Source: Marine Biology Ch. 1, Intro to Marine Biology (Florida International University)
Difficulty: Introductory | Prerequisites: None. Basic familiarity with chemistry concepts (atoms, electrons) is helpful but not required.
This chapter lays the chemical foundation for everything that follows in marine biology. Before you can understand how ocean currents distribute nutrients, why coral reefs bleach, or how deep-sea organisms survive crushing pressures, you need to understand what water actually is and why it behaves the way it does. The material here covers particle theory, atomic structure, and the three types of chemical bonding most relevant to ocean chemistry: covalent, ionic, and hydrogen bonds. If you are coming into the course cold, this is where you start.
Water (H₂O) is a polar covalent compound whose bent molecular shape creates partial charges, allowing hydrogen bonds to form between molecules. These hydrogen bonds give water its unusual properties, including its ability to act as the ocean's primary solvent. The three bond types you need for this course are covalent (shared electrons, strong), ionic (transferred electrons, attraction between opposite charges), and hydrogen (weak electrostatic attraction between molecules).
Atom
The smallest unit of an element that retains the properties of that element. Made up of protons (+), neutrons (neutral), and electrons (-). In simple terms, atoms are the building blocks of all matter. Everything in the ocean, from salt to whale bone, is made of atoms.
Compound
A substance formed when two or more different elements bond together in a fixed ratio. Water is a compound: two hydrogen atoms to one oxygen atom (2H:1O). Think of it as a recipe with exact proportions. Change the ratio and you get a different substance entirely.
Emergent properties
New characteristics that a compound gains through bonding, which neither of its individual elements possessed alone. In simple terms, when atoms bond to form a compound, the result behaves differently from the raw ingredients. Water can dissolve salt; isolated hydrogen and oxygen atoms cannot.
Covalent bond
A chemical bond in which two atoms share one or more pairs of electrons. Both atoms contribute electrons and share them equally (or near-equally). Think of it as a mutual arrangement: both atoms use the shared electrons to become stable. These bonds are strong and require significant energy to break.
Ionic bond
A chemical bond formed by the electrostatic attraction between two oppositely charged ions, where one atom has transferred an electron to another. Think of it as one atom giving an electron away entirely, creating a positive ion and a negative ion that then attract each other.
Ion
An atom that has gained or lost an electron from its outer shell, giving it a net electrical charge. Losing an electron makes the ion positive (cation); gaining one makes it negative (anion). In simple terms, ions are atoms that are no longer electrically neutral because they have handed off or picked up an electron.
Hydrogen bond
A weak bond between two molecules caused by the electrostatic attraction between a hydrogen atom in one molecule and an electronegative atom (oxygen, nitrogen, or fluorine) in another. Think of it as a gentle tug between neighbouring water molecules. Individually weak, but collectively responsible for many of water's unusual properties.
Polarity
An uneven distribution of electrical charge across a molecule, creating a positive end and a negative end. In simple terms, a polar molecule has a lopsided charge. Water is polar because oxygen hogs the shared electrons, leaving hydrogen slightly positive and oxygen slightly negative.
Electronegativity
A measure of how strongly an atom attracts shared electrons towards itself within a bond. Think of it as how greedy an atom is for electrons. Oxygen is more electronegative than hydrogen, which is why water molecules end up polar.
Solvent
A substance capable of dissolving other substances (solutes) to form a solution. In simple terms, water is called the "universal solvent" because its polarity lets it pull apart and surround many different compounds.
Matter exists in three states, each defined by the energy level and spacing of its particles.
Solids: Low energy, densely packed particles that vibrate in place. Water as a solid (ice) is unusual: it forms a crystalline lattice that spaces particles further apart, making ice less dense than liquid water.
Liquids: Moderate energy, particles less tightly packed and free to move past one another. Liquid water has more energy than ice but less than steam.
Gases: Highest energy, large spaces between particles. Gas molecules collide frequently, transferring energy. These collisions are what allow gaseous water to absorb and distribute heat.
Atoms are the smallest unit of an element and are composed of protons (positive charge), electrons (negative charge), and neutrons (no net charge).
The nucleus (protons + neutrons) sits at the centre; electrons orbit in shells around the nucleus.
An element's bonding behaviour depends on the number of electrons in its outer shell and how many it needs to share, gain, or lose to become stable.
Two atoms share a pair of electrons. Each atom contributes one electron to the shared pair, and both use the pair equally.
Found between non-metal elements
Among the strongest bond types; require substantial energy to break
Water (H₂O) is the most common covalent compound
Covalent compounds can exist in any state: solid, liquid, or gas
Common covalent compounds in the ocean: water, carbon dioxide (CO₂), and molecular oxygen (O₂)
Formed through the transfer (not sharing) of electrons. One atom loses an electron and becomes a positive ion (cation); the other gains that electron and becomes a negative ion (anion). The bond is the electrostatic attraction between these opposite charges.
If an atom loses an electron, the resulting ion is positive
If an atom gains an electron, the resulting ion is negative
Sodium chloride (table salt, NaCl) is the classic example: sodium donates an electron to chlorine
A weak electrostatic attraction between a hydrogen atom in one molecule and an electronegative atom (oxygen, nitrogen, or fluorine) in a neighbouring molecule.
