Biology 2e Chapter 2: The Chemical Foundation of Life

Atoms as Building Blocks of Molecules and Matter

Fundamental Nature of Matter: - Life is entirely composed of matter. - Matter is defined as anything that occupies space and has mass. - Elements represent unique forms of matter characterized by: 1. Specific chemical properties. 2. Specific physical properties.
Elements and the Living World: - Elements are substances that cannot be broken down into smaller substances through ordinary chemical means. - Chemical Symbols: Each element is designated by a unique one or two-letter chemical symbol. - Examples: Sulfur is represented as $S$ ; Calcium is represented as $Ca$ . - The Four Essential Elements: Four elements comprise the vast majority of living organisms: - Carbon ( $C$ ) - Oxygen ( $O$ ) - Hydrogen ( $H$ ) - Nitrogen ( $N$ )

Atomic Structure and Subatomic Particles

The Atom: The smallest unit of matter that retains all of the chemical properties of an element.
Atomic Regions: - Nucleus: Located at the center of the atom; it contains protons and neutrons. - Outermost Region: The area that holds electrons in orbit around the nucleus.
Subatomic Particles: Protons, neutrons, and electrons are the collective constituents of atoms.
Key Properties Summary: - Proton: - Charge: $+1$ - Mass: $1\,amu$ - Location: Nucleus - Neutron: - Charge: $0$ - Mass: $1\,amu$ - Location: Nucleus - Electron: - Charge: $-1$ - Mass: $0\,amu$ - Location: Orbitals within the outermost region
Visual Depiction: Helium ( $He$ ) serves as a classic example of how these subatomic particles are arranged in a simplified depiction.

Atomic Number, Atomic Mass, and Isotopes

Atomic Standards: - Atoms of a specific element have a standard number of protons and electrons. - Atomic Number: Defined as the number of protons in an atom. Each element has a unique, distinct atomic number. - Atomic Mass: The total mass of an atom, roughly equal to the sum of protons and neutrons. - Units: Expressed in atomic mass units ( $amu$ ). - Calculation of Neutrons: $\text{Atomic Mass} - \text{Atomic Number} = \text{Number of Neutrons}$ . - Note: Electrons are not included in atomic mass calculations due to their negligible mass.
Isotopes: - Forms of an element that contain different numbers of neutrons and therefore possess different mass numbers. - Carbon Example: Carbon has an atomic number of $6$ . It has two stable isotopes with mass numbers of $12$ and $13$ . Carbon-12 has an atomic mass of approximately $12.11$ . - Hydrogen Example: - $^{1}H$ : $0$ neutrons. - $^{2}H$ : $1$ neutron. - $^{3}H$ : $2$ neutrons.
The Periodic Table: Displays the atomic mass and atomic number for every element. Typically, the atomic number is placed above the chemical symbol, and the approximate atomic mass is placed below it.

Electron Shells and the Bohr Model

Neutral Atoms: In a neutrally charged atom, the number of protons equals the number of electrons, which is equal to the atomic number.
The Bohr Model: An early conceptual model showing the nucleus at the center with electrons in circular orbits at specific distances (energy levels).
Electron Distribution: - Orbits are referred to as electron shells or energy levels. - Electrons occupy the lowest available energy shell (closest to the nucleus) first. - Shell filling sequence: Shell $1n$ fills first, followed by $2n$ , then $3n$ , etc.
Valence Shells and Stability: - Valence Shell: The outermost electron shell. - Stability is achieved when the valence shell is completely filled. - Group 18 Elements: These elements naturally possess full valence shells and are highly stable. - Octet Rule: For the first two outer shells, a full shell consists of $8$ electrons. - Group 1 Elements (e.g., $H$ , $Li$ , $Na$ ): Can achieve stability by losing their single outer electron. - Group 17 Elements: Can achieve stability by gaining one additional electron to complete the octet.

Chemical Reactions and Bonding

Defining Chemical Reactions: These represent changes in the distribution of electrons between participating atoms. - Reactants: Substances present at the beginning of the reaction. - Products: Substances formed at the conclusion of the reaction. - Example Reaction: $2H_2O_2 \rightarrow 2H_2O + O_2$ .
Reaction Types: - Irreversible: Proceeds in one direction until reactants are exhausted ( $\rightarrow$ ). - Reversible: Reactants convert to products, but products can also convert back into reactants ( $\rightleftharpoons$ ). - Example: $HCO_3^{-} + H^{+} \rightleftharpoons H_2CO_3$ .
Chemical Bonds: The attractive forces that link atoms together to form molecules.

