HS Chemistry - Essentials

Isotopes & Ions

Overview of The Page

This page will cover:

• What are isotopes? How are they written?
• What are ions? How are they written?

Isotopes and ions are important to learning about atoms.

Isotopes

Isotopes are atoms of an element that contain the same number of protons but different numbers of neutrons. Since they have the same number of protons, their atomic number remains the same, and they are still atoms of the same element. However, with different number of neutrons, they have different atomic masses.

Isotopes of an element are often denoted as [Element Name]-[Mass Number]. For example, Carbon-12 is an isotope of Carbon with a mass of 12 amu. Since Carbon is element #6, it has 6 protons, which means Carbon-12, with a mass of 12 amu, has $12 \- 6$ neutrons, or 6 neutrons. Similarly, Carbon-14 is an isotope of Carbon with a mass of 14 amu, and since it also has 6 protons, it has $14 \- 6$ neutrons, or 8 neutrons.

Different isotopes of an element are available in different amounts. For example, Hydrogen-1 $1 proton, 0 neutrons$ is the most abundant isotope of Hydrogen, while Hydrogen-2 atoms $1, proton, 1 neutron, also known as deuterium$ makes up roughly 1% of all Hydrogen atoms, and Hydrogen-3 $1 proton, 2 neutrons, also known as tritium$, is only available in very tiny amounts $known as trace amounts$.

It seems as though each element could have an infinite number of isotopes, as the number of neutrons does not affect what element the atom is. However, if an atom has too many neutrons in its nucleus, it becomes unstable, and loses the excess neutrons. Thus, only stable isotopes are considered when counting isotopes.

Given that an element can have multiple isotopes, it is important to note what isotope one is talking about. The [Element Name]-[Mass Number] notation is one method of doing so. However, the following notation is also used:

In this notation, X is the element symbol, A is the atomic mass of the isotope, and Z is the atomic number of the element. In this notation, Carbon-12, for example, would be written as:

And Carbon-14 would be written as:

After taking into account the relative abundance of each isotope to other isotopes of the same element, the masses of the different stable isotopes are then averaged to give a single number that is used as the average atomic mass for that element. This is why atomic masses on the Periodic Table are often shown as decimals.

And that covers isotopes for now.

Ions

Ions are atoms of an element that contain the same number of protons and neutrons, but different numbers of electrons. They thus have the same atomic number and atomic mass, but different charges.

In ionic bonding $covered more in [Unit 3's Ionic Bonding page](../Unit-3/2-Ionic-Bonding.md)$, ionic bonds are formed between metals and non-metals. The non-metal turns into an anion $an ion with a negative charge$, and the metal turns into a cation $an ion with a positive charge$. The metal is left with a positive charge, and the non-metal is left with a negative charge. They are then attracted to one another, and bond.

When an atom $usually a non\-metal$ is turned into an anion, it gains and adds electrons to its valence shell until the valence shell is completed. Cl $Chlorine$ for example, with 7 valence electrons, only needs one more electron to complete its valence shell. Thus, in an ionic reaction, Cl will gain one more electron to get 8 valence electrons, becoming Cl^-. After this, Cl will not take in any more electrons because its outer shell is full $octet rule$. Thus, the only stable Cl anion is Cl^-.

When an atom $usually a metal$ is turned into a cation, it loses electrons from its valence shell until the outermost shell is empty, at which point the shell below it, which is full, will become the new outermost shell. Na $Sodium$, for example, has one valence electron in its outer shell. Thus, in an ionic reaction, Na will lose its electron and become Na⁺. After this, Na will not lose any more electrons because it now has a full outer shell $octet rule$. Thus, the only stable Na cation is Na⁺.

But why does it happen this way? Why do metals become cations, while non-metals become anions? Why not the other way?

Metals, with few valence electrons $less than 4$, tend to have lower ionization energies for their valence electrons, while non-metals, with many valence electrons $more than 4$, tend to have higher ionization energies for their valence electrons. The ionization energy is the amount of energy required to remove an electron from an atom. The 1^st ionization energy is the amount required to remove the first electron from the atom $e.g. turning Al into Al^\+^$, the 2^nd ionization energy is the amount of energy required to remove the second electron $e.g. turning Al^\+^ into Al^2\+^$, and so on. This is covered more in Electronegativity & Ionization Energy.

Since cations and anions are oppositely charged particles, they can bond to one another to form ionic compounds (covered more in Ionic Bonding). And since ions are charged particles, free-moving ions can conduct electricity.

And that covers ions for now.