Science 4th QT

• Atoms

  • Building blocks of matter, everything on Earth is made up of it.
• Subatomic Particles of an Atom

  1. Neutrons (N)
  2. Protons (+)
  3. Electrons (-)
• Protons (+)

  • located inside the Nucleus of an atom
  • positive charge
  • mass of one
• Neutrons (N)

  • located in the Nucleus of an atom
  • has no charge
  • mass of one
• Electrons (-)

  • orbit the nucleus of an atom
  • located at the atoms shells/orbits
  • negative charge
  • are very small (basically no mass)
• In a Neutral Atom, there are the same number of Protons and Electrons.
• It is the Electromagnetic Force of Attraction between the positive protons in the Nucleus and the negative electrons orbiting around the nucleus that holds the atom together.
• Electromagnetic Force of Attraction (EMFA)

  • force of attraction or repulsion between all electrically charged particles. 
• Remember!

  • Atomic Number = Number of Protons = Number of Electrons
  • Mass Number = Number of Protons + Number of Neutrons
  • Number of Electrons = Number of Protons (for a neutral atom)
• Atomic Number

  • includes the number of Protons and Neutrons, they are two largest particles in the atoms.
  • Mass of the atoms is located in the Nucleus
  • Atomic Mass Number = Protons + Neutrons
• Compounds

  • distinct group of atoms held together by chemical bonds.
• Ions

  • group of atoms with a positive and negative charge.
  • does not have the same number of Electrons and Protons
• + = Cations

  -= Anions
• Isotopes

  • each of two or more forms of the same element that contain equal numbers of protons but different numbers of neutrons in their nuclei.
  • differ in relative atomic mass but not in chemical properties.
• Solid Sphere Model - J. Dalton 1803

  • Atoms are tiny balls that can't be broken and are all made of the same material.
  • Helped explain how different chemicals mix together and what makes them different.
  • Couldn't explain differences in atomic mass within an element (isotopes)
  • Couldn't explain the behavior of atoms in chemical reactions.
• John Dalton 1766-1844

  • British chemist and physicist
  • Proposed the Solid Sphere Model in the early 19th century
  • Shifted from philosophical ideas to scientific theory
  • First atomic model based on experimental evidence and quantitative observations
  • paved the way for the development of modern atomic theories
• Plum Pudding Model - J.J. Thomson 1897

  • Atoms are like plum pudding. with tiny positive charges scattered throughout a cloud of negative electrons.
  • Helped explain why atoms have a neutral charge overall and why they emit light when they collide with each other.
  • Couldn't explain why electrons didn't collapse into the positive sphere.
  • failed to predict the distribution and arrangement of electrons.
  • Lacked explanation for the nucleus and its positive charge.
• Joseph John Thomson 1856-1940

  • English physicist known for his work on the nature of electrons.
  • Proposed the Plum Pudding Model in the late 19th century.
  • Discovered electrons as distinct particles.
  • Shifted understanding from indivisible atom to subatomic particles.
  • Paved the way for future exploration of atomic structure.
• Nuclear Model - E. Rutherford 1911

  • Atoms have a nucleus with a positive charge and most of the mass, surrounded by electrons that orbit like planets.
  • Explains why particles can pass through or bounce off atoms, and is the basis of our current understanding of the atomic structure.
  • Didn't explain the stability of the nucleus against electrostatic repulsion.
  • Lacked details about electron orbits and energy levels.
  • Didn't incorporate the principles of quantum mechanics.
• Ernest Rutherford 1871-1937

  • New Zealand born physicist known for his contributions to nuclear physics.
  • Introduced the Nuclear Model in the early 20th Century
  • First model to propose a central, massive nucleus.
  • Explained the behavior of positively charged alpha particles in the gold foil experiment.
  • Laid the groundwork for understanding atomic structure and radioactivity.
• Planetary Model - N. Bohr 1913

  • Electrons orbit the nucleus of an atom in specific energy levels or shells.
  • Explains why atoms emit light and why they absorb certain colors of light.
  • Explains the stability of atoms and why they don't fall apart.
  • Limited explaining the Hydrogen atom.
  • Couldn't account for the behavior of multi-electron atoms.
  • Didn't incorporate the wave-like nature of electrons.
• Niels Bohr 1885-1962

  • Danish physicist known for his pioneering work in atomic structure.
  • Proposed the Planetary Model in the early 20th century.
  • Explained atomic structure spectra with precision.
  • Introduced the concept of quantized energy levels.
  • Bridged classical physics with emerging quantum mechanics.
• Quantum Model - E. Schrodinger 1920s

