Overview: A Chemical Connection to Biology

Biology and Chemistry

Living organisms are governed by the fundamental laws of physics and chemistry. Biology is a multidisciplinary science, integrating concepts from chemistry to explain life processes. Life is organized hierarchically, from atoms to molecules, and then to cells, with emergent properties appearing at each level.

• Emergent properties: New characteristics arise when simple components combine into more complex structures.

• Transition from nonlife to life: Occurs as molecules organize into cells.

Concept 2.1: Matter, Elements, and Compounds

Matter and Its Forms

Matter is anything that occupies space and has mass. It is composed of elements and compounds.

• Element: A pure substance that cannot be broken down by chemical reactions.

• Compound: A substance made of two or more elements in a fixed ratio, with properties different from its constituent elements.

Emergent Properties of Compounds

When elements combine, the resulting compound often has new properties.

• Example: Sodium (a metal) and chlorine (a gas) combine to form sodium chloride (NaCl), an edible salt.

The Elements of Life

Only a subset of the 92 naturally occurring elements are essential for life.

• Major elements: Carbon, hydrogen, oxygen, and nitrogen make up 96% of living matter.

• Other essential elements: Calcium, phosphorus, potassium, and sulfur.

• Trace elements: Required in minute quantities (e.g., iron, iodine).

Concept 2.2: Atomic Structure and Properties

Atoms and Subatomic Particles

Atoms are the smallest units of matter retaining the properties of an element.

• Subatomic particles: Protons (positive), neutrons (neutral), electrons (negative).

• Atomic nucleus: Dense core of protons and neutrons.

• Electron cloud: Electrons move around the nucleus.

Atomic Number and Atomic Mass

• Atomic number: Number of protons in the nucleus; also equals the number of electrons in a neutral atom.

• Mass number: Sum of protons and neutrons.

• Atomic mass: Approximate total mass of an atom.

Isotopes

Isotopes are atoms of the same element with different numbers of neutrons.

• Stable isotopes: Do not decay.

• Radioactive isotopes: Decay spontaneously, emitting particles and energy.

• Example: Carbon-12, Carbon-13 (stable); Carbon-14 (radioactive).

Energy Levels of Electrons

Electrons have potential energy based on their distance from the nucleus.

• Electron shells: Discrete energy levels where electrons reside.

• Energy transitions: Electrons absorb or release energy to move between shells.

Electron Distribution and the Periodic Table

The arrangement of electrons in shells determines an atom’s chemical properties.

• Periodic table: Shows electron distribution for each element.

• Valence electrons: Electrons in the outermost shell; determine chemical reactivity.

Electron Orbitals

Orbitals are three-dimensional regions where electrons are likely to be found.

• Each shell: Contains a specific number of orbitals.

• Maximum per orbital: 2 electrons.

Concept 2.3: Chemical Bonds and Molecular Formation

Covalent Bonds

Covalent bonds involve the sharing of valence electrons between atoms.

• Single bond: One pair of electrons shared.

• Double bond: Two pairs of electrons shared.

• Bonding capacity (valence): Number of covalent bonds an atom can form.

Electronegativity and Types of Covalent Bonds

Electronegativity is an atom’s ability to attract electrons in a bond.

• Nonpolar covalent bond: Electrons shared equally.

• Polar covalent bond: Electrons shared unequally, creating partial charges.

Ionic Bonds and Ionic Compounds

Ionic bonds form when electrons are transferred from one atom to another, creating ions.

• Cation: Positively charged ion.

• Anion: Negatively charged ion.

• Ionic compound: Formed by the attraction between cations and anions (e.g., NaCl).

Weak Chemical Interactions

• Hydrogen bonds: Weak attractions between a hydrogen atom covalently bonded to an electronegative atom and another electronegative atom.

• Importance: Stabilize DNA, protein structure, and water properties.

Van der Waals Interactions

• Temporary partial charges: Created by uneven electron distribution.

• Collective strength: Many weak interactions together can be strong (e.g., gecko adhesion).

Molecular Shape and Function

The shape of a molecule is determined by the arrangement of its atoms’ orbitals and is crucial for its biological function.

• Hybridization: s and p orbitals combine to form specific shapes.

• Recognition: Molecules with similar shapes can bind the same biological receptors.

Summary Table: Types of Chemical Bonds

Key Terms and Concepts

• Element: Pure substance, cannot be broken down chemically.

• Compound: Substance of two or more elements in fixed ratio.

• Atom: Smallest unit of an element.

• Isotope: Atom with same protons, different neutrons.

• Covalent bond: Shared electrons.

• Ionic bond: Transferred electrons, attraction of ions.

• Hydrogen bond: Weak attraction involving hydrogen and electronegative atoms.

• Van der Waals interaction: Weak, temporary attractions.

• Valence electrons: Outermost shell electrons, determine reactivity.

• Electronegativity: Atom’s pull on shared electrons.

Example Exam Questions

• How many protons, neutrons, and electrons are in an isotope of titanium with mass number 48?

• What numbers must be placed as coefficients in the blanks for the chemical reaction below in order to ensure that matter is conserved? Fe3O4 + __C → __Fe + __CO

• Which bond type is most similar to a person carefully walking on a slippery sidewalk?

Additional info: This guide expands on brief points with academic context, definitions, and examples for clarity and completeness.