Atoms, Elements, Molecules, and Chemical Bonds

Core Concepts

Matter and atoms

Matter is anything that occupies space and has mass. All matter is composed of atoms, the smallest units of an element that retain the element's chemical properties. Atoms are incredibly small — roughly 0.1 nanometers in diameter.

Atomic structure

An atom consists of three types of subatomic particles:

• Protons: Positively charged particles located in the nucleus. Each proton has a mass of approximately 1 atomic mass unit (amu).

• Neutrons: Neutral (uncharged) particles located in the nucleus. Each neutron has a mass of approximately 1 amu.

• Electrons: Negatively charged particles that orbit the nucleus in regions called electron shells. Electrons have negligible mass compared to protons and neutrons.

In a neutral atom, the number of electrons equals the number of protons, so the overall charge is zero.

Atomic number and mass number

• Atomic number: The number of protons in an atom's nucleus. This defines the element. Every carbon atom has 6 protons; every oxygen atom has 8.

• Mass number: The total number of protons plus neutrons in the nucleus. Carbon-12 has 6 protons and 6 neutrons, so its mass number is 12.

Isotopes

Isotopes are atoms of the same element that differ in the number of neutrons. Carbon-12 has 6 neutrons; carbon-13 has 7 neutrons; carbon-14 has 8 neutrons. All are carbon (6 protons), but their mass numbers differ. Some isotopes are radioactive and decay over time, which makes them useful in scientific dating and medical imaging. Biologically, isotopes of the same element behave nearly identically in chemical reactions.

Elements

An element is a substance that cannot be broken down into simpler substances by chemical means. Each element consists of atoms with a specific number of protons. About 25 of the 92 naturally occurring elements are essential for life. The most abundant elements in living organisms are:

• Oxygen (O): Component of water and most organic molecules.

• Carbon (C): The backbone of all organic molecules.

• Hydrogen (H): Component of water and all organic molecules; participates in hydrogen bonding.

• Nitrogen (N): Component of proteins and nucleic acids.

• Calcium (Ca): Structural component of bone; signaling molecule.

• Phosphorus (P): Component of nucleic acids, ATP, and phospholipids.

• Potassium (K): Major ion in cellular fluids; nerve signaling.

• Sulfur (S): Component of some amino acids and proteins.

• Sodium (Na): Major ion in extracellular fluid; nerve signaling.

• Magnesium (Mg): Component of chlorophyll; enzyme cofactor.

Trace elements — iron, iodine, zinc, copper, and others — are required in very small quantities but are essential for specific biological functions.

Molecules and compounds

A molecule consists of two or more atoms held together by chemical bonds. A compound is a substance composed of two or more different elements in a fixed ratio. All compounds are molecules, but not all molecules are compounds — O2 is a molecule but not a compound because it contains only one element.

Electrons and chemical bonding

Electrons are arranged in shells (energy levels) around the nucleus. The first shell holds up to 2 electrons; the second and third shells hold up to 8 electrons each (for the elements most relevant to biology). Atoms are most stable when their outermost shell is full — a principle called the octet rule for elements in the second row of the periodic table.

Valence electrons are the electrons in the outermost shell. They determine an atom's chemical behavior. Atoms with incomplete outer shells tend to bond with other atoms to achieve full outer shells.

Types of chemical bonds

Covalent bonds: A covalent bond forms when two atoms share one or more pairs of valence electrons. Covalent bonds are the strongest type of chemical bond in biological systems and are the primary bonds holding organic molecules together.

• Nonpolar covalent bonds: Electrons are shared equally between atoms with similar electronegativity (electron-attracting power). Examples: C—H bonds, O=O in O2.

• Polar covalent bonds: Electrons are shared unequally because one atom attracts electrons more strongly (has higher electronegativity). The more electronegative atom carries a partial negative charge, and the less electronegative atom carries a partial positive charge. Example: O—H bonds in water.

Ionic bonds: An ionic bond forms when one atom transfers one or more electrons to another atom, creating ions (charged atoms or molecules). The resulting positive and negative ions attract each other electrostatically. Example: Na+ and Cl− in table salt (NaCl). In biological systems, ionic bonds are weaker than covalent bonds and can be disrupted by water.

