The Bohr Model of the Atom
Experimental Evidence
Theories are based on observations. What did Bohr observe that made it necessary for him to develop a new model for the structure of the atom?
Bohr noticed that when white light was passed through a prism, the visible spectrum of light was obtained. In contrast to this, Bohr saw that when light of a particular colour (like from the hydrogen discharge tube) was passed through a prism, a line spectrum was obtained. A line spectrum contains distinct lines of a particular colour.
To Bohr, the line spectra phenomenon showed that atoms could not emit energy continuously, but only in very precise quantities (he described the energy emitted as quantized). Because the emitted light was due to the movement of electrons, Bohr suggested that electrons could not move continuously in the atom (as Rutherford had suggested) but only in precise steps. Bohr hypothesized that electrons occupy specific energy levels. When an atom is excited, such as during heating, electrons can jump to higher levels. When the electrons fall back to lower energy levels, precise quanta of energy are released as a specific wavelengths (lines) of light.
The Bohr Model of the Atom
The atom consists of two main parts:
• The nucleus
• The electron cloud
The nucleus is the centre of the atom and contains the protons and neutrons. The nucleus comprises most of the mass of the atom but little of its volume. The electron cloud surrounds the nucleus. The electron cloud contains electrons. The electrons are arranged in levels around the nucleus. There is a maximum capacity for electrons at each of these levels.
Electron Capacity
The electrons are arranged in levels around the nucleus. There is a maximum capacity for electrons at each of these levels.
Level (n) / Electron Capacity (2n2)
1 / 2(1)2 = 2
2 / 2(2)2 = 8
3 / 2(3)2 = 18
4 / 2(4)2 = 32
5 / 2(5)2 = 50
6 / 2(6)2 = 72
7 / 2(7)2 = 98
The levels get closer together as you move further and further from the nucleus. "n" is the outermost level where electrons can be and still be considered a part of the atom.
An electron is at infinity if the electron is removed from the atom.
Why Do Spectral Lines Occur?
Ground State vs. Excited State
Electrons are normally found in the lowest available energy orbitals. This is known as the ground state.
If for some reason, an electron has absorbed energy and moved to an orbital of higher energy, then the electron is in an excited state.
The electrons in an excited state will eventually lose their extra energy and return to their ground state. When they release their excess energy, they emit electromagnetic radiation. When this excited electron returns to the second energy level in the atom, the electromagnetic radiation is in the form of visible light of a particular colour.
Much of your chemistry will deal with the elements with atomic numbers from 1 to 20. Let's examine where the electrons are in atoms of these elements.
We will have to remember the electron capacity at each level.
Lewis Symbols or Diagrams
Elemental properties and reactions are determined only by electrons in the outer energy levels. Electrons in completely filled energy levels are ignored when considering properties. Simplified Bohr diagrams which only consider electrons in outer energy levels are called Lewis Symbols.
A Lewis Symbol consists of the element symbol surrounded by "dots" to represent the number of electrons in the outer energy level as represented by a Bohr Diagram. The number of electrons in the outer energy level is correlated by simply reading the Group number. Lewis symbols for oxygen, fluorine, and sodium are given in the diagram.
Lewis Symbols for the elements of the second period. Correlate the number of dots with the group number.
Valence Electrons
Electron(s) located in the outermost energy level of the atom are called the valence electrons.
Why do you think that the valence electrons are of particular interest to the chemist?
Look at your periodic table in which you wrote the location of all the electrons of each of the first twenty elements. Can you remember a quick way to determine?
· the number of valence electrons that an atom of an element has?
· the level at which the valence electrons are found (the valence level)?
Lewis Structures of Atoms
Chemists draw Lewis structures to show the valence electrons in an atom.
To draw a Lewis structure you:
a. write the symbol for the element
b. locate the element on the periodic table and determine the number of valence electrons it has. For elements in the principal groups IA to VIIIA the number of valence electrons equals the group number.
c. Imagine that the symbol for the element has four sides, like sides of a square.
d. Place dots on each side of the square, one per side before doubling up.
e. Each side can have at most 2 dots.
f. When two dots are on the same side, place them as a pair.
When looking at the spectral lines of elements other than hydrogen, scientists wondered why more than one spectral line of a particular colour occurred. For example: In the spectrum for lithium there were 4 red lines, 1 orange line and 2 blue lines. If Bohr's model was accurate, then more than one line of the same colour would not be possible.
