Biochemistry 101: Atoms, Ions, Bonding, Water, and Basic Biochemical Concepts

Atoms, Elements, and the Basis of Biochemistry

Biochemistry focuses on chemistry as it applies to biology and living systems; connects chemistry to physics and biology.
A chemical element is the simplest form of matter with unique chemical properties. Elements are identified by their atomic number, which equals the number of protons in the nucleus.
- Example elements discussed: Hydrogen, Magnesium, Oxygen, Carbon, Iron.
The terms atom and element are interchangeable in this context; atoms are the building blocks that define elements.
Subatomic building blocks of atoms:
- Protons: positive charge, located in the nucleus;
- Neutrons: neutral charge, located in the nucleus;
- Electrons: negative charge, orbit around the nucleus; determine chemical reactivity.
The nucleus contains protons and neutrons; electrons are outside the nucleus and drive chemical behavior.
Charges of the building blocks:
- Proton: +1
- Neutron: 0
- Electron: −1
Isotopes:
- Atoms of the same element with varying numbers of neutrons; chemically similar because outer electrons (valence electrons) are the same, but they decay at different rates.
- Isotopes undergo decay via ionizing radiation, which is the energy release from changes in the nucleus.
Mass and charge considerations:
- Mass is largely due to protons and neutrons; electrons contribute negligibly to atomic mass.
- The nucleus is the core; electrons live in regions called orbitals around the nucleus.
Atomic models:
- Early models depicted electrons orbiting the nucleus like planets around the sun (simplified); not fully accurate.
- The quantum mechanical model presents the nucleus with protons and neutrons and probability distributions for electron locations; electrons exist in orbitals with certain probabilities.
Valence electrons:
- The outermost electrons (in the outer shell) are the ones that participate in chemical reactions.
- The term “valence” refers to those outer-shell electrons that determine chemical behavior.
What is the chemical identity of an element? The number of protons in the nucleus (atomic number). The total number of protons defines the identity of the element (e.g., Na, Cl, O, C).
Common chemical intuition aids memory:
- Water and sugars (e.g., glucose) are given as examples of molecular formulas:
- Water: $ext{H}_2 ext{O}$
- Glucose: $ext{C}6 ext{H}{12} ext{O}_6$
A quick periodic-table reminder:
- Six most abundant elements in the human body by typical teaching: Oxygen (O), Carbon (C), Hydrogen (H), Nitrogen (N), Calcium (Ca), Phosphorus (P).
- Symbols: O, C, H, N, Ca, P.
- Trace elements exist beyond these six (e.g., magnesium, iron) and are present at lower levels.
Organic vs inorganic distinction:
- Organic compounds contain carbon (by definition in this course).
- Inorganic compounds generally do not contain carbon (e.g., many minerals and salts).
Definition of ions and ionic context:
- An ion is a charged particle formed by gaining or losing electrons (electrons, not protons, are the focus here).
- Cation: positively charged ion (e.g., Na⁺).
- Anion: negatively charged ion (e.g., Cl⁻).
- Example of ion formation: Sodium gives up one electron to chlorine, producing Na⁺ and Cl⁻, which attract each other to form a compound.
- Sodium chloride: $ext{NaCl}$ (table salt).
Ionic bonding in brief:
- Result of electron transfer between a cation and an anion.
- Ionic bonds are relatively weak in water and readily dissociate into ions in aqueous solutions (electrolytes).
- Electrolytes conduct electricity when dissolved in water (e.g., Na⁺, Cl⁻ from table salt).
Covalent bonding and bond strength:
- Covalent bonds involve sharing electrons between atoms.
- Compared to ionic bonds, covalent bonds are generally much stronger and do not dissociate easily in water.
- Double and triple covalent bonds (e.g., O=O, H–O–H in water) involve sharing more electrons.
Polar vs nonpolar covalent bonds:
- Polar covalent bonds occur when electrons are shared unequally due to differences in electronegativity; one end becomes slightly negative (electronegative) and the other slightly positive (electropositive).
- Oxygen is more electronegative than hydrogen in H₂O, leading to partial charges: $ext{O} ext{δ}^- ext{ and } ext{H} ext{δ}^+$ .
- This unequal sharing leads to a polar covalent bond in water.
Water as a special case:
- Within a single water molecule, bonds between hydrogen and oxygen are covalent, with polar covalent character due to unequal sharing.
- Across water molecules, hydrogen bonds form between the electropositive hydrogen of one molecule and the electronegative oxygen of another: this is a hydrogen bond.
- Hydrogen bonding is crucial for many properties of water and is foundational for life.
Polar covalent bonds and the concept of electronegativity/electropositivity:
