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Chapter #9. Atomic Theory Quantum Model of the Atom. Physics Review. J.J. Thompson determined all mater contains electrons with cathode ray tube experiments. Rutherford demonstrated that atoms contain mostly empty space by bombarding metals with alpha particles (helium nuclei)
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Chapter #9 Atomic Theory Quantum Model of the Atom
Physics Review • J.J. Thompson determined all mater contains electrons with cathode ray tube experiments. • Rutherford demonstrated that atoms contain mostly empty space by bombarding metals with alpha particles (helium nuclei) • Previously we talked about the plumb pudding model of an atom and Rutherford’s planetary model. • Then along came Bohr who applied Planck’s the ideas of energy quantization to Rutherford’s planetary model of an atom.
Evidence for Quantization of Energy The circular orbits of the Bohr model are characterized bythe principle quantum number, n, which has positive integer values. e.g.1, 2, 3….
About Energy Energy is anything that has the capacity to do work. Energy Examples: Food; how much energy for a heart beat? Gasoline; how many $ goes to heat and how much to work? Dynamite; what can it move? Electricity; what can it move? Apple on a tree; what can it move?
The Bohr Model of an Atom Rutherfordproposed the electrons were located in orbits around the nucleus similar to planets around a star. This is sometimes called the planetary model, which is not the modern day model of an atom. Li atom 3p, 3n, 3e-
The Bohr Model of an Atom Neils Bohr developed a mathematical model based on Rutherford's proposal using Planck’s quantized energy levels. In Bohr’s model electrons could only have fixed Potentialenergy levels.
Classification of Electromagnetic Radiation White light is the collection of all of the colors in the rainbow. When white light is passed through a glass prism the waves are bent by the glass. The shortest (most energetic) are bent the most.
Potential and Kinetic Energy Potential energy; energy due to position above the ground. Kinetic energy; energy due to matter in motion. Electromagneticenergy; energy possessing both particle and wave properties traveling at the speed of light. Units of energy, joule, or calorie.
Properties of Waves Wavelength (m) Amplitude (m) Speed 3.0X108 m/s Energy (j) Frequency(1/s, Hz)
Electromagnetic Radiation 100 10-2 102 10-4 10-8 10-7 10-10 10-12 Micro Waves (Radar) Infrared rays Radio and Television waves Uv-rays Visible rays X- rays Gamma rays Increasing wave length in meters Increasing energy
Line Spectrum Evidence for Quantization of Energy The red-orange light from hydrogen gas passes through a prism to form a line spectra. Each different colored light has its own unique energy. This is called an emission spectra
Additional Emission Line Spectra Note: Each element has its own unique emission spectra (Finger print)
Emission vs. Absorption Spectra Early on physicists theorized that light emitted by the sun should be a continuous spectrum. They were troubled by black lines when observing the suns spectrum through a glass prism. Bunsen (Bunsen burner) and Kirchoff studied emission spectra from emission tubes containing various gaseous elements. Interesting the wavelengths of the black lines match know wave lengths of different elements. This lead physics to conclude that gases surround the sun that absorb the emitted wave lengths of the sun.
What is Quantization Quantization means no in between Energy levels in atoms are quantized. Anything that comes in units such as: stairs television channels, gears, and bookshelves are quantized. A turtle on stairs may only be at specific heights. Its potential energy is quantized.
Quantization Are you quantized right now? Coming to class you increase your potential energy one quanta (step) at a time.
Evidence for Quantization of Energy The circular orbits of the Bohr model are characterized bythe principle quantum number, n, which has positive integer values. e.g.1, 2, 3….
Evidence for Quantization of Energy In terms of the Bohr model absorption and emission looks like this.
Evidence for Quantization of Energy Electrons move between energy levels by absorbing and emitting energy in the form of light. We call the lowest energy level the ground state. The higher energy level is called the excited state.
Evidence for Quantization of Energy The Bohr model works well for the hydrogen atom which has only one electron but performs poorly for more complex atoms. This led to the development of the current quantum mechanical model describing the arrangement of electrons in atoms. Unfortunately for us this model is more complex than that developed by Neils Bohr.
Quantum Numbers • The location of an electron in an atom can be described in order of precision by its: • shell (a positive integer given the symbol n = 1,2,3,….) • subshell (Orbital) (designated by letters s, p, d or f) • orientation (orbital symmetrical on x, y, or z axis) • Spin (clockwise or counter clockwise) +1/2 Like an electrons Social Security number: 4px
Quantum Numbers As the value for n of a shell increases its energy and distance from the nucleus increases. This is similar to the Bohr model. Each shell has a number of subshells equal to its value for n (up to a maximum of 4). e.g. A shell with n = 1 will have one subshell (s) A shell with n = 2 will have two subshells (s,p) A shell with n = 3 will have three subshells (s,p,d) A shell with n = 4 will have four subshells (s,p,d,f) A shell with n = 5 will have four subshells (s,p,d,f,g)
Quantum Numbers Each subshell is designated with the letter s, p, d or f. The subshells are named by putting the value of n in front of the symbol for the subshell. e.g. a p subshell in the second shell is named 2p The subshells vary in order of energy s < p < d < f. The difference in energy between subshells is much smaller than the difference in energy between shells. Just like the space between buildings is smaller than the space between streets.
