Chemistry is all over the world around us, whether we realize it or not. It’s in the air we breathe, the things we eat, the products we consume, and it’s even within us! As we’ve gained a deeper understanding of the science, we’ve established a relatively universal way to communicate it on paper. This has made chemistry a language of its own. Thus, being able to read and interpret written chemistry is critical to deeply understanding and problem-solving in chemistry. We use a variety of symbols to represent elements, compounds, molecules, ions, isotopes, and so on. To some chemistry learners, the symbology may be confusing. Introductory chemistry courses often gloss over how to properly read written chemistry. Some instructors assume the topic to be self-explanatory, but we will not do that here. We will go through the fundamental symbols and their meaning with the goal of boosting your understanding of the language of chemistry!

Important Background

Fundamental Particles

Understanding the fundamental particles is key to understanding chemistry. They are what compose atoms and are thus the building blocks for all chemical species from element to molecule and so forth. The most basic particles are protons, neutrons, and electrons. By convention, we have assigned each of these particles a particular charge. Protons are positively charged, neutrons are neutral, and electrons are negatively charged. Protons and neutrons directly make up the nucleus, while electrons cloud the nucleus.

Protons are essentially the chemical equivalent of a fingerprint. If the number of protons changes, so does the element’s identity. For example, sodium (Na) has 11 protons. If a proton were added, our sodium is now magnesium (Mg). If a proton were removed, our sodium is now neon (Ne). However, additions and removals of protons are nuclear processes and are not relevant for early learners. The number of neutrons and electrons does not change chemical identity. The number of neutrons an atom has is relevant when discussing isotopes, atoms that have the same number of protons but a different number of neutrons. The number of electrons an atom has is relevant in discussing chemical reactions and ions, atoms with an unequal number of protons and electrons. We will discuss isotopes and ions in later sections.

In summary:

• The fundamental particles are:
  • Protons: positively charged particles that sit in the nucleus
  • Neutrons: uncharged/neutral particles that sit in the nucleus
  • Electrons: negatively charged particles that cloud the nucleus

Symbology

For Elements

To represent elements, the chemicals that comprise the periodic table, we use either a single capital letter (ex. P for phosphorus) or a capital letter followed by a lowercase letter (ex. Na for sodium). The letters used are unique to each individual element and are chosen for a variety of different reasons. But no matter how they are derived, the symbol you see on the periodic table borrows one or two letters from the element’s name or a word related to the element. Let’s consider Cu, copper’s elemental symbol. It’s derived from the element’s Latin name, “cuprum”. Some elements may be more obvious or intuitive to the average English speaker, such as P for phosphorus.

In summary…

•  An element is a substance made of a single atom. Each square on the periodic table is occupied by a single element.
• To represent elements, we use either of the following:
  • a) A single capital letter
    • Examples:
      • P for phosphorous
      • W for tungsten
  • b) A capital letter followed by a lowercase letter
    • Examples
      • Cl for chlorine
      • Na for sodium

For Compounds and Molecules

Compounds are chemicals composed of two or more atoms of different elements. These elements are bonded together by shared electrons. Some elements in compounds exist in nice 1:1 ratios. A great example is table salt or sodium chloride, which has the chemical formula NaCl. However, the combining ratios are not always this simple. So how do chemists account for this? We use subscripts! Let’s consider calcium chloride with the chemical formula CaCl₂. Ca, having no subscript, means there is only one calcium atom in a single molecule of calcium chloride. Cl, having a subscript of two, means there are two chlorine atoms in a single molecule of calcium chloride. A more lengthy example is the chemical formula of glucose, C₆H₁₂O₆. Carbon and oxygen both have subscripts of six, meaning there are six atoms of both elements in a single glucose molecule. Similarly, hydrogen’s subscript of twelve corresponds to twelve hydrogen atoms in a single glucose molecule.

The rules are identical for molecules, which are chemicals composed of two or more atoms of elements that may or may not be the same. That being said, all compounds are molecules but not all molecules are compounds. Great examples are homonuclear diatomics, molecules composed of two atoms of the same element. Oxygen gas, O₂, is a great example. Heteronuclear diatomics, such as NaCl, are molecules composed of two atoms of different elements.

