
$$CH_3Cl + OH \rightarrow Cl^- + CH_3OH$$
Because of these reactions, carbon can form many compounds. The most common type of carbon compound is a hydrocarbon.
A hydrocarbon is a compound containing both carbon and hydrogen.
We currently know approx. 1 million hydrocarbon compounds and that number increases every year! While carbon can form a number of inorganic (non-C-H) molecules, there are a lot fewer of those. Carbon, in total, forms more compounds than all the other elements, combined. That's a lot of compounds!
Carbon can also form long, sturdy C-C chains. In fact, the largest chain on record was 6,000 carbons long!
Uses of carbon
Carbon has a a lot of uses. Here is just a small chunk of its uses:
- Used in fossil fuels (hydrocarbons).
- Used in pencils (graphite).
- Used in jewelry (diamond).
- Breathed in by plants and converted to oxygen (photosynthesis) (CO2).
- Used in high tensile strength materials (fullerene/buckyball).
- 12% of the human body (elemental carbon).
There are a lot more uses, but these give you a good look into how widespread carbon is. They don't call it a building block of life for nothing!
The carbon cycle is a summary of how carbon is transferred.
Fig. 4: Carbon cycle diagram. Wikimedia commons.
Processes like fossil fuel burning and land use, pump carbon into the atmosphere. The decomposition of organic life also releases carbon, which is absorbed by the ground. Processes like photosynthesis and the sea-surface gas exchange take in carbon. The ocean is a major reservoir of carbon and holds 50x more carbon than our atmosphere.
Carbon - Key takeaways
- An allotrope is a version of an element with a different physical structure. The bonding of the atoms within the molecule differs, leading to different physical properties.
- Carbon's most common allotropes are graphite and diamond.
- Carbon has 6 protons, (usually) 6 neutrons, and 6 electrons.
- Isotopes are different forms of one element that have varying numbers of neutrons.
- Valence electrons are the electrons in the outermost ring/the highest energy level. These electrons are used in bonding.
- In a covalent bond, the electrons are shared between the two elements.
- Carbon has 4 valence electrons, so it can make four covalent bonds.
- Carbon is very versatile and can form long, sturdy C-C chains.
- A hydrocarbon is a compound containing both carbon and hydrogen.
- Carbon has many uses, such as for fuel, photosynthesis, and forming organic life.
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