Chemistry of Life: Elements, Bonding, and Biochemical Reactions

Fundamental Concepts of Matter and Elements in Living Organisms

Matter: Defined as anything in the universe that possesses mass and/or volume.
Elements: Matter is composed of elements, which are pure substances that cannot be broken down using ordinary chemical or physical techniques.
Atoms: The atom is the smallest particle of an element that retains its chemical and physical properties.
Organic Compounds: All organic compounds contain the element Carbon ( $C$ ). Carbon is uniquely suited for building complex molecules due to its 4 valence electrons ( $e^-$ ), which allow for diverse bonding possibilities, including single, double, and triple bonds.

Elemental Composition of Living Organisms

Primary Elements (96%): The vast majority of a living organism's mass is composed of four elements:
- Carbon ( $C$ )
- Hydrogen ( $H$ )
- Oxygen ( $O$ )
- Nitrogen ( $N$ )
Secondary Elements (4%): These elements are essential but present in much smaller proportions:
- Calcium ( $Ca$ )
- Phosphorus ( $P$ )
- Potassium ( $K$ )
- Sulfur ( $S$ )
- Sodium ( $Na$ )
- Chlorine ( $Cl$ )
- Magnesium ( $Mg$ )
Trace Elements (< 0.1%): Necessary in minute quantities for physiological function:
- Iodine ( $I$ )
- Iron ( $Fe$ )

Atomic Structure and Isotopes

Subatomic Particles:
- Protons ( $p^+$ ): Carry a positive charge; located in the nucleus.
- Neutrons ( $n^0$ ): Carry no charge (neutral); located in the nucleus.
- Electrons ( $e^-$ ): Carry a negative charge; orbit the nucleus in energy shells.
Electrical Neutrality: Atoms typically have no net charge, meaning the number of protons equals the number of electrons ( $\#p^+ = \#e^-$ ).
Atomic Constants:
- Atomic Number: Equal to the number of protons ( $\#p^+$ ).
- Mass Number: The sum of protons and neutrons ( $\#p^+ + \#n$ ) within the nucleus. Electrons are excluded from this calculation because their mass is considered negligible.
Isotopes: Forms of a single element that differ in their neutron count, resulting in different mass numbers. Since they maintain the same number of protons and electrons, isotopes behave identically in chemical reactions.

Radioisotopes: Definition, Decay, and Applications

Definition: A radioisotope is a radioactive form or isotope of an element.
Instability and Decay: The nuclei of certain isotopes are unstable. They undergo a decay process where they break down, transforming the radioisotope into an atom of a different element.
Carbon-14 ( $^{14}C$ ) Example:
- Structure: 8 neutrons, 6 electrons, and 6 protons.
- Process: As it decays, it gives off particles and energy. Specifically, 1 neutron splits into a high-energy electron and a proton.
- Result: After the split, the isotope has 7 neutrons, 7 electrons, and 7 protons, effectively transforming $^{14}C$ into Nitrogen-14 ( $^{14}N$ ).
Scientific and Medical Uses:
- Dating: Used for dating organic materials, rocks, and fossils due to their steady, predictable rate of decay.
- Radioactive Tracers: Emitted particles serve as signals to:
  - Trace pathways as substances move through cells.
  - Visualize internal body structures for medical diagnosis.
  - Track chemical reactions to gain a deeper understanding of biological processes.

Electron Arrangements and Chemical Stability

Chemical Properties: The arrangement of electrons determines the chemical properties of an atom.
Valence Electrons: These are the electrons located in the atom's outermost energy shell (the valence shell).
Stability Rules: Atoms react by transferring or sharing valence electrons to achieve chemical stability, typically reaching a state of an octet (8 valence electrons) or a duet (2 valence electrons for smaller atoms).

Chemical Bonding: Ionic vs. Covalent

Ionic Bonds: These result from the attraction between two oppositely charged atoms or molecules. Electrons are transferred between atoms to create ions:
- Cation: A positively charged ion ( $+$ ).
- Anion: A negatively charged ion ( $-$ ).
- Solubility: Ionic compounds are often soluble in polar solvents, such as water ( $H_2O$ ).
Covalent Bonds: These result from the sharing of one or more pairs of electrons between atoms. The strength of the bond is determined by the electronegativity of the participating atoms.

Bond Polarity and Electronegativity

Electronegativity ( $EN$ ): A measure of an atom's attraction to shared electrons. If one atom is more electronegative, the electrons will reside closer to it rather than being equidistant.
Electronegativity Difference ( $\Delta EN$ ): Used to determine the specific bond type:
- Polar Covalent Bond: Uneven sharing of electrons where $\Delta EN \geq 0.41$ .
- Nonpolar Covalent Bond: Equal sharing of electrons where $\Delta EN \leq 0.40$ .

