Biochemistry — Elements of Life and Basic Bonding Concepts (Study Notes)

Page 2: Elements of Life

• Learning Objective ENE-1.A: Describe the composition of macromolecules required by living organisms.
• Main idea: Life relies on a set of key elements and the macromolecules they form (carbohydrates, proteins, lipids, nucleic acids).
• Context: CHON are the primary building blocks; trace elements are required in smaller amounts but are essential for homeostasis.

Page 3: Biochemistry Intro Question

• Question: What are atoms and why are they important to living organisms?
• Core idea: Atoms are the fundamental units of matter that compose all biological molecules and structures; understanding atoms explains how molecules form, interact, and function in biology.

Page 4: Atoms — The Smallest Unit Biochemists Care About

• Answer: The smallest stable unit of matter that retains the characteristics of its element.
• Structure info:
  • Nucleus contains protons and neutrons; nucleus is positively charged by protons.
  • Protons: positively charged.
  • Neutrons: neutral.
• Significance: Determines the identity and properties of the element; foundational for bonding and reactivity.

Page 5: Electrons and Orbitals

• Electrons: negatively charged particles that surround the nucleus in orbitals.
• Each orbital has a different energy level.
• Energy relationship:
  • Closer to the nucleus = lower energy level.
  • Farther from the nucleus = higher energy level.
• Implication: Energy levels influence bonding possibilities and chemical behavior.

Page 6: Valence Electrons and the Octet Tendency

• Valence electrons: the electrons in the outermost (valence) shell.
• Key role: Used for making different types of bonds.
• Octet rule (stability): Most elements strive to have 8 electrons in their valence shell.
• Consequence: Drives bonding patterns and molecular stability.

Page 7: Ions — Anions and Cations

• Atoms are usually neutral: #electrons = #protons.
• Ions: non-neutral atoms with a net charge.
• Anions: negatively charged (more electrons than protons).
• Cations: positively charged (more protons than electrons).
• Two types of ions: cations and anions.
• Practical note: Ion formation is crucial for reactions in biology (electrostatic interactions, salt balance, neurotransmission, etc.).

Page 8: Elements of Life — CHON and Trace Elements

• Main elements of life: Carbon (C), Hydrogen (H), Oxygen (O), Nitrogen (N) – collectively CHON.
• Trace elements: all other essential elements present in small amounts in the body.
• Despite being trace, they are critical for homeostasis and survival.
• Without trace elements, organisms would die due to failed biological processes.

Page 9: The Cycle(s) of Life

• Concept: Life operates in cycles (carbon, nitrogen, phosphorus, etc.) connecting organisms and the environment.
• (Note: The slide title indicates cycles, but no further details are given in this page excerpt.)

Page 10: The Importance of Carbon

• Carbon is the base of all life.
• Plant role: Plants draw carbon from the atmosphere during photosynthesis.
• Carbon incorporation: Carbon is incorporated into carbohydrates and serves as the main source of biomass in ecosystems.
• Biomolecule construction: Organisms build carbohydrates, proteins, nucleic acids, and lipids using carbon.
• Recycling: Decomposers recycle carbon back into the environment when organisms die.
• Carbon-depleted areas: Organisms there die because they cannot synthesize necessary biomolecules without carbon.

Page 11: The Importance of Nitrogen

• Nitrogen is essential: "No nitrogen, no GAINS" (mnemonic reference).
• Nitrogen cycle: Inorganic nitrogen is fixed from the atmosphere by bacteria and decomposers; plants absorb fixed nitrogen into the food web.
• Role in biology: Nitrogen is used to make proteins and nucleic acids.
• Recycling: Nitrogen is recycled back into the environment by decomposers.
• Nitrogen-depleted areas: Organisms cannot synthesize proteins or nucleic acids without nitrogen.

Page 12: The Importance of Phosphorus

• Phosphorus is used to build nucleic acids and certain phospholipids.
• Phosphorus-depleted environments lead to failure to make nucleic acids and phospholipids (major component of cell membranes).
• Note: Referred to as the “wild card” due to its essential and sometimes limiting role in biology.

Page 13: Practice — Aquarium Metabolism Reflection

• Prompt: Think about how the chemicals secreted by the fish are used by the plant in this aquarium.
• Concept focus: Nutrient cycling between organisms via the microbial loop and plant uptake.

Page 14: Practice Question — Nutrient Pathways in an Aquarium Model

• Question: Which statement best describes how molecules released by fish become nutrients for plants?
• Options (paraphrased):
  • A: CO₂ from fish is converted by bacteria into oxygen atoms used to make water.
  • B: O₂ from fish is converted by bacteria into ammonia for lipids/fatty acids.
  • C: Nitrites from fish are converted by bacteria into CO₂ for carbohydrates.
  • D: Ammonia from fish is converted by bacteria into nitrates for proteins and nucleic acids.
• Correct interpretation: D is the accurate pathway—ammonia is oxidized by bacteria to nitrite and then to nitrate, which plants use to synthesize proteins and nucleic acids.
• Context: The model depicts plant, fish, and bacteria interactions in the aquarium ecosystem.

