1. Quantum Theory and the Atom
Electrons in Atoms and the Periodic Table

2. Learning Goals
Compare the Bohr and quantum mechanical models of the atom.
Explain the impact of de Broglie’s wave particle duality and the Heisenberg uncertainty principle on the current view of electrons in atoms.
Identify the relationships among a hydrogen atom’s energy levels, sublevels, and atomic orbitals.

3. Bohr’s Model of the Atom
Einstein’s theory of light’s dual nature accounted for several unexplainable phenomena, but it did not explain why atomic emission spectra of elements were discontinuous.

4. Bohr’s Model of the Atom
In 1913, Niels Bohr, a Danish physicist working in Rutherford’s laboratory, proposed a quantum model for the hydrogen atom that seemed to answer this question.
This model correctly predicted the frequency lines in hydrogen’s atomic emission spectrum.

6. Bohr’s Model of the Atom
The lowest allowable energy state of an atom is called its ground state.
When an atom gains energy, it is in an excited state.

7. Bohr’s Model of the Atom
Bohr suggested that an electron moves around the nucleus only in certain allowed circular orbits.

8. Bohr’s Model of the Atom
Each orbit was given a number, called the quantum number.
Bohr orbits are like steps of a ladder, each at a specific distance from the nucleus and each at a specific energy.

9. Bohr’s Model of the Atom
Hydrogen’s single electron is in the n = 1 orbit when it is in the ground state.
When energy is added, the electron moves to the n = 2 orbit.

10. Bohr’s Model of the Atom
The electron releases energy as it falls back towards the ground state.

11. Bohr’s Model of the Atom
Bohr’s model explained the hydrogen’s spectral lines, but failed to explain any other element’s lines.
For this and other reasons, the Bohr model was replaced with a more sophisticated model called the quantum-mechanical or wave- mechanical model.  

12. Quantum Mechanical Model
Louis de Broglie (1892–1987) hypothesized that particles, including electrons, could also have wavelike behaviors.
Electrons do not behave like particles flying through space.
We cannot, in general, describe their exact paths.

13. Quantum Mechanical Model
Heisenberg showed it is impossible to take any measurement of an object without disturbing it.
The Heisenberg uncertainty principle states that it is fundamentally impossible to know precisely both the velocity and position of a particle at the same time.

15. Quantum Mechanical Model
The only quantity that can be known is the probability for an electron to occupy a certain region around the nucleus.

16. Quantum Mechanical Model
Schrödinger treated electrons as waves in a model called the quantum mechanical model of the atom.
Schrödinger’s equation applied equally well to elements other than hydrogen (unlike Bohr’s model).

17. Quantum Mechanical Model
The quantum mechanical model makes no attempt to predict the path of an electron around the nucleus.
Bohr orbits were replaced with quantum-mechanical orbitals.

18. Quantum Mechanical Model
Orbitals are different from orbits in that they represent probability maps that show a statistical distribution of where the electron is likely to be found.

19. Quantum Mechanical Model
In the quantum-mechanical model, a number and a letter specify an orbital.
The lowest-energy orbital is called the 1s orbital.
It is specified by the number 1 and the letter s.

20. Hydrogen’s Atomic Orbitals
The number is called the Principal quantum number (n) and it indicates the relative size and energy of atomic orbitals.
n specifies the atom’s major energy levels, called the principal energy levels.

21. Hydrogen’s Atomic Orbitals
Energy sublevels are contained within the principal energy levels.

22. Hydrogen’s Atomic Orbitals
Each energy sublevel relates to orbitals of different shape.
s, p, d, f
s, p, d
s, p s

23. Hydrogen’s Atomic Orbitals
s sublevel:

24. Hydrogen’s Atomic Orbitals
p sublevel:

25. Hydrogen’s Atomic Orbitals
d sublevel:

26. Hydrogen’s Atomic Orbitals
f sublevel:

28. Hydrogen’s Atomic Orbitals
Orbitals are sometimes represented by dots, where the dot density is proportional to the probability of finding the electron.
The dot density for the 1s orbital is greatest near the nucleus and decreases farther away from the nucleus.
The electron is more likely to be found close to the nucleus than far away from it.

30. Hydrogen’s Atomic Orbitals

31. Hydrogen’s Atomic Orbitals
At any given time, hydrogen’s electron can occupy just one orbital.
When hydrogen is in the ground state, the electron occupies the 1s orbital.
When the atom gains a quantum of energy, the electron is excited to one of the unoccupied orbitals.