BIOLOGY 2e Chapter 2: THE CHEMICAL FOUNDATION OF LIFE

The Composition and Nature of Matter in Biological Systems

Matter and Elemental Basics
- Life is composed of matter, which is defined as anything that occupies space and has mass.
- Elements represent unique forms of matter that possess specific chemical and physical properties.
- Elements cannot be broken down into smaller substances through ordinary chemical means.
Essential Elements for Living Organisms
- Each element is designated by a chemical symbol consisting of one or two letters. Examples include Sulfur ( $S$ ) and Calcium ( $Ca$ ).
- The four most common elements found in living organisms are:
  - Carbon ( $C$ )
  - Oxygen ( $O$ )
  - Hydrogen ( $H$ )
  - Nitrogen ( $N$ )
Comparative Distribution of Elements The following table compares the mass percentage of elements across different environments:
- Oxygen ( $O$ ):
  - Life (Humans): $65 \%$
  - Atmosphere: $21 \%$
  - Earth’s Crust: $46 \%$
- Carbon ( $C$ ):
  - Life (Humans): $18 \%$
  - Atmosphere: trace amounts
  - Earth’s Crust: trace amounts
- Hydrogen ( $H$ ):
  - Life (Humans): $10 \%$
  - Atmosphere: trace amounts
  - Earth’s Crust: $0.1 \%$
- Nitrogen ( $N$ ):
  - Life (Humans): $3 \%$
  - Atmosphere: $78 \%$
  - Earth’s Crust: trace amounts

Atomic Structure and Sub-atomic Particles

The Atom
- The atom is the smallest unit of matter that retains all the chemical properties of an element.
- Atoms are the fundamental building blocks of elements.
Atomic Regions
- Nucleus: Located at the center of the atom; it contains protons and neutrons.
- Outermost Region: Holds electrons in orbit around the nucleus.
Key Properties of Sub-atomic Particles
- Protons:
  - Charge: $+1$
  - Mass: $1\,amu$ (atomic mass unit)
  - Location: Nucleus
- Neutrons:
  - Charge: $0$ (neutral)
  - Mass: $1\,amu$
  - Location: Nucleus
- Electrons:
  - Charge: $-1$
  - Mass: $0\,amu$ (negligible mass)
  - Location: Orbitals (orbiting the nucleus)
- Example: In a Helium ( $He$ ) atom, these sub-atomic particles are arranged with protons and neutrons in the center and electrons in the surrounding orbitals.

Atomic Measurements and Isotopic Variation

Atomic Number and Mass
- Atomic Number: Defined as the number of protons in an atom. Each element has a distinct and unique atomic number.
- Atomic Mass: The mass of the atom, which is roughly equal to the sum of the number of protons and neutrons. It is expressed in atomic mass units ( $amu$ ).
- Electrons are excluded from atomic mass calculations due to their negligible mass.
- The number of neutrons can be calculated using the formula: $\text{Neutrons} = \text{Atomic Mass} - \text{Atomic Number}$ .
Isotopes
- Isotopes are forms of an element that have the same number of protons but different numbers of neutrons, resulting in different mass numbers.
- Carbon Example: Carbon has an atomic number of $6$ . Carbon-12 has a mass number of $12$ (6 protons, 6 neutrons), while its stable isotope Carbon-13 has a mass number of $13$ (6 protons, 7 neutrons). The approximate average atomic mass of Carbon is $12.11\,amu$ .
- Hydrogen Example:
  - $^1H$ (Protium): $0$ neutrons.
  - $^2H$ (Deuterium): $1$ neutron.
  - $^3H$ (Tritium): $2$ neutrons.
The Periodic Table
- The periodic table displays the atomic number and atomic mass for each element.
- The atomic number is placed above the element symbol, while the approximate atomic mass is placed below it.

Applications of Radioisotopes in Research

Radiometric Dating
- Radioisotopes are unstable isotopes that emit neutrons, protons, and electrons as they decay to reach a more stable form.
- Carbon Dating: Scientists use the decay of Carbon-14 ( $^{14}C$ ) into Nitrogen-14 ( $^{14}N$ ) to estimate the age of carbon-containing remains (e.g., a pygmy mammoth).
- This method is effective for objects less than approximately $50,000$ years old.
- For older objects, scientists analyze other isotopes such as Uranium (which decays to Lead), Potassium, or Rubidium.

