Atomic Structure & Ions Flashcards

Learning Outcomes

• Subatomic Particles: Comprehensive identification and understanding of protons, neutrons, and electrons, including their relative masses, charges, and locations within the atom.
• Atomic Structure & Periodic Table: Rationalizing the arrangement of electrons in energy levels/shells and linking valence shell electron counts to group positions in the periodic table.
• Isotopes & Atomic Mass: Differentiating between radioactive and non-radioactive isotopes, and calculating relative atomic mass ($A_r$) from isotopic abundance data.
• Ion Formation: Explaining the mechanisms of cation and anion formation, their chemical representations, and the governing principles behind trends in ionic radii.

Fundamental Concepts of Atomic Structure

• Definition of an Atom: The atom is the smallest fundamental unit of matter that retains the chemical identity and properties of an element.
• Subatomic Composition: Atoms are comprised of three principal subatomic particles:
  • Protons: Positively charged particles residing within the central nucleus.
  • Neutrons: Electrically neutral particles residing within the central nucleus.
  • Electrons: Positively/negatively charged fundamental particles revolving around the nucleus in specified quantum energy levels (shells).

• Detailed Characterization of Subatomic Particles:
  • Proton ($p$ or $p^+$):
    • Location: Confined strictly within the atomic nucleus.
    • Relative Charge: $+1$
    • Relative Mass: $1$
    • Function: Determines the chemical identity of an element (Proton Number / Atomic Number, $Z$). All atoms of a specific element possess an identical proton count. In a neutral atom, positive nuclear charge from protons is balanced by an equal number of surrounding negative electrons.
  • Neutron ($n$ or $n^0$):
    • Location: Confined strictly within the atomic nucleus (present in all atomic species except simple hydrogen, ${_1^1}\text{H}$).
    • Relative Charge: $0$ (Electrically neutral).
    • Relative Mass: $1$ (Approximately equivalent to the mass of a proton).
    • Function: Functions as the "nuclear glue" that stabilizes the nucleus by counteracting the strong electrostatic repulsion between positively charged protons.
  • Electron ($e^-$):
    • Location: Resides outside the nucleus within discrete quantum shells, subshells, and three-dimensional atomic orbitals.
    • Relative Charge: $-1$
    • Relative Mass: Approximately $\frac{1}{1840}$ (frequently treated as negligible / $0$ in basic mass calculations).
    • Function: Primary determinant of atomic reactivity, chemical bonding, valence behaviors, and molecular geometry.

Historical Evolution of Atomic Theory

• Democritus (460 B.C. – 370 B.C.):

  • First proposed the conceptual existence of fundamental particles, coining the term atomos (meaning indivisible or uncuttable).
  • Hypothesized that matter is composed of indivisible and indestructible particles moving through empty space.
  • Philosophical in origin; lacked empirical validation via the scientific method and failed to explain chemical reactions or atomic interactions.

• Comparison: Ancient Philosophy vs. Modern Scientific View:

  • Discreteness of Matter:
    • Democritus' View: Matter consists of distinct, individual particles rather than a continuous substance.
    • Modern Scientific View: Correct. Matter is discrete and composed of distinct atoms, ions, and molecules.
  • Indivisibility:
    • Democritus' View: Atoms are absolute fundamental units that cannot be broken down.
    • Modern Scientific View: Incorrect. Atoms consist of subatomic particles (protons, neutrons, electrons) and can be split into smaller entities via nuclear fission and transmutation reactions.
  • Indestructibility:
    • Democritus' View: Atoms are permanent, eternal, and unchangeable.
    • Modern Scientific View: Partially Incorrect. Atoms can undergo radioactive decay, transform into different elements, or undergo mass-energy conversion ($E = mc^2$).
  • Chemical Behavior:
    • Democritus' View: Macroscopic properties are driven by physical geometric shapes (e.g., sharp/spiky atoms for bitter liquids, smooth atoms for sweet liquids).
    • Modern Scientific View: Incorrect. Chemical behavior is governed by valence electron configurations, electronegativity, and electrostatic interactions.

• Chronological Milestone Summary:

  • John Dalton: Formulated modern Atomic Theory.
  • J.J. Thomson: Discovered the electron using cathode ray tube experiments.
  • Ernest Rutherford: Discovered the atomic nucleus (via Gold Foil Experiment) and identified the proton.
  • Niels Bohr: Developed the planetary atomic model, establishing that electrons travel in quantized energy shells surrounding the nucleus.
  • James Chadwick: Proved the existence of the neutron.

• Dalton's Atomic Theory:

  • All matter is comprised of tiny, definite particles called atoms.
  • Atoms are indivisible and indestructible.
  • All atoms belonging to a specific element share identical mass and physical properties.
  • Atoms of different elements contain different masses.
  • Atoms of different elements combine in fixed, whole-number ratios to form chemical compounds.

