Notes on Atoms, Elements, and the Periodic Table (Pages 1–5)

Origins of chemical symbols

Elements are referenced by chemical symbols derived from their names in English, Latin, or Greek.
Examples:
- Oxygen: symbol O, from the English word "oxygen".
- Gold: symbol Au, from the Latin word "aurum" meaning "gold".

Elements in the Animal Body (Table 2.1)

Major elements (make up ~96% of body mass in animals; shading in the table indicates importance):
- Oxygen (O)
- Atomic number: Z = 8
- Body mass percentage: ext{Body Mass ( ext{%)} } o ext{65.0}
- Function: Necessary for cellular respiration; component of water
- Carbon (C)
- Atomic number: Z = 6
- Body mass percentage: 18.5
- Function: Component of organic molecules
- Hydrogen (H)
- Atomic number: Z = 1
- Body mass percentage: 9.5
- Function: Component of water and organic molecules
- Nitrogen (N)
- Atomic number: Z = 7
- Body mass percentage: 3.3
- Function: Component of all proteins and nucleic acids
Minor elements
- Calcium (Ca)
- Atomic number: Z = 20
- Body mass percentage: ~1.0%
- Function: Principal component of bones and teeth; required for muscle contraction, nerve impulse transmission, and blood clotting
- Phosphorus (P)
- Atomic number: Z = 15
- Body mass percentage: ~0.4%
- Function: Component of backbone of nucleic acids; important in energy transfer and respiration; ion influences pH of fluids
- Potassium (K)
- Atomic number: Z = 19
- Body mass percentage: trace to minor (~0.3–0.4%)
- Function: Principal positive ion within cells; important in nerve function
- Sulfur (S)
- Atomic number: Z = 16
- Body mass percentage: ~0.3%
- Function: Component of most proteins and various enzymes
- Sodium (Na)
- Atomic number: Z = 11
- Body mass percentage: ~0.2%
- Function: Important positive ion in extracellular fluid; important in nerve function
- Chlorine (Cl)
- Atomic number: Z = 17
- Body mass percentage: ~0.2%
- Function: Component of many energy-transferring processes and molecules
- Magnesium (Mg)
- Atomic number: Z = 12
- Body mass percentage: ~0.1%
- Function: Component of many energy-transferring enzymes
Trace elements
- Silicon (Si)
- Atomic number: Z = 14
- Body mass percentage: ~0.1%
- Function: Component of some enzymes
- Aluminum (Al)
- Atomic number: Z = 13
- Body mass percentage: ~0.1%
- Function: Component of some enzymes
- Iron (Fe)
- Atomic number: Z = 26
- Body mass percentage: ~0.1%
- Function: Critical component of hemoglobin
- Manganese (Mn)
- Atomic number: Z = 25
- Body mass percentage: ~0.1%
- Function: Needed for fatty acid synthesis
- Fluorine (F)
- Atomic number: Z = 9
- Body mass percentage: ~0.1%
- Function: Component of bones and teeth
- Vanadium (V), Chromium (Cr)
- Atomic numbers: Z = 23 ext{ (V)}, 24 ext{ (Cr)}
- Body mass percentage: ~0.1%
- Function: Component of some enzymes; Cr involved in glucose metabolism
The periodic table (Fig. 2.4) provides: symbol, atomic number, and atomic weight for each element; groups show similar properties; left side is mostly metallic elements, right side contains inert gases.
Shading in Fig. 2.4:
- Red: major elements
- Blue: minor elements
- Yellow: trace elements

The Periodic Table: Key Concepts

Purpose: to give important information about each element: chemical symbol, atomic number, and atomic weight.
Organization:
- Elements with similar properties are grouped together.
- Metallic elements are on the left; inert (noble) gases are on the right.
Major elements make up 96% of animal body mass (as indicated by red shading in Fig. 2.4).