- Polar covalent bonds arise when electrons spend more time near the more massive atom (in water, more time near oxygen).
- Resulting charges create a dipole: electronegative end (partial negative) vs electropositive end (partial positive).
- The term polar covalent bond captures the idea of two different ends with opposite partial charges.
Hydrogen bonding and its importance:
- Hydrogen bonds are weaker than covalent bonds but numerous and cooperative, giving water many of its unique properties (e.g., high boiling point relative to size, cohesion, solvent capabilities).
Osmosis and electrolytes (practice applications):
- Salts like NaCl, KCl dissociate in water into ions and can act as electrolytes.
- Osmosis is the movement of water across a semipermeable membrane driven by solute (salt) gradients.
- Example: Salt on a slug dehydrates it by drawing water out due to osmotic effects; water moves toward the higher salt concentration.
Electrolyte balance and clinical relevance:
- Electrolyte balance is essential for physiological function, especially for nerve and muscle activity.
- Imbalances can lead to severe outcomes (e.g., cardiac or nerve failure).
- Common clinical driver of electrolyte imbalance: dehydration, especially with chronic illness and diarrhea, affecting water and salt balance and kidney function.
Free radicals and antioxidants (biochemical relevance):
- Free radical: an unstable, highly reactive particle with an unusual number of electrons; can damage molecules and cells.
- Common example: reactive oxygen species such as superoxide anion (e.g., O₂⁻).
- Antioxidants neutralize free radicals: selenium, vitamins A, C, carotenoids (e.g., beta-carotene in carrots).
Basic molecular concepts:
- Molecule: two or more atoms bonded together.
- Water (H₂O) is a molecule composed of two hydrogen atoms and one oxygen atom.
- A compound is a molecule composed of two or more different elements (e.g., H₂O is a compound; O₂ is not).
- Molecular formula examples: $ext{H}2 ext{O}, ext{C}6 ext{H}{12} ext{O}6$ .
Isotopes and radiation in clinical context:
- Isotopes are chemically similar but decay at different rates, making them useful in medical imaging and therapy.
- Ionizing radiation types: ultraviolet (UV), X-rays, alpha, beta, gamma rays.
- Radiation dose units and concepts:
- The Sievert (Sv) is the unit used to measure radiation dose in terms of biological effect.
- Common practical units: $1 ext{ Sv} = 1000 ext{ mSv}$ ; a typical public exposure limit is on the order of tens of millisieverts per year; exposure of around 5 Sv at once is fatal.
- Physical half-life vs biological half-life:
- Physical half-life: time required for 50% of a radioactive isotope to decay to a stable state, denoted as $t_{1/2}^{ ext{phys}}$ .
- Biological half-life: time required for the body to metabolize or eliminate half of the substance, denoted as $t_{1/2}^{ ext{bio}}$ .
- Natural sources of ionizing radiation include cosmic rays and radon gas; radon levels vary by geological conditions and can be monitored.
- Medical history context: Marie Curie (Polish origin) contributed to early work on radioactivity; coined the term radioactivity; discovered polonium and radium; trained doctors in early X-ray use; died from radiation exposure.
Key historical and foundational connections:
- Understanding atoms and ions underpins how drugs interact with biological systems, how electrolytes govern physiology, and how radiation interacts with matter in medicine.
- The quantum mechanical view of atoms sets the stage for quantum-level medicine and quantum computing in future research.
Quick reference recaps (for exam-style prompts):
- Six most abundant elements in the human body: $ext{O}, ext{C}, ext{H}, ext{N}, ext{Ca}, ext{P}$ .
- Isotopes: chemically similar; different neutron numbers; different decay rates; used in imaging/therapy; concept of physical vs biological half-life.
- Ions and bonds:
- Ionization: transfer of electrons leading to ions.
- Ionic bond: electrostatic attraction between cation and anion; salts dissolve/disassociate in water into ions; electrolytes conduct electricity.
- Covalent bond: sharing of electrons; can be polar (electrons not shared equally) or nonpolar.
- Polar covalent bonds create partial charges; water is a classic example with an electronegative oxygen end and an electropositive hydrogen end.
- Hydrogen bonds: between molecules (e.g., between water molecules); enable water’s unique properties and biological processes.
- Water chemistry and life: water’s polarity and hydrogen bonding are central to cellular processes, nutrient transport, and macromolecular interactions.
- Clinical relevance: dehydration and electrolyte imbalance can threaten cardiac and nervous system function; osmosis and electrolyte balance are foundational in patient care.
If you want, I can turn these notes into a condensed study sheet with suggested exam questions and quick-run answers.

