Notes: Atoms, Bonds, Water, pH, and Applications

Atoms and Subatomic Particles

Body composition note from lecture: about 5% of body weight is made up of key elements (the teacher mentions carbon, oxygen, nitrogen, hydrogen, phosphorus, and sulfur as part of the body’s makeup).
Atoms are the smallest unit of an element that still displays the element’s properties.
An element and its atoms share the same name (e.g., carbon, hydrogen, oxygen).
Subatomic particles:
- Protons: positive charge
- Neutrons: neutral (no charge)
- Electrons: negative charge
Location in the atom:
- Protons and neutrons are in the nucleus.
- Electrons orbit the nucleus in electron shells (orbit around the nucleus).
Chemical bonds are largely driven by electrons, not protons or neutrons; electrons in the outer shell (valence shell) determine bonding.
Which subatomic particle is responsible for chemical bonds? Electrons (in the outer shell) initiate interactions that form bonds.
Atomic symbols: one- or two-letter abbreviations (H for hydrogen, Na for sodium, C for carbon).
Atomic number (Z): equal to the number of protons in an atom.
Mass number (A): equal to the sum of protons and neutrons, i.e., $A = Z + N$ where N is the number of neutrons.
In a neutral atom, the number of protons equals the number of electrons (Z = number of electrons) so the net charge is zero.
Carbon example (to connect Z, A, and neutral state):
- Atomic number $Z=6$ (protons) and the mass number shown in some contexts is $A=12$ for carbon-12.
- Therefore, the number of neutrons $N = A - Z = 12 - 6 = 6.$
- In a neutral carbon atom, protons = electrons = 6.
Periodic table basics:
- Arrangement: horizontal rows are periods (increasing atomic number).
- Vertical columns are groups.
- Elements in the same group tend to share chemical bonding characteristics.
Isotopes:
- Definition: atoms of the same element that differ in the number of neutrons (and thus in mass number).
- Examples with carbon:
- Carbon-12: $Z=6, A=12, N=6$
- Carbon-13: $Z=6, A=13, N=7$
- Carbon-14: $Z=6, A=14, N=8$
- Isotopes may be radioactive and undergo spontaneous decay; can be helpful or harmful depending on context.
- Uses of radioactive isotopes:
- Diagnostics/medical imaging (e.g., PET scans) to visualize organs or detect tumors.
- Therapy to deliver radiation to cancer cells.
- Research into drug metabolism and how biological processes occur.
- Industrial uses such as leak detection in pipes.
- Carbon-14 dating to determine the age of organic materials in archaeology.

Energy Shells and the Bohr Model

Electrons are held in orbit around the positively charged nucleus by electrostatic attraction; energy is required to keep electrons in their shells.
Bohr-like energy shells describe electron location: electrons occupy discrete energy levels (shells).
Shell capacity and filling order:
- First energy shell can hold up to $2$ electrons.
- Each additional shell can hold up to $8$ electrons (in this simplified model relevant to biologically significant elements).
- Lower shells are filled before electrons populate higher shells.
Examples and implications:
- Hydrogen (Z = 1, e⁻ = 1): its single electron resides in the first shell; the first shell has capacity for 2, so hydrogen is chemically reactive.
- Helium (Z = 2, e⁻ = 2): the first shell is full; helium is inert and does not form bonds under normal conditions.
- Neon (Z = 10): first shell holds 2, second shell holds 8; the outer (valence) shell is filled, so neon does not readily form bonds.
Valence shell:
- The outermost shell determines many of an atom’s chemical properties.
- The octet rule: most stable when the outermost shell has 8 electrons. Exception: one-shell atoms like helium are stable with 2 electrons.
Additional examples for common elements (valence electron counts):
- Hydrogen: valence shell = 1 electron
- Carbon: atomic number 6, valence shell electrons = 4
- Nitrogen: atomic number 7, valence shell electrons = 5
- Oxygen: atomic number 8, valence shell electrons = 6

Molecules, Compounds, and Chemical Formulas

Molecule: two or more atoms bonded together; the smallest unit that retains the properties of the substance.
Compound: a molecule that contains at least two different elements bonded together.
Formula gives the number of each type of atom in a molecule (e.g., glucose): $ext{C}6 ext{H}{12} ext{O}_6$
- This indicates 6 carbon atoms, 12 hydrogen atoms, and 6 oxygen atoms in a glucose molecule.

