JG

2.1 Elements and Atoms: The Building Blocks of Matter

Matter and Weight

  • The substance of the universe—from a grain of sand to a star—is matter.
  • Matter is defined as anything that occupies space and has mass.
  • Mass vs. weight:
    • Mass is the amount of matter contained in an object and is constant regardless of location (Earth, Moon, space).
    • Weight is the mass under the influence of gravity; it varies with gravity.
    • Example: An object weighs less on the Moon than on Earth because lunar gravity is weaker; a piece of cheese would weigh about a pound on Earth but only a few ounces on the Moon.
  • Relationship summary: Mass is intrinsic; weight is gravity-dependent.

Elements and Compounds

  • All matter is composed of 92 fundamental substances called elements.
  • An element is a pure substance that cannot be created or broken down by ordinary chemical means.
  • Elements versus compounds:
    • Elements cannot be broken down into simpler substances by chemical means.
    • Compounds are substances composed of two or more elements joined by chemical bonds and always in the same relative amounts.
  • Calcium (Ca) as an example element essential to the human body; calcium is obtained from environment (via diet) and comes in elemental form in foods like cheese.
  • In the human body, the most abundant elements are arranged as: Oxygen (O), Carbon (C), Hydrogen (H), Nitrogen (N) (Figure 2.2).
  • All elements in the body are derived from the foods we eat and the air we breathe.
  • Glucose as a representative compound:
    • Always composed of the same three elements: carbon, hydrogen, and oxygen, in fixed ratios.
    • In glucose, there are always 6 carbon, 6 oxygen, and 12 hydrogen units:
    • Formula: ext{C}6 ext{H}{12} ext{O}_6
  • Elements in nature rarely occur alone; they combine to form compounds.
  • Atoms are the smallest unit of an element that retains its properties; examples show atoms are extremely small (the period at the end of the sentence is millions of atoms wide).

Atoms and Subatomic Particles

  • An atom is the smallest quantity of an element that retains the element’s unique properties.
  • Subatomic particles:
    • Protons (p): positively charged, located in the nucleus.
    • Neutrons (n): electrically neutral, located in the nucleus.
    • Electrons (e): negatively charged, orbit the nucleus at high speeds (near the speed of light) in regions around the nucleus.
  • Charge and mass:
    • Protons and neutrons contribute to atomic mass; electrons contribute negligibly to mass (electrons have ~1/2000 the mass of a proton or neutron): me \approx \frac{mp}{2000}
  • Nucleus and electron cloud:
    • The nucleus contains protons and neutrons; electrons are attracted to the positively charged nucleus.
    • This attraction gives the atom structural stability; electrons are held near the nucleus by this attraction.
  • Models of atomic structure:
    • Planetary model (fixed orbits) is a simplification.
    • Electron cloud model shows electrons occupying regions of space at various distances from the nucleus over time.

Atomic Number and Mass Number

  • Protons determine the identity of the element; the atomic number (Z) is the number of protons in the nucleus.
  • In a neutral atom, the number of electrons equals the number of protons; hence, Z also identifies the typical number of electrons.
  • Neutrons and protons can vary between isotopes of the same element.
  • Mass number (A) is the sum of protons and neutrons in the nucleus:
    • A = Z + N where N is the number of neutrons.
  • Examples:
    • Carbon (C): Z = 6 (six protons). The most common carbon has N = 6 neutrons, so A = 12 (C-12).
    • Uranium (U): Z = 92, A = 238 (U-238); it has 146 neutrons (N = 146).
  • Isotopes share the same Z (same element) but have different N (neutron count).
  • The periodic table reflects elements organized by Z; it provides the chemical symbol, atomic number, and mass number, and groups elements by similar chemical behavior.

Isotopes

  • Isotope: one of the different forms of an element distinguished by a different number of neutrons while keeping the same number of protons.
  • Hydrogen isotopes (Figure 2.5):
    • Protium (^{1}H): 1 proton, 0 neutrons (most abundant hydrogen).
    • Deuterium (^{2}H or D): 1 proton, 1 neutron.
    • Tritium (^{3}H): 1 proton, 2 neutrons.
  • C isotopes: Commonly discussed isotopes include C-12 (6p, 6n), C-13 (6p, 7n), and C-14 (6p, 8n).
  • Heavy isotopes: Contain more neutrons than the most common form; tend to be unstable and radioactive.
  • Radioactive isotopes (radioisotopes): Nuclei that readily decay, emitting subatomic particles and electromagnetic energy.
  • Half-life: Time required for half of any size sample of an isotope to decay; e.g., Tritium has a half-life of about 12 years: T_{1/2} \approx 12 \text{years}.
  • Uses and cautions: Radioisotopes can be harmful with excessive exposure but have important medical applications when used carefully and controlled (e.g., diagnostic imaging, treatment).

