The Chemical Basis of Life: Elements, Atoms, and Water Chemistry

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Elements, Atoms, and Compounds

2.1 Organisms Are Composed of Elements, Usually Combined into Compounds

All living organisms are made of matter, which consists of chemical elements. These elements combine in specific ways to form compounds essential for life.

Matter is anything that occupies space and has mass.
About 25 elements are essential for human life, but four elements—oxygen (O), carbon (C), hydrogen (H), and nitrogen (N)—make up about 96% of the mass of most living organisms.
A compound is a substance consisting of two or more different elements combined in a fixed ratio (e.g., NaCl for table salt).
Emergent properties are new characteristics that arise when elements combine to form compounds (e.g., table salt is edible, though its components sodium and chlorine are not safe alone).

Example: Table salt (NaCl) illustrates emergent properties because its characteristics differ from those of its constituent elements.

Atoms and Subatomic Particles

2.3 Atoms Consist of Protons, Neutrons, and Electrons

Atoms are the smallest units of matter that retain the properties of an element. They are composed of three types of subatomic particles.

Protons (positive charge) and neutrons (no charge) are located in the nucleus.
Electrons (negative charge) orbit the nucleus in electron shells.
The atomic number is the number of protons in an atom and defines the element.
The mass number is the sum of protons and neutrons in the nucleus.
The atomic mass (measured in daltons or atomic mass units) is approximately equal to the mass number.
Isotopes are atoms of the same element with different numbers of neutrons.

Example: A nitrogen atom with 7 protons and 9 neutrons (radioactive isotope) has an atomic number of 7 and a mass number of 16.

Table: Isotopes of Carbon

Isotope	Protons	Neutrons	Mass Number
Carbon-12	6	6	12
Carbon-13	6	7	13
Carbon-14	6	8	14

Chemical Bonds

2.5 The Distribution of Electrons Determines an Atom’s Chemical Properties

The arrangement of electrons in an atom’s electron shells determines how it interacts with other atoms.

Electrons occupy electron shells at varying distances from the nucleus.
An atom with an incomplete outer shell tends to share, gain, or lose electrons, forming chemical bonds.
Ionic bonds result from the transfer of electrons between atoms.
Covalent bonds involve the sharing of electrons between atoms.

Example: Hydrogen forms 1 covalent bond, oxygen forms 2, nitrogen forms 3, and carbon forms 4.

2.6 Covalent Bonds Join Atoms into Molecules Through Electron Sharing

Nonpolar covalent bonds: Electrons are shared equally (e.g., H2).
Polar covalent bonds: Electrons are shared unequally, resulting in partial charges (e.g., H2O).
Polarity refers to the separation of electrical charge within a molecule.

Example: In water, oxygen is more electronegative than hydrogen, so electrons are pulled closer to oxygen, making the molecule polar.

2.7 Ionic Bonds are Attractions Between Ions of Opposite Charge

An ion is an atom or molecule with an electrical charge due to the loss or gain of electrons.
Cations are positively charged ions; anions are negatively charged ions.
Ionic bonds form when oppositely charged ions attract each other (e.g., Na+ and Cl– in table salt).

Example: The crystal structure of table salt is held together by the attraction between Na+ and Cl– ions.

2.8 Hydrogen Bonds are Weak Bonds Important in the Chemistry of Life

Hydrogen bonds are weak attractions between a hydrogen atom (already covalently bonded to a more electronegative atom) and another electronegative atom.
In water, hydrogen bonds form between the slightly positive hydrogen of one molecule and the slightly negative oxygen of another.
Hydrogen bonds are crucial for the structure and function of biological molecules (e.g., DNA, proteins).

Example: Neighboring water molecules are held together by hydrogen bonds.

2.9 Chemical Reactions Make and Break Chemical Bonds

Chemical reactions involve the breaking and forming of chemical bonds, converting reactants into products.
Matter is neither created nor destroyed in chemical reactions; it is rearranged.

Example: The reaction of hydrogen and oxygen to form water:

Water’s Life-Supporting Properties

2.10 Hydrogen Bonds Make Liquid Water Cohesive

Cohesion: The tendency of molecules of the same kind to stick together (important for water transport in plants).
Adhesion: The clinging of one substance to another (e.g., water to plant cell walls).
Surface tension: A measure of how difficult it is to break the surface of a liquid, caused by cohesion.

Example: Sweat forms droplets on skin due to water’s high surface tension.

2.11 Water’s Hydrogen Bonds Moderate Temperature

Thermal energy: Energy from the random movement of atoms and molecules.
Heat: Transfer of thermal energy from a warmer to a cooler body.
Temperature: Measures the intensity of heat.
Evaporative cooling: As water evaporates, the surface cools because the hottest molecules leave first.

Example: High humidity reduces the effectiveness of evaporative cooling, making hot days feel hotter.

2.12 Ice Floats Because It Is Less Dense than Liquid Water

Water is unique in that its solid form (ice) is less dense than its liquid form.
Hydrogen bonds stabilize and keep water molecules farther apart in ice than in liquid water.
This property allows ice to float, insulating bodies of water and protecting aquatic life in winter.

Example: Freezing water expands and can crack rocks and boulders.

2.13 Water is the Solvent of Life

A solution is a homogeneous mixture of two or more substances.
Solvent: The dissolving agent (water in biological systems).
Solute: The substance dissolved.
Water’s polarity allows it to dissolve many substances, forming aqueous solutions.

Example: Blood is an aqueous solution because water dissolves salts, proteins, and other molecules.

Acids, Bases, and pH

2.14 The Chemistry of Life Is Sensitive to Acidic and Basic Conditions

In water, a small fraction of molecules dissociate into H+ (hydrogen ions) and OH– (hydroxide ions).
The pH scale measures how acidic or basic a solution is, ranging from 0 (most acidic) to 14 (most basic).
A buffer is a substance that minimizes changes in pH by accepting or donating H+ ions.

Example: A solution at pH 4 is 100,000 times more acidic than one at pH 9.

Acids and Bases

An acid increases the H+ concentration of a solution.
A base reduces the H+ concentration, often by accepting H+ or releasing OH–.

The pH Scale

At 25°C, the product of H+ and OH– concentrations in water is constant:

pH is defined as the negative logarithm of the H+ concentration:

For a neutral solution: , so pH = 7.

Buffers

Most living cells maintain an internal pH close to 7.
Buffers are acid-base pairs that reversibly bind H+ to resist pH changes.
Major biological buffer systems include:
- Carbonic Acid/Bicarbonate Buffering System
- Phosphate Buffering System
- Protein Buffering System

Example: The carbonic acid/bicarbonate system helps maintain blood pH.

Additional info: The carbonic acid/bicarbonate buffer system is represented by the equilibrium: