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 biological processes. Understanding the nature of matter, elements, and compounds is essential for studying how living organisms function at the molecular level.

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

What is Matter?

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

Elements and Compounds

Element: A substance that cannot be broken down to other substances by chemical reactions.
Compound: A substance consisting of two or more different elements combined in a fixed ratio.
Compounds have emergent properties that are different from those of their constituent elements.
Example: Sodium (Na) is a soft metal, and chlorine (Cl) is a poisonous gas, but together they form sodium chloride (NaCl), or table salt, which is edible and essential for life.

Elements of Life

About 20–25% of the 92 natural elements are required for life; these are called essential elements.
Carbon (C), hydrogen (H), oxygen (O), and nitrogen (N) make up about 96% of living matter.
The remaining 4% consists mainly of calcium (Ca), phosphorus (P), potassium (K), and sulfur (S).
Trace elements are required by organisms in very small amounts (e.g., iron, iodine, zinc).

Table: Major Elements in the Human Body

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%
Trace elements	-	<0.01%

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

Atoms and Subatomic Particles

Atom: The smallest unit of matter that retains the properties of an element.
Atoms are composed of subatomic particles:
- Protons: Positive charge (+1)
- Neutrons: No charge (0)
- Electrons: Negative charge (–1)
Protons and neutrons are located in the atomic nucleus; electrons move around the nucleus in electron shells.

Atomic Number and Atomic Mass

Atomic number: Number of protons in the nucleus of an atom; defines the element.
Mass number: Sum of protons and neutrons in the nucleus.
Atomic mass: The atom’s total mass, approximately equal to the mass number (measured in daltons).
Formula:

Isotopes

Isotopes: Atoms of the same element with the same number of protons but different numbers of neutrons.
Radioactive isotopes: Unstable isotopes that decay spontaneously, releasing particles and energy.
Applications: Radioactive isotopes are used as tracers in medical diagnostics and research.

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

Chemical Bonds

Atoms with incomplete valence shells can share or transfer valence electrons, resulting in chemical bonds.
Covalent bonds: Involve the sharing of pairs of valence electrons between atoms.
Ionic bonds: Involve the transfer of electrons from one atom to another, resulting in oppositely charged ions that attract each other.

Covalent Bonds

Single covalent bond: Sharing of one pair of electrons (e.g., H–H).
Double covalent bond: Sharing of two pairs of electrons (e.g., O=O).
Molecule: Two or more atoms held together by covalent bonds.
Valence: The bonding capacity of an atom, usually equal to the number of unpaired electrons in its valence shell.

Electronegativity and Bond Polarity

Electronegativity: The attraction of an atom for the electrons in a covalent bond.
Nonpolar covalent bond: Electrons are shared equally between atoms.
Polar covalent bond: Electrons are shared unequally, resulting in partial charges on atoms (e.g., H2O).

Ionic Bonds and Ions

Ion: An atom or molecule with a net electric charge due to the loss or gain of electrons.
Cation: Positively charged ion (loss of electrons).
Anion: Negatively charged ion (gain of electrons).
Ionic compound: A compound formed by ionic bonds (e.g., NaCl).
Ionic compounds are often found as crystals and dissociate easily in water.

Weak Chemical Interactions

Weak bonds, such as hydrogen bonds and van der Waals interactions, play important roles in the structure and function of biological molecules.
Hydrogen bond: Forms when a hydrogen atom covalently bonded to one electronegative atom is attracted to another electronegative atom (commonly oxygen or nitrogen).
Van der Waals interactions: Weak attractions between molecules or parts of molecules that result from transient local partial charges.

Concept 2.4: Chemical Reactions Make and Break Chemical Bonds

Chemical Reactions

Chemical reaction: The making and breaking of chemical bonds, leading to changes in the composition of matter.
Reactants: Starting materials in a chemical reaction.
Products: Resulting materials from a chemical reaction.
Example: Formation of water:
Photosynthesis: An important chemical reaction in which light energy is used to convert carbon dioxide and water into glucose and oxygen:

Summary Table: Types of Chemical Bonds

Bond Type	Description	Relative Strength	Example
Covalent	Sharing of electron pairs between atoms	Strong	H2O, O2
Ionic	Transfer of electrons; attraction between oppositely charged ions	Strong (in dry conditions)	NaCl
Hydrogen	Attraction between a hydrogen atom and an electronegative atom	Weak	Between water molecules
Van der Waals	Weak attractions due to transient local charges	Very weak	Gecko's toe hairs on surfaces

Key Learning Objectives

Identify the basic unit of matter and the subatomic particles that make them up.
Identify the essential elements in the human body.
Distinguish between and discuss the biological importance of polar covalent bonds and hydrogen bonds.
Understand how to identify valence electrons.
Determine the difference between covalent and ionic bonds.