Chemical Bonds

Chemical bonds are interactions between atoms that contain energy and hold substances together.
Bond formation is a chemical reaction driven by electron interactions in valence shells.
Types of bonds:
- Ionic bonds: formed when electrons are transferred from one atom to another; oppositely charged ions attract and form a bond.
- Covalent bonds: formed when atoms share electrons; each atom seeks a full octet in its valence shell.
- Nonpolar covalent bonds: electrons shared equally between atoms.
- Polar covalent bonds: electrons shared unequally, creating partial charges.
- Hydrogen bonds: weakest type of bond; a weak attraction between a slightly positive hydrogen atom in one molecule and a slightly negative atom (often O, N) in another molecule.
Electronegativity:
- Definition: the ability of an atom to attract electrons in a chemical bond.
- In water, oxygen is highly electronegative relative to hydrogen, leading to polar covalent bonds and partial charges (δ⁺ on H, δ⁻ on O).
Ionic bond example (from lecture): Sodium and chlorine form table salt (NaCl).
- Sodium (Na): Z = 11, protons/electrons = 11.
- Chlorine (Cl): Z = 17, protons/electrons = 17.
- Sodium tends to lose 1 electron (forming Na⁺, electron count becomes 10; protons remain 11).
- Chlorine tends to gain 1 electron (forming Cl⁻, electron count becomes 18; protons remain 17).
- Result: Na⁺ and Cl⁻ attract to form an ionic bond (NaCl).
Covalent bond types:
- Nonpolar covalent: electrons shared equally (e.g., H–H in H₂, O–O in O₂).
- Polar covalent: electrons shared unequally, creating dipoles (e.g., H–O in H₂O).
Water as an example of a polar molecule:
- Oxygen is more electronegative than hydrogen, leading to a partial negative charge on oxygen and partial positive charges on hydrogens.

Water: Properties and Significance in Biology

Water is a polar molecule that serves as a universal solvent because of its polarity; polar substances dissolve readily in water (e.g., table salt, NaCl).
Solutions: solvent (water) plus solute (substance dissolved in water).
Water as solvent example: NaCl dissolves in water due to polar interactions with ions (Na⁺ and Cl⁻).
Nonpolar substances (e.g., oils, waxes) do not dissolve well in water due to lack of charge separation.
Cohesion: water molecules stick to each other via hydrogen bonds, helping form droplets.
Adhesion: water molecules cling to other surfaces, enabling capillary action (e.g., water uptake in plant roots and transport in stems).
High heat capacity: water can absorb a large amount of thermal energy with only a small change in temperature, helping organisms maintain stable internal temperatures.
High heat of vaporization: significant energy is required to break hydrogen bonds so water can evaporate; sweating helps cool the body as water absorbs heat during evaporation.
Density and phase behavior:
- Ice is less dense than liquid water due to the lattice arrangement of hydrogen-bonded water molecules in the solid form, causing ice to float on liquid water.
- When water freezes, its molecules spread out into a lattice, making ice less dense than liquid water.
Water as a solvent in biology and medicine:
- Blood and many bodily fluids are water-based.
- Water dissolves salts and other polar molecules essential for physiology.

