Life and water

Chapter 2: Life, Chemistry, and Water

Why It Matters

• Living organisms are assemblies of atoms and molecules linked through chemical bonds.

• Chemistry governs the behavior of both living and nonliving entities.

• Grasping the connection between the structure of substances and their properties is foundational in biology.

2.1 The Organization of Matter

• An element is defined as a pure substance that cannot be decomposed into simpler substances by ordinary chemical or physical processes.

• Matter is composed of various elements and their combinations.

• There are 92 naturally occurring elements on Earth, with additional artificial elements created in laboratories.

25 Key Elements

• Four primary elements (carbon, hydrogen, oxygen, nitrogen) account for over 96% of the mass of living organisms.

• Seven additional elements (calcium, phosphorus, potassium, sulfur, sodium, chlorine, magnesium) make up nearly 4%.

• Some trace elements (e.g., iodine) are essential in minute amounts (below 0.01%).

Proportions of Different Elements

Atoms and Molecules

• Elements consist of atoms, the smallest unit retaining the properties of that element.

• Atoms combine in definite ratios to form molecules identified by chemical formulas (e.g., CO2).

Compounds

• Molecules formed by the combination of different atoms are referred to as compounds.

• The properties of compounds differ from those of their constituent elements (e.g., water is formed from flammable gases hydrogen and oxygen).

Common Elements in Living Organisms

Study Break 2.1

• Distinguish between:

  1. An Element and an Atom

  2. A Molecule and a Compound

2.2 Atomic Structure

• Each element comprises one unique type of atom, comprising an atomic nucleus surrounded by electrons.

• The nucleus consists of positively charged protons and generally also has neutrons (except hydrogen).

Atomic Structure - Examples

A. Hydrogen

• 1 proton and 1 electron.

B. Carbon

• 6 protons, 6 neutrons, 2 inner electrons, 4 valence electrons.

Isotopes

• Isotopes are variants of an element that differ in neutron count but retain the same number of protons.

• They do not alter chemical reactions of the element in biological contexts.

Isotopes of Hydrogen and Carbon

• Hydrogen Isotopes:

  • 2H (Deuterium): 1 proton, 1 neutron; mass number = 2.

  • 3H (Tritium): 1 proton, 2 neutrons; mass number = 3.

  • 1H: 1 proton; mass number = 1.

• Carbon Isotopes:

  • 12C: 6 protons, 6 neutrons; mass number = 12.

  • 13C: 6 protons, 7 neutrons; mass number = 13.

  • 14C: 6 protons, 8 neutrons; mass number = 14.

Isotopes in Research

• Radioisotope decay helps estimate ages of organic materials, rocks, or fossils.

• Isotopes function as tracers in biological studies (e.g., 14C, 32P).

Focus On Applied Research

• Iodine-123 is used to scan thyroid glands, concentrating in the gland.

Polarity

• Nonpolar Covalent Bond: Electrons are shared equally.

• Polar Covalent Bond: Electrons are unevenly shared; one atom carries a partial negative charge (δ-) while another has a partial positive charge (δ+).

Water, A Polar Molecule

• In water (H2O), the oxygen atom forms polar covalent bonds with two hydrogen atoms, making the molecule asymmetric and strongly polar.

Symmetry

• Polar Groups: Hydroxyl (—OH), amino (—NH), and sulfhydryl (—SH) groups are typically asymmetrically located in biological molecules.

• Nonpolar Groups: Carbon-hydrogen bonds are symmetrically arranged, making them nonpolar.

Polar and Nonpolar Associations

• Polar molecules attract other polar molecules, separating nonpolar substances into clumps.

• Nonpolar associations reduce surface area exposure to the polar environment.

Hydrophilic and Hydrophobic

• Hydrophilic: Molecules that readily associate with water.

• Hydrophobic: Molecules that are excluded by water.

Hydrogen Bonds

• Formed between partially positive hydrogen atoms and partially negative atoms (e.g., in H—O, N—H bonds).

• Can be intramolecular or intermolecular and are individually weak but collectively strong.

Stabilizing Effect of Hydrogen Bonds

• Hydrogen bonds maintain structure in large biological molecules, such as proteins.

• Begin to break at temperatures exceeding 45°C.

Van der Waals Forces

• Weak interactions occur between nonpolar molecules as electrons create temporary charge zones, influencing molecular shapes.

Chemical Reactions

• Form or break chemical bonds between atoms or molecules.

• Enzymes facilitate reactions in biological systems.

Chemical Equations

• Represent reactions with reactants on the left and products on the right.

• Balanced equations maintain equal numbers of each type of atom throughout the reaction (e.g., Photosynthesis: 6 CO2 + 6 H2O → C6H12O6 + 6 O2).

Study Break 2.2

• Question Topics:

  1. Formation of Ionic Bonds

  2. Formation of Covalent Bonds

  3. Understanding Electronegativity and its relation to bond types

  4. Definition of Chemical Reaction

2.4 Hydrogen Bonds and the Properties of Water

• Hydrogen bonds create a lattice in water that defines its density, heat absorption, cohesion, and surface tension properties.

Water and Temperature

• Liquid water resists individual molecule escape due to its hydrogen-bond lattice, remaining liquid between 0°C and 100°C.

• High specific heat and heat of vaporization enable temperature regulation.

Cohesion and Adhesion

• Cohesion: Water molecules' tendency to stick to each other due to hydrogen bonding.

• Adhesion: Water molecules' ability to adhere to other surfaces via hydrogen bonds with charged/polar groups.

Surface Tension

• Water's surface tension allows it to resist separation due to uneven bonding at the air-water interface.

Effects of Surface Tension

• Creates droplets and supports small insects or objects on water.

Polar and Nonpolar Environments

• Water lattice creates distinct environments critical to cell organization, excluding nonpolar molecules.

Biological Membranes

• Formed from lipids possessing dual polarity; polar heads interact with water while nonpolar tails form a bilayer barrier.

Water as a Solvent

• Water coats and reduces attraction between polar and charged molecules, forming hydration layers that prevent reassociation.

Acids and Bases

• Acids: Proton donors increasing H+ concentration in solutions (e.g., HCl → H+ + Cl–).

• Bases: Proton acceptors that reduce H+ concentration (e.g., NaOH → Na+ + OH–).

pH Scale

• Measures acidity or alkalinity of solutions; ranges from 0 (strong acids) to 14 (strong bases) with 7 being neutral.Examples:

• 0: Hydrochloric acid,

• 7: Pure water,

• 14: Sodium hydroxide.