Chapter 2: The Chemical Context of Life – Study Notes

Chapter 2: The Chemical Context of Life

Concept 2.1: Elements and the Periodic Table

The chemical basis of life is founded on a small set of elements, most of which are essential for biological processes. The periodic table organizes all known elements by their atomic structure and properties.

• Element: A substance that cannot be broken down into other substances by chemical reactions.

• SPONCH elements: Sulfur, Phosphorus, Oxygen, Nitrogen, Carbon, Hydrogen make up ~98% of body weight in living things.

• Trace elements: Elements required in minute quantities (e.g., Fe, Zn, I).

Element	Symbol	Percentage of Body Mass
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%

Concept 2.2: Atomic Structure

Atoms are the smallest units of matter that retain the properties of an element. Their structure determines chemical behavior.

• Atomic number: Number of protons in the nucleus.

• Mass number: Number of protons plus neutrons.

• Electrons: Negatively charged particles found in energy levels (shells) around the nucleus. Their mass is ~1/1800 of a proton and is usually ignored in atomic mass calculations.

• Electron shells: Layers around the nucleus where electrons reside; each shell has a different energy level.

• Potential energy: Electrons in higher energy levels have greater potential energy due to their position.

• Electron movement: Electrons can move between shells; moving farther from the nucleus requires input of energy.

Electron Energy Levels

• Energy is absorbed when electrons move to higher shells; energy is lost when they fall to lower shells.

• Analogy: Like a ball bouncing down stairs, electrons lose energy as they move closer to the nucleus.

Electron Configuration and Orbitals

• Valence shell: The outermost shell, determines chemical behavior.

• Octet rule: Atoms are most stable when their valence shell is full (usually 8 electrons).

• Orbitals: 3D spaces where electrons are likely to be found; each shell contains several orbitals (e.g., s, p).

• Inert (noble) gases: Have full valence shells and are chemically nonreactive.

Element	Electron Configuration	Reactivity
Neon	1s22s22p6	Inert
Oxygen	1s22s22p4	Reactive

Concept 2.2: Electronegativity

Electronegativity is a measure of an atom's ability to attract electrons in a chemical bond. The difference in electronegativity between atoms determines the type of bond formed.

• Electronegativity values: Higher values indicate stronger attraction for electrons.

• Key biological elements:

  Element	Electronegativity
  O	3.5
  N	3.0
  C	2.5
  H	2.1
  S	2.5
  P	2.1
  Na	0.9
  Cl	3.0

• Bond type: Absolute value of electronegativity difference determines bond type (covalent, ionic).

Concept 2.3: Chemical Bonding

Atoms form chemical bonds to achieve stable electron configurations. The main types of bonds in biology are covalent, ionic, hydrogen, and van der Waals interactions.

Covalent Bonds

• Definition: Sharing of electron pairs between atoms.

• Single, double, triple bonds: Number of shared electron pairs; triple bonds are strongest, single are weakest.

• Bond strength: Quantified by energy required to break the bond. $\text{C--C: 82.7 kcal/mol}$ $\text{C=C: 143.9 kcal/mol}$ $\text{C≡C: 200 kcal/mol}$ $\text{O--H single covalent: 110 kcal/mol}$ $\text{O--H hydrogen bond: 4.5 kcal/mol}$

Nonpolar vs. Polar Covalent Bonds

• Nonpolar covalent bond: Electrons shared equally; occurs when electronegativity difference < 2.0 (e.g., H--H, C--H).

• Polar covalent bond: Electrons shared unequally; occurs when electronegativity difference is larger (e.g., O--H, C--O).

• Partial charges: Polar bonds create regions of partial positive and negative charge (dipoles).

Geometry and Polarity

• Molecular geometry: Determines overall polarity; linear geometry can cancel dipoles (nonpolar), bent geometry adds dipoles (polar).

• Example: Water (H2O) is polar due to bent geometry and electronegativity difference between O and H.

Ionic Bonds

• Definition: Transfer of electrons from one atom to another, creating ions (cations and anions).

• Example: Sodium chloride (NaCl) forms when Na donates an electron to Cl, resulting in Na+ and Cl-.

• Ionic bond strength: Strong in air, weak in water due to solvation.

Hydrogen Bonds

• Definition: Weak attraction between a hydrogen atom in one polar molecule and an electronegative atom in another.

• Example: Hydrogen bonding between water molecules, or between water and ammonia.

Van der Waals Interactions

• Definition: Weak attractions due to transient dipoles in nonpolar molecules.

• Example: Instantaneous dipoles in methane molecules.

Hybridization and Molecular Shape

• Hybrid orbitals: Atomic orbitals combine to form new shapes, influencing molecular geometry.

• Tetrahedral geometry: Common in molecules like methane (CH4) and water (H2O).

• Bond angles: Water has a bond angle of 104.5°, methane has 109.5°.

Summary Table: Types of Chemical Bonds

Bond Type	Mechanism	Strength	Example
Covalent	Electron sharing	Strong	H2, O2, CH4
Ionic	Electron transfer	Strong (air), weak (water)	NaCl
Hydrogen	Attraction between polar molecules	Weak	H2O, NH3
Van der Waals	Transient dipoles	Very weak	CH4

Example Applications

• Photosynthesis: Electron movement between shells is essential for energy capture in plants.

• Water properties: Hydrogen bonding gives water its unique characteristics (cohesion, adhesion, high specific heat).

Additional info: These notes expand on the provided slides and handwritten content, adding definitions, examples, and tables for clarity and completeness.