2 Section 13.1 Models of the Atom OBJECTIVES:Summarize the development of atomic theory.Explain the Quantum Mechanical model and the theory that electrons form an electron “cloud”.
3 Greek Idea Democritus and Leucippus Matter is made up of solid indivisible particlesJohn Dalton - one type of atom for each element
4 J. J. Thomson’s Model Discovered electrons Atoms were made of positive stuffNegative electron floating around“Plum-Pudding” model
5 Ernest Rutherford’s Model Discovered dense positive piece at the center of the atom- nucleusElectrons would surround itMostly empty space“Nuclear model”
6 Niels Bohr’s ModelHe had a question: Why don’t the electrons fall into the nucleus?Move like planets around the sun.In circular orbits at different levels.Have different energies and therefore orbit at different levels.Cannot exist between orbits.A quanta is the energy needed to jump to a higher energy level- “Quantum Leap”“Planetary model”
8 The Quantum Mechanical Model Erwin Schrodinger derived an equation that described the energy and position of the electrons in an atomThings that are very small behave differently from things big enough to see.The quantum mechanical model is a mathematical solutionIt is not like anything you can see.
9 The Quantum Mechanical Model Has energy levels for electrons.Orbits are not circular.They are not even ovals, they are random three-dimensional shapes.It can only tell us the probability of finding an electron a certain distance from the nucleus.
10 The Quantum Mechanical Model The atom is found inside a blurry “electron cloud”Think of fan blades spinning fast.Electrons are moving so fast that they create this kind of blur. Except that they are moving in a 3D space.
11 Atomic OrbitalsPrincipal Quantum Number (n) = the energy level of the electron.Within each energy level, the complex math of Schrodinger’s equation describes several shapes.These are called atomic orbitals - regions where there is a high probability of finding an electron.Sublevels- like theater seats arranged in sections
12 Summary # of shapes Max electrons Starts at energy level s 1 2 1 p 3 6 d51037144f
13 By Energy Level First Energy Level Second Energy Level only s orbitalonly 2 electronsSecond Energy Levels and p orbitals are available8 total electronsThird energy levels, p, and d orbitals18 total electronsFourth energy levels,p,d, and f orbitals32 total electrons
14 By Energy LevelAny more than the fourth and not all the orbitals will fill up.You simply run out of electronsThe orbitals do not fill up in a neat order.The energy levels overlapLowest energy fill first.
15 Section 13.2 Electron Arrangement in Atoms OBJECTIVES:Apply the aufbau principle, the Pauli exclusion principle, and Hund’s rule in writing the electron configurations of elements.
16 Section 13.2 Electron Arrangement in Atoms OBJECTIVES:Explain why the electron configurations for some elements differ from those assigned using the aufbau principle.
18 Electron Configurations The way electrons are arranged in atoms.Aufbau principle- electrons enter the lowest energy first.This causes difficulties because of the overlap of orbitals of different energies.Pauli Exclusion Principle- at most 2 electrons per orbital - different spins
19 Electron Configuration Hund’s Rule- When electrons occupy orbitals of equal energy they don’t pair up until they have to.Let’s determine the electron configuration for PhosphorusNeed to account for 15 electrons
20 Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p The first two electrons go into the 1s orbitalNotice the opposite spinsonly 13 more to go...
21 Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s The next electrons go into the 2s orbitalonly 11 more...
22 Increasing energy The next electrons go into the 2p orbital 3d4d5d7p6d4f5fThe next electrons go into the 2p orbitalonly 5 more...
23 Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s The next electrons go into the 3s orbitalonly 3 more...
24 Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s The last three electrons go into the 3p orbitals.They each go into separate shapes3 unpaired electrons= 1s22s22p63s23p3
25 Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Now do Oxygen.
26 Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s 2p
31 Orbitals fill in order Lowest energy to higher energy. Adding electrons can change the energy of the orbital.Half filled orbitals have a lower energy.Makes them more stable.Changes the filling order
32 Write these electron configurations Titanium - 22 electrons1s22s22p63s23p64s23d2Vanadium - 23 electrons1s22s22p63s23p64s23d3Chromium - 24 electrons1s22s22p63s23p64s23d4 expectedBut this is wrong!!
33 Chromium is actually: 1s22s22p63s23p64s13d5 Why? This gives us two half filled orbitals.Slightly lower in energy.The same principal applies to copper.
34 Copper’s electron configuration Copper has 29 electrons so we expect: 1s22s22p63s23p64s23d9But the actual configuration is:1s22s22p63s23p64s13d10This gives one filled orbital and one half filled orbital.Remember these exceptions: d4, d9
35 Section 13.3 Physics and the Quantum Mechanical Model OBJECTIVES:Calculate the wavelength, frequency, or energy of light, given two of these values.
36 Section 13.3 Physics and the Quantum Mechanical Model OBJECTIVES:Explain the origin of the atomic emission spectrum of an element.
37 If the light is not white By heating a gas with electricity we can get it to give off colors.Passing this light through a prism does something different.
38 Atomic Spectrum Each element gives off its own characteristic colors. Can be used to identify the atom.How we know what stars are made of.
39 These are called discontinuous spectra, or line spectra unique to each element.These are emission spectraThe light is emitted given offSample 13-2 p.375
40 Explanation of atomic spectra When we write electron configurations, we are writing the lowest energy.The energy level, and where the electron starts from, is called it’s ground state- the lowest energy level.
41 Changing the energyLet’s look at a hydrogen atom
42 Changing the energyHeat or electricity or light can move the electron up energy levels (“excited”)
43 Changing the energyAs the electron falls back to ground state, it gives the energy back as light
44 Changing the energy May fall down in steps Each with a different energy
46 UltravioletVisibleInfraredFurther they fall, more energy, higher frequency.This is simplifiedthe orbitals also have different energies inside energy levelsAll the electrons can move around.
47 The physics of the very small Quantum mechanics explains how the very small behaves.Classic physics is what you get when you add up the effects of millions of packages.Quantum mechanics is based on probability
48 Heisenberg Uncertainty Principle -It is impossible to know exactly the location and velocity of a particle.The better we know one, the less we know the other.Measuring changes the properties.Instead, analyze interactions with other particles
49 More obvious with the very small To measure where a electron is, we use light.But the light moves the electronAnd hitting the electron changes the frequency of the light.
50 Before After Photon changes wavelength Photon Moving Electron Electron Changes VelocityMoving ElectronFig , p. 382