Electron Shells and Stability

  • Atoms are not independent; they interact with other atoms to form or break down substances.
  • Electron shells: Regions around the nucleus where electrons reside at distinct energy levels.
  • Electron shell capacity:
    • The first shell can hold up to 2 electrons.
    • Each subsequent shell can hold up to 8 electrons.
    • Atoms can have from 1 to 5 electron shells depending on the number of electrons.
  • Examples of electron configurations by shell count:
    • Hydrogen (Z = 1): 1 electron (first shell not full).
    • Helium (Z = 2): 2 electrons fill the first shell.
    • Lithium (Z = 3): 2 electrons fill the first shell; 1 electron in the second shell.
    • Carbon (Z = 6): 2 electrons fill the first shell; 4 electrons in the second shell (second shell not full).
    • Neon (Z = 10): fills both first and second shells (2 in the first, 8 in the second).
  • The number of electrons in the outermost shell is the valence shell; the electrons in this shell are the valence electrons.
  • Stability and reactivity:
    • An atom is most stable when its valence shell is full.
    • If the valence shell is not full, the atom is reactive and tends to gain, lose, or share electrons to achieve a full valence shell.
  • Octet rule (outermost shell stability):
    • Most atoms are most stable with 8 electrons in their valence shell (except hydrogen and helium, which are stable with 2 electrons in their single shell).
    • Example: Oxygen has 6 electrons in its valence shell and commonly gains 2 electrons to reach 8 (O with 8 electrons in valence shell).
  • Bonding examples:
    • Water formation: two hydrogen atoms share electrons with oxygen to form H–O–H, completing valence shells.
    • Methane formation: carbon tends to form four covalent bonds with hydrogen to satisfy its valence requirements (CH4).
  • Terminology:
    • The term hydrogen reflects its contribution to water: hydro- = water; -gen = maker (Hydrogen = water maker).
  • Periodic table placement and shells:
    • Hydrogen and Helium reside at the sides of the top row because they have only one electron shell.
    • Elements larger than He require a second shell; Li (Z = 3) has two shells: 2 electrons in the first shell, 1 in the second.
    • The second shell can hold up to 8 electrons.
    • Carbon with 6 electrons fills its first shell (2) and partially fills the second shell (4).
    • Neon (10 electrons) fills both first and second shells.
    • Elements in the second row (Li to Ne) have two electron shells.
  • Valence electrons and periodic table groups:
    • Elements in a single column (group) have the same number of valence electrons, which participate in chemical reactions.

Periodic Table and Chemical Relationships

  • The periodic table is arranged by atomic number and shows:
    • Chemical symbol for each element.
    • Atomic number (Z).
    • Mass number (A).
    • Grouping by propensity to react with other elements.
  • The table captures the fact that the number of protons equals the number of electrons in neutral atoms; the total number of protons and electrons equals the atomic number and determines chemical behavior.
  • 92 naturally occurring elements; additional larger, unstable elements have been discovered experimentally.

Examples and Everyday Relevance

  • Glucose as an example of a compound: Always composed of C, H, and O in fixed ratios (C6H12O6).
  • Water (H2O) and methane (CH4) as canonical examples of molecules formed by sharing electrons to satisfy octet rule.
  • The arrangement of elements in the body (O, C, H, N) and how they derive from diet and air underpins physiology and biochemistry.

Medical Applications of Isotopes (Career and Imaging Context)

  • Radioisotopes in medicine:
    • Radioisotopes emit subatomic particles and electromagnetic energy that can be detected by imaging techniques.
    • A key application is PET scanning using radioactive glucose to reveal metabolically active tissues: cancerous tissues often show high glucose uptake and appear as bright spots on PET images (Figure 2.6).
  • Interventional radiology:
    • Interventional radiologists treat disease using minimally invasive techniques involving radiation.
    • They use radioisotopes to diagnose and treat conditions, including tumors that are difficult to remove surgically.
    • Radioembolization: a procedure where radioactive seeds are delivered via needle into blood vessels feeding a tumor (e.g., liver tumors) to disrupt blood supply and kill nearby tumor cells.
  • Practical and ethical considerations:
    • Radioisotope exposure carries health risks; benefits must be weighed against risks.
    • Controlled use in medical settings can reduce the need for invasive surgery and chemotherapy in select cases, lowering cost and recovery time.

Connections and Key Takeaways

  • Matter consists of elements and compounds; understanding atomic structure explains chemical behavior.
  • Atomic number (Z) identifies the element; mass number (A) and neutron count (N) define isotopes.
  • Electron configuration, especially valence electrons, determines reactivity and bonding patterns (octet rule).
  • Isotopes differ in neutron number; some are stable, others radioactive with characteristic half-lives.
  • Real-world relevance includes nutrition, physiology, and modern medical imaging and therapy using radioisotopes.
  • Always distinguish mass (intrinsic) from weight (gravity-dependent).

Important Equations and Numbers (LaTeX)

  • Mass number: A = Z + N
  • Electron mass approximation: me \approx \frac{mp}{2000}
  • Octet rule (conceptual): valence shell aims to contain 8 electrons (except H and He, which strive for 2 electrons in their single shell)
  • Half-life concept (example): T_{1/2} \approx 12\text{ years} \quad (\text{for Tritium})
  • Common molecular formulas:
    • Water: \text{H}_2\text{O}
    • Methane: \text{CH}_4
  • Glucose formula: \text{C}6\text{H}{12}\text{O}_6
  • Isotope notation examples: Protium ≈ ^1H, Deuterium ≈ ^2H, Tritium ≈ ^3H; Carbon isotopes as ^{12}C, ^{13}C, ^{14}C

References to Figures (from OpenStax pages)

  • Figure 2.2: Elements in the human body from most abundant to least: O, C, H, N (plus others such as Na, Fe).
  • Figure 2.3: Two models of atomic structure for He and C (planetary vs electron cloud).
  • Figure 2.4: The Periodic Table of the Elements—showing symbols, atomic numbers, and mass numbers.
  • Figure 2.5: Isotopes of Hydrogen (Protium, Deuterium, Tritium).
  • Figure 2.6: PET Scan showing metabolic activity via glucose uptake.
  • Figure 2.7: Electron shells illustrating hydrogen, carbon, and neon configurations.