• Discovery of the Electron (J.J. Thomson, 1897):

  • Utilized a high-voltage cathode ray tube containing a vacuum pump, cathode (negative electrode), anode (positive electrode), and electric deflection plates.
  • Observed that the emitted cathode ray beam was consistently deflected towards the positively charged plate ($+\text{plate}$).
  • Concluded that cathode rays consist of streams of negatively charged subatomic particles, later named electrons.

• Determination of Electron Mass (Robert Millikan, 1909):
  • Utilized the Oil Drop Apparatus, spraying microscopic oil droplets through an atomizer into a chamber with ionizing radiation, a microscope, and charged electric plates ($+\text{plate}$ and $-\text{plate}$).
  • Measured the precise charge of an electron and derived its absolute mass to be $9.11 \times 10^{-28}\,\text{g}$.
  • Established that an electron has one unit of negative elementary charge and a mass equivalent to approximately $\frac{1}{1840}$ of a hydrogen atom.

• Discovery of the Atomic Nucleus (Ernest Rutherford, 1909–1911):
  • Experiment Setup: Fired positively charged alpha particles ($\alpha$-particles) from a radioactive source inside a lead shield block at an ultrathin sheet of gold foil surrounded by a circular fluorescent detection screen.
  • Observations:
    • The vast majority of alpha particles passed straight through the gold foil undeflected.
    • A small fraction of alpha particles were deflected at large angles.
    • An extremely small minority (roughly 1 in 8000) bounced directly backwards toward the source.
  • Deductions (1911):
    • The atom consists mostly of empty space.
    • The positive charge and almost all the mass of the atom are concentrated in an extremely dense, tiny central core called the nucleus.

• Discovery of the Proton (Ernest Rutherford, 1919–1920):
  • In 1919, Rutherford bombarded pure nitrogen gas (${_7^{14}}\text{N}$) with high-energy alpha particles (${_2^4}\text{He}$).
  • Observed the artificial transmutation of nitrogen into oxygen (${_8^{17}}\text{O}$) along with the emission of positively charged hydrogen nuclei (${_1^1}\text{H}$).
  • Proved that the hydrogen nucleus is a subatomic building block present within all atomic nuclei, officially naming this particle the proton in 1920.

• Discovery of the Neutron (James Chadwick, 1932):
  • Bombarded beryllium (${_4^9}\text{Be}$) targets with alpha particles (${_2^4}\text{He}$), causing Beryllium to emit a highly penetrating, neutral radiation beam.
  • Directed this unknown radiation stream at paraffin wax (a substance rich in hydrogen atoms/protons).
  • The neutral radiation collided with protons in the paraffin wax, knocking protons out at high speeds.
  • Demonstrated that this unknown beam consisted of uncharged subatomic particles with a relative mass virtually identical to protons, proving the existence of the neutron.

Subatomic Mass Comparison and Atomic Notation

• Relative Mass Relationships:

  • Protons and neutrons have equal relative mass: $\text{Mass of Neutron} \approx \text{Mass of Proton} = 1\,\text{amu}$.
  • Electrons possess negligible mass: $\text{Mass of Electron} = \frac{1}{1840} \times \text{Mass of Proton}$.
  • Consequently, nearly $100\%$ of an atom's mass is concentrated in the nucleus.

• Standard Nuclide Notation:

  • Represented generally as ${_{Z}^{A}\text{X}}$, where:
    • $\text{X}$ = Chemical symbol of the element.
    • $Z$ = Proton Number (Atomic Number), defining the element's identity and number of nuclear protons.
    • $A$ = Nucleon Number (Mass Number), defining the combined sum of protons and neutrons ($A = Z + n$).
  • Worked Example (Carbon-12, ${_{6}^{12}\text{C}}$):
    • Element: Carbon
    • Mass Number ($A$): $12$
    • Atomic Number ($Z$): $6$
    • Proton Count ($Z$): $6$
    • Electron Count (Neutral atom): $6$
    • Neutron Count ($A - Z$): $12 - 6 = 6$

Electronic Configuration and Periodic Table Organization

• Energy Shell Model:

  • Electrons populate specific spherical energy levels or shells designated by principal quantum numbers ($n = 1, 2, 3, \dots$) progressing outward from the central nucleus.
  • Maximum Electron Capacity Formula: The theoretical maximum number of electrons occupying shell $n$ is given by:         $\text{Capacity} = 2n^2$
    • For Shell 1 ($n = 1$): $2(1)^2 = 2\,\text{electrons}$
    • For Shell 2 ($n = 2$): $2(2)^2 = 8\,\text{electrons}$
    • For Shell 3 ($n = 3$): $2(3)^2 = 18\,\text{electrons}$

• Valence Electrons and Chemical Reactivity:

  • Electrons occupying the outermost shell are termed valence electrons.
  • Valence electrons determine an element's bonding behavior and reactivity.
  • Atoms lose, gain, or share valence electrons to attain a stable noble gas configuration (Octet Rule, generally 8 valence electrons).

