2.1 Elements and Atoms: The Building Blocks of Matter

Elements and Atoms: The Building Blocks of Matter

Everything in anatomy and physiology comes back to chemistry, and chemistry starts here: with elements and atoms. The structure of atoms determines how they bond, which determines the molecules your body can build, which determines how your cells function. This section covers atomic structure, isotopes, electron configurations, and the basics of chemical bonding.

Elements and Atoms

Relationships in Atomic Structure

Matter is anything that has mass and occupies space. All matter is composed of elements, which are fundamental substances that cannot be broken down into simpler substances by ordinary chemical means. About 26 elements are found in the human body, though just four of them (oxygen, carbon, hydrogen, and nitrogen) make up roughly 96% of your body mass.

Each element is made up of a single type of atom. An atom is the smallest unit of an element that still retains that element's chemical properties.

Atoms are composed of three subatomic particles:

• Protons: positively charged, located in the nucleus
• Neutrons: no charge, located in the nucleus
• Electrons: negatively charged, orbit the nucleus in energy shells

When two or more elements chemically combine in a specific ratio, they form a compound. Compounds have properties that differ from those of their component elements. For example, hydrogen is a flammable gas and oxygen supports combustion, but together as water ($H_2O$) they form a liquid that extinguishes fire. Other biologically important compounds include carbon dioxide ($CO_2$) and glucose ($C_6H_{12}O_6$).

Atomic Number vs. Mass Number

• Atomic number (Z) = the number of protons in an atom's nucleus. This number is unique to each element and defines its identity. Carbon always has 6 protons, so its atomic number is 6. Change the number of protons and you change the element.
• Mass number (A) = the total number of protons + neutrons in the nucleus. Carbon-12 has a mass number of 12 (6 protons + 6 neutrons).
• Atomic mass is the weighted average mass of all naturally occurring isotopes of an element. This is the decimal number you see on the periodic table (carbon's atomic mass is approximately 12.01, not a whole number, because it accounts for the relative abundance of each isotope).

The atomic number tells you what element you're looking at. The mass number tells you which version (isotope) of that element you have.

Isotopes and Elemental Properties

Isotopes are atoms of the same element that have different numbers of neutrons. They share the same atomic number but have different mass numbers. Carbon, for example, has three naturally occurring isotopes: carbon-12, carbon-13, and carbon-14. All three have 6 protons, but they have 6, 7, and 8 neutrons respectively.

How isotopes affect properties:

• Chemical properties stay essentially the same across isotopes, because chemical behavior depends on electrons, not neutrons.
• Physical properties like mass differ slightly between isotopes.
• Some isotopes are radioactive, meaning their nuclei are unstable and undergo radioactive decay, emitting particles and energy. Carbon-14 is a radioactive isotope of carbon.

Radioactive isotopes (radioisotopes) have direct medical applications. They're used in PET scans and other medical imaging, in radiation therapy for cancer treatment, and in radiometric dating techniques.

Electron Shells and Atomic Behavior

Electrons don't just float randomly around the nucleus. They occupy specific energy levels called electron shells, labeled K, L, M, N (or 1, 2, 3, 4), with K being closest to the nucleus.

Each shell holds a maximum number of electrons:

For this course, the most important concept here is valence electrons, the electrons in the outermost shell. Valence electrons determine an atom's chemical behavior and reactivity.

Stability and reactivity:

Atoms are most stable when their outermost shell is full. This is sometimes called the octet rule because most atoms need 8 electrons in their valence shell to be stable (the exception is the K shell, which is full with just 2). Noble gases like helium (He), neon (Ne), and argon (Ar) already have full valence shells, which is why they rarely react with anything.

Atoms with partially filled valence shells are reactive. They will gain, lose, or share electrons to reach a stable configuration. For instance, sodium has 1 valence electron and tends to lose it, while chlorine has 7 valence electrons and tends to gain one.

Types of chemical bonds:

• Ionic bonds form when one atom transfers electrons to another. This typically happens between metals (which lose electrons) and nonmetals (which gain them). Sodium chloride ($NaCl$) is a classic example: sodium gives up its one valence electron to chlorine, and both atoms end up with full outer shells.
• Covalent bonds form when atoms share electrons. This typically happens between two nonmetals. In water ($H_2O$), oxygen shares electrons with two hydrogen atoms. Other examples include methane ($CH_4$) and molecular oxygen ($O_2$).

Atomic Structure and Periodic Trends

The periodic table organizes elements by increasing atomic number, and its layout reflects patterns in atomic structure. Elements in the same column (group) have the same number of valence electrons, which is why they behave similarly in chemical reactions.

Two periodic trends worth knowing:

• Electronegativity is a measure of how strongly an atom attracts electrons in a chemical bond. Oxygen and nitrogen are highly electronegative, which matters a lot for the polar bonds found in biological molecules like water.
• Atomic radius is the distance from the nucleus to the outermost electron shell. Atomic radius generally increases as you move down a group (more shells) and decreases as you move across a period from left to right (stronger nuclear pull on electrons).

Within electron shells, electrons occupy specific orbitals (s, p, d, f), and the arrangement of electrons across these orbitals is called an atom's electron configuration. For this course, the key takeaway is that electron configuration determines how an atom will bond with others, which in turn determines the structure and function of every molecule in your body.