IONS
7.1
Valence Electrons, The Octet Rule, and
formation of Cations and Anions
Think about it…
• The most common example of an ionic compound is Sodium
Chloride (table salt)
• What is it about sodium and chlorine atoms that cause them to
combine?
• Why is the formula unit for sodium chloride NaCl and not
Na2Cl? or NaCl2 ?
Valence Electrons
• Elements within each group of the periodic table behave
similarly because they have the same number of valence
electrons
• Valence Electrons – the electrons in the highest
occupied energy level of an element’s atoms.
• This can be determined by looking at the element’s
electron configuration
• For representative elements (not the d or f block), the number
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of valence electrons is the same as the group number
Ex. Group 1A elements (H, Li, Na, K, etc.) have 1 valence
electron.
Carbon and Silicon (in Group 4A) have 4 valence electrons.
Nitrogen and Phosphorous have 5 valence electrons. They are
in group 5A.
If you count noble gases as group 0, they are the exception.
They have 8 valence electrons.
Another way to figure out the number of valence electrons is to
count the position of the element in it’s row on the periodic
table. For example, Lithium is 1st in in it’s row. It has one
valence electron. Fluorine is 7th in it’s row. It has 7 valence
electrons.
Dot structures
• Valence electrons are usually the only electrons used in
chemical reactions.
• Sometimes it’s helpful to see what electrons are doing in
chemical reactions by drawing models.
• We show these electrons using electron dot structures (or
also called Lewis Structures)
• Electron Dot Structures – diagrams that show valence
electrons as dots surrounding the symbol of the element
• Ex. Sulfur
• Electron Configuration
• 1s2 2s2 2p6 3s2 3p4
• Valence Electrons
• 6
• Dot Structure is shown at right
• When writing a dot structure:
• Dots are placed around the element’s symbol with a maximum of 2
dots on each side
• The exact location of the dots is not critical, but we put on dot on
each side first before pairing them up. It’s good to show them in
pairs usually, but there are exceptions when you draw several
together in a bond. We’ll look at those later.
Write the electron dot structures for the following. You
can check your book for the answers.
• Fluorine
• Magnesium
• Phosphorous
• Krypton
• Oxygen
• Nitrogen
• Chlorine
The Octet and Duet Rules
• We know that Noble Gases like Neon and Argon are
unreactive.
• This was explained by Gilbert Lewis who came up with the
Octet Rule in 1916 as part of the Lewis Theory of Bonding
• Octet Rule – Atoms tend to achieve the electron configuration
of a noble gas when forming compounds
• Octet – set of 8 electrons in the highest energy orbital
• Atoms with an octet are stable!
• You made electron dot structures on the last page for atoms that follow
the octet rule
• Hydrogen, lithium, Beryllium and Helium are exceptions – their
most stable electron configuration is known as a duet (2
electrons in their highest energy level)
●He● (this is the electron dot structure for helium, showing the how it
follows the duet rule.)
Ion Formation: Cations
• Metals tend to lose/give away electrons to form cations. This
leaves an octet or duet in the next-lowest energy level
• Example: Sodium
• Electron configuration: 1s2 2s2 2p6 3s1
• Valence electrons: 1 (in the 3s orbital)
• Lewis Structure: Na●
• Sodium tends to lose this electron, becoming Na+
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Na Na+ + e-
• New electron configuration: 1s2 2s2 2p6
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This is the same electron configuration as Neon
• New valence electrons: 8 (in the 2s and 2p orbitals)
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New Lewis Structure: (brackets, no dots, charge on outside)
• For transition metals, the charges of cations may vary. They
sometimes form exceptions to the octet rule. We’ll learn about
transition metal cations at a later time.
Formation of compounds: Cations
• Example: How does magnesium become more stable
when forming a compound?
• Magnesium gives away 2 valence electrons to attain the
same electron configuration as Neon (an octet)
• Mg  Mg2+ + 2e• 1s2 2s2 2p6 3s2  1s2 2s2 2p6
Ion Formation: Anions
• Nonmetals tend to take on electrons to form anions. This
also creates an octet or duet in the next-lowest energy
level.
• Example: Fluorine
• Electron configuration: 1s2 2s2 2p7
• Valence electrons: 7 (in the 2p orbital)
• Lewis Structure: (see right )
• Chlorine tends to gain one electron, becoming Cl•
Cl + e-  Cl-
• New electron configuration: 1s2 2p6
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This is the same electron configuration as Neon
• New valence electrons: 8 (in the 2s and 2p orbitals)
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New Lewis Structure: (brackets, full set of dots, charge on outside)
Formation of Compounds: Anions
• Example: How does chlorine become stable when forming
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a compound?
Chlorine gains one electron to gain the same electron
configuration as argon
Cl + e-  Cl1s2 2s2 2p6 3s2 3p5  1s2 2s2 2p6 3s2 3p6
The ions produced when halogens gain electrons are
called halide ions.
• All halide ions have charges of 1• All halide ions have a full octet in their outermost shell
Dot structures and bonding
• In the Lewis Theory of Bonding:
• A chemical bond involves the sharing or transfer of electrons to
attain stable electron configurations (outer shells with 8 electrons)
for all bonding atoms
• If the electrons are transferred, the bond is an ionic bond
• If the electrons are shared, the bond is a covalent bond.
• The electrons in the bonding atoms are arranged so that all atoms
involved get stable electron configurations.
• Below is an example of an ionic bond. Covalent bonds will be
shown later.