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 chemical context of life is essential for comprehending how living organisms function at the molecular level.

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

Definition and Composition of Matter

Matter is anything that takes up space and has mass.
All 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 elements in a fixed ratio.
Compounds have emergent properties that are different from those of their constituent elements.

Example: Emergent Properties of a Compound

Sodium (Na): A soft, reactive metal.
Chlorine (Cl): A poisonous, greenish gas.
Sodium chloride (NaCl): Table salt, a safe edible compound formed from Na and Cl.

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

Atomic Structure

Atom: The smallest unit of matter that retains the properties of an element.
Atoms are composed of subatomic particles:
- Protons: Positive charge
- Neutrons: No charge
- Electrons: Negative charge

Atomic Number and Mass Number

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 in the nucleus.
Number of neutrons: Mass number – Atomic number.

Isotopes

Isotopes: Different atomic forms of the same element, with the same number of protons but different numbers of neutrons.
Some isotopes are radioactive and decay spontaneously, emitting energy.

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.
Covalent bonds: Sharing of a pair of valence electrons by two atoms.
Ionic bonds: Transfer of electrons from one atom to another, resulting in oppositely charged ions that attract each other.
Hydrogen bonds: Weak attractions between a hydrogen atom covalently bonded to one electronegative atom and another electronegative atom.
Van der Waals interactions: Weak attractions due to transient local partial charges.

Types of Covalent Bonds

Single bond: Sharing of one pair of electrons (e.g., H–H).
Double bond: Sharing of two pairs of electrons (e.g., O=O).
Polar covalent bond: Electrons are shared unequally, resulting in partial charges (e.g., in water, H2O).
Nonpolar covalent bond: Electrons are shared equally.

Ions and Ionic Compounds

Cation: Positively charged ion.
Anion: Negatively charged ion.
Salts: Compounds formed by ionic bonds (e.g., NaCl).

Concept 2.4: Chemical Reactions Make and Break Chemical Bonds

Chemical Reactions

Chemical reactions involve the making and breaking of chemical bonds.
Reactants: Starting materials in a chemical reaction.
Products: Final materials in a chemical reaction.
All chemical reactions are reversible to some extent.
Chemical equilibrium: The point at which the forward and reverse reaction rates are equal.

Concept 2.5: Hydrogen Bonding Gives Water Properties That Help Make Life Possible

Structure and Polarity of Water

Water (H2O) is a polar molecule, with oxygen having a partial negative charge and hydrogens having partial positive charges.
Hydrogen bonds form between the partially positive hydrogen of one water molecule and the partially negative oxygen of another.

Emergent Properties of Water

Cohesion: Water molecules stick together due to hydrogen bonding, aiding in the transport of water against gravity in plants.
Adhesion: Water molecules stick to other substances.
Surface tension: A measure of how difficult it is to stretch or break the surface of a liquid; water has a high surface tension.
Moderation of temperature: Water absorbs and releases heat with only slight temperature changes due to its high specific heat.
Evaporative cooling: As water evaporates, it cools the surface, helping regulate temperature in organisms and environments.
Expansion upon freezing: Ice is less dense than liquid water because hydrogen bonds keep water molecules further apart in solid form, allowing ice to float.
Versatility as a solvent: Water dissolves many substances due to its polarity, making it the solvent of life.

Acids, Bases, and pH

Acid: Substance that increases the hydrogen ion (H+) concentration of a solution.
Base: Substance that reduces the hydrogen ion concentration, often by accepting H+ or releasing OH−.
pH scale: Measures the concentration of H+ in a solution; ranges from 0 (most acidic) to 14 (most basic), with 7 being neutral.
Buffer: Substance that minimizes changes in pH by accepting or donating H+ ions.

Example: Buffer in Human Blood

Carbonic acid (H2CO3): Acts as a buffer by dissociating into bicarbonate (HCO3−) and H+, helping maintain blood pH.

Acidification: A Threat to Our Oceans

Burning fossil fuels increases atmospheric CO2, which dissolves in seawater to form carbonic acid, lowering ocean pH (ocean acidification).
Lower pH reduces carbonate ion concentration, threatening organisms that build shells or reefs from calcium carbonate.

Table: Comparison of Key Chemical Bonds

Bond Type	Definition	Relative Strength	Example
Covalent	Sharing of electron pairs between atoms	Strongest	H2O, O2
Ionic	Transfer of electrons, attraction between oppositely charged ions	Strong (in dry conditions), weaker in water	NaCl
Hydrogen	Attraction between a hydrogen atom and an electronegative atom	Weak (individually)	Between water molecules
Van der Waals	Weak attractions due to transient local charges	Weakest	Gecko feet adhesion

Key Equations

Calculating number of neutrons:
Dissociation of water:
pH calculation:

Summary

Life depends on the unique chemical and physical properties of water and other molecules.
Understanding atomic structure, chemical bonds, and the behavior of water is foundational for studying biology.