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Atomic Structure

Fundamentals of Atomic Theory
Understanding Subatomic Particles
Mass Spectrometry
Atomic Structure and Isotopes
Electronic Configuration
Prep Questions

Fundamentals of Atomic Theory
Definition of an Atom
  • Definition:
    An atom is the smallest unit of matter that retains the identity of a chemical element. It consists of a nucleus containing protons and neutrons, surrounded by electrons in shells.
Structure of an Atom:
  • Nucleus:
    • The nucleus is the central part of the atom, containing protons and neutrons.
    • Protons are positively charged particles.
    • Neutrons are neutral particles.
    • The nucleus is dense and contains most of the atom's mass.
  • Electron Cloud:
    • Electrons are negatively charged particles that orbit the nucleus in energy levels or shells.
    • The arrangement of electrons in shells determines the atom's chemical properties.
Historical Development of Atomic Theory
  • Dalton's Atomic Theory (1803):
    • Proposed by John Dalton.
    • Main points:
      • All matter is composed of tiny, indivisible particles called atoms.
      • Atoms of the same element are identical in mass and properties.
      • Atoms of different elements have different masses and properties.
      • Atoms combine in simple whole-number ratios to form compounds.
      • Atoms cannot be created or destroyed in chemical reactions.
  • Thomson's Plum Pudding Model (1897):
    • Proposed by J.J. Thomson.
    • Described the atom as a sphere of positive charge with negatively charged electrons embedded in it, like raisins in a pudding.
    • Discovered the electron through his experiments with cathode rays.
  • Rutherford's Nuclear Model (1911):
    • Proposed by Ernest Rutherford.
    • Conducted the gold foil experiment, which led to the discovery of the nucleus.
    • Found that most of the atom is empty space with a small, dense, positively charged nucleus.
    • Electrons orbit the nucleus, similar to planets orbiting the sun.
  • Bohr's Planetary Model (1913):
    • Proposed by Niels Bohr.
    • Suggested that electrons move in fixed orbits or energy levels around the nucleus.
    • Each orbit has a specific energy level.
    • Electrons can jump between orbits by absorbing or emitting energy in fixed quanta.
Understanding Subatomic Particles
Protons
  • Properties:
    • Charge: +1
    • Mass: Approximately 1 atomic mass unit (amu)
    • Location: Nucleus
  • Role in the Atom:
    • Determines the element's identity (atomic number).
    • Contributes to the mass of the atom.
Neutrons
  • Properties:
    • Charge: 0 (neutral)
    • Mass: Approximately 1 amu
    • Location: Nucleus
  • Role in the Atom:
    • Stabilizes the nucleus by reducing repulsive forces between protons.
    • Contributes to the mass of the atom.
Electrons
  • Properties:
    • Charge: 0 (neutral)
    • Mass: Approximately 1 amu
    • Location: Nucleus
  • Role in the Atom:
    • Involved in chemical bonding and reactions.
    • Determines the atom's chemical properties and reactivity.
Mass Spectrometry
Definition of a Mass Spectrometer
A mass spectrometer is an analytical device used to measure the mass-to-charge ratio of ions. It helps in identifying the composition of a sample by generating a spectrum of the masses of its constituent ions.
Working of a Mass Spectrometer
  • Ionization:
      The sample is ionized, usually by losing an electron to form positive ions.
      Methods of ionization include electron impact, chemical ionization, and electrospray ionization.
  • Acceleration:
      The ions are accelerated by an electric field to have the same kinetic energy.
  • Deflection:
    • The accelerated ions are deflected by a magnetic field. The degree of deflection depends on their mass-to-charge ratio.
    • Lighter ions or ions with a higher charge are deflected more.
  • Detection:
      The deflected ions are detected, and a mass spectrum is produced showing the relative abundance of each ion.
Applications of Mass Spectrometry
  • Identifying the Composition of Unknown Substances:
      Used in chemistry and biochemistry to identify unknown compounds.
  • Determining Isotopic Composition of Elements:
      Helps in determining the isotopic distribution in a sample.
  • Applications in Various Fields:
    • Environmental science: Detecting pollutants.
    • Forensic science: Analyzing substances found at crime scenes.
    • Medicine: Drug testing and metabolic studies.
Atomic Structure and Isotopes
Definition of an Element
  • Definition:
    An element is a pure substance consisting of only one type of atom, characterized by its atomic number (number of protons).
  • Examples:
    Hydrogen (H), Carbon (C), Oxygen (O), etc.
  • Periodic Table:
    Elements are arranged in the Periodic Table based on their atomic number.
    Provides valuable information about the properties of elements, including atomic number, atomic mass, and electron configuration.
Explanation of Isotopes
  • Definition:
      Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons.
  • Examples:
      Carbon-12 (12C) and Carbon-14 (14C) are isotopes of carbon.
  • Typical Example:
    • Carbon-12: 6 protons, 6 Neutrons
    • Carbon-14: 6 protons, 8 neutrons
  • Notation:
      Isotopes are often written with the element symbol, the atomic number as a subscript, and the mass number as a superscript. For example, 126 C and 146 C
Applications of Isotopes
  • In Medicine:
      Radioactive isotopes are used in medical imaging and cancer treatment (e.g., iodine-131 in thyroid treatment).
  • In Archaeology:
      Carbon dating using Carbon-14 to determine the age of ancient artifacts.
  • Other Applications:
    • Isotopes in environmental studies to trace chemical pathways and sources.
    • Isotopes in agriculture to study nutrient uptake in plants.
Electronic Configuration
Definition of Electronic Configuration
  • Definition:
    The arrangement of electrons in the energy levels or shells around the nucleus of an atom.
  • Importance:
    Determines the chemical properties and reactivity of an element.
Rules for Writing Electronic Configurations
  • Aufbau Principle:
      Electrons fill the lowest energy orbitals first before filling higher energy ones.
  • Pauli Exclusion Principle:
      No two electrons in an atom can have the same set of four quantum numbers.
  • Hund's Rule:
      Electrons will fill degenerate orbitals (orbitals of the same energy) singly before pairing up.
Writing Electronic Configurations
  • Practice Examples:
    • Hydrogen (H): 1s¹
    • Helium (He): 1s²
    • Lithium (Li): 1s² 2s¹
    • Beryllium (Be): 1s² 2s²
    • Boron (B): 1s² 2s² 2p¹
    • Carbon (C): 1s² 2s² 2p²
    • Nitrogen (N): 1s² 2s² 2p³
    • Oxygen (O): 1s² 2s² 2p⁴
    • Fluorine (F): 1s² 2s² 2p⁵
    • Neon (Ne): 1s² 2s² 2p⁶
  • Pauli Exclusion Principle:
    • Sodium (Na): 1s² 2s² 2p⁶ 3s¹
    • Magnesium (Mg): 1s² 2s² 2p⁶ 3s²
    • Chlorine (Cl): 1s² 2s² 2p⁶ 3s² 3p⁵
Importance of Electronic Configuration
  • Chemical Bonding:
    • Determines how atoms bond with each other to form molecules.
    • Valence electrons (electrons in the outermost shell) play a key role in bonding.
  • Reactivity:
    • Elements with similar electronic configurations exhibit similar chemical properties.
    • Noble gasses have stable electronic configurations and are generally unreactive.
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