1. Atoms and Elements Fundamentals

Biochemistry: Chemistry applied to biology and living systems.
Chemical Element: Simplest form of matter with unique chemical properties, defined by its atomic number (number of protons).
Atom: Building block defining an element.
Subatomic Particles:
- Protons: Positive charge ( $+1$ ), in nucleus.
- Neutrons: Neutral charge ( $0$ ), in nucleus.
- Electrons: Negative charge ( $-1$ ), orbit nucleus, determine chemical reactivity.
Isotopes: Atoms of the same element with varying numbers of neutrons; chemically similar (same valence electrons) but decay at different rates via ionizing radiation (energy release from nuclear changes).
Valence Electrons: Outermost electrons; participate in chemical reactions and determine chemical behavior.

2. Chemical Identity and Key Elements

Chemical Identity: Defined solely by the number of protons (atomic number).
Six Most Abundant Elements in Human Body: Oxygen (O), Carbon (C), Hydrogen (H), Nitrogen (N), Calcium (Ca), Phosphorus (P).
Organic vs. Inorganic Compounds:
- Organic: Contains Carbon (by course definition).
- Inorganic: Generally does not contain Carbon (e.g., minerals, salts).

3. Ions and Chemical Bonds

Ion: Charged particle formed by gaining (anion) or losing (cation) electrons.
- Cation: Positively charged ion (e.g., $\text{Na}^+$ ).
- Anion: Negatively charged ion (e.g., $\text{Cl}^-$ ).
Types of Chemical Bonds:
- Ionic Bond: Formed by electron transfer between cation and anion (electrostatic attraction).
  - Relatively weak in water; readily dissociate into ions (electrolytes) in aqueous solutions ( $\text{e.g., NaCl} \to \text{Na}^+ + \text{Cl}^-$ ).
  - Electrolytes conduct electricity in water.
- Covalent Bond: Involves sharing electrons between atoms; generally stronger than ionic bonds and do not easily dissociate in water.
  - Polar Covalent Bond: Unequal sharing of electrons due to electronegativity differences (e.g., $\text{H}_2\text{O}$ : Oxygen is $\text{δ}^-$ , Hydrogen is $\text{δ}^+$ ).
  - Nonpolar Covalent Bond: Equal sharing of electrons.
- Hydrogen Bond: Weaker intermolecular bond between an electropositive hydrogen of one molecule and an electronegative atom (e.g., Oxygen) of another molecule.
  - Crucial for water's unique properties (high boiling point, cohesion, solvent).

4. Special Topics and Clinical Relevance

Molecules vs. Compounds:
- Molecule: Two or more atoms bonded together (e.g., $\text{H}2\text{O}, \text{O}2$ ).
- Compound: A molecule composed of two or more different elements (e.g., $\text{H}2\text{O}$ is a compound; $\text{O}2$ is not).
Osmosis: Movement of water across a semipermeable membrane driven by solute gradients.
Electrolyte Balance: Essential for physiological function (nerve, muscle activity); imbalances (e.g., from dehydration) can be severe.
Free Radicals: Unstable, highly reactive particles with an unusual number of electrons (e.g., superoxide anion $\text{O}_2^-$ ); can damage cells.
- Antioxidants: Neutralize free radicals (e.g., selenium, vitamins A/C, carotenoids).
Isotopes and Radiation:
- Used in medical imaging and therapy due to their decay rates.
- Ionizing Radiation: UV, X-rays, alpha, beta, gamma rays.
- Sievert (Sv): Unit for radiation dose (biological effect); $1 \text{ Sv} = 1000 \text{ mSv}$ .
- Physical Half-life ( $t_{1/2}^{\text{phys}}$ ): Time for 50% of an isotope to decay.
- Biological Half-life ( $t_{1/2}^{\text{bio}}$ ): Time for body to eliminate 50% of a substance.
- Marie Curie: Pioneer in radioactivity research (coined term, discovered polonium/radium), died from radiation exposure.

5. Quick Recap for Exams

Elements (Human Body): O, C, H, N, Ca, P.
Isotopes: Same element, different neutrons; chemically similar; different decay rates; clinical uses (imaging/therapy); physical vs. biological half-life.
Ions: Charged particles from electron transfer (gain/loss).
Ionic Bonds: Electron transfer; form ions; electrolytes (conduct electricity); dissociate in water.
Covalent Bonds: Electron sharing; strong.
- Polar: Unequal sharing (e.g., water).
- Nonpolar: Equal sharing.
Hydrogen Bonds: Intermolecular bonds (e.g., between water molecules); crucial for water properties and life.
Water Chemistry & Life: Polarity and hydrogen bonding are central to biological processes.
Clinical Relevance: Dehydration, electrolyte imbalance, osmosis are vital in patient care.