  • Electrons exist as a probable wave-like pattern around the nucleus, not in a specific orbit.
  • Explains why electrons act like particles and waves, and is the foundation of our understanding of atomic structure and widely used in modern physics.
  • Mathematical complexity of the model.
  • Requires advanced mathematics to calculate electron probabilities.
  • Doesn't provide a simple visual representation of atomic structure.
• Erwin Schrodinger 1887-1961

  • Austrian physicist renowned for his contributions to quantum mechanics.
  • Proposed the Quantum Model in the 1920s.
  • Quantum mechanics provides a comprehensive understanding of electron behavior.
  • Schrodinger's model successfully explains multi-electron atoms.
  • Quantum mechanics is the foundation of modern atomic theory.
• If Atoms are Solid Spheres, what would happen if you tried to cut one in half?

  In the Solid Sphere Model, atoms were believed to be indivisible, so you couldn't actually cut one in half. It was thought that atoms were the smallest, fundamental building blocks of matter.
• How would you describe the distribution of "raisins" (electrons) in this atomic pudding?

  In the Plum Pudding Model, the "raisins" (electrons) were thought to be scattered throughout the positive "pudding" (atom). So, the distribution of electrons was assumed to be relatively uniform within the atom.
• What did Rutherford's Gold Foil experiment reveal about the atomic Nucleus?

  Rutherford's Gold Foil experiment revealed that most of the atom's mass is concentrated in a small, positively charges nucleus at the center. This discovery overturned the idea of a uniformly distributed positive charge.
• Why might electrons be compared to planets in a solar system withing Bohr's model?

  Bohr's Planetary model drew an analogy between electrons orbiting the nucleus and planets orbiting the sun. It simplified the complex behavior of electrons, suggesting they had quantized energy levels like planets have orbits.
• How does Schrodinger's cat relate to the concept of electron probability clouds in the Quantum Model?

  Schrodinger's cat is a thought experiment in Quantum Mechanics. not directly related to electrons. However, it illustrates the probabilistic nature of quantum systems, similar to how electrons are described by probability clouds in the Quantum Model. The cat is in a superposition of states (both alive and dead) until observed, much like electrons can exist in multiple states until measured.
• Periodic Table of Elements

  • is a method of showing the chemical elements in a table with the elements arranged in order of increasing atomic number.
• Johannes Wolfgang Dobereiner 1817-1829

  • grouped elements based on similarity.
  • "Triads"
• John Newlands 1864

  • arranged the 62 known elements in order of increasing atomic weights, noted that after interval of eight elements similar physical/chemical properties reappeared.
  • Law of Octaves
• Dmitri Ivanovich Mendeleev 1869

  • proposed arranging elements by atomic weights and properties (Lothar Meyer independently reached similar conclusion but published results after Mendeleev).
  •  Mendeleev’s periodic table of 1869 contained 17 columns with two partial periods of seven elements each.
  • Periodical Law
  • "The properties of elements are periodic functions of their atomic masses."
• Henry Moseley 1913

  • An English physicist
  • He assigned a whole number to the size of nuclear charge of an atom and he called this atomic number.
  • able to derive the relationship between x-ray frequency and number of protons.
  • When Moseley arranged the elements according to increasing atomic numbers and not atomic masses, some of the inconsistencies associated with Mendeleev’s table were eliminated. The modern periodic table is based on Moseley’s Periodic Law (atomic numbers).
• Groups

  • are vertical columns on the periodic table.
  • are numbered 1A to 8A from left to right across the periodical table (excluding the transition metals, lanthanide series, metals, and actinide series metals.)
  • are numbered 3B to 12B from left to right across the periodic table for the transition metals only.
• Periods

  • are horizontal rows on the periodic table.
  • Periods are numbered 1 to 7, top down on the periodic table.
• Atomic Size

  • based on atomic radius
  • the atomic radius increases as you go down and/or go to the left
• Ionization Energy

  • energy needed to remove a valence electron
  • the ionization increases as you go up and/or go to the right
• Electron Affinity

  • the energy change that occurs when an electron is added to a gaseous atom.
  • the electron affinity increases as you go up and/or go to the right
• Electron Negativity

  • the relative ability of a covalently bonded atom to attract shared electrons
  • the electron negativity increases as you go up and/or right