Development of Modern Atomic Theory

1. Development of Modern Atomic Theory From Democritus to Schrodinger

2. Democritus • Greek (lived from 460 to 370 B.C.) • Named the atom “atomos” meaning indivisible The Atom Song Video

3. John Dalton • Dalton developed the 1st version of modern Atomic Theory in 1803 • All matter is made of small particles called atoms. • Atoms of a given element are identical in properties; atoms of different elements differ in properties. • Atoms cannot be subdivided, created, or destroyed. • Atoms of different elements combine in simple whole number ratios to form compounds. • In chemical reactions, atoms are combined, separated, or rearranged. • Problems with Dalton’s Model = atoms do have smaller parts (subatomic particles/quarks), atoms of the same element can have different mass, atoms of different elements can have the same mass (isotopes) (Other parts are still correct)

4. J.J. Thomson • J.J. Thomson concluded that the atom had negatively charged particles (1897) • He is credited with discovering the electron • His model looked like “plum pudding” with a large positive “pudding” and small, evenly distributed negative charges (“plums”)

5. Ernest Rutherford • Determined that the atom contained a small, dense center of positive charge = the nucleus (1911) • Very small volume compared to size of the atom • Credited with discovery of protons (1919) – positively charged particles and nucleus • Suggested that electrons moved around like bees around a hive • Missing behavior of electrons = how the electrons are distributed in space and why they were not drawn to the positive nucleus.

6. Overview of wave motion • Wavelength (λ) = distance between corresponding points on adjacent waves. • Frequency (v) = number of waves that pass thru a given point.

7. Planck and Einstein • Max Planck (1900) determined particles can move as waves and the energy of the particle is related to wave’s frequency • E=h (as frequency increases, energy increases) • In 1905, Einstein determined Electromagnetic Spectrum, thus light, is wave-like, but it is also a stream of particles. • Einstein named the particles (which have no mass) photons. • The energy of the photon depends on the frequency of the radiation

8. Hydrogen Line Emission Spectrum • Scientists conducted studies passing an electric current through a vacuum tube with hydrogen and saw pink light. • Already knew that when an excited atom (an atom with higher energy) returns to its ground state, it gives off the energy as light, but thought that it should be continuous spectrum • Looks continuous to the eye, however, if this light is shined through a prism, it is separated into 4 visible lines

9. Quantum Theory • This new discovery (of the lines) began the new atomic theory called quantum theory. • When an excited hydrogen atom falls to a lower level, it emits a photon of radiation. • Energy of photon = the difference in energy between the two levels. • Because only specific frequencies of light are emitted, electrons only exist in specific energy states, not just random locations.

10. Excited vs. ground state Ground State = closer to nucleus, less energy Excited State = further from nucleus, more energy

11. The Bohr Model of the Hydrogen Atom • Niels Bohr’s model (1913) explained the line emission spectrum: • Electrons can only circle the nucleus in allowed paths or orbits. • An electron is in the lowest state in the orbit closest to the nucleus and higher in orbits farther from the nucleus. • As electrons drop from higher states/excited to lower/ground states, energy (as a photon) is emitted = emission of energy/light • As electrons move up from lower/ground to higher/excited states , they absorb energy = absorption of energy • Bohr’s model had some things right (fixed energy levels and can move between levels) and some things wrong (circle the nucleus in fixed paths)

12. Heisenberg • Heisenberg uncertainty principle (1927) states that it is impossible to determine simultaneously both the position and velocity of an electron or any other particle

13. Erwin Schrodinger • Erwin Schrodinger developed an equation that treats electrons as waves using the dual wave-particle theory (1926) • Together with Heisenberg’s uncertainty principle, laid foundation for quantum theory • Quantum theory describes mathematically the wave properties of electrons and other very small particles • Main point = electrons to not travel around the nucleus in orbits like Bohr suggested, they exist in specific regions called orbitals. • Orbitals are not exact location, but a 3D region around the nucleus that indicates the probable location of an electron.

14. Irene and Frederic Joilot-Curie and James Chadwick • Joilot-Curies discovered the neutron • James Chadwick (1932) identified it as a new particle and named the neutron • At this point, the model of the atom correctly explained the nucleus and the behavior of electrons!

15. Electron Configuration • The arrangement of electrons in any element’s atom • Describes the orbitals for electrons of each element • A single electron configuration exists for each different element • Atoms assume arrangements that have the lowest possible energy = ground-state electron configurations. • Electron Configuration Activity • Large numbers = principle quantum number or energy level for electrons • Whole Numbers • 1-8 for current elements • The lower the energy level, the closer an electron is to the nucleus

16. The sublevels • Letters = sublevels for each energy level • Sublevel s has 1 orbital • 2 maximum electrons, found in all energy levels • Sublevel p has 3 orbitals • 6 maximum electrons, not found in energy level 1 • Sublevel d has 5 orbitals • 10 maximum electrons, not found in energy levels 1-2 • Sublevel f has 7 orbitals • 14 maximum electrons, not found in energy levels 1-3 • Subscripts represent # of electrons in each sublevel

17. Determining Configurations • Energy level corresponds to the row except for sublevel d (one lower) and sublevel f (two lower) • Divide Periodic Table into 4 sublevel areas and move to a new sublevel as you move across • A sublevel is full when it hits the max # of electrons for the sublevel • Standard vs. noble gas notation • 1s22s22p63s2 or [Ne] 3s2 = Mg • 1s22s22p63s23p64s23d3 or [Ar] 4s23d3 = V • Noble Gas in brackets represents the configuration of the noble gas, the rest of the configuration is written after. • Don’t forget that all electrons are still represented!

18. Orbital Notation • Use arrows rather than subscripts to represent electrons • Show each orbital, not just each sublevel with lines (Remember, s has two orbitals, p has 3, d has 5, f has 7) • Each orbital can hold a max of 2 electrons – one up and one down • Always fills one in each first orbital of equal energy and then a second (everyone gets one helping first – Hund = helping) • Hund’s rules: Orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron and all electrons in singly occupied orbitals must have the same spin state.

19. What else do you learn from electron configuration? • Highest energy levels (highest number) • Types of electrons (what type – s,p,d,f) • Number of outer level electrons (how many in highest level?) • Practice: - 1st 3 rows in packet together