1. 5.2 – Quantum Theory and the Atom
Chapter 5
5.2 – Quantum Theory and the Atom

2. Bohr’s Model Danish physicist, Niels Bohr proposed model for H atom
Proposed H atom has only certain allowable energy states
Lowest allowable energy state is ground state
Atom gains energy, it’s in an “excited” state
Related H atom’s energy states to electrons w/i atom
Suggested electron in H atom moves around nucleus in certain circular orbits
Smaller orbit => lower energy state

3. Bohr’s Model

4. Bohr’s Model
Bohr assigned a # n (quantum #) to each orbit for calculations
First orbit is that closest to the nucleus
n=1
For H, atom is in ground state when the single electron is in the n=1 orbit and does not radiate energy
If outside E (energy) is added, the electron will move to higher orbit
Atom is now in excited state

5. Bohr’s Model Side note: Limit’s of Bohr’s Model:
Orbitals are not evenly spaced apart
Limit’s of Bohr’s Model:
Explained H well
Failed to explain other elements
Didn’t fully explain chemical behavior of atoms

6. Quantum Mechanical Model of the Atom
French grad student Louis Broglie proposed a fix to Bohr’s Model
Compared Bohr’s electron orbits to waves

7. Quantum Mechanical Model of the Atom
Quantum Mechanical Model – model where electrons are treated as waves
AKA – “Mechanical Model”
Like Bohr’s – limits electrons E to certain values
Unlike Bohr’s – does NOT attempt to describe electron path

8. Hydrogen Atomic Orbitals
Boundary of atomic orbital is unclear => orbital does not have exact shape
Principal Quantum # - first quantum # of atomic orbitals (n)
Indicates relative size and energy of atomic orbitals
As n increases, orbital grows and electron spends more time away from nucleus increasing E
Each major energy level is called Principal Energy Level

9. Hydrogen Atomic Orbitals
Principal Energy Level
Lowest E level is assigned a principal quantum # of 1
Up to 7 E levels for H
Energy Sublevels
Principal E level 1 – 1 sublevel
Principal E level 2 – 2 sublevels
Principal E level 3 – 3 sublevels and so on…

10. Hydrogen Atomic Orbitals
Sublevels are labeled s, p, d or f according to shape of atom’s orbital
s orbitals are spherical
p orbitals are dumb-bell shaped

11. Hydrogen Atomic Orbitals
Sublevels are labeled s, p, d or f according to shape of atom’s orbital
d and f orbitals are not all the same

12. Hydrogen Atomic Orbitals
Each orbital will contain at most, 2 electrons
Single sublevel in princ. E level 1 corresponds with 1s orbital
S sublevel has 1 orbital
2 sublevels in princ. E level 2 are 2s and 2p
Note: 2s is larger in size than 1s
p sublevels have 3 orbitals of E
Princ. E level 3 has 3 sublevels 3s, 3p and 3d
d sublevels have 5 orbitals of energy
4th sublevel consists of 4s, 4d, 4p, and 4f
f sublevels have 7 orbitals of E

13. Atomic Orbitals

14. 5.3 – Electron Configuration

15. Ground State Electron Configuration
Electron Configuration – arrangement of electrons in an atom
Low E systems are more stable => atoms tend to arrange themselves in such a way to give them lowest E possible (ground state electron configuration).
3 Rules/Principles:
Aufbau Principle
Pauli Exclusion Principle
Hund’s Rule

16. Aufbau Principle Aufbau Principle:
Electrons occupy lowest energy orbital
Diagram gives sequence (p. 156)
Sequence of energy sublevels go: s p d f

17. Pauli’s Exclusion Pauli’s Exclusion:
Electrons in orbitals are represented by arrows in boxes
Each electron has associated spin (like a top)
Can only spin in 1 of 2 directions
Arrows point up or down to represent spin
Empty box indicates unoccupied orbital
Box with 1 arrow implies one electron
Box with arrow up and down means orbital is full
**MAX of 2 arrows can occupy a box (must have opposite spin)

18. Pauli’s Exclusion

19. Hund’s Rule Hund’s Rule:
Single electron with same spin must occupy each equal- energy orbital before additional electrons with opposite spins can occupy the same orbital

20. Electron Arrangement 2 ways to represent electron configuration:
Orbital Diagrams (ex: C)
Electron Configuration Notation
Describes principal energy level and energy sublevel associated with each orbital
Superscripts represent # of electrons in orbital
Ex: C = 1s22s22p2

21. Electron Arrangement Noble Gas Notation
Method of representing electron configuration of noble gases
Noble gases are elements in last column of periodic table
Usually have 8 electrons in outermost orbital and are unusually stable
Notation uses bracketed symbols
Ex: Helium: 1s2 or [He]
Ex: Neon: 1s22s22p6 or [Ne]
Ex: Sodium:
[Ne]3s1

22. Exceptions
Use the aufbau diagram to write correct ground-state electron configurations for all elements up to #23
For elements after that, look at a sublevel diagram to see the order in which the orbitals are filled (there is an increased stability of half-filled and filled sets of s and d orbitals)
Periodic Table Method

23. Sublevel Diagram

24. Valence Electrons
Valence Electrons – electrons in atom’s outermost orbitals
Determine chemical properties of an element
Ex: S [Ne]3s23p4
6 valence electrons
Ex: Cs [Xe]6s1
1 valence electron

25. Electron Dot Structure
Electron Dot Structure – Consists of elements symbol and is surrounded by dots to represent valence electrons.
These electrons are involved in forming chemical bonds
Place one dot on each of the 4 sides of the symbol (any sequence is fine)
Once all single spots are filled, begin pairing up the electrons
Ex: Carbon

26. Electron Dot
Ex: O
Ex: F
Ex: N