Atom Structure, Isotopes, and the Mole
Atomic Constituents: Protons, Neutrons, and Electrons
Atoms are composed of three main subatomic particles: protons, neutrons, and electrons.
Mass/weight approximations (as used in this course):
Proton: weight ~ 1 atomic mass unit (amu); charge +1.
Neutron: weight ~ 1 amu; charge 0.
Electron: weight ≈ 0 amu (often approximated as 0 in simple models); charge −1.
Notation of weights and charges (simplified):
mp \approx 1\ \text{u},\quad qp = +1
mn \approx 1\ \text{u},\quad qn = 0
me \approx 0\ \text{u},\quad qe = -1
Practical simplification: we round weights to whole numbers and focus on relative contributions rather than exact fractions.
In simple atoms, the nucleus contains protons and neutrons; electrons orbit around the nucleus in a region called the electron cloud.
The Nucleus and Electron Cloud
Nucleus: central region containing protons (positive charge) and neutrons (neutral).
Electron cloud: electrons orbit rapidly around the nucleus and form a dynamic region of negative charge.
The electrons move so fast that their positions form a surrounding region rather than a fixed orbit; sometimes described with the idea that rapid spinning creates a barrier-like effect, giving the appearance of a solid plane for interaction purposes (an analogy used in the lecture).
The nucleus is relatively stable; electrons are organized around it in shells/levels.
Hydrogen atom (simplest atom) usually has:
1 proton in the nucleus and typically 1 electron surrounding it.
Simplified view: one proton, one electron; optional neutron(s) may be present in some isotopes.
Atomic Number, Mass Number, and Isotopes
Elements are defined by the number of protons in the nucleus, which is the atomic number Z.
In a neutral atom, the number of electrons equals the atomic number (protons balance electrons).
Hydrogen example: Z = 1, so 1 proton and typically 1 electron in neutral hydrogen.
When protons are added, electrons are added to balance the charge, creating a new element.
Example progression:
Helium: Z = 2 (2 protons, 2 electrons)
Lithium: Z = 3 (3 protons, 3 electrons)
Neon: Z = 10 (2 electrons in the inner shell, 8 in the second shell in the simplified model)
The mass number A is the total number of protons and neutrons: A = Z + N,
where N\approx\text{number of neutrons} and each typically weighs ~1 amu.Isotopes: atoms with the same Z (same element) but different numbers of neutrons N. Examples with hydrogen:
Hydrogen-1 (protium): 1 proton, 0 neutrons, A = 1
Hydrogen-2 (deuterium): 1 proton, 1 neutron, A = 2
Hydrogen-3 (tritium): 1 proton, 2 neutrons, A = 3
Radioisotopes: isotopes that emit radiation as they decay (e.g., tritium, carbon-14, iodine-125). Radiation can damage molecules like DNA but has clinical uses (e.g., imaging).
Iodine-125 imaging of the thyroid: radioactive iodine is taken up by the thyroid, allowing imaging of thyroid tissue for tumors or pathology.
Carbon-14 dating uses the radioisotope carbon-14 to estimate age.
Chlorine example: Chlorine (Cl) has Z = 17 and electrons = 17 in a neutral atom; neutrons vary by isotope (e.g., Cl-35 with N = 18, Cl-37 with N = 20), which is why the atomic weight is not an integer.
Atomic Weight, Isotopes, and the Periodic Table
Atomic number Z equals the number of protons and, in a neutral atom, the number of electrons.
The mass number A = Z + N is an integer that varies with isotopes.
Atomic weight (or atomic mass) on the periodic table is a weighted average of the isotopic masses, not a whole number due to the mixture of isotopes in nature.
Example intuition: natural hydrogen is mostly \text{H-1} but includes some \text{H-2} and a small amount of \text{H-3}, which raises the average atomic weight slightly above 1.0.
The Dalton (atomic mass unit, amu) serves as the unit for these weights; 1 Dalton ≈ the mass of a proton or neutron.
The atomic weight on the periodic table is effectively the average mass per atom in atomic mass units, then numerically equal to the same value in grams per mole (g/mol) when using the mole concept.
