Chemistry and the Chemical Basis of Life

Introduction to Chemistry in Biology

Chemistry is fundamental to understanding biological processes because all living things are composed of chemical substances. The study of chemistry involves examining the composition, structure, properties, and reactiveness of matter. Chemical compounds serve as the building blocks of life, forming the molecules that make up cells and tissues.

• Composition: What substances are made of.

• Structure: How atoms are arranged within a substance.

• Properties: Characteristics such as melting point, density, and solubility.

• Reactivity: How substances interact and change during chemical reactions.

Atoms: The Basic Unit of Matter

Atomic Structure

An atom is the smallest unit of matter that retains the properties of an element. Atoms are composed of three main subatomic particles:

• Protons: Positively charged particles located in the nucleus; determine the atom's identity.

• Neutrons: Neutral particles (no charge) also found in the nucleus; contribute to atomic mass.

• Electrons: Negatively charged particles that orbit the nucleus; determine chemical reactivity.

Key Point: In a neutral atom, the number of protons equals the number of electrons.

Elements and the Periodic Table

Definition and Properties of Elements

An element is a pure substance consisting of only one type of atom. Each element is defined by its number of protons, known as the atomic number.

• Atomic Number (Z): The number of protons in an atom's nucleus. Determines the element's identity.

• Mass Number (A): The sum of protons and neutrons in the nucleus.

• Chemical Symbol: One- or two-letter abbreviation for the element (e.g., O for oxygen).

Example: Oxygen has 8 protons (atomic number 8) and typically 8 neutrons, giving it a mass number of 16.

Note: The atomic mass on the periodic table is often a decimal because it represents the weighted average of all naturally occurring isotopes.

Isotopes and Radioactivity

Isotopes

Isotopes are atoms of the same element that have the same number of protons but different numbers of neutrons, resulting in different mass numbers. Isotopes of an element have nearly identical chemical properties but may differ in stability.

• Example: Carbon-12 (6 protons, 6 neutrons) and Carbon-14 (6 protons, 8 neutrons).

Radioactive Isotopes (Radioisotopes)

Some isotopes are unstable and undergo radioactive decay, emitting particles and energy. These radioisotopes can be used in medical diagnosis and treatment, as well as in biological research (e.g., tracing metabolic pathways).

• Example: Phosphorus-32 decays by converting a neutron into a proton and an electron.

Electron Shells and Valence Electrons

Electron Arrangement

Electrons are arranged in shells (energy levels) around the nucleus. Each shell can hold a specific maximum number of electrons:

• First shell: 2 electrons

• Second shell: 8 electrons

• Third shell: 8 or 18 electrons (depending on the element)

• Fourth shell: 18 or 32 electrons

The valence electrons are those in the outermost shell and are responsible for chemical bonding and reactivity.

Stability and Reactivity of Atoms

Full vs. Incomplete Outer Shells

Atoms are most stable when their outermost electron shell is full. Atoms with incomplete outer shells are chemically reactive and tend to form bonds to achieve stability.

• Example: Noble gases like helium, neon, and argon have full outer shells and are chemically inert.

• Example: Potassium has one valence electron and is highly reactive.

Bohr and Lewis Structures

Representing Atoms

• Bohr Model: Shows all electrons in their respective shells around the nucleus.

• Lewis Structure: Shows only the valence electrons as dots around the chemical symbol.

Example: Beryllium (Be) with atomic number 4: 2 electrons in the first shell, 2 in the second shell.

Chemical Bonds

Types of Chemical Bonds

Atoms form chemical bonds to achieve stable electron configurations. The main types of bonds in biology are:

• Nonpolar Covalent Bond: Electrons are shared equally between atoms with similar electronegativity (e.g., H2, O2).

• Polar Covalent Bond: Electrons are shared unequally due to a difference in electronegativity, resulting in partial charges (e.g., H2O).

• Ionic Bond: Electrons are transferred from one atom to another, creating oppositely charged ions that attract each other (e.g., NaCl).

• Hydrogen Bond: A weak bond formed between a slightly positive hydrogen atom in one polar molecule and a slightly negative atom in another molecule (e.g., between water molecules, between DNA strands).

Electronegativity and Bond Formation

Electronegativity is an atom's ability to attract electrons in a chemical bond. The difference in electronegativity between atoms determines the type of bond formed:

• Nonpolar Covalent: Difference < 0.4

• Polar Covalent: Difference between 0.4 and 1.8

• Ionic: Difference > 1.8

Table: Comparison of Bond Types

Examples and Applications

• Water (H2O): Polar covalent bonds between hydrogen and oxygen create partial charges, allowing hydrogen bonding and unique properties like cohesion and high specific heat.

• Sodium Chloride (NaCl): Ionic bond forms when sodium donates an electron to chlorine, resulting in Na+ and Cl- ions.

• DNA: Hydrogen bonds hold the two strands of the DNA double helix together.

Summary

• Atoms are the fundamental units of matter, composed of protons, neutrons, and electrons.

• Elements are defined by their atomic number (number of protons).

• Isotopes have the same number of protons but different numbers of neutrons.

• Electron arrangement, especially valence electrons, determines chemical reactivity.

• Chemical bonds (covalent, ionic, hydrogen) are essential for forming molecules and compounds in living organisms.