Page 15: Electronegativity

• Definition: Electronegativity is the measure of how strongly atoms attract bonding electrons to themselves.
• Interpretation: It indicates how much an atom will pull electrons toward itself in a bond.
• Determinants: Primarily determined by the number of electrons in the valence shell.
• Trend: The closer an atom is to having eight electrons in its valence shell, the more electronegative it tends to be.

Page 16: Electronegative Elements You Need to Know

• Top trio context: Fluorine, Oxygen, Nitrogen are among the most electronegative elements.
• Relative electronegativity in biology:
  • Fluorine is the most electronegative element among common elements, but it is a trace element in biology and not commonly used in biomolecules.
  • Oxygen is more electronegative than nitrogen.
  • Nitrogen is less electronegative than oxygen.
• Taken together: F > O > N within the typical biologically relevant set; fluorine’s high electronegativity is mathematically represented, but its biological role is limited compared to O and N.

Page 17: Electropositivity

• Definition: Measure of an element’s ability to donate electrons and form positive ions.
• Common characteristic: Elements with 1 or 2 electrons in their valence shells.
• Correlation: They are not very electronegative, i.e., they tend to lose electrons rather than attract them.

Page 18: Electrons and Bonding

• Topic title: Electrons and Bonding (intro to how atoms bond and how electrons are involved in bonding).

Page 19: Covalent Bonds vs. Ionic Bonds

• Ionic bonds (concept as stated): Transfer of valence electrons from a metal to a non-metal; often weaker and can dissociate in water.
• Covalent bonds (concept as stated): Electrons are shared between atoms; energy is stored in covalent bonds and released when broken; typically stronger and central to most biology.
• Note on slide accuracy: The provided slide contains inaccuracies (e.g., statement that energy is stored in covalent bonds while describing ionic bond formation). Correct definitions are provided above: covalent bonds involve electron sharing; ionic bonds involve electron transfer and are often weaker in aqueous environments.

Page 20: Polarity

• Definition: Polarity arises from unequal sharing of electrons across a covalent bond.
• Cause: Large differences in electronegativity between bonded atoms.
• Consequence: Creates partial charges on atoms, leading to dipole moments in molecules.

Page 21: Polarity — Charge Distribution

• Polar molecules have an overall neutral net charge.
• Partial charges:
  • Partially negative on the more electronegative atom due to electron density drawn toward it.
  • Partially positive on the less electronegative atom.
• Implication: Polar molecules interact via dipole-dipole interactions and hydrogen bonding, affecting solubility and interactions in biological systems.

Page 22: Hydrogen Bonds

• Definition: Weak attraction between a hydrogen atom attached to a highly electronegative atom (O, N, or F) and another electronegative atom (O, N, or F).
• Why O, N, F?: They are highly electronegative; hydrogen attached to them carries a partial positive charge; the other electronegative atom has partial negative charge.
• Result: Hydrogen bonds help stabilize structures (e.g., DNA double helix, protein folding) and mediate interactions in water and biomolecules.

Page 23: Bonds and Molecular Shape

• Key idea: The way atoms bond determines the molecule’s shape and geometry.
• Importance: Shape dictates chemical properties and biological function.
• Takeaway: Structure, shape, and chemical properties are intimately linked and will recur across biology.

Page 24: Laws of Conservation

• Energy conservation: Energy is always conserved in a reaction; energy lost is typically released as heat.
• Atomic conservation: The number and types of atoms are conserved in a reaction (atoms are not created or destroyed).
• Bond conservation: The total number of bonds is conserved in a reaction (you cannot create or destroy bonds out of nowhere; bonds may break and form, but their total count remains the same in a closed process).

Page 25: Any Questions?

• Open forum for questions and clarification about the material.

Notes and tips for study:

• Remember the CHON elements and why trace elements matter despite their low abundance.
• Understand how carbon, nitrogen, and phosphorus form the backbone of essential biomolecules (carbohydrates, proteins, nucleic acids, lipids).
• Practice the nitrogen cycle concept depicted in the aquarium model: ammonia -> nitrite -> nitrate via bacteria, with plants incorporating nitrates.
• Distinguish covalent vs. ionic bonds and how polarity arises from electronegativity differences.
• Be able to explain hydrogen bonds and their role in biomolecular structure and interactions.
• Keep in mind the conservation laws; they are universal in chemistry and biology and underpin metabolic pathways and reaction stoichiometry.

LaTeX examples to remember:

• Valence octet: $8$ electrons in the valence shell.
• Carbon in biomolecules: $ext{C}$ as a central element in carbohydrates, proteins, nucleic acids, lipids.
• Common molecules: $H<em>2O$, $CO</em>2$, $NH<em>3$, $NO</em>3^-$, etc.