Electron Configuration and the Bohr Model

Standard Electron Arrangement
- In an atom with a neutral charge, the number of protons equals the number of electrons.
- The Bohr model depicts electrons in circular orbits (electron shells or energy levels) at specific distances from the nucleus.
- Electrons occupy the lowest available energy shell ( $1n$ ) first before filling shells further away ( $2n, 3n$ , etc.).
The Octet Rule and Stability
- The outermost shell is known as the valence shell.
- The most stable atomic configuration occurs when the valence shell is filled.
- Octet Rule: For the first few shells, stability is achieved with eight electrons in the outer shell.
- Group 18 Elements: Known for having full valence shells and high stability.
- Group 1 Elements (e.g., $H, Li, Na$ ): Can achieve stability by losing an outer electron.
- Group 17 Elements: Can achieve stability by gaining one additional electron.
Electron Orbitals and Quantum Mechanics
- Modern science recognizes that electrons do not follow simple planet-like orbits.
- Orbitals: Complex-shaped regions where there is a high probability of finding an electron at any given time.
- Subshells within Bohr shells have unique shapes:
  - S subshell: Spherical shape, contains one orbital.
  - P subshell: Dumbbell shaped, contains three orbitals.
  - Other subshells include d and f subshells.

Chemical Reactions and Molecular Formation

Nature of Chemical Reactions
- Chemical reactions involve changes in the distribution of electrons between atoms.
- Reactants: Substances present at the start of the reaction.
- Products: Substances formed at the conclusion of the reaction.
Reaction Types
- Irreversible Reactions: Proceed in one direction until reactants are exhausted. Example: $2H_2O_2 \rightarrow 2H_2O + O_2$ .
- Reversible Reactions: Products can be converted back into reactants. Example: $HCO_3^- + H^+ \rightleftharpoons H_2CO_3$ .
Chemical Bonds
- A chemical bond is the attractive force linking atoms together to form molecules.
- Covalent Bonds: Formed when atoms share electrons. A water molecule ( $H_2O$ ) forms when two Hydrogens and one Oxygen share electrons.
- Double Bonds: Occur when more than one set of electrons is shared, such as in an Oxygen molecule ( $O_2$ ).
- Ionic Bonds: Occur when atoms transfer electrons. Metals typically lose electrons while nonmetals gain them to achieve an octet.
Polarity in Covalent Bonds
- Polar Covalent Bonds: Electrons are shared unequally because one nucleus has higher electronegativity (a stronger pull on electrons). In water, Oxygen is more electronegative than Hydrogen, creating partial charges ( $\delta$ ).
- Non-Polar Covalent Bonds: Electrons are shared equally. Example: Methane ( $CH_4$ ).
- Molecular Shape: Both bond type and shape determine polarity. Carbon Dioxide ( $CO_2$ ) contains polar bonds but is a nonpolar molecule due to its linear shape.

Intermolecular Forces and the Properties of Water

Bond Strength
- Ionic and covalent bonds are strong and require significant energy to break.
- Weaker interactions include Hydrogen bonds and Van der Waals interactions.
Weak Interactions
- Hydrogen Bonds: Interactions between the partial positive charge ( $\delta^+$ ) of a Hydrogen atom and the partial negative charge ( $\delta^-$ ) of a more electronegative atom on a different molecule. These frequently occur between water molecules.
- Van der Waals Interactions: Very weak attractions occurring between molecules in close proximity due to temporary changes in electron density.
Water’s Biological Importance
- Water comprises approximately $60\% \text{ to } 70\%$ of the human body.
- It is the most critical molecule for life on Earth because it is polar and can form hydrogen bonds.
- In a water molecule, the Oxygen pulls shared electrons closer, resulting in a slightly negative charge for Oxygen ( $\delta^-$ ) and a slightly positive charge for Hydrogen ( $\delta^+$ ). This allows the Hydrogen of one molecule to form a weak bond with the Oxygen of an adjacent molecule.