You may be wondering, how do we decide which subscripts to use? It really depends on the nature of our molecule and the intramolecular forces at work! These are the forces we observe between atoms within a molecule, and their relative strength based on Pauling’s scale of electronegativity. Those topics are important to understanding the underlying stoichiometry but not in reading the chemistry.

A final important note is that you generally want the simplest whole number ratio unless you are working with an organic compound or biomolecule. The prior example C₆H₁₂O₆ is not in the simplest whole number ratio, but it is a biomolecule and thus is an exception to this rule.

In summary…

• Let’s make an important distinction between compounds and molecules:
  • All compounds are molecules but not all molecules are compounds.
    • Compound: chemicals composed of two or more atoms of different elements
    • Molecule: chemicals composed of two or more atoms of the same OR different elements
• How do we express how many atoms of each element are present in a given compound?
  • One atom
    • We do not place any subscript next to the elemental symbol within the compound
      • Ex: NaCl
        • Sodium (Na) and chlorine (Cl) are in a 1:1 ratio. We don’t express it as Na₁Cl₁ because it is already assumed with the presence of each element.
  • Two or more atoms
    • We attach a subscript next to the elemental symbol within the compound
      • Ex: C₆H₁₂O₆ (glucose)
        • There are six carbon, six oxygen, and twelve hydrogen atoms in a molecule of glucose. We know this because of the numerical subscripts that come after each elemental symbol.

For Ions

So far, we have only discussed neutral atoms. These are atoms that possess an equal number of protons and electrons. We do not consider neutrons since they are not relevant to charge. While many atoms in our universe are neutral, not all of them are. This is where ions come into play. Ions are chemical species that possess an unequal number of protons and electrons. If a compound has more protons than electrons, it is a cation and positively charged. Metals commonly make cations. If a compound has more electrons than protons, it is an anion and negatively charged. Non-metals commonly make anions.

How we express ionic charges is with numerical superscripts followed by a plus (+) or minus (-) signs. Ca²⁺ is a great example of a cation with two electrons fewer than the number of protons, while Cl^– is a great example of an anion with one electron greater than the number of protons. Some students find the positive and negative charges counterintuitive since a negative charge corresponds to an excess of electrons relative to protons, while a positive charge corresponds to a lack of electrons relative to protons. However, this is purely by convention and has to do with the assigned charges of the fundamental particles. Electrons are negatively charged, so it follows that an excess of negatively charged particles confers a negative overall charge on a given atom.

Another important distinction to make is monatomic versus polyatomic ions. Monatomic ions are ions that are formed by a single element containing an excess or a lack of electrons. Sometimes, compounds may also have an excess or a lack of electrons. These ionic compounds are called polyatomic ions. We represent their charges the same way as a monatomic ion. The only difference is that the superscript follows the last element in the ionic compound. For example, a sulfate ion has the formula SO₄^2-.

In summary…

• Ions are chemical species with an excess OR a lack of electrons relative to the number of protons, giving them a net charge that is not 0
  • So how do we represent these charges?
    • With a superscript and positive (+) or negative (-) sign
      • Ex:
        • Lithium ion, Li⁺
        • Oxide, O^2-
• There are several subclasses of ions. They are based on type of charge and how many atoms are present in the given ion.
  • Charge:
    • Cation: an ion with fewer electrons than protons, giving it a net positive charge
      • How are cations denoted? With a positive (+) sign
      • Ex:
        • Chromium (III) ion, Cr³⁺
    • Anion: an ion with more electrons than protons, giving it a net negative charge
      • How are anions denoted? With a negative (-) sign
      • Ex:
        • Iodide, I^–
  • Number of atoms:
    • Monatomic: an ion containing a single atom/element
      • Ex:
        • Nitride, N^3-
    • Polyatomic: an ion containing more than one atom
      • How do we represent polyatomic ions? With the charge superscript following the last element in the formula
      • Ex:
        • Carbonate ion, CO₃^2-
        • Ammonium ion, NH₄⁺