Easily broken individually, but collectively powerful in large numbers
In water, hydrogen bonds form because of the molecule's polarity
A water molecule consists of two hydrogen atoms covalently bonded to one oxygen atom. The molecule's overall shape is bent (roughly 104.5°), caused by electron repulsion around the oxygen.
Oxygen is more electronegative than hydrogen, so it pulls shared electrons closer to itself
This creates a partial negative charge (δ-) on the oxygen side and partial positive charges (δ+) on each hydrogen
The positive end of one water molecule attracts the negative end of another, forming a hydrogen bond
This polarity is what makes water an effective solvent: it can surround and pull apart ionic compounds and interact with other polar molecules
Water molecule formula: H₂O (2 hydrogen : 1 oxygen)
Molecular geometry: Bent shape, approximately 104.5° bond angle
Charge distribution: Oxygen carries a partial negative charge (δ-); each hydrogen carries a partial positive charge (δ+)
Covalent bond in water: Each O-H bond is a shared electron pair
Hydrogen bond: The δ+ hydrogen of one water molecule aligns with the δ- oxygen of an adjacent molecule
Ice crystal structure: When water freezes, hydrogen bonds lock molecules into a hexagonal lattice that is less dense than liquid water (this is why ice floats)
Water's polarity and hydrogen bonding are the reason salt dissolves when it enters the ocean. The partially positive hydrogens surround chloride ions (Cl⁻) while the partially negative oxygens surround sodium ions (Na⁺), pulling the crystal apart. This same mechanism drives nutrient transport throughout marine ecosystems.
Ice floating rather than sinking is not a trivial detail. If ice were denser than liquid water, lakes and oceans would freeze from the bottom up, killing most aquatic life. The insulating layer of surface ice is what allows organisms beneath it to survive polar winters.
Students often think hydrogen bonds and covalent bonds are similar in strength. They are not. Covalent bonds (within a molecule) are strong and require substantial energy to break. Hydrogen bonds (between molecules) are weak and break easily.
Students sometimes assume ionic bonds involve sharing electrons. They do not. Ionic bonding is a transfer: one atom gives up an electron entirely, the other accepts it.
A common error is thinking ice is denser than liquid water because solids are "usually" denser. Water is the exception. Its crystalline structure spaces molecules further apart, making ice less dense.
Students confuse polarity with ionic character. A polar molecule like water still has covalent bonds; the electrons are shared, just unequally. In an ionic bond the electron is transferred outright.
⚠️ Know the difference between covalent, ionic, and hydrogen bonds. Be able to identify each by how electrons are handled (shared, transferred, or attracted between molecules).
⚠️ Understand why ice floats. The crystalline structure of ice and its lower density compared to liquid water is a frequently tested concept.
⚠️ Be able to explain how water's polarity arises from oxygen's higher electronegativity and the bent molecular shape.
⚠️ Remember that water is the most common covalent compound and functions as a solvent because of its polarity.
⚠️ Exam terminology note: refer to the loss/gain of electrons in ionic bonding precisely. Loses electron = positive ion. Gains electron = negative ion. Do not mix these up.
True or false: Hydrogen bonds are stronger than covalent bonds.
Fill in the blank: When an atom loses an electron, it becomes a ______ ion.
True or false: Ice is denser than liquid water.
Fill in the blank: Water's bent molecular shape is caused by electron ______.
True or false: Water is a polar molecule because oxygen is more electronegative than hydrogen.
Answers: 1. False 2. Positive 3. False 4. Repulsion 5. True
Q: What type of bond holds the atoms within a single water molecule together?
A: Covalent bonds. Each oxygen-hydrogen bond is a shared pair of electrons.
Q: Explain why water molecules form hydrogen bonds with each other.
A: Oxygen is more electronegative than hydrogen, pulling shared electrons closer and creating a partial negative charge on the oxygen side and partial positive charges on each hydrogen. The positive end of one molecule attracts the negative end of another, forming a hydrogen bond.
Q: Why does ice float on liquid water?
A: When water freezes, hydrogen bonds lock molecules into a crystalline lattice that spaces them further apart than in liquid water. This makes ice less dense than liquid water, so it floats.
Q: Distinguish between ionic and covalent bonding.
A: In covalent bonding, atoms share electrons. In ionic bonding, one atom transfers an electron to another, creating oppositely charged ions that attract each other.
Q: If an atom gains an electron, what charge does the resulting ion carry?
A: Negative. Gaining an electron adds a negative charge, making the ion an anion.
This material connects directly to Chapter 1.2 (Solubility in Water), where water's role as a polar solvent explains how salts dissolve in the ocean and how salinity is established. It also underpins Chapter 1.3 (Density and Pressure), since temperature, salinity, and pressure all affect how water molecules pack together.
Beyond this course, the principles here appear in any discipline that deals with aqueous chemistry: environmental science, biochemistry, and chemical oceanography all build on the same foundation of polarity, bonding, and solvent behaviour.
Water properties, particle theory, states of matter, atomic structure, protons, electrons, neutrons, covalent bond, ionic bond, hydrogen bond, polarity, electronegativity, emergent properties, compound, solvent, ice density, crystalline structure, bent molecular shape, water chemistry, marine chemistry, BIO 101, intro marine biology, Chapter 1