Atomic Structure and Ions

Atomic Structure

• Atoms of different elements contain varying quantities of subatomic particles
• The nucleus is composed of neutrons and protons
• Proton number is equivalent to the quantity of protons
• Nucleon number is the sum of protons and neutrons

Identifying Elements

• Each element has a unique proton number
• The proton number serves as an element's identifier
  • Example: Oxygen has a proton number of 8
  • Example: Sodium has a proton number of 11
• In neutral atoms, the proton number matches the number of electrons
• Number of neutrons can be found by subtracting the proton number from the nucleon number
• Nucleon number is also known as mass number
• Relative atomic mass is approximately equal to its nucleon number
  • Example: Chlorine has 17 protons and 18 neutrons
  • Therefore:
    • Proton number = 17
    • Nucleon number = 35 (17 + 18)

Subatomic Particles in Ions

• Ions arise when atoms gain or lose electrons
• Atoms losing electrons form positive ions
• Atoms gaining electrons form negative ions
• During ion formation, the count of protons and neutrons remains constant
• Formation of positive ions leads to a decrease in electron count
• Formation of negative ions increases the electron count
• Metal atoms tend to lose electrons, resulting in positive ions
  • Example: Sodium (proton number 11, nucleon number 23)
    • Sodium atom: 11 protons, 12 neutrons (23 - 11), 11 electrons
    • Sodium ion: 11 protons, 12 neutrons, 10 electrons (lost 1)
• Non-metal atoms gain electrons to form negative ions
  • Example: Oxygen (proton number 8, nucleon number 16)
    • Oxygen atom: 8 protons, 8 neutrons (16 - 8), 8 electrons
    • Oxide ion: 8 protons, 8 neutrons, 10 electrons (gained 2)

Standard Notations of Elements

• Each element is identified by a name and a symbol
• Mass number (nucleon number) = Z + N (number of neutrons)
• Atomic number = number of protons
• Chemical symbol denotes the element

Elements and Symbols

• Each symbol contains one or two letters
• Two-letter symbols capitalize the first letter and lowercase the second
• Some elements' symbols come from their Latin names
  • Sodium (Na, from "natrium")
  • Potassium (K, from "kalium")

Atomic Structure and Electron Arrangement

• Electrons arrange in shells, starting nearest the nucleus
• Each shell holds a specific maximum number of electrons
• Max electrons for atoms with proton numbers 1-20:
  • 1st shell (n=1): 2 electrons
  • 2nd shell (n=2): 8 electrons
  • 3rd shell (n=3): 8 electrons
• New shells fill only when the previous shell is full

Writing Electron Arrangements

• Step 1: Find the atom's proton number
• Step 2: Find the atom's number of electrons (same as proton number in neutral atoms)
• Step 3: Arrange electrons in shells, starting with n=1, filling each to capacity before moving to the next
  • Example: Sulfur (proton number 16)
    • 16 electrons
    • 1st shell (n=1): 2 electrons
    • 2nd shell (n=2): 8 electrons
    • 3rd shell (n=3): 6 electrons
    • Electron arrangement: 2.8.6

Valence Electrons

• Valence electrons are in the outermost occupied electron shell (valence shell)
• Number of valence electrons comes from the electron arrangement
  • Example: Magnesium (proton number 12, arrangement 2.8.2)
  • Number of valence electrons = 2

Isotopes

• Isotopes are atoms of an element with the same proton number but different nucleon numbers
  • Example: Hydrogen isotopes (protium, deuterium, tritium)
  • All have proton number 1, but different nucleon numbers
• Isotopes can be written in several forms
  • Uranium-235 (U-235)
  • Uranium-238 (U-238)

Differences in Isotopes

• Different nucleon numbers
• Different number of neutrons
• Different physical properties (due to neutron count)

Similarities in Isotopes

• Same number of protons
• Same number of electrons
• Same chemical properties (due to electron count)

Calculating Relative Atomic Masses for Elements with Isotopes

• Relative atomic mass is the average mass of isotopes
• Step 1: Identify all known isotopes
• Step 2: Identify the abundance of each isotope
• Step 3: Calculate relative atomic mass:
  • Relative atomic mass = [(% Isotope 1 × Isotopic mass of Isotope 1) + (% Isotope 2 × Isotopic mass of Isotope 2) + (% Isotope 3 × Isotopic mass of Isotope 3)] / 100
  • Example: Chlorine has two isotopes:
    • Chlorine-35 (75% abundance)
    • Chlorine-37 (25% abundance)
  • Relative atomic mass of chlorine = [(75 × 35) + (25 × 37)] / 100 = 35.5

Uses of Isotopes

• Radioactive isotopes decay and emit radiation

Agriculture

• Phosphorus-32 tracks phosphorus uptake in plants

Medicine

• Cobalt-60 (gamma rays) treats cancer (chemotherapy)
• Iodine-131 treats thyroid disease

Power Source

• Uranium-235 is a common nuclear power station fuel

Archeology

• Carbon-14 estimates age of bones, wood, fossils

Industry

• Sodium-24 detects leaks in gas/oil pipes and ventilation
• Cobalt-60 gamma radiation preserves food by killing bacteria