(negative of spectral lines for printing purporses)
Record in your notebook the line spectrum of three elements having more than one line of a particular colour. Many elements show this characteristic. Choose any 3 elements that do this. Be sure to label your diagrams of the line spectra that you choose. Ensure that you clearly record which element belongs to each of your recorded line spectra. (Use coloured pencils to assist with spectral lines)
Element 1) ______
Element 2) ______
Problems-Explanations
Why would the line spectrum for an element contain more than one red line?
Bohr's model of the atom could not explain why there is more than one line of a particular colour in the line spectrum of various elements. His model did explain why one line of a particular colour occurred. Do you remember why?
Refresh your memory.
Bohr's model of the atom needed to be modified in order to explain the occurrences of more than one line of a particular colour in the line spectrum of elements.
What was this explanation?
The Modified Model of the Atom
• The atom still consisted of a nucleus and an electron cloud.
• The electron cloud still contained levels on which electrons would be located.
• The energy levels for electrons were subdivided into sublevels: s, p, d and f. The sublevels were called orbitals.
• The first energy level has only one sublevel or orbital, the 1s orbital
• The second energy level has two types of sublevels or orbitals, the 2s and three 2p orbitals (making 4 sublevels in total).
• The third energy level has three types of sublevels or orbitals, the 3s, three 3p, and five 3d orbitals.
• The fourth energy level has four types of orbitals: the 4s, three 4p, five 4d and seven 4f orbitals.
• Each of the fifth, sixth, and seventh energy levels has similar sublevels or orbitals as the fourth energy level.
• Each orbital can hold 2 electrons at most.
• The electrons within an orbital or sublevel have opposite spins.
• For any orbitals having the same energy, one electron goes into each orbital before you double up.
• The order of electrons filling into sublevels is in order of lowest energy to highest energy.
• The sublevels overlap as you move further from the nucleus.
• The order of filling electrons is not as straightforward as it may seem. This is because the energy levels start to overlap as you move further from the nucleus.
Orbital Representation
Orbital Representation is a way for chemists to show where the electrons in an atom are located. To show the orbital representation, chemists use the atomic orbital chart sometimes called an Aufbau diagram. This atomic orbital chart can be designed using lines, squares or circles to indicate each of the orbitals (sublevels).
Explanation
More Dead Chemists!
In 1926, Wolfgang Pauli (1900-1958) determined the assignments of electrons in various orbitals. Known as the Pauli exclusion principle, it formed the basis of the electronic structure in atoms. What this principle states is that each atomic orbital holds at most two electrons. The ground-state electronic structures are then obtained following Pauli's principle: by arranging the atomic orbitals in order of increasing energy and filling them one electron at a time starting from the lowest-energy level. This process of starting at the bottom and working up is commonly known as the Aufbau principle, Aufbau being derived from German, meaning "building-up". A group of orbitals with the same energy (such as the three different 2p orbital) make up a sub-shell, and a group of sub-shells with similar energy compromise a shell.
Friedrich Hund (1896-1997) suggested that if more than one orbital in an orbit is available, electrons will fill empty orbitals before pairing in one of them. This is known today as Hund's rule.
To summarize, the ground-state electron configuration is obtained simply by "filling-up" orbitals of the lowest energy in accord with the Pauli exclusion principle and Hund's rule.
MOLECULE LEWIS DIAGRAMS
A Lewis diagram depicts a mmolecule using an element symbol to represent the nucleus and core electrons of each atom. Valence electrons are represented by lines for electron pair bonds and dots for unbonded electrons.
The following procedure can be followed to derive Lewis diagrams for most molecules.
1. Find the total number of electrons:
Tabulate the total number of outer energy level electrons for all atoms in the molecule. For each atom, read the group number.
2. Draw a first tentative structure:
The element with the least number of atoms is usually the central element. Draw a tentative molecular and electron arrangement attaching other atoms with single bonds as the first guess. Single bonds represented with a line represent 2 electrons
3. Add electrons as dots to get octets around atoms:
When counting electrons for the octet around an atom, count both electrons in a bond for each atom and any lone pair electrons. Hydrogen, of course, gets only 2 electrons.
4. Count the total number of electrons in the final structure to see if the total agrees with the number tabulated in step #1. If not, then move a lone pair of electrons into a double bond. Or add more lone pairs of electrons.
5. Cycle through steps 3 and 4 several times until you get it right by trial and error.