Atoms and Subatomic Particles

An atom is the smallest unit of an element that retains the element’s properties (Fig. 2.6).
Subatomic particles:
- Protons: positively charged; mass ~1 atomic mass unit; located in the nucleus.
- Neutrons: electrically neutral; mass ~1 atomic mass unit; located in the nucleus.
- Electrons: negatively charged; very small mass; orbit the nucleus in regions called electron clouds (Fig. 2.7).
Nuclear composition and atomic weight:
- The nucleus contains protons and neutrons; together they determine the atomic weight of the atom.
- Protons and neutrons are the heaviest particles; each has a mass of about 1 amu.
Electron behavior:
- Electrons are extremely light and their mass does not contribute significantly to the atom’s weight.
- Electrons exist in motion around the nucleus and their exact positions are best described statistically via electron clouds.
- The negative charge of electrons balances the positive charge of protons in neutral atoms.
- In neutral atoms, the number of protons equals the number of electrons, giving no net charge.
Ionization:
- If an atom loses or gains electrons, it becomes charged (an ion).
- An ion has a net electrical charge due to the imbalance between protons and electrons.
Isotopes:
- Atoms can have different numbers of neutrons; atoms with the same number of protons but different neutrons are isotopes.
- Example: Carbon normally has 6 protons and 6 neutrons (carbon-12), but some isotopes have more neutrons (e.g., carbon-14).
- Carbon-14 is radioactive and emits energy as it decays to a stable element; the rate of decay is used to date fossils.
The nucleus and electron arrangement are represented by two common models:
- Planetary model: Protons and neutrons in the nucleus with electrons orbiting like planets around the sun. This model is not physically accurate but aids basic understanding.
- Orbital model: A three-dimensional electron cloud representing the most likely locations of electrons at any given time; reflects probability distributions rather than fixed orbits.
Equality of particles in neutral atoms:
- The number of protons in the nucleus equals the number of electrons in a neutral atom: Z = E, where Z is the atomic number and E is the number of electrons.
Fundamental relationships (definitions and simple equations):
- Atomic number: Z = ext{number of protons}
- Mass number (nucleon count): A = Z + Nn, where Nn is the number of neutrons
- Isotopes differ in neutron number (same Z, different N_n)
- Charge of an ion: q = Z - E
- Atomic weight is determined by protons and neutrons in the nucleus
- For radioactive decay (example: carbon-14): Activity A = \lambda N, half-life t_{1/2} = \frac{\ln 2}{\lambda}

Isotopes and Carbon-14: Practical Implications

Isotopes differ in neutron number but have the same chemical properties because they have the same number of protons (the same element).
Radioactive isotopes (e.g., carbon-14) decay at a constant rate (characterized by the decay constant \lambda); this decay can be used for dating fossils and geological samples.
Carbon-14 dating relies on measuring the ratio of carbon-14 to carbon-12 in a sample and applying the decay law to estimate age.

Electron Shells and Energy Levels

Electrons occupy electron shells around the nucleus; the shell a given electron resides in is determined by its energy level.
First shell (closest to nucleus) contains electrons with the lowest energy; second shell contains electrons with higher energy (as energy increases with distance from the nucleus).
The electron cloud concept represents the most probable locations of electrons around the nucleus rather than fixed orbits.

Connections to Foundational Principles and Real-World Relevance

Understanding the origin of chemical symbols helps in reading scientific literature and chemical nomenclature.
Knowledge of major vs trace elements clarifies why certain minerals are essential for life and how elemental balance affects physiology (e.g., ions for nerve impulses, energy transfer in cells).
The periodic table’s organization reflects recurring chemical properties, enabling predictions about element behavior in reactions and compounds.
Atomic structure underpins biochemistry: the way atoms bond and the way electrons are arranged determines molecular structure and function (e.g., water as a molecule of H and O, organic molecules built from carbon frameworks).
Isotopes and radioactive decay provide tools for dating and tracing biological and geological processes; they also illustrate how tiny changes at the subatomic level have large-scale implications.

Ethical, Philosophical, and Practical Implications

The ability to date fossils and determine environmental history using isotopes raises questions about privacy, environmental management, and historical interpretation.
Understanding ionization and electron behavior informs medical and technological applications (e.g., imaging, radiation therapy, materials science).
The simplifications in teaching models (planetary vs orbital) highlight the importance of refining scientific models as evidence and technology advance; this reflects the broader philosophical point that scientific understanding evolves with better models.

Summary of Key Concepts (Quick Reference)

Symbol origins: O from English, Au from Latin "aurum".
Major elements in animal bodies: O (Z=8, ~65%), C (Z=6, ~18.5%), H (Z=1, ~9.5%), N (Z=7, ~3.3%).
Minor/trace elements include Ca, P, K, S, Na, Cl, Mg, Si, Al, Fe, Mn, F, V, Cr, etc., with varying small body-mass contributions.
Atomic nucleus contains protons (positive) and neutrons (neutral); electrons (negative) orbit the nucleus in electron clouds.
Atomic number: Z = ext{number of protons}; Mass number: A = Z + N_n; Neutral atom: Z = E (electrons).
Ions: charged species with q = Z - E.
Isotopes: same Z, different Nn; carbon-14 as a radioactive isotope used for dating with decay rate \lambda and half-life t{1/2} = \frac{\ln 2}{\lambda}.
Models of the atom: planetary (simplified) vs orbital (electron cloud) representations.
Electron shells: energy levels determine shell; first shell has the lowest energy, second shell higher energy.
The periodic table organizes elements by properties; major elements are red, minor blue, trace yellow in the provided figure.