Atoms and Molecules

FAS Astronomers Blog, Volume 32, Number 7.

We all were taught that things are made up of molecules, which, in turn, are made up of atoms. Atoms are composed of electrons (with a negative charge) and a nucleus, which contains protons (with a positive charge) and neutrons (with no charge). Electrons are thought to be fundamental particles. Protons and Neutrons are not and are composed of Quarks. For more on this see The Standard Model of Particle Physics. 

There are atoms and molecules, as well as elements and compounds. You can think of a chemical element as a specific type of atom. A molecule is a group of atoms bonded together. A compound is something composed of more than one element. Not all molecules are compounds. A molecule of Oxygen (O₂) is composed of two Oxygen atoms, but it is only composed of one element and is therefore not a compound.

Each chemical element is assigned an “atomic number”, which corresponds to the number of protons and electrons in a neutral atom. The number of protons defines what the atom is – that is, which element it is. Hydrogen (H) has one proton, Helium (He) two protons, Lithium (Li) three protons, all the way up to Oganesson (Og) with 118 protons. 

There are different versions of each element, called isotopes, with a varying number of neutrons. For lighter elements (Hydrogen through Sulfur), the number of neutrons tends to equal the number of protons for the most common isotope. For heavier elements, the number of neutrons generally exceed the number of protons. Isotopes are referred to using the total number of nucleons (protons plus neutrons). For example, Uranium 235 (with 92 protons and 143 neutrons) and Uranium 238 (with 92 protons and 146 neutrons). Some isotopes, such as Hydrogen, have specific names, for example, Deuterium (Hydrogen 2) and Tritium (Hydrogen 3). 

Each isotope has an atomic mass, which is roughly equal to the number of protons plus neutrons. The atomic mass unit of measure is an “unified atomic mass unit” (u), Dalton, or amu. 1 u is defined as 1/12 the atomic mass of carbon 12. In other words, it is defined so that carbon 12 has an atomic mass of 12. 

Each element has an atomic weight, which is a weighted average of the atomic masses for all the element’s isotopes as they exist in nature. For example, Hydrogen’s atomic weight, 1.007, is determined from the relative atomic masses and abundance of neutral hydrogen (1 u), deuterium (2 u), and tritium (3 u). 

One of the more confusing terms in Chemistry is the mole. Wilhelm Ostwald first used the term mole, which comes from the German word for molecule. A mole is a way to easily represent a very large number of things. A mole is basically 6.0221 x 10²³ things, just as a dozen is 12 things. For example, 6.02 X 10²³ grains of sand are a mole of sand. 

The number 6.02 x 10²³ is called Avogadro’s number (N_A). Jean Baptiste Perrin named it after Amedeo Avogadro, who first theorized that a specific volume of gas was made up of a given number of particles. Experimentally, Avogadro’s number is determined from Carbon 12. Twelve grams of Carbon 12 contain 6.02 x 10²³ atoms. So, in Chemistry, the atomic weight/atomic mass can be expressed in grams per mole (g/mol). For example, Hydrogen has 1 g/mol, Helium has 4 g/mol, Carbon 12 has 12 g/mol and so on. 

The advantage of this comes from the ease in which moles allow us to work with chemical formulas. We could express things in grams or number of atoms; however, different atoms and molecules have different masses (in grams) and different numbers of atoms. Moles are easier. Water gives us a simple example: 2H₂ + O₂ = 2H₂O; four moles of hydrogen plus two moles of oxygen results in two moles of water. This is easier to work with than using grams; 4.028 g of hydrogen (4 x 1.007 g/mol) plus 31.998 grams of oxygen (2 x 15.999 g/mol) results in 36.030 grams of water (2 x 18.015 g/mol).

By the way, October 23 is “Mole Day.” It is, of course, to be celebrated from 6:02 am to 6:02 pm. 

Neutral atoms have the same number of electrons and protons. However, atoms can gain or lose electrons; in which case they are referred to as ions. 

• A positive ion (cation) has lost one or more electrons. Elements to the left in the periodic table have a small valence number and tend to lose electrons.
• A negative ion (anion) has gained one or more electrons. Elements to the right in the periodic table have many valence electrons and tend to add electrons. 