For Ionic Compounds

Sometimes, individual ions bond together to form ionic compounds, a subclass of compounds. While they follow many of the rules that generic compounds do, they have some important conventions we must consider. One such convention is that we write the cation before the anion. Previously, we gave the example of NaCl in the “For Compounds” section. The sodium ion has a charge of positive one (Na⁺) and is hence why it comes before the chloride ion, which has a charge of negative one (Cl^–). Additionally, just as we did with elements in generic compounds, we use subscripts after each ion in the formula to denote how many of each are present. For example, CaCl₂. There is only one calcium ion (Ca²⁺), while there are two chloride ions (Cl^–).

But what about ionic compounds containing polyatomic ions? We have to treat polyatomic ions like monatomic ions in a formula, meaning that the elements that compose them are inseparable and should be seen as a single unit. If there is only one molecule of a given polyatomic ion in the ionic compound, it is written as a unit and placed accordingly if it is cationic or anionic. A good example is calcium carbonate, CaCO₃. Both ions are in a 1:1 ratio. The calcium ion (Ca²⁺) comes first since it is cationic, unlike the anionic carbonate ion (CO₃^2-). If there is more than one molecule of a given polyatomic ion in the ionic compound, the entire polyatomic ion is placed in parentheses and followed by the appropriate numerical subscript. A good example is calcium nitrate, Ca(NO₃)₂. The ions are in a 1:2 ratio. The calcium ion (Ca²⁺) comes first since it is cationic, unlike the anionic nitrate ion (NO₃^–).

In summary…

• Ionic compounds are a subclass of compounds with some additional rules for reading and interpretation.
  • Rules:
    • 1) Cations are written before anions
      • Ex:
        • The one Ca²⁺ ion becomes before the two Cl^– ions in CaCl₂
    • 2) Each ion is followed by a numerical subscript to denote how many there are in the ionic compound’s formula
      • Example:
        • CaCl₂ contains one Ca²⁺ ion and two Cl^– ions
    • 3) Polyatomic ions must be treated as units and represented as such when there are one or more present in the ionic compound’s formula
      • Example #1: Only one polyatomic ion of a given type present in the formula
        • CaCO₃ contains one Ca²⁺ ion and one CO₃^2- ion
      • Example #2: More than one polyatomic ion of a given type present in the formula
        • Ca(NO₃)₂ contains one Ca²⁺ ion and two NO₃^– ions

For Isotopes

Sometimes, neutral atoms have an unequal number of neutrons relative to protons. Two atoms that have the same number of protons but not the same number of neutrons are called isotopes. Isotopes have the same chemical identity on the atomic level since protons are unique to a given element. Thus, if two atoms are isotopes, they are the same element. We represent isotopes using the following generic nuclear notation form, ^A_ZX. A represents the number of nucleons, the total number of protons and neutrons, while Z represents the number of protons and X is the chemical symbol of the element. Since the number of protons directly determines the isotope’s chemical identity, some chemists will omit Z altogether. This omission results in the generic form ^AX, where A is the number of nucleons and X is the chemical symbol of the element.

Let’s consider carbon-14, which is written as ¹⁴₆C. Carbon’s elemental symbol is C, contains six protons, and has a total of fourteen nucleons. Six of those nucleons are carbon’s protons, while neutrons contribute the remaining eight. An isotope of carbon-14 is carbon-12. Carbon-12 is represented as ¹²₆C and has an equal number of protons and neutrons.

In summary…

• Isotopes are two atoms that have the same number of protons but a different number of neutrons
• Isotopes are written using nuclear notation, which takes the following generic form:
  • ^A_ZX
    • Where:
      • A = the number of nucleons, the total number of protons and neutrons
      • Z = the number of protons
      • X = the chemical symbol of the element
    • Note!
      • Since the number of protons is unique to each individual element, chemists sometimes omit Z in nuclear notation
    • Ex: Carbon-14
      • ¹⁴₆C
        • A = 14, 6 protons and 8 neutrons
        • Z = 6 protons
        • X = C, representing carbon