• Periodic Table Group Correlations:

  • The Periodic Table organizes elements into vertical columns called Groups based on the number of valence (outermost) electrons.
  • Example (Aluminium, ${_{13}^{27}\text{Al}}$):
    • Proton count = $13$
    • Electron arrangement = $2, 8, 3$
    • Outermost (valence) electrons = $3$
    • Group Assignment: Placed in Group III because it has 3 outermost electrons.
  • Group Summary Table:

Isotopes and Relative Atomic Mass ($A_r$)

• Definition of Isotopes: Atoms of the same chemical element possessing the same number of protons (identical atomic number $Z$) but different numbers of neutrons (different mass number $A$).

• Examples of Isotopic Systems:

  • Carbon Isotopes:
    • Carbon-12 (${_{6}^{12}\text{C}}$): 6 protons, 6 neutrons, 6 electrons (Stable, most abundant).
    • Carbon-13 (${_{6}^{13}\text{C}}$): 6 protons, 7 neutrons, 6 electrons (Stable).
    • Carbon-14 (${_{6}^{14}\text{C}}$): 6 protons, 8 neutrons, 6 electrons (Unstable, radioisotope).
  • Uranium Isotopes:
    • Uranium-235 (${_{92}^{235}\text{U}}$) and Uranium-238 (${_{92}^{238}\text{U}}$) occur naturally in the Earth's crust with long half-lives. Uranium-234 (${_{92}^{234}\text{U}}$) forms as an intermediate decay product.

• Relative Atomic Mass ($A_r$):

  • Definition: The ratio of the weighted average mass of one atom of an element compared to $\frac{1}{12}\text{th}$ of the mass of a single atom of Carbon-12 (${^{12}\text{C}}$).
  • Explains why atomic weights on periodic tables are non-integer values (decimals).
  • Calculation Formula:$A_r = \sum \left( \frac{\text{Isotope Mass} \times \text{Percentage Abundance}}{100} \right)$

• Sample Relative Atomic Mass Calculations:

  • Example 1: Chlorine ($\,\text{Cl}$)
    • Isotopes: ${_{17}^{35}\text{Cl}}$ ($75\%$ relative abundance) and ${_{17}^{37}\text{Cl}}$ ($25\%$ relative abundance).
    • Calculation:$A_r(\text{Cl}) = \left(35 \times \frac{75}{100}\right) + \left(37 \times \frac{25}{100}\right)$$A_r(\text{Cl}) = 26.25 + 9.25 = 35.5$
  • Example 2: Rhenium ($\,\text{Re}$)
    • Isotopes: ${^{187}\text{Re}}$ ($62.60\%$ relative abundance) and ${^{185}\text{Re}}$ ($37.40\%$ relative abundance).
    • Calculation:$A_r(\text{Re}) = \left(187 \times \frac{62.60}{100}\right) + \left(185 \times \frac{37.40}{100}\right)$$A_r(\text{Re}) = 117.062 + 69.19 = 186.252\,\text{amu}$
  • Example 3: Lithium ($\,\text{Li}$)
    • Isotopes: $\text{Li-6}$ (Mass = $6.015\,\text{amu}$, Abundance = $7.42\%$) and $\text{Li-7}$ (Mass = $7.016\,\text{amu}$, Abundance = $92.58\%$).
    • Calculation:$A_r(\text{Li}) = \left(6.015 \times \frac{7.42}{100}\right) + \left(7.016 \times \frac{92.58}{100}\right)$$A_r(\text{Li}) = 0.446313 + 6.4954128 = 6.9417\,\text{amu}$

Mass Spectrometry

• Functions and Applications:

  • Determining the relative atomic mass ($A_r$) of elements.
  • Determining the relative molecular mass ($M_r$) of chemical compounds.
  • Identifying naturally occurring isotopes, their relative isotopic masses, and fractional abundances.

• Instrumentation Components & Operational Steps:

  1. Vaporization Chamber / Sample Inlet: Solid/liquid element sample is vaporized into gaseous atoms using a heater furnace.
  2. Ionization Chamber: Gaseous atoms are bombarded with a high-energy electron beam ($e^-$) to knock off outer electrons, forming positive gaseous ions ($\text{M}^+$).
  3. Acceleration Plates / Electric Field: High-voltage accelerator plates create an electric field that accelerates positive ions into a focused beam.
  4. Deflection Chamber / Magnetic Sector: Powerful electromagnet applies a perpendicular magnetic field that deflects positive ions along a curved path according to their mass-to-charge ratio ($\frac{m}{z}$):
     • Light Ions: Deflected too strongly ("Too Light").
     • Heavy Ions: Deflected insufficiently ("Too Heavy").
     • Target Ions: Deflected precisely along the trajectory curve to impact the collector slit ("Just Right").
  5. Ion Detector & Recorder: Detects arriving electrical current from positive ions, outputting a mass spectrum showing signal intensity versus mass-to-charge ratio ($\frac{m}{z}$).

