Elements and Atoms_ The Building Blocks of Matter

Learning Objectives

  • By the end of this section, you will be able to:

    • Discuss the relationships between matter, mass, elements, compounds, atoms, and subatomic particles

    • Distinguish between atomic number and mass number

    • Identify the key distinction between isotopes of the same element

    • Explain how electrons occupy electron shells and their contribution to an atom’s relative stability

Matter and Mass

  • Matter: Anything that occupies space and has mass.

    • Ranges from a grain of sand to a star.

  • Mass vs Weight:

    • Mass: Amount of matter in an object, constant in any environment.

    • Weight: Mass affected by gravity, thus variable.

      • Example: Cheese weighs less on the moon than Earth due to reduced gravity.

Elements and Compounds

  • Elements: Pure substances that cannot be broken down by ordinary chemical means.

    • 92 fundamental elements in nature; examples include calcium (Ca), oxygen (O), sodium (Na), iron (Fe).

  • Compounds: Substances composed of two or more elements joined by chemical bonds.

    • Example: Glucose, composed of carbon, hydrogen, and oxygen in a specific ratio (C6H12O6).

Atoms and Subatomic Particles

  • Atoms: Smallest unit of an element that retains its unique properties.

    • Atoms are incredibly small.

  • Subatomic Particles:

    • Protons: Positively charged; determine the atomic number (element identity).

    • Neutrons: Neutral; contribute to atomic mass.

    • Electrons: Negatively charged; equal in number to protons in neutral atoms, and contribute minimally to mass.

Atomic Structure

  • Models of Atomic Structure:

    • Planetary Model: Electrons in fixed orbits.

    • Electron Cloud Model: Electrons move rapidly in regions of space around the nucleus.

  • Electrical Charges:

    • Protons (p+): Positive charge

    • Electrons (e−): Negative charge

    • Neutrons: No charge

  • Stability: The attraction between protons and electrons creates structural stability.

Atomic and Mass Number

  • Atomic Number: Unique number of protons in an atom, determines the element.

    • Example: Carbon has an atomic number of 6 (6 protons).

  • Mass Number: Sum of protons and neutrons in the nucleus.

    • Example: Common carbon form (12C) has 6 protons and 6 neutrons (mass number 12).

  • Periodic Table: Organized list of elements by atomic number, displaying symbols and mass numbers.

Isotopes

  • Isotope: Varieties of an element differing in neutron count.

    • Example: Carbon has isotopes 12C, 13C, and 14C (with respective neutron counts).

  • Heavy Isotopes: Isotopes with more than usual neutrons, often unstable and radioactive.

  • Radioactive Isotopes: Nuclei decay, emitting particles; half-life indicates decay rate, e.g., tritium (12 years).

Medical Applications of Radioisotopes

  • Career Connection: Interventional Radiologists: Use radioisotopes for minimally invasive treatments.

    • Example: Treating liver tumors with radioembolization, using radioactive seeds to disrupt blood supply.

  • PET Scanning: Using low doses of radioactive glucose to detect active tissues in the body, revealing potential tumors.

Electron Behavior and Chemical Reactions

  • Atoms react to form compounds, dictated by electron behavior in electron shells:

    • Each shell has defined energy levels.

    • First shell holds 2 electrons, subsequent shells hold 8.

    • Example: Hydrogen (1 electron), Helium (2 electrons).

  • Valence Electrons: Electrons in the outermost shell governing reactivity.

    • Full valence shell (8 electrons) = stable atom.

    • Unfull valence shell = reactive atom, seeks to fill through bonds.

  • Octet Rule: Atoms will gain, lose, or share electrons to achieve a filled valence shell.

    • Example: Oxygen (6 valence electrons) seeks 2 more, often bonding with 2 hydrogen atoms to form H2O.