1. Bohr and
Introduction
To Electron
Structure
Part I

2. Big Picture Let’s look at three consecutive elements
Chlorine – atomic # 17
Argon – atomic # 18
Potassium - atomic # 19
How are they similar and different?

3. Chlorine Argon Potassium yellow-green gas at room temperature
Highly reactive
Argon
Gas at room temperature
Gas used in light bulbs
Extremely unreactive
Potassium
Solid metal at room temperature

4. Essential questions What accounts for the element’s differences?
What role do electrons play in all of this?

5. Let’s start by looking back at our atomic theory of the atom

6. Rutherford’s Model Described the nucleus, but…
It did not explain similarity and differences in element’s chemical properties
It did not explain why the how the electrons are arranged around the nucleus.
Nor why the negatively charged electrons did not fall into the positively charged nucleus and collapse the atom.

7. Emission Spectra
In the early 1900’s, scientists observed that different elements emitted visible light when heated.
Analysis of this emitted light reveals that chemical behavior is related to electron arrangement.
(flame test demo)

8. Emission Spectra
Atoms can absorb energy, but they must eventually release it
When atoms emit energy, it is released in the form of light
Atoms don’t absorb or emit all colors
The spectrum of colors emitted can identify the element

9. Emission Spectra Lithium chloride Red Sodium chloride Yellow
Potassium chloride
Violet
Calcium chloride
Red-orange
Strontium chloride
Bright red

10. Line Spectra = specific wavelengths are emitted; characteristic of atoms

11. The Bohr Model of the Atom
Bohr postulated that electrons do not fall into the nucleus because they can only travel in certain allowable orbits or energy levels.
Proposed a model of the atom that explained the light given off by the heated or excited atoms.
When atoms are excited, electrons jump up energy levels and then emit light as they fall down to their ground states.

12. Bohr Model of Atom: Electron Orbits
In the Bohr Model, electrons travel in orbits or energy levels around the nucleus
The farther the electron is from the nucleus the more energy it has.

13. The Bohr Model of the Atom: Orbits and Energy
Each orbit (energy level) has a specific amount of energy
Energy of each orbit is symbolized by n, with values of 1, 2, 3 etc; the higher the value the farther it is from the nucleus and the more energy an electron in that orbit has
Tro's Introductory Chemistry, Chapter 9

14. The Bohr Model of the Atom: Energy Transitions
Electrons can move from a lower to a higher (farther from nucleus) energy level by absorbing energy
When the electron moves from a higher to a lower (closer to nucleus) energy level, energy is emitted from the atom as a photon of light
Tro's Introductory Chemistry, Chapter 9

15. The Bohr Model of the Atom Ground and Excited States
Ground state – atoms with their electrons in the lowest energy level possible; this lowest energy state is the most stable.
Excited state – a higher energy state; electrons jump to higher energy levels by absorbing energy
Atom is less stable in an excited state; it will release the extra energy to return to the ground state

16. The limits of Bohr’s model
The Bohr Model very accurately predicts the spectrum of hydrogen with its one electron
It is inadequate when applied to atoms with many electrons
It has did not explain the chemical behavior of atoms.
Even though the model is incorrect, it laid the groundwork for future atomic models.

17. Ever wonder how glow sticks work?
Similar to how we use fire to excite our atoms earlier, glow sticks use a chemical reaction to excite electrons
The chemical reaction causes the electrons to rise to a higher energy level
When the electrons return to their ground state, they emit light in the form of chemiluminescence.

18. Check for Understanding
Now work with a partner to complete:
Electron Energy and Light
Bohr Atomic Models Questions