Describe how electrons are arranged in energy levels, write electron configurations and Lewis dot structures, and explain why valence electrons determine chemical behavior (MA STE HS-PS1-1, patterns of electrons).

What this topic is asking

Standard HS-PS1-1 is explicitly about the patterns of electrons in the outermost energy level, and using them to predict an element's properties. To do that you must know how electrons fill energy levels, how to count the valence (outer) electrons, and why those outer electrons control how an atom bonds and reacts. Massachusetts chemistry builds the whole periodic table and all of bonding on this single idea: chemistry is the behavior of valence electrons.

Energy levels and electron configuration

Electrons are not scattered randomly; they occupy energy levels (also called shells) at set distances from the nucleus. Lower levels are closer to the nucleus and lower in energy, and they fill first. For the elements covered in a first chemistry course, a simple shell model works well:

• First shell: holds up to 2 electrons.
• Second shell: holds up to 8 electrons.
• Third shell: holds up to 8 electrons for the main-group elements you meet first.

So sodium (11 electrons) is written 2, 8, 1 by shell, and chlorine (17 electrons) is 2, 8, 7. This shell notation is the level of detail Massachusetts high school chemistry expects; the full subshell notation (1s, 2s, 2p, and so on) is an AP and college extension. The key habit is to fill from the inside out and stop when you have placed all the electrons.

Valence electrons

For a main-group element, the number of valence electrons equals the group number's ones digit: group 1 has 1, group 2 has 2, group 13 has 3, up to group 18 with 8 (except helium with 2). That is why elements in the same column behave alike, a pattern explored in the periodic table and periodic trends. Sodium has 1 valence electron, so it tends to lose 1 and form $\text{Na}^+$. Chlorine has 7, so it tends to gain 1 and form $\text{Cl}^-$.

The octet rule

The octet rule states that atoms gain, lose, or share electrons in order to reach a full outer shell of 8 electrons, the stable arrangement of the noble gases. (The exception is the smallest atoms, which aim for 2, the full first shell, like helium.) The octet rule is the single best predictor of behavior in a first chemistry course:

• Atoms with 1 to 3 valence electrons tend to lose them, becoming positive ions.
• Atoms with 5 to 7 valence electrons tend to gain electrons, becoming negative ions.
• Atoms with a full octet (the noble gases) are stable and very unreactive.

Lewis dot diagrams

A Lewis dot diagram (electron dot diagram) shows an element's symbol surrounded by dots for its valence electrons only. You place dots one at a time on the four sides before pairing them. Oxygen, with 6 valence electrons, has two lone pairs and two single dots, showing it needs 2 more electrons to complete its octet. These diagrams make bonding visual and are used constantly in Module 2.

Try this

Q1. Write the electron arrangement (by shell) for an atom with 15 electrons and state its number of valence electrons. [2]

• Cue. 2, 8, 5; it has 5 valence electrons (this is phosphorus).

Q2. Why do all group 1 elements form $1+$ ions? [2]

• Cue. They each have 1 valence electron, which they lose to reach a full outer shell, giving a $1+$ charge.