Shells
Valence Electrons
The electrons in the outermost shell control most chemical bonding. If that shell is nearly empty or nearly full, the atom is more likely to react.
Chemistry · Lesson 02
Molecular Bonding Lab
The periodic table tells you which electrons sit on the outside of an atom. Bonding explains what those outer electrons do next. In this lab, drag atoms together and watch whether electrons transfer or overlap, then connect that behavior to crystal salts, polar molecules, and the reason water is liquid at room temperature.
Big Idea
Atoms react because outer shells want to become more stable.
Most atoms in this lesson are trying to reach 8 valence electrons, which is called the octet rule. Hydrogen is the special case: its first shell is full with 2 electrons. Metals tend to give electrons away more easily. Nonmetals tend to pull electrons toward themselves.
Stability
Octet Rule
Main-group atoms often become more stable when their outer shell reaches 8 electrons. Hydrogen follows a duet rule and becomes stable with 2.
Pull
Electronegativity
Electronegativity measures how strongly an atom attracts electrons in a bond. A large difference often leads to ionic transfer. A smaller difference usually leads to sharing.
Outcome
Ionic vs Covalent
Metal + nonmetal often becomes ionic. Two nonmetals usually become covalent. Those different structures help explain why substances have different physical properties.
Atomic Structure
Electron shell diagrams
Each ring is an electron shell. The outermost ring is the valence shell — the electrons that actually do the bonding in the lab below. Compare how full or empty that outer ring is for every element in this lesson.
Sodium (Na)
2, 8, 1
Magnesium (Mg)
2, 8, 2
Chlorine (Cl)
2, 8, 7
Oxygen (O)
2, 6
Hydrogen (H)
1
Sodium's outer shell holds just 1 electron — easy to lose. Chlorine's outer shell holds 7 of 8 — one electron away from full. That imbalance is exactly why sodium and chlorine react so readily.
Trend
Electronegativity across the periodic table
Electronegativity rises left-to-right across a period (more protons pull on the same shell) and rises bottom-to-top within a group (electrons sit closer to the nucleus). Compare F (3.98) above Cl (3.16) in the same column.
Lower pull · metals lose electrons Higher pull · nonmetals gain electrons
Interactive Lab
Atom Snapper
Choose a pair, then drag the atoms close together or use the Snap together button. Metal + nonmetal pairs should transfer electrons. Two nonmetals should overlap and share electrons.
Drag atoms together
The lab simplifies Lewis-style bonding on purpose. Use it to watch the decision: transfer or share. Then read the chemistry notes on the right for the full explanation.
Formal Notation
Lewis dot structures
This is the textbook shorthand for what you just dragged together above. Each dot is a valence electron, and the same four pairs are shown before and after bonding.
Na gives its 1 valence electron to Cl. Both become charged ions, shown in brackets with their charge.
Mg gives up 2 electrons to O. Each ion now has a full outer shell of 8.
H and Cl share one pair of electrons. No brackets — this is a molecule, not a pair of ions.
Identical atoms share the pair equally — a nonpolar covalent bond.
Properties
Why salt is a crystal but water is a liquid
Bonding behavior does not stop once the atoms connect. The way particles arrange in bulk changes melting point, state of matter, brittleness, and conductivity.
Ionic Example
Salt forms a crystal lattice.
In sodium chloride, positive and negative ions lock into a repeating three-dimensional pattern. That rigid ionic lattice makes salt a brittle crystal with a high melting point.
Covalent Example
Water is made of separate H2O molecules.
Inside each water molecule, electrons are shared covalently. Between molecules, hydrogen bonding pulls them together strongly enough for a liquid, but not into the same rigid kind of ionic crystal lattice that salt forms.
Rule of Thumb
Electron arrangement predicts physical behavior.
If atoms become charged ions and build a lattice, expect crystal-like behavior. If atoms stay inside discrete molecules, expect lower melting points and states like gases or liquids more often.