HL

Atomic Structure and Periodic Table Fundamentals

Subatomic Particles: masses, charges, and relative scales

  • Subatomic particles to know: protons, neutrons, electrons.
  • Relative masses:
    • Protons and neutrons have similar masses; conventionally both are treated as 1 atomic mass unit (amu).
    • Electron mass is much smaller: about m_e \approx 0.0005\ \text{amu}, i.e. roughly 1/1836 of a proton mass.
    • In practice, electrons contribute negligible mass to the overall atomic mass compared to protons and neutrons.
  • Electric charges:
    • Proton: +1\,e
    • Electron: -1\,e
    • Neutron: neutral (0 charge)
  • Summary about mass vs charge:
    • Protons and neutrons determine mass; electrons determine charge distribution.
    • Atoms are overall electrically neutral when the number of protons equals the number of electrons.

Atomic identity vs mass: what determines the element

  • The identity (which element it is) is determined by the number of protons in the nucleus.
  • This number is called the atomic number, denoted by Z.
  • You cannot change the number of protons without changing the element (e.g., 6 protons defines carbon; 7 defines nitrogen).
  • The nucleus contains protons (positive charge) and neutrons (neutral mass).
  • Electrons move around the nucleus and provide the charge balance.
  • The nucleus is the center of mass for mass; electrons orbit the nucleus.

Dalton’s Atomic Theory and the historical progression

  • Dalton’s contributions (early view of atoms):
    • Everything is made of atoms.
    • Atoms are indestructible (indivisible in Dalton’s sense).
    • Atoms of one element are identical in mass and properties; atoms of different elements differ.
    • Atoms combine in simple whole-number ratios to form compounds (e.g., H2O, not H2O0.5).
  • Dalton’s view used the metaphor of marble-like indivisible spheres; no knowledge of subatomic structure.
  • Later experiments revealed more structure: electrons, protons, neutrons, nucleus, etc.

Thomson’s plum pudding model and the electron

  • JJ Thomson (cathode ray tube) discovered the electron and its charge-to-mass ratio.
  • He proposed a model where electrons are embedded in a positively charged ‘pudding’ (the atom as a uniform sphere with embedded electrons).
  • This model could not explain all observations and was later superseded by Rutherford’s nuclear model.

Millikan’s oil drop experiment: electron charge

  • Robert Millikan measured the elementary charge e by balancing gravitational and electric forces on charged oil drops.
  • From this, and Thomson’s charge-to-mass ratio experiments, the mass of the electron was estimated along with its charge.
  • Result: precise measurement of the electron’s charge: e = 1.602 imes10^{-19}\ \text{C} (historic value; used to determine electron mass via charge-to-mass ratio).

Rutherford: the nuclear (central) atom

  • Rutherford validated a new atomic model: the nuclear atom.
  • Key idea: most of the atom is empty space; a dense, positively charged nucleus sits at the center.
  • Protons reside in the nucleus; electrons move around the nucleus in largely empty space.
  • Early mass discrepancy led to the realization that there was more to the nucleus than protons alone; the missing mass was later attributed to neutrons (Chadwick).

Chadwick: the neutron

  • James Chadwick identified the neutron, a neutral particle in the nucleus, which accounts for the mass not explained by protons alone.
  • Neutrons contribute to the atomic mass but have no charge.

Key relationships and definitions (quick recap)

  • Atomic number: Z = ext{number of protons}
    • Determines the element (identity).
    • Symbol often shown as a subscript or directly on the periodic table as the element’s Z.
  • Mass number: A = Z + N where N is the number of neutrons.
  • Nuclear composition: nucleus contains Z protons and N = A - Z neutrons.
  • Neutral atom: number of electrons equals number of protons, i.e., electron count = Z.
  • Electron mass is negligible compared to protons and neutrons, but electrons are responsible for chemical behavior (bonding, charge).

The periodic table: structure and key terms

  • Periods vs Groups:
    • Periods: horizontal rows; across a period, proton number increases; atoms gain protons and electrons.
    • Groups: vertical columns; elements in a group share similar chemical properties.
  • Four main pieces of information on most periodic tables:
    • Chemical symbol (one or two letters; sometimes derived from Latin).
    • Atomic number Z (number of protons).
    • Element name.
    • Atomic mass (average atomic mass or atomic weight) in amu; sometimes mass number is shown for isotopes but not as the standard table value.
  • Atomic mass vs Mass number:
    • Atomic mass: weighted average of all isotopes’ masses in amu.
    • Mass number A: integer equal to protons + neutrons for a specific isotope, i.e., A = Z + N.

