1. Electrons in Atoms

2. The Atom & Unanswered Questions
How are electrons arranged in the space around the nucleus?
Why aren’t electrons (negative) pulled into the nucleus (positive)?
To understand more about the atom, we have to understand the nature of light.

3. White Light
White light contains all of the colors of the visible light spectrum.
A prism will separate white light into a continuous spectrum of colors according to energy.
Increasing Energy

4. Electrons and Light
Scientists observed certain elements released visible light when heated in a flame.
Analysis of this light revealed that an element’s chemical properties are related to the arrangement of it’s electrons.

5. Atomic Emission Spectra The set of colors in the light released by atoms of an element.
If the light is passed through a prism, the atomic emission spectrum is produced.
A spectrum is unique to each element, like a fingerprint.
An element can be identified by the emission spectrum.

7. Bohr’s Model of a Hydrogen Atom
Electrons move around the nucleus in certain circular orbits.
Smaller orbits mean lower energy
Larger orbits mean higher energy

8. Bohr’s Model of a Hydrogen Atom
When energy is added from an outside source…
electrons move to higher-energy orbits and the atom is in an excited state.
When e- return to their lower-energy orbit, or “ground” state, the energy is lost and given off as light of a particular color.

9. Bohr’s Model of the Atom
Unlike the rungs of a ladder, hydrogen’s orbits are not evenly spaced.
Electrons can only gain or lose energy in specific amounts - enough to move up or down orbit(s).

10. The only problem…
The Bohr model worked for hydrogen, however, it was not valid for atoms with two or more electrons.

11. Where are Electrons? The Heisenberg Uncertainty Principle It is impossible to know precisely both the speed and position of a particle at the same time.
It’s impossible to take a measurement of an electron without disturbing it.
It’s impossible to assign fixed paths for electrons like in Bohr’s model.
All we know is the probability for an electron to occupy a certain region around the nucleus.

12. Where are Electrons?
Schrodinger used complex math equations to predict the most probable location of an electron (not the path!)
The equations worked for ALL elements - not just hydrogen

13. Atomic Orbitals
Orbitals are drawn to represent a 3-D region of space with a 90% probability of containing electrons.

14. Describing Electrons in Atoms
The major energy levels of the atom are numbered 1,2,3,4…
As the number increases the orbitals become larger and the energy increases.
Within the energy levels, there are sublevels.
Energy level 1 has a single sublevel
As you move farther from the nucleus (higher energy level), the number of sublevels increases.

15. Describing Electrons in Atoms
In each sublevel there are atomic orbitals (Schrodinger)…
Sublevels are labeled s, p, d, or f according to the shape of the orbitals.

16. Orbital Shapes s orbital – spherical One orbital
Exists at EVERY energy level
LOWEST energy orbital at any energy level

17. Orbital Shapes p orbital - dumb-bell shaped
Set of three orbitals aligned along the x, y, and z axis
Exists at every energy level starting at the second

18. Orbital Shapes d orbital f orbital
set of 5 orbitals, beginning at the 3rd energy level
f orbital
set of 7 orbitals, beginning at the 4th energy level

19. Orbital Shapes Energy Level Number of sublevels Types of orbitals 1 s2
s, p 3
s, p, d 4
s, p, d, f

20. Electron Configuration
The arrangement of electrons in an atom is called the electron configuration.
Electrons want to be stable…
the most stable, lowest-energy arrangement of electrons is the element’s GROUND-STATE ELECTRON CONFIGURATION
3 Rules:

21. Aufbau Principle Each electron occupies the lowest energy orbital available.
The Aufbau diagram shows the sublevels from lowest energy to highest energy.
Each box represents an orbital (s,p,d,f)

22. Hund’s Rule Single e- must occupy each equal-energy orbital before a second e- can occupy the same orbital.
Ex. P-orbitals

23. The Pauli Exclusion Principle A maximum of 2 e- can occupy an orbital but only if the e- have opposite spins.
e- in orbitals can be represented by arrows in boxes.
Each e- has a spin.
e- can only spin in one direction ↑ or ↓
unoccupied orbital
↑ orbital with one e-
↑↓ filled orbital (2 e-)

24. Electron Arrangement.
There are three ways to describe the arrangement of electrons in an atom.
Electron Configuration Notation: uses energy level, sublevel, and includes a superscript representing #e- in the orbital.
Ex. Carbon: 1s22s22p2
2. Orbital diagrams:
Ex. Carbon:

25. Electron Arrangement
3. Noble-Gas Notation: short-hand way to write electron configurations for elements with a large atomic number!
Write the noble gas (8A) symbol from the previous period (row) in brackets and write the electron configuration after.
Ex. Na is [Ne]3s1
Sr is [Kr]5s2

26. Electron Arrangement

27. Electron Arrangement