Chemistry: Bonding and Reactions

Everything you can touch, taste, breathe, or drink is made of atoms — and almost always, atoms bonded together into molecules. How atoms bond determines all the properties of matter. How bonds break and re-form is what chemistry is mostly about. If you understand chemical bonding and reactions, you understand why water is wet, why fire burns, why DNA works, and why batteries hold charge.

Atoms bond to achieve a more stable arrangement of electrons. Most atoms are stable when their outermost shell (valence shell) is full — usually 8 electrons (the "octet rule"), or 2 for hydrogen and helium.\n\nThree main ways to achieve this:\n\n1. **Give up electrons** (to become a positive ion)\n2. **Take electrons** (to become a negative ion)\n3. **Share electrons** (covalent bond)\n\nThe specific choice depends on the atoms involved and their affinity for electrons (called **electronegativity**).

**Ionic bonds** form when one atom transfers electrons to another. The atom that gives up electrons becomes positively charged (a cation). The atom that gains them becomes negatively charged (an anion). Opposite charges attract — and hold each other in a crystal lattice.\n\nClassic example: sodium chloride (table salt, NaCl).\n- Sodium (Na) has 1 electron in its outer shell. Giving it up → stable, +1 charge.\n- Chlorine (Cl) has 7 electrons in its outer shell. Gaining one → stable, −1 charge.\n- Na⁺ and Cl⁻ attract in a crystal structure. That's salt.\n\nIonic compounds: typically solids at room temperature, high melting points, dissolve in water, conduct electricity when dissolved or melted. (Salt water conducts electricity because Na⁺ and Cl⁻ ions can move freely.)

**Covalent bonds** form when atoms SHARE electrons rather than transferring them. Each shared pair of electrons counts as one bond.\n\n- **Single bond**: 1 shared pair (like H−H in hydrogen gas)\n- **Double bond**: 2 shared pairs (like O=O in oxygen gas)\n- **Triple bond**: 3 shared pairs (like N≡N in nitrogen gas — the strongest kind of covalent bond)\n\nMost molecules in biology — water (H₂O), glucose, DNA, proteins, fats — are held together by covalent bonds. Covalent compounds tend to be gases, liquids, or low-melting solids, and generally don't conduct electricity.\n\n**Polar covalent bonds** form when the two atoms have different electronegativities — sharing electrons, but unequally. Water is the classic example: oxygen pulls the electrons more than hydrogen does, so the oxygen side has a slight negative charge and the hydrogen side has a slight positive charge. That polarity is why water is such a universal solvent.

**Metallic bonds** — in metals, atoms pool their outer electrons into a shared "electron sea." This explains why metals conduct electricity (free-moving electrons), are malleable and ductile (atoms can slide past each other), and are shiny.\n\n**Intermolecular forces** — forces BETWEEN molecules (not within them). These include:\n- **Hydrogen bonds** — especially between water molecules. Responsible for water's high boiling point, ice's density anomaly, and DNA's double helix structure.\n- **Van der Waals forces** — weak attractions that hold nonpolar liquids together.\n\nIntermolecular forces are weaker than covalent bonds, but they determine whether a substance is solid, liquid, or gas at room temperature.

A **chemical reaction** is when bonds break and re-form, changing one set of molecules into a different set.\n\nFour basic types:\n\n- **Synthesis**: A + B → AB (simple combination)\n- **Decomposition**: AB → A + B (one compound splits)\n- **Single displacement**: A + BC → AC + B (one element replaces another)\n- **Double displacement**: AB + CD → AD + CB (partners swap)\n\nCombustion is a special case: a fuel reacts with oxygen to produce CO₂ and water (if complete).\n\nExample — combustion of methane (natural gas):\n**CH₄ + 2O₂ → CO₂ + 2H₂O + energy**\n\nThe atoms are conserved (balanced on both sides) — 1 C, 4 H, 4 O. The bonds have changed. Energy is released because the new bonds are more stable than the old ones.

In every chemical reaction:\n\n- **Mass is conserved** — all atoms present at the start are present at the end (just rearranged).\n- **Energy is conserved** — total energy doesn't disappear or appear, though it can change form (chemical to thermal, for example).\n\nThis is why chemical equations must be **balanced**: the same number and type of atoms on both sides.\n\nReactions can be **exothermic** (release energy — combustion, cellular respiration) or **endothermic** (absorb energy — photosynthesis, ice melting). Chemical energy is stored in the bonds themselves.

Chemical bonding and reactions are the basis of every material and every living process. Understanding them opens up biochemistry (how cells work), materials science (how to design new substances), pharmacology (how drugs affect the body), and environmental science (how pollutants transform in nature). Chemistry is the central science — it connects physics to everything alive.