Symbolism and historical naming quirks

  • Symbols often reflect common names or Latin/Greek roots (e.g., Na from natrium, Fe from ferrum, K from kalium, Pb from plumbum, Hg from hydrargyrum).
  • Many elements bear names honoring scientists or places (e.g., Einsteinium Es, Curium Cm, Polonium Po, Americium Am).
  • The first element discovered by a person often influenced its naming; sometimes we use a name reflecting properties (e.g., Bromine from Greek bromos meaning ‘stench’).
  • Some symbols are not intuitive (e.g., Mg for magnesium vs Mn for manganese; Hg for mercury).

Metalloids and the staircase on the periodic table

  • Metalloids (semimetals) sit along the staircase boundary between metals and nonmetals:
    • Common metalloids: boron (B), silicon (Si), germanium (Ge), arsenic (As), antimony (Sb), tellurium (Te), polonium (Po), astatine (At).
  • Metalloids possess properties intermediate between metals and nonmetals; behavior can depend on the reaction conditions.
  • The staircase may be drawn with a bold line in some tables to indicate these elements.

Metals vs Nonmetals: core properties

  • Metals (the vast majority of the table):
    • Excellent conductors of heat and electricity.
    • Ductile (can be drawn into wires) and malleable (can be hammered into sheets).
    • Tend to lose electrons to form positively charged ions (cations).
    • Generally shiny (lustrous).
  • Nonmetals (right side of the table):
    • Poor conductors of heat and electricity (insulators) and highly variable in state (solids, liquids, gases).
    • Not easily malleable or ductile; many are gases at room temperature.
    • Tend to gain electrons to form negatively charged ions (anions) during reactions.
  • Notable exception: carbon is not a metal—despite its central biological role and abundance in life, it behaves chemically like a nonmetal.
  • A classroom reminder: “Hydrogen” is often shown in Group 1 but is not a metal; its placement is more about historical grouping than metal-like behavior.

Practical connections and implications

  • Why the order matters: knowing the order of discovery and experiments helps connect the development of atomic theory to modern understanding.
  • The progression from Dalton to Thomson to Millikan to Rutherford to Chadwick shows how experimental evidence refined the model of the atom.
  • The concept of the nucleus and mass concentration in the nucleus explains why most of the atom is empty space yet has most of the mass concentrated in a tiny region.
  • The periodic table’s structure is not arbitrary: patterns in properties appear in periods and groups, guiding predictions about reactivity and bonding.
  • Atomic number vs atomic mass vs mass number: these terms have distinct meanings but are often conflated in casual use; students must distinguish:
    • Z = ext{number of protons} (identity of the element).
    • A = Z + N (mass number for a given isotope).
    • Atomic mass (weighted average of isotopes) in amu.

Quick formula sheet (LaTeX)

  • Electron mass: m_e \approx 9.109 \times 10^{-31}\ \text{kg} \approx 0.0005486\ \text{amu}
  • Proton/neutron mass: mp \approx mn \approx 1.6726 \times 10^{-27}\ \text{kg} \approx 1\ \text{amu}
  • Mass number: A = Z + N
  • Neutral atom electron count: N_e = Z
  • Atomic mass (weighted): \text{Atomic mass} = \sumi (fi \cdot Ai) where fi is the natural abundance of isotope i and A_i is its mass number.
  • Atomic number definition: Z = \text{number of protons}
  • Nuclear composition: \text{nucleus} = Z\text{ protons} + N\text{ neutrons}, \, N = A - Z

Representative exam-style question (based on the lecture)

  • Question: Which element is expected to have similar properties to magnesium (Mg)?
    • Magnesium is in Group 2 (alkaline earth metals). Elements in the same group share similar reactivity and chemical behavior.
    • Possible correct answers include: Be (beryllium), Ca (calcium), Sr (strontium), Ba (barium), Ra (radium).
    • NOTE: Be, Mg, Ca, Sr, Ba, Ra all have similar trends in reactivity, especially with water (alkaline earth metals), though reactivity increases down the group.

Quick recap of key takeaways

  • The nucleus houses protons (positive charge) and neutrons (neutral mass); electrons orbit in largely empty space.
  • Atomic number Z defines the element; mass number A defines a particular isotope via A = Z + N.
  • The periodic table organizes elements into periods and groups; groups reflect similar chemical properties.
  • Metals vs nonmetals vs metalloids have characteristic properties that drive bonding and material behavior.
  • Historical experiments (Dalton, Thomson, Millikan, Rutherford, Chadwick) progressively unveiled the subatomic world and led to the modern atomic model.
  • Hydrogen’s placement is special and not a perfect fit in Group 1 as a metal; carbon is not a metal despite its central role in life.
  • Always distinguish between atomic mass (weighted average) and mass number (specific isotope).