Acids, Bases, pH, and Buffers

pH: a measure of hydrogen ion concentration in a solution; water can ionize to produce H⁺ and OH⁻ ions in equilibrium.
Acids: substances that dissociate in water to release hydrogen ions (H⁺).
Bases: substances that accept hydrogen ions or release hydroxide ions (OH⁻) or otherwise decrease H⁺ concentration.
pH scale: ranges from 0 to 14; 0–6.9 is acidic; 7 is neutral; 7.1–14 is basic (alkaline).
Common acidic substances discussed: vinegar, soda, lemon juice, stomach acid, tomatoes, black coffee, urine.
Common basic/alkaline considerations: discussions of alkaline water and alkaline diets; note that practical health claims should be evaluated critically.
Buffers: chemicals or mixtures that help maintain pH within normal physiological ranges; crucial for the health of organisms.
Blood pH: typically around $pH ext{ (blood)} ightarrow 7.4$ (slightly basic).
- Acidosis: blood pH drops below $7.0$ , which can lead to serious problems such as infection, organ failure, and even death if severe.
- Alkalosis: blood pH rises above $7.8$ , which can trigger hyperventilation and other symptoms.
The body has buffering mechanisms to prevent dangerous pH changes and maintain homeostasis, analogous to a thermostat that maintains a target temperature.

Real-World Applications and Contexts

Radioisotopes in medicine and industry:
- Diagnostics (e.g., PET scans) using radioisotopes to image organs and detect tumors.
- Therapy delivering targeted radiation to cancer cells.
- Research to study drug metabolism and physiological processes.
- Industrial uses such as leak detection in piping systems.
- Carbon-14 dating to determine the age of organic materials in archaeology.
Practical notes from the lecture:
- The distinction between ionic and covalent bonds has practical implications for properties like solubility, reactivity, and the structure of biomolecules.
- Water’s properties (high heat capacity, high heat of vaporization, cohesion, adhesion, solvent abilities, and ice being less dense than liquid water) profoundly influence biology and climate.
- The octet rule and the idea of stable valence shells help explain why atoms form certain bonds and shapes, which in turn determine the structure and function of biomolecules such as glucose ( $ext{C}6 ext{H}{12} ext{O}_6$ ).

Quick Reference: Key Formulas and Concepts

Atomic and mass numbers:
- $Z = ext{number of protons} = ext{number of electrons in a neutral atom}$
- $A = Z + N$ where $N = ext{neutrons}$
Isotope relationships (example with carbon):
- Carbon-12: $Z=6, ext{ } N = 6, ext{ } A = 12$
- Carbon-13: $Z=6, ext{ } N = 7, ext{ } A = 13$
- Carbon-14: $Z=6, ext{ } N = 8, ext{ } A = 14$
Electron shell capacities (Bohr-like model):
- First shell: up to $2$ electrons
- Second and subsequent shells: up to $8$ electrons (in this simplified biologically relevant context)
Valence and the octet rule:
- Outer (valence) shell stability often achieved with $8$ electrons; helium is an exception with only $2$ electrons in its shell.
Bond types summarized:
- Ionic: transfer of electrons → ions attract to form a bond; typically between a metal and a nonmetal.
- Covalent: sharing of electrons; can be nonpolar (equal sharing) or polar (unequal sharing) depending on electronegativity differences.
- Hydrogen bonds: weak intermolecular attractions between a partially positive hydrogen and a negatively charged atom like O or N in another molecule.
Water-specific properties:
- Polar solvent; universal solvent for many substances
- High heat capacity and high heat of vaporization
- Cohesion and adhesion support droplets and capillary action
- Ice is less dense than liquid water due to hydrogen-bonded lattice structure
Acids, bases, pH, and buffers:
- Acid: dissociates to release H⁺ in solution
- Base: accepts H⁺ or releases OH⁻; pH indicates acidity/basicity
- pH scale ranges 0–14, with 7 neutral; below 7 acidic; above 7 basic
- Buffers stabilize pH to maintain homeostasis (e.g., blood around $pH ightarrow 7.4$ )

// End of notes — include connections to broader foundational principles (atomic structure, chemical bonding, thermodynamics of reactions, and homeostasis) and be prepared to explain specific examples (NaCl formation, H₂O polarity, carbon dating, PET imaging) on exam.