Heavier atoms, particularly those in the actinide series, transform into one another through radioactive decay, when the composition of the nucleus changes. However, this is more in the realm of Physics, and we’ll leave it to another day.

Much of the information about the elements and their structure can be found in the periodic table. Here the elements are grouped by their characteristics. Each row (period) has elements with the same electron structure. Each column (group or family) has elements with similar characteristics and are said the have the same number of valence electrons (the electrons in an atom’s outer shell). Noble gases such as Helium, Neon, Argon, and elements in group 18 have outer shells that are complete, and, as such, are very stable. Others have incomplete outer shells and bond together to form molecules.

The more common molecules found in nature form from the elements in the main part of the periodic table (columns 1, 2, and 13-17). Carbon readily bonds with elements on either side forming a vast array of complex organic molecules. 

For many printable periodic tables, see the collection by Science Notes.

Electronegativity measures the attraction of the electrons (negative charge) to the nucleus (positive charge). It is typically expressed as a dimensionless number from 0.79 to 3.98 and it determines the ability of an element to form a chemical bond. Electronegativity is the greatest for elements in the top right of the periodic table (excluding the noble gases), which have the most valence electrons, and as, such, the most electrical attraction to the positive nucleus. Fluorine (F⁹), with a valence of 7, has the greatest electronegativity at 3.98 and easily gains an electron to fill its outer shell. Science Notes has a nice chart showing the electronegativity by element.

Closely related to electronegativity is the ionization energy, which is the energy required to remove an electron from a gaseous atom. The required energy increases as the electronegativity increases – moving right and up the periodic table.

Molecules are created through chemical bonds (the transfer or sharing of electrons) to completely fill or empty their valence/outer electron shell. Electron shells are composed of one or more subshells (s², p⁶, d¹⁰, f¹⁴), each holding a different number of electrons. Note that other than the s subshell, the last subshell could be empty. For example, Lithium, with three electrons, has a structure 1s²2s¹p⁰.  

• Hydrogen and Helium, the lightest atoms, have only two electrons (s²) in their outer shell.
• The remaining elements in the periodic table main group (columns 1-2 and 13-18/1A-8A) have room for 8 electrons (s²p⁶) in their outer shell (the octet rule). 
  • Elements in columns 1 and 2 have a configuration of s^xp⁰, where x = 1 or 2.
  • Elements in columns 13-18 have a configuration of s²p^y, where y = 1 to 6.
• The transitional and inner transitional metals (periodic table columns 3-12/1B-8B, along with the Lanthanide and Actinide series) are more complicated and can have more than 8 electrons in their outer shell.

Science Notes has three periodic tables showing information about electron shells (Electron Shells, Outer Electron Orbits & Electron Structure).

There are several primary types of chemical bonds.

• Ionic bonds are when electrons are transferred between atoms to empty or fill the outer shell. They take place primarily between a metal and nonmetal. For example, lithium fluoride is formed when lithium (with one valence electron) loses an electron to fluorine (with seven valence electrons). These bonds form when the electronegativity is greater for the atom receiving the electron (stronger electronic bond) than for the one losing the electron (weaker electronic bond).
• Covalent bonds are when electron pairs are shared between two atoms. They take place primarily between two nonmetals. For example, Oxygen (with six valence electrons) shares electrons with two hydrogen atoms (one valence electron each) to form water (H₂O). These bonds form when the electronegativity is similar between the two atoms.
  • Non-Polar (Covalent) bonds are when there is an even sharing of electrons. In this case, the electronegativity of the two atoms is similar. 
  • Polar (Covalent) bonds are when there is an uneven sharing of electrons. Electrons tend to “gather” more toward one atom resulting in a slightly negative charge, leaving the other atom with a slight positive charge. These bonds are formed when there is a slight difference in electronegativity between the atoms.
• Metallic bonds are formed when electrons in the outer shell are shared (delocalized) among a “electron sea” or “lattice” of ions. They take place between two metals. The shared electrons are free to move about making metals excellent electrical and thermal conductors. Because the bonds can be easily broken, metals are malleable and can be formed into thin wires (ductile).

There are also free radicals, such as the hydroxyl radical (HO), which have one or more unpaired valence electron. In this case the Oxygen atom has five rather than six electrons. Biradicals, such as nitrogen oxide (NO₂), also have one or more unpaired electron.