Definition
Ionic bond: forms when a metal transfers electron(s) to a non-metal, creating oppositely charged ions that attract — metals lose electrons easily, non-metals gain them easily.
Covalent bond: forms when two non-metal atoms share electron pairs instead of transferring them.
Rule
Ionic compounds: high melting points (strong attraction throughout a rigid crystal lattice), usually soluble in water, conduct electricity when molten or dissolved.
Covalent compounds: generally low melting points, often insoluble in water, don't conduct electricity (no free charged particles).
Example 1 — why NaCl forms a crystal
Briefly describe how ions arrange themselves to form a NaCl crystal.
- Each Na ion is surrounded by Cl ions and vice versa, in a regular repeating (lattice) pattern.
- This maximizes attraction between opposite charges and minimizes repulsion between like charges.
Answer: A regular ionic lattice of alternating Na and Cl
Definition
In a metal, positive metal ions sit in a "sea" of freely moving, delocalized electrons — this shared electron sea is what's called the metallic bond.
Remember
This sea of free electrons explains metals' properties directly: they conduct electricity and heat well (electrons carry charge/energy freely), and are malleable/ductile (layers of ions can slide past each other while the electron sea keeps holding them together).
Example 2 — why water is a liquid but H₂S is a gas
Explain why is a liquid at room temperature while is a gas.
- Oxygen is far more electronegative than sulfur, so O–H bonds are strongly polar and water molecules form hydrogen bonds with each other.
- molecules, with much weaker intermolecular forces, need far less energy to separate into a gas.
Answer: Hydrogen bonding in water gives it a much higher boiling point than
Remember
The same reasoning explains why HF is a liquid while HCl is a gas: HF hydrogen-bonds strongly (F is highly electronegative), HCl does not.
Definition
Graphite: carbon atoms in flat, layered sheets — each carbon bonds to only others, leaving one free (delocalized) electron per atom that can move along the layer, making graphite electrically conductive and slippery (a lubricant, since the layers slide over each other).
Diamond: each carbon bonds to others in a rigid 3D network — no free electrons, so it doesn't conduct electricity, but the strong network makes it extremely hard.
Quick Recap
- ●Ionic: metal + non-metal, electron transfer, high melting point, conducts when molten/dissolved
- ●Covalent: non-metal + non-metal, shared electrons, low melting point, usually doesn't conduct
- ●Metallic bond: positive ions in a sea of free electrons — explains conductivity and malleability
- ●Hydrogen bonding (as in HO, HF) raises boiling point far above similar molecules without it
- ●Graphite: layers, free electrons, conducts + lubricates. Diamond: rigid 3D network, no free electrons, very hard, insulator