Types of Chemical Bonds

Covalent Bonds: - Formed when two or more atoms share electrons. - Bond Strength: Strong. - Single Bonds: Two hydrogens and one oxygen share electrons to form a water molecule ( $H_2O$ ). - Double Bonds: Two sets of electrons are shared, such as the bond between oxygen atoms in an $O_{2}$ molecule.
Ionic Bonds: - Formed when atoms either give up or gain electrons to achieve an octet, resulting in an attraction between oppositely charged ions. - Bond Strength: Strong. - Metals vs. Nonmetals: Metals typically lose electrons while nonmetals gain them. - Example: Sodium Chloride ( $NaCl$ ): - Sodium ( $Na$ ): Highly reactive, explodes in contact with water. - Chlorine ( $Cl$ ): A poisonous, deadly gas. - Sodium Chloride ( $NaCl$ ): A stable compound used as a food preservative and flavoring agent.
Polar vs. Non-polar Covalent Bonds: - Polar Covalent Bonds: Electrons are shared unequally and are attracted more to one nucleus than the other due to differences in electronegativity. - Example: Water ( $H_2O$ ). Oxygen has higher electronegativity than hydrogen. - Non-Polar Covalent Bonds: Electrons are shared equally between atoms. - Example: Methane ( $CH_4$ ). Carbon shares electrons with four hydrogens equally. - Example: Carbon Dioxide ( $CO_2$ ) is nonpolar because its molecular shape offsets the polar covalent bonds. - Determinants of Polarity: Both bond type and molecular shape determine if a molecule is polar or nonpolar.
Hydrogen Bonds: - An attraction between the slightly positive ( $\delta+$ ) hydrogen atom of one molecule and the slightly negative ( $\delta-$ ) atom (usually oxygen) of another molecule. - Bond Strength: Weak.

The Unique Properties of Water

Criticality for Life: - Water makes up $60$ to $70\%$ of the human body. - Its essential nature is derived from its polarity and ability to form hydrogen bonds.
States of Water: - Liquid: Hydrogen bonds are constantly being made, broken, and remade. - Gas (Steam): Heating increases kinetic energy, causing hydrogen bonds to break completely and molecules to escape into the air. - Solid (Ice): Lowered temperatures allow for the maintenance of a crystalline lattice structure.
Density of Ice: Hydrogen bonding makes ice less dense than liquid water because the lattice structure keeps molecules further apart, allowing ice to float.
Heat Properties: - Specific Heat Capacity: The amount of heat $1\,g$ of a substance must absorb to raise its temperature by $1^{\circ}C$ . Water has a high specific heat, meaning it takes a long time to heat up or cool down. - Heat of Vaporization: The energy required to change $1\,g$ of a liquid to a gas. Water has a high heat of vaporization.
Solvent Properties: - Water is a versatile solvent; it can dissolve ions and other polar molecules. - Solute: The substances dissolved or mixed into the water. - Dissociation and Spheres of Hydration: When $NaCl$ is mixed in water, the atoms break off to form ions ( $Na^+$ and $Cl^-$ ), and water molecules form spheres of hydration around them.

Cohesion, Adhesion, and Surface Tension

Cohesion: Water molecules at the liquid-gas interface stick together due to hydrogen bonding. - Surface Tension: The capacity of a substance to withstand being ruptured under tension/stress (e.g., a needle floating on water or a water strider staying afloat).
Adhesion: Adaptation through the attraction between water molecules and other molecules. - Capillary Action: Observed in glass tubes where adhesive forces between water and the glass exceed the cohesive forces between water molecules, causing the water to move up the tube.

pH, Buffers, Acids, and Bases

Ionization of Water: A small percentage of water molecules dissociate into hydrogen ions ( $H^+$ ) and hydroxide ions ( $OH^-$ ). - Equation: $H_2O(l) \rightleftharpoons H^+(aq) + OH^-(aq)$ .
Concentrations in Pure Water: - $[H^+] = 1 \times 10^{-7}\,mol/dm^3$ . - $[OH^-] = 1 \times 10^{-7}\,mol/dm^3$ .
Defining pH: - pH is the negative of the base 10 logarithm of the hydrogen ion concentration: $\text{pH} = -\log_{10}([H^+])$ . - Neutral: pH of $7$ . - Acidic: High $H^+$ concentration; $pH < 7$ . - Alkaline (Basic): High $OH^-$ concentration; $pH > 7$ .

Carbon: The Basis of Life

Macromolecules: Carbon is the key component of proteins, carbohydrates, lipids, and nucleic acids.
Covalent Versatility: Carbon can form covalent bonds with up to four different atoms, serving as a "backbone." - Carbon has $4$ electrons in its outer shell. - It forms $4$ covalent bonds to satisfy the octet rule.
Hydrocarbons: Molecules consisting only of carbon and hydrogen (e.g., methane). - Covalent bonds in hydrocarbons store significant energy, which is released when burned (powering cars and heating homes).
Methane Geometry: Methane ( $CH_4$ ) adopts a tetrahedral geometry where the four hydrogen atoms are spaced exactly $109.5^{\circ}$ apart.

Additional Information and Credits

Dialogue & Credits: - Source material credit to Rao, A., Fletcher, S., Ryan, K., Tag, A. and Hawkins, A. from the Department of Biology, Texas A&M University. - Visual reference: Amoeba Sisters "Paramecium Parlor" regarding water's electronegativity. - Visual Molecular Dynamics (VMD) software used for ice lattice images (Jane Whitney and Carlos Ponte).