1. Electron Configuration & Orbital Writing Pre AP Chemistry
Electrons in Atoms
Electron Configuration & Orbital Writing
Pre AP Chemistry

2. Energy Levels in Atoms
Because Rutherford’s model could not explain some of the properties of elements, his students continued to study the atom. Niels Bohr used Rutherford’s model to create a new atomic model.
Niels Bohr proposed a model in which the electrons move around the nucleus. He theorized that the electron orbits the nucleus. He also theorized that the orbits were different energy levels that the electrons travel in and can be excited to a high energy level and that the electrons did not lose energy and fall into the nucleus. (The weakness in Rutherford’s theory.)
A quantum of energy is the amount of energy required to move an electron from its present energy level to the next higher one. (Also called a quantum leap)

3. B. The Quantum Mechanical Model
Erwin Schrödinger related the amplitude of the electron wave, Y (psi), to any point in space around the nucleus. His equation treated the electron as a wave and developed an equation to describe this behavior. The quantum mechanical model comes from the mathematical solutions to Schrödinger equation.
The quantum mechanical model does not define an exact path for the electron to take around the nucleus but instead estimates a probability of finding the electron in a certain position. Since the volume occupied by an electron is somewhat vague, it is better to refer to an electron cloud.

4. Atomic Models

5. Electron Configuration /Orbitals
Electrons are arranged around the nucleus in specific regions called shells.
Electrons of lower energy are located near the nucleus and electrons of higher energy are in the highest shell and are called valence electrons.
Each shell is broken down into subshells or orbitals (s, p, d, f) which are three dimensional regions around the nucleus that indicate the probable location where electrons are likely to occur.
Elements located in the same group on the periodic table have their highest energy electrons arranged in the same way and that makes them chemically similar.

6. Electron Configuration
the phone number or address of an atom.
1st = energy level of the e-.
represented by rows ( Period) on the periodic table
1 = lowest, 7 = highest
2nd = shape of sub-level
4 different shapes…found on each energy level
s, p, d, f
3rd = orientation of orbital in space.
- number of electrons, each element represents 1
electron
- s has 1 orbital with a total of 2 electrons
- p has 3 orbitals with a total of 6 electrons
- d has 5 orbitals with a total of 10 electrons
- f has 7 orbitals with a total of 14 electrons

7. Single atom orbital

8. Electron Configurations
In e- configurations, you must represent (in order), the:
energy level # (electron shell), subshell orbital (s,p,d or f) and e- # for the atom.
Always read left to right, top to bottom.

9. II. Electron Arrangement in Atoms A. Electron Configurations

10. Orbital Filling- Aufbau Principle
Aufbau principle Electrons enter at the lowest energy level Some energy levels overlap into the adjacent principal energy level.

11. Orbital Filling- Pauli Exclusion Principle
Pauli exclusion principle Spectral data shows that only 2 electrons can exist in the same orbital Electrons behave as if they were spinning about their own axis When electrons occupy the same orbital – they are said to spin in opposite directions (assign +1/2 and – 1/2).

12. Orbital Filling- Hund’s Rule
Hund’s Rule Also with the principle, you must have all orbital filled with one electron before you can add the other electron with opposite spin to the orbital All elements would like to have a completely filled orbital and the maximum number of electrons that can exist in a filled orbital is eight

13. 1s1 Hydrogen letter = the orbital location
first number =the energy level.
letter = the orbital location
superscript = number of electrons

14. 1s2 Helium Electron shell (energy level) Type of subshell orbitals
Number of electrons

15. 1s22s1 Lithium Electron shell (energy level) Type of subshell orbitals
Number of electrons

16. 1s22s2 Beryllium Electron shell (energy level)
Type of subshell orbitals
Number of electrons

17. Period 2 Boron = 1s2 2s2 2p1 Carbon = 1s2 2s2 2p2
Nitrogen = 1s2 2s2 2p3
Oxygen = 1s2 2s2 2p4
Fluorine = 1s2 2s2 2p5
Neon = 1s2 2s2 2p6
Remember that the superscript numbers must equal the atomic number