Quantum Numbers • As mentioned earlier each shell has a number of subshells equal to its value for n. • Therefore for: • n = 1 there will be one subshell 1s • n = 2 there will be two subshells 2s and 2p • n = 3 there will be three subshells 3s, 3p and 3d • n = 4 there will be four subshells 4s, 4p, 4d and 4f • n = 5 there will be four subshells 5s, 5p, 5d and 5f
Quantum Numbers Each subshell is made up of one or more orbitals. An orbital is a volume of space where an electron is likely to be found. What is the orbital for books called? Fish? Cars? It is important not to confuse an orbit (a circular path on which an electron moves in the Bohr model) with an orbital. They are two very different things.
Quantum Numbers An s subshell has one orbital which is spherically shaped. If you were to measure where the electron was within an s subshell many, many times and plot the results on a graph you would get something like this.
Quantum Numbers The p-orbital is next in energy after the s-orbital. The p-orbital has a dumbbell shape with electrons located either side of the nucleus in tear drop shaped lobes. There are three types of p-orbitals all having identical shape and energy directed along the x, y and z axis.
Quantum Numbers Next in energy are the d-orbitals of which there arefive all with the sameenergy. The different d-orbitals do have different shapes as well orientation. These are followed by the sevenf-orbitals.
Quantum Numbers The way in which electrons are organized into shells, subshells and orbitals in an atom is called the electronic configuration. The electronic configuration of an atom can be determined using the“Aufbau rule”also known as the“building up principle”. Aufbau comes from the German meaning construction although it was the Danish physicist Neils Bohr who came up with the idea !!
Quantum Numbers • The Aufbau Principle states that: • “The orbitals of lower energy are filled in first with the electrons and only then the orbitals of high energy are filled.” • What is the lowest energy orbital of an atom? 1s orbital What is the third lowest energy orbital of an atom? 2p orbital
Quantum Numbers • As we have seen previously for p, d and f subshells there are multiple orbitals with the same energy, called degenerateorbitals. • In particular: • p subshells have three orbitals with the same energy • d subshells have five orbitals with the same energy • f subshells have seven orbitals with the same energy • Each of these orbitals may accommodate a maximum of twoelectrons.
Quantum Numbers If there are multiple orbitals with the same energy how do we decide which orbital to put an electron? We use Hund’s rule which states: “Electrons fill degenerate orbitals one at a time before doubling up in the same orbital”
p subshell px py pz p subshell px py pz Quantum Numbers Using Hund’s rule how would we put three electrons in a p subshell ?
p subshell px py pz p subshell px py pz Quantum Numbers When we do put two electrons in one orbital then they obey the Pauli exclusion principle. “only electrons with opposite spin can occupy the same orbital”
2s subshell 1s subshell Quantum Numbers How would we use our rules to “build up” the electron configuration of a Li atom? Li has Z = 3 so has 3 e-. We can write this in shorthand as 1s22s1
2p subshell 2s subshell 1s subshell Quantum Numbers How would we use our rules to “build up” the electron configuration of a N atom? N has Z = 7 so has 7 e-. We can write this in shorthand as 1s22s22p3
Quantum Numbers Beyond the 3p subshell the orbitals don’t fill in an obvious way. For example the 4s level lies lower in energy than the 3d.
Diagonal Rule There is an easy way to remember the sequence of the energies of the subshells.
Quantum Numbers So now we have everything we need to determine the electronic configuration of any atom What is the electronic configuration of a Na atom? Start here Z = 11, Na has 11 e- 3s1 1s2 2s2 2p6
Quantum Numbers The electrons in the highest energy shell of an atom are (those furthest from the nucleus) are called the valence electrons. These electrons are very important as when atoms interact with each other it is through their valence electrons.
PERIODIC TRENDS • Atomic size • Increases from top to bottom (gaining a new outer shell) • Decreases left to right in a period (Increase of protons attracts electrons stronger, thus contracting the element. • Ease of ionization (relative ease of losing electrons) • The larger the atom, the further electrons are from nucleus, and the less they are held by nucleus. • Metals and nonmetals • Metallic character (increases going away from metals) • Nonmetallic character
How many electrons fill the following sublevels? a. The 3rd level. b. The 4th level. c. The 5th level