Isotopic Properties and Radioisotope Applications

• Comparison of Isotope Properties:

  • Chemical Properties: Identical. Isotopes contain an identical proton number and electronic configuration (same valence electron arrangement), leading to identical chemical behavior.
  • Physical Properties: Slightly Different. Variations in mass numbers cause subtle differences in physical properties such as density, melting point, boiling point, and rates of gaseous diffusion.

• Classification of Isotopes:

  • Non-radioactive (Stable): Possess an optimal, balanced proton-to-neutron ratio within the nucleus.
  • Radioactive (Radioisotopes / Unstable): Possess an unstable nuclear balance due to excessive neutrons, deficient neutrons, or an excessively large atomic nucleus.

• Key Radioisotope Applications:

  • Americium-241 (${^{241}\text{Am}}$): Used as an alpha-particle ionizing source in household ionization smoke detectors.
  • Carbon-14 (${^{14}\text{C}}$): Used in radiocarbon dating to determine the age of ancient organic remains.
  • Technetium-99m (${^{99m}\text{Tc}}$): Most widely utilized gamma-emitting medical radioisotope tracer in diagnostic imaging.
  • Cobalt-60 (${^{60}\text{Co}}$): Emits gamma radiation utilized in external beam radiation therapy for cancer treatment and industrial sterilization of surgical instruments.
  • Uranium-235 (${^{235}\text{U}}$): Used as nuclear fuel in commercial nuclear power reactors, undergoing controlled nuclear fission to generate heat energy for electricity production.

Formation of Ions and Trends in Ionic Size

• Atomic Neutrality vs. Charged Ions:

  • Neutral Atom: Electrically neutral entity where proton count equals electron count ($\text{Protons} = \text{Electrons}$).
  • Ion: A charged chemical species formed when electrons are transferred (added or lost) from a neutral atom, making proton count unequal to electron count ($\text{Protons} \neq \text{Electrons}$).

• Cations (Positive Ions):

  • Formation Mechanism: Elements with low electronegativity and low ionization energy (metals) readily lose valence electrons.
  • Charge Imbalance: Total proton count in the nucleus exceeds electron count, leaving a net positive charge.
  • General Formation Equation:$\text{M}_{(g)} \rightarrow \text{M}_{(g)}^{n+} + n\text{e}^-$
  • Example (Sodium, $\,\text{Na}$):
    • Sodium Atom Configuration: $2, 8, 1$ ($11\,p^+$, $11\,e^-$)
    • Loss of $1\,e^-$ gives Sodium Ion ($\text{Na}^+$) Configuration: $2, 8$ ($11\,p^+$, $10\,e^-$)

• Anions (Negative Ions):

  • Formation Mechanism: Elements with high electronegativity and exothermic electron affinity (non-metals) readily capture additional valence electrons.
  • Charge Imbalance: Total electron count exceeds proton count, leaving a net negative charge.
  • General Formation Equation:$\text{X}_{(g)} + n\text{e}^- \rightarrow \text{X}_{(g)}^{n-}$
  • Example (Oxygen, $\,\text{O}$):
    • Oxygen Atom Configuration: $2, 6$ ($8\,p^+$, $8\,e^-$)
    • Gain of $2\,e^-$ gives Oxide Ion ($\text{O}^{2-}$) Configuration: $2, 8$ ($8\,p^+$, $10\,e^-$)

• Trends in Ionic Radii:

  • Cation Size vs. Parent Atom: Cations are smaller than their neutral parent atoms. Loss of valence shell electrons reduces electron-electron shielding and increases the effective nuclear charge ($Z_{\text{eff}}$) pulled per remaining electron. Higher positive charges yield progressively smaller ionic radii.
  • Anion Size vs. Parent Atom: Anions are larger than their neutral parent atoms. Addition of electrons into the valence shell increases mutual inter-electronic repulsion without altering nuclear positive charge, expanding the electron cloud. Higher negative charges yield progressively larger ionic radii.

Concept Map Summary

• Summary Flow Relationships:
  • Atoms contain Sub-atomic particles (Protons, Neutrons, Electrons).
  • Protons determine the Proton Number ($Z$).
  • Protons + Neutrons together constitute the Nucleon Number ($A$).
  • Electrons dictate the Electronic Configuration.
  • Isotopes have identical proton numbers but differing neutron numbers, determining the weighted Relative Atomic Mass ($A_r$).