18. Write out the correct e- configuration for Na (Z = 11)
1s2 2s2 2p6 3s1
Write out the correct e- configuration for Mg (Z = 12)
1s2 2s2 2p6 3s2
Write out the correct e- configuration for P (Z = 15)
1s2 2s2 2p6 3s2 3p5
‘d’ orbitals are so large that they reach into the next energy level. Therefore, the 1st ‘d’ orbital belongs to the 3rd energy level…even though we don’t see it until the 4th row of the P.T.!
Write out the correct e- configuration for Br (Z = 35)
1s2 2s2 2p6 3s2 3p5 4s2 3d10 4p5
‘f’ orbitals are so large that they reach into the next 2 energy levels. Therefore, the 1st ‘f’ orbital belongs to the 4th energy level…even though we don’t see it until the 6th row of the P.T.!

19. Orbital Filling

20. Electron Arrangement in Atoms Exceptional Electron Configurations
Filled sublevels are more stable than partial filled or half-filled sublevels.
But sometimes half-filled may be more stable than other configurations.
Example: Iron, Copper, and Chromium
1s22s22p63s23p64s23d5
1s22s22p63s23p64s13d10 This is more stable because the higher energy sublevel is completely filled.
*1s22s22p63s23p64s13d5 3d and 4s are very close together indeed in the transition series, so filled shell or half filled shell stability (Cr, 4s13d5) is enough to tip the balance. *

21. Atomic Emission Spectra and the Quantum Mechanical Model Light and Atomic Emission Spectra
This energy consist of variation in electric and magnetic fields taking place in a regular, repeating fashion (Electromagnetic energy), Light is a form of electromagnetic radiation
If you plot the strength of the variation against time, the graph shows “waves” of energy. The number of waves peaks that occur in a unit of time is called the frequency of the wave (Greek letter v and units are Hertz (Hz)).
The distance between the peaks is the wavelength (Greek letter λ) and the amplitude of a wave is the height from the maximum displacement from zero.
These characteristics of waves are related by
c= λv
where c is the speed of light which is 3.0 x m/s.

22. Atomic Emission Spectra and the Quantum Mechanical Model Light and Atomic Emission Spectra
The wavelengths of light can separate into a spectrum of color. This is part of the visible spectrum. There are two types of spectrums. a. Adsorption spectrum. b. Emission spectrum.
Adsorption spectrum is when the energy gained by the excited electron is is absorbed so that it is missing in visible spectrum.
Emission spectrum is when the excited electrons lose the energy and it is emitted at specific points on the visible spectrum that appear as single lines on a detector.

23. The Quantum Concept and Photons
Consider the electron of a hydrogen atom in its lowest energy level, or ground state. The quantum numbers represent the different energy states. The difference between these energy states corresponds to the lines in the hydrogen spectrum.
With more complex atoms more than one electron is present and the interaction between electrons make solution to the equation impossible because electrons have the same charge.
It is possible to approximate the electronic structure of a multi-electron atom. This approximation is made by first calculating the various energy states and quantum numbers.
It is assumed that the various electrons in multi-electron atom occupy the same energy states without affecting each other.

24. Atomic Emission Spectra and the Quantum Mechanical Model The Quantum Concept and the Photoelectric Effect
Max Planck used Bohr’s theory to develop his hypothesis. He assumed that energy is given off in packets called quanta or photons instead of a steady stream.
He stated that the amount of energy given off is related to the frequency of light (v - Greek letter nu). He thought a quantum energy was equal to
E = hv
where h is the constant 6.63 x J/Hz (Hz = Hertz).
Albert Einstein proposed that light could be described as a quanta of energy that behaved as if they were particles. The dual wave-particle behavior is called the photoelectric effect. In the photoelectric effect, metals eject electrons when light shines on them. The frequency and the wavelength of the light determine if the photoelectric effect will occur. The light quanta are called photons.

25. Heisenberg Uncertainty Principle
Werner Heisenburg refined ideas about atomic structure. He stated that it is impossible to know the exact position and momentum of an electron in an atom. Using the equation for momentum, he proposed that mv = p where m is mass and p is momentum.
The uncertainty of position and momentum are related to Planck’s constant
∆p ∆x > h
where p is momentum and x is position (∆ = change).
Because h is constant, ∆ p and ∆ x are inversely proportional to each other.

26. Heisenberg Uncertainty Principle
Impossible to know exactly where an electron is at any moment.