Molecular Polarity and Physical Properties

Polar Molecules:
- Dipoles: Covalent bonds can create partial charges ( $\delta^+$ or $\delta^-$ ). The atom attracting electrons more strongly becomes $\delta^-$ , while the other becomes $\delta^+$ .
- Characteristics: If partial charges are distributed non-uniformly, the whole molecule has polarity.
- Solubility: Polar molecules attract and align with other polar molecules and are generally soluble in water (e.g., Water and $CH_3OH$ ).
Nonpolar Molecules:
- Occur when all bonds are nonpolar OR when polar bonds are distributed symmetrically so their charges cancel out.
- Shape Factors: Molecular shape is a major determinant. Symmetrical molecules with no lone pairs on the central atom (where all outer atoms are identical) are nonpolar.
- Solubility: These have very low solubility in polar liquids (e.g., Oil and fat in water).
Example Cases:
- O-F bond: Polar covalent ( $4.0 - 3.5 = 0.5$ ; $0.5 \geq 0.41$ ).
- $SO_3$ molecule: Nonpolar because electrons are shared equally, charges cancel out, there are no lone pairs on the central atom, and all outer atoms are the same.

Intermolecular Forces (Van der Waals Forces)

Intermolecular Forces (IMF) are forces of attraction between two molecules. They are generally weak and occur when molecules are in close proximity.

Dipole-Dipole Forces: Permanent attractions between the $\delta^+$ and $\delta^-$ ends of two different polar molecules (e.g., attraction between $H-Cl$ molecules).
Hydrogen Bonding: A specific, strong attraction between a Hydrogen atom bonded to Nitrogen, Oxygen, or Fluorine ( $N, O, F$ ) and a lone electron pair on an $N, O, \text{ or } F$ atom of another molecule. This is stronger than standard dipole-dipole forces.
London Dispersion Forces: Attractions between the electrons of one molecule and the protons of another. These exist between all particles due to temporary dipoles formed by random electron movement. Strength increases with the number of electrons and molecule length.

Physical Property Influence: IMFs determine melting point, boiling point, state of matter, and surface tension. Molecules that are polar, larger, or have complex shapes typically exhibit higher melting/boiling points and are more likely to be liquid or solid at room temperature.

Major Types of Chemical Reactions

All chemical reactions involve breaking and forming bonds, rearranging atoms and ions. Four types are critical in biology:

Dehydration Synthesis (Condensation):
- A large molecule is formed from smaller subunits by removing water.
- An $-OH$ group and an $-H$ atom are removed from two reactant molecules to form $H_2O$ , joining the reactants.
- Anabolic: This is a building reaction that requires an input of energy.
Hydrolysis:
- Water is used to break down larger molecules into smaller ones.
- The bond in the reactant is broken, and the $-OH$ and $-H$ from water are attached to the resulting products.
- Catabolic: This is a breakdown reaction that releases energy.
Neutralization:
- An acid and a base react to form water and an ionic compound (a salt).
- Example: $H_2SO_4(aq) + Ca(OH)_2(aq) \rightarrow CaSO_4(s) + 2H_2O(l)$ .
Oxidation-Reduction (Redox):
- Involves the transfer of electrons.
- Oxidation: Loss of electrons.
- Reduction: Gain of electrons.
- Mnemonics: "OIL RIG" (Oxidation It Loses, Reduction It Gains) or "LEO the lion says GER" (Loss of Electrons is Oxidation, Gain of Electrons is Reduction).
- Agents: The oxidizing agent is the entity being reduced; the reducing agent is the entity being oxidized.
- Sample Reaction: $2NH_3 + 3Br_2 \rightarrow N_2 + 6HBr$ . In this reaction, Nitrogen is oxidized (losing $e^-$ ) and Bromine is reduced (gaining $e^-$ ).

Questions & Discussion

Q1: One atom has 6 protons and a mass number of 13. Another has 6 protons and a mass number of 15. Identify them and explain the difference.
- Response: Both are isotopes of Carbon ( $C$ ). The difference is the number of neutrons ( $7$ vs $9$ ).
Q2: List the three common isotopes of hydrogen and define radioisotopes and their uses.
- Response: Radioisotopes are unstable isotopes used in research and medicine for dating and as trackers.
Q3: An atom has 8 electrons, 6 of which are valence. Identify it and draw the model.
- Response: The element is Oxygen ( $O$ ). The first shell (2 $e^-$ ) and second shell (6 $e^-$ ) are occupied.
Q4: How do bonding arrangements affect molecular shape?
Q5: Compare ionic and covalent bonds.
Q6: How does atomic composition and shape affect polarity?
Q7: What effects do polarity, size, and shape have on physical properties and intermolecular forces?
Q8: Differentiate between polar and non-polar covalent bonds.
Q9: In an $N-H$ bond, which atom are electrons closer to?
- Response: Electrons are closer to Nitrogen because it is more electronegative than Hydrogen.
Q10: Explain Oxygen's role in polarity, bond shape, and redox reactions.
Q11: How do hydrogen bonds produce attractive forces and influence water's properties?
Q12: Describe the relationship between dehydration and hydrolysis.
- Response: They are opposite reactions; one builds by removing water, the other breaks down by adding water.
Q13: Describe redox and whether reduction can occur independently of oxidation.
- Response: No, they must occur together because electrons lost by one molecule must be gained by another.