Example: for hydrogen, the atomic weight is about \approx 1.01\ \text{g/mol}; for carbon, about 12.01\ \text{g/mol}; the idea is that 1 mole (Avogadro’s number) of atoms weighs the atomic weight in grams.
The mole is a fundamental counting unit in chemistry.
The Mole and Avogadro’s Number
A mole is a specific count: N_A = 6.02\times 10^{23} entities (atoms, molecules, etc.).
One mole of hydrogen atoms weighs approximately 1.01\ \text{g}.
One mole of carbon atoms weighs approximately 12.01\ \text{g}.
In general:
If you could count out exactly N_A atoms of an element and place them on a scale, their total mass would equal the element’s atomic weight in grams.
This underpins the link between atomic weight (in amu) and molar mass (in g/mol).
Electron Energy Levels and Shell Capacities
Electrons occupy energy shells (or energy levels) around the nucleus.
Shell capacities (simplified, as described in the lecture):
First shell (innermost) holds 2 electrons.
Second shell holds 8 electrons.
Third shell holds 8 electrons (per the simplified course model).
Fourth shell can hold 18 electrons (not deeply discussed in this course).
Filling order (simplified):
Hydrogen: 1 electron goes into the first shell.
Helium: 2 electrons fill the first shell.
Lithium: 3 electrons fill the first shell (2) and begin filling the second shell (1 of 8).
Neon: 10 electrons total (2 in the first shell, 8 in the second shell).
The capacity and arrangement explain how adding protons (and thus electrons) moves you to new elements and changes the electron configuration.
The first electron shell is closest to the nucleus due to electrostatic attraction; outer shells are farther away and require more energy to maintain electrons at that distance.
Practical Examples and Common Elements in Human Tissue
The periodic table organizes elements by increasing atomic number; the chemical symbol is shown in the table (e.g., H for hydrogen).
Common elements relevant to human tissue include:
Hydrogen (H), Oxygen (O), Carbon (C), Nitrogen (N)
Calcium (Ca), Potassium (K), Sodium (Na)
Chlorine (Cl), Iron (Fe), Iodine (I)
In isotopes, the neutron number can vary, leading to different mass numbers for the same element.
Example: Boron example in the lecture shows a nucleus with 5 protons and a mixture of neutrons, causing the atomic weight to be around 10–11, illustrating why the weight is not a fixed integer.
Imaging and Radioisotopes: Practical Implications
Radioisotopes emit radiation as they decay; this is the molecularly damaging aspect but also provides clinical utility.
Nuclear imaging uses radioisotopes to visualize tissues and function; e.g., radioactive iodine uptake by the thyroid for imaging thyroid health.
Older imaging techniques (e.g., nuclear medicine imaging with radioisotopes) coexisted with ultrasound; some studies still rely on radioisotopes for certain imaging tasks.
Notable isotopes mentioned:
Tritium (H-3): a radioisotope of hydrogen that decays and emits radiation.
Carbon-14 (C-14): used in radiometric dating.
Iodine-125 (I-125): used to image thyroid tissue due to the thyroid’s uptake of iodine.
Important caution: radioisotopes can be harmful due to radiation exposure; their use requires justified clinical benefits and safety measures.
Quick Concept Check: Formulas and Key Relationships
Atomic number and electron count in neutral atoms:
Z =#\text{protons} = #\text{electrons (neutral atom)}.
Mass number (nucleon count):
A = Z + N, where N = \text{neutrons}.
Isotopes: same Z, different N; different mass numbers A.
Atomic weight vs. molar mass:
Atomic weight (amu) is the weighted average of isotopic masses.
The molar mass in g/mol equals the atomic weight in amu (numerical value).
Avogadro’s number (mole definition):
N_A = 6.02\times 10^{23}.
For a mole of atoms, the mass in grams equals the atomic weight in amu:
e.g., 1\text{ mole of H atoms} \approx 1.01\ \text{g}.
e.g., 1\text{ mole of C atoms} \approx 12.01\ \text{g}.