Hydrogen bonds: A hydrogen bond is a weak attraction between a partially positive hydrogen atom (covalently bonded to an electronegative atom such as O or N) and a partially negative atom (usually O or N) in another molecule or another part of the same molecule. Individual hydrogen bonds are weak, but many hydrogen bonds together can be strong — as in the DNA double helix or the structure of proteins.

Ions

An ion is an atom or molecule that has gained or lost one or more electrons, acquiring a net electrical charge. A cation is positively charged (lost electrons); an anion is negatively charged (gained electrons). Na+ (sodium ion), K+ (potassium ion), Ca2+ (calcium ion), and Cl− (chloride ion) are critical for nerve signaling, muscle contraction, enzyme function, and fluid balance.

Bond strength comparison

In biological systems:

• Covalent bonds: strong (do not break under normal cellular conditions without enzymes)

• Ionic bonds: moderate (can be disrupted by water and changes in pH)

• Hydrogen bonds: individually weak, collectively strong

ELI Example

Atoms are like neighbors in an apartment building. Each apartment (shell) has limited spots. Unstable neighbors with empty spots find ways to fill them: sharing a roommate equally (nonpolar covalent), unequally (polar covalent), transferring a roommate (ionic), or just waving across windows (hydrogen bond).

Do Not Confuse

Lab Link

Understanding elements and bonding is essential for the biomolecule testing laboratory activities in Chapter 27. The chemical tests for carbohydrates, proteins, and lipids rely on specific bonding interactions between reagents and functional groups. Recognizing which elements are present in each type of biomolecule helps you predict and interpret test results.

High-Yield Memory Anchors

• Protons = identity of element. Electrons = bonding behavior.

• Covalent = sharing electrons. Ionic = transferring electrons. H-bond = partial charge attraction.

• Valence electrons = outermost shell = determine bonding.

• O, C, H, N = ~96% of living matter.

• Polar = unequal sharing; partial charges matter for biology.

Quick Check

Q1 (Foundational): An atom has 8 protons, 8 neutrons, and 8 electrons. What is its atomic number? Its mass number? What element is it?

Q2 (Application): Oxygen has 8 electrons. How many valence electrons does oxygen have? Based on the octet rule, how many covalent bonds does oxygen typically form?

Q3 (Comparison/Reasoning): Compare the O—H bond in a water molecule with a hydrogen bond between two water molecules. Which is stronger? Which involves electron sharing? Which can be disrupted by heating water to boiling?

Quick Check Answers

A1: Atomic number = 8 (number of protons). Mass number = 16 (8 protons + 8 neutrons). The element is oxygen.

A2: Oxygen has electron configuration 2, 6 — meaning 2 electrons in the first shell and 6 in the second (valence) shell. Oxygen has 6 valence electrons. To complete its outer shell to 8, it needs 2 more electrons, so it typically forms 2 covalent bonds (as in H2O, where oxygen shares electrons with two hydrogen atoms).

A3: The O—H covalent bond within a water molecule is far stronger than a hydrogen bond between two water molecules. The O—H bond involves electron sharing and is not broken by boiling water. Hydrogen bonds between water molecules are weak attractions between the partially positive H of one water and the partially negative O of another. These hydrogen bonds are disrupted when water boils (water molecules separate into vapor), but the covalent O—H bonds within each water molecule remain intact.

Chapter Summary

Atoms are the building blocks of matter. Valence electrons determine bonding: covalent (sharing), ionic (transfer), hydrogen (partial-charge attraction). The most abundant biological elements — H, C, N, O, P, S — form the foundation of all biological molecules.

Common Mistakes

Mistake: "Hydrogen bonds are strong bonds like covalent bonds."

Reality: Individual hydrogen bonds are weak — about 1/20 the strength of a covalent bond. However, many hydrogen bonds together can collectively produce a stable structure, as in the DNA double helix.

Mistake: "Ionic compounds like NaCl exist as individual NaCl molecules in water."

Reality: In water, NaCl dissociates into separate Na+ and Cl− ions. The ions are surrounded by water molecules, not paired as intact NaCl units.

Mistake: "Protons determine chemical behavior."

Reality: Chemical behavior — bonding — is determined by electrons, specifically valence electrons. The number of protons defines the element, but the electrons determine how the atom interacts with other atoms.