Chapter 2: The Chemical Context of Life – Study Notes

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Chapter 2: The Chemical Context of Life

Introduction

This chapter explores the fundamental chemical principles that underlie all biological processes. Understanding the nature of matter, atomic structure, and chemical bonding is essential for studying life at the molecular level.

Concept 2.1: Matter Consists of Chemical Elements in Pure Form and in Combinations Called Compounds

Definition of Matter

Matter is anything that takes up space and has mass.
All living organisms are composed of matter.

Elements and Compounds

An element is a substance that cannot be broken down to other substances by chemical reactions.
A compound is a substance consisting of two or more elements combined in a fixed ratio.
Compounds have characteristics different from those of their constituent elements.
Example: Sodium (Na) and chlorine (Cl) are elements; when combined, they form sodium chloride (NaCl), a compound with different properties.

Elements in the Human Body

The human body is primarily composed of a few key elements, with others present in trace amounts.

Element	Symbol	Percentage of Body Mass (including water)
Oxygen	O	65.0%
Carbon	C	18.5%
Hydrogen	H	9.5%
Nitrogen	N	3.3%
Calcium	Ca	1.5%
Phosphorus	P	1.0%
Potassium	K	0.4%
Sulfur	S	0.3%
Sodium	Na	0.2%
Chlorine	Cl	0.2%
Magnesium	Mg	0.1%

Additional info: Trace elements such as iron (Fe), zinc (Zn), and iodine (I) are essential in very small amounts for proper biological function.

Concept 2.2: An Element’s Properties Depend on the Structure of Its Atoms

Atomic Structure

An atom is the smallest unit of matter that retains the properties of an element.
Atoms are composed of subatomic particles:
- Protons (positive charge, located in the nucleus)
- Neutrons (no charge, located in the nucleus)
- Electrons (negative charge, orbiting the nucleus in electron shells)
The nucleus contains protons and neutrons; electrons form a "cloud" around the nucleus.
Proton and neutron masses are nearly identical and are measured in daltons.

Atomic Number and Atomic Mass

The atomic number is the number of protons in an atom’s nucleus and defines the element.
The mass number is the sum of protons and neutrons in the nucleus.
Atomic mass is the atom’s total mass, approximately equal to the mass number.
Example: Oxygen has 8 protons (atomic number 8) and typically 8 neutrons (mass number 16).

Isotopes

Isotopes are atoms of the same element with different numbers of neutrons.
Radioactive isotopes decay spontaneously, releasing particles and energy.
Example: Carbon-12 and Carbon-14 are isotopes of carbon; Carbon-14 is radioactive.

Electron Distribution and Chemical Properties

The chemical behavior of an atom is determined by the distribution of electrons in electron shells.
The valence shell is the outermost electron shell; electrons in this shell are called valence electrons.
Elements with a full valence shell are chemically inert (unreactive).
The periodic table arranges elements by atomic number and electron configuration.

Concept 2.3: The Formation and Function of Molecules Depend on Chemical Bonding Between Atoms

Chemical Bonds

Atoms with incomplete valence shells can share or transfer valence electrons, resulting in chemical bonds.
Chemical bonds are attractions that hold atoms together in molecules or compounds.

Covalent Bonds

A covalent bond is the sharing of a pair of valence electrons by two atoms.
A single covalent bond involves one pair of shared electrons; a double covalent bond involves two pairs.
Molecules are formed when two or more atoms are held together by covalent bonds.
Structural formulas represent atoms and bonding (e.g., H—H for hydrogen gas).
Molecular formulas show the number and type of atoms (e.g., H2).

Electronegativity and Bond Polarity

Electronegativity is an atom’s attraction for electrons in a covalent bond.
In a nonpolar covalent bond, electrons are shared equally between atoms.
In a polar covalent bond, one atom is more electronegative and pulls electrons closer, resulting in partial charges.
Example: In water (H2O), oxygen is more electronegative than hydrogen, making the bonds polar.

Ionic Bonds

Sometimes, atoms strip electrons from their bonding partners, forming ions (charged atoms or molecules).
A cation is a positively charged ion; an anion is a negatively charged ion.
An ionic bond is the attraction between a cation and an anion.
Example: Sodium (Na) donates an electron to chlorine (Cl), forming Na+ and Cl-, which combine to form NaCl (table salt).
Ionic compounds (salts) often form crystalline structures.

Hydrogen Bonds

A hydrogen bond forms when a hydrogen atom covalently bonded to one electronegative atom is attracted to another electronegative atom.
Common in water (H2O) and important in the structure of DNA and proteins.
Example: The attraction between the hydrogen of one water molecule and the oxygen of another.

Summary Table: Types of Chemical Bonds

Bond Type	Description	Example
Covalent	Sharing of electron pairs between atoms	H2, O2, H2O
Ionic	Transfer of electrons from one atom to another, forming ions	NaCl
Hydrogen	Attraction between a hydrogen atom and an electronegative atom	Between water molecules

Additional info: Understanding these chemical principles is foundational for studying biological molecules and processes, such as metabolism, genetics, and cellular structure.