Energy Changes in Chemical Reactions

Ever wondered why some reactions feel hot whilst others feel cold? It's all about whether energy is being released or absorbed during the process.

Exothermic reactions are the ones that give out energy to their surroundings - think of burning wood or hand warmers. These happen because making new bonds between atoms releases energy. You'll spot these reactions because they warm up their surroundings.

Endothermic reactions do the opposite - they absorb energy from around them, making things feel colder. This happens because breaking bonds between atoms requires energy input. Ice packs use endothermic reactions to cool injuries.

Remember: Bond making = energy out (exothermic), Bond breaking = energy in (endothermic)

Giant Covalent Structures

These massive molecular structures explain why diamonds are so incredibly hard and why graphite makes excellent pencil lead. Giant covalent structures are built from huge numbers of non-metal atoms all connected by strong covalent bonds in regular, repeating patterns.

The main examples you'll encounter are diamond, graphite, and silicon dioxide. Because breaking all those strong covalent bonds requires loads of energy, these materials have very high melting and boiling points.

Most giant covalent structures don't conduct electricity because they lack charged particles. However, graphite is the fascinating exception to this rule.

Allotropes are simply different structural forms of the same element - like how diamond and graphite are both pure carbon but arranged completely differently.

Carbon's Amazing Allotropes

Carbon shows off by forming two completely different structures with opposite properties. Understanding these helps explain why the same element can be both the hardest natural substance and a soft writing material.

Diamond has each carbon atom bonded to four others in a rigid 3D network. This creates an incredibly strong structure with a very high melting point. Since all electrons are locked in bonds, diamond can't conduct electricity - making it an excellent insulator.

Graphite takes a different approach, with each carbon bonded to only three others, forming layers. These layers have weak forces between them, so they slide over each other easily - perfect for pencils! The spare electron from each carbon can move freely, allowing graphite to conduct electricity and heat.

Key difference: Diamond = 4 bonds per carbon (insulator), Graphite = 3 bonds per carbon (conductor)

Ionic Bonding Basics

Ionic bonding happens when metals meet non-metals and electrons get transferred from one to the other. Common examples include sodium chloride (table salt) and magnesium oxide.

Metals always lose electrons to become positive cations, whilst non-metals gain these electrons to become negative anions. This creates oppositely charged ions that attract each other strongly.

These ions arrange themselves in regular 3D lattices held together by strong electrostatic forces. This gives ionic compounds high melting and boiling points. Solid ionic compounds can't conduct electricity because the ions are fixed in place, but they conduct when molten or aqueous (dissolved in water) because the ions can move freely.