Atomic Structure and Periodic Trends

Atomic Structure and Models

Early Ideas About the Atom

• Ancient Greeks (Democritus) proposed the idea of an ultimate indivisible particle called "atomos".

Early Quantitative Ideas

• Law of Conservation of Mass: Mass is neither created nor destroyed in a chemical reaction.
• Law of Definite Proportions: Pure substances always contain the same proportions of elements by mass.

Dalton’s Theory

• Matter is composed of indestructible atoms.
• Atoms of one element cannot be converted into other elements.
• Atoms of a given element are identical in mass and properties.
• Atoms combine in simple whole number ratios.
• Dalton's theory incorporated the law of conservation of mass and the law of definite proportions.

Avogadro's Contribution

• Suggested elements could exist as molecules (e.g., $O_2$) rather than single atoms.

Limitations of Dalton’s Theory

• Could not explain why elements combine in specific ratios.
• Did not address the electrical nature of particles.

Discovery of the Electron

• Humphry Davy proposed an electromotive force holding elements together.
• Michael Faraday: The amount of substance produced during electrolysis is proportional to the quantity of electricity.
• Sir William Crookes: Cathode rays are negatively charged particles.

Thomson's Model

• Thomson determined the charge-to-mass ratio ($e/m$) of electrons: $1.8 × 10^{11} C kg^{−1}$.
• Proposed the "plum pudding" model: electrons embedded in a sphere of positive charge.

Rutherford and the Nuclear Atom

• Discovered alpha (α), beta (β), and gamma (γ) radiation.

Gold Foil Experiment

• Most alpha particles passed through undeflected.
• A few particles were deflected slightly.
• Very few particles were deflected backwards.

Conclusions

• Atom is mostly empty space.
• Positive charge concentrated in a tiny, dense nucleus.
• Electrons revolve around the nucleus.
• The number of positive charges in the nucleus is the atomic number ($Z$).

Sub-Atomic Particles

• Proton: Positive charge, located in the nucleus.
• Neutron: Neutral charge, located in the nucleus.

Atomic Number and Mass Number

• Atomic Number ($Z$): Number of protons in the nucleus.
• Mass Number: Total number of protons and neutrons in the nucleus.
• Number of neutrons = Mass number ($A$) − Atomic number ($Z$).

Isotopes

• Atoms of the same element with different numbers of neutrons.
• Same atomic number ($Z$) but different mass numbers ($A$).

Relative Atomic Mass and Isotopic Mass

• Relative Atomic Mass (RAM): Mass of an atom relative to $^{1/12}$ the mass of carbon-12.
• Relative Isotopic Mass: Mass of a specific isotope relative to carbon-12.

The Bohr Model

• Electrons revolve around the nucleus in circular orbits.
• Electrons in specific orbits have specific energies.
• Only certain orbits are allowed.
• $ΔE = (E<em>{higher} − E</em>{lower}) = hν$ (Bohr frequency condition).
• $E=hv$ (Energy of a photon).

Successes of the Bohr Model

• Accurately calculated wavelengths of spectral lines for hydrogen.
• Predicted the radius of the $n = 1$ orbit of hydrogen, which was $5.3 × 10^{−11} m$.
• Predicted the ionization energy of hydrogen, which was $2.18 × 10^{−18} J$.

Limitations of the Bohr Model

• Did not explain why electrons absorb or emit radiation when moving between energy levels.
• Did not accurately predict spectra of multi-electron atoms.
• Assumed electrons occupy specific orbits at specific distances.

Quantum Atom

Matter and Waves

• Louis de Broglie: Material particles have wave properties: $λ = h/p$.

Uncertainty Principle

• Werner Heisenberg: It's impossible to know both position and momentum of a particle accurately.

Pauli Exclusion Principle

• No two electrons in the same atom can have the same four quantum numbers.

Aufbau Principle

• Fill the energy levels in order when building up electrons around an atom.

Developing the Periodic Table

Dobereiner’s Triads

• Grouped elements with similar properties in triads.

Mendeleev's Periodic Table

• Arranged elements in order of increasing atomic weights.

Modern Periodic Table

• Elements are arranged in order of increasing atomic number.
• Horizontal rows are periods, vertical columns are groups.
• Elements in the same group have similar electronic configurations. ### Blocks
  • s-block: Groups I and II (1 and 2), configuration $ns^1$ and $ns^2$ respectively
  • p-block: Groups III-VIII (13-18), electronic configurations ranging from $ns^2 np^1$ to $ns^2 np^6$
  • d-block: Groups IIIB–IIB (3–12) configuration $(n−1)d^1 ns^2$ to $(n−1)d^{10} ns^2$
  • f-block: configurations $(n−2)f^1 (n−1)d^{10} ns^2$ to $(n−2)f^{14} (n−1)d^{10} ns^2$

Periodic Properties

• Atomic Radius: Generally increases down a group and decreases across a period.
• Ionic Radius: Positive ions (cations) are smaller than neutral atoms, negative ions (anions) are larger.

Ionization Energy

• Energy required to remove an electron from a gaseous atom: $A(g) → A^+(g) + e^−$
• Increases across a period, decreases down a group.

Successive Ionization Energies

• Removal of successive electrons requires more energy
• The greatest difference occurs after removal of valence layer electrons.

Electron Affinity

• Energy change when an electron is added to a neutral atom: $X + e^− → X^−$
• Increases across a period and decreases down a group.

Electronegativity

• The tendency of an atom to pull electrons in a bond towards itself.
• Electronegativity increases across a period and decreases down a group.

Summary of periodic trends

• Atomic radius decreases across a period and increases down a group (with some exceptions)
• The first ionization energy increases across a period and decreases down a group (with some exceptions)