2 Section 13.1 Models of the Atom OBJECTIVES:Summarize the development of atomic theory.
3 Section 13.1 Models of the Atom OBJECTIVES:Explain the significance of quantized energies of electrons as they relate to the quantum mechanical model of the atom.
4 J. J. Thomson’s Model Discovered electrons Negative electron floating around“Plum-Pudding” model
5 Ernest Rutherford’s Model Discovered dense positive piece at the center of the atom- nucleus“Nuclear model”
6 Niels Bohr’s Model Move like planets around the sun. In circular orbits at different levels.Amounts of energy separate one level from another.“Planetary model”
7 Bohr’s planetary model electron cannot exist between energy levels, just like you can’t stand between rungs on ladderQuantum of energy required to move to the next highest level
8 The Quantum Mechanical Model Since the energy of an atom is never “in between” there must be a quantum leap in energy.Erwin SchrodingerMathematical solution
9 The Quantum Mechanical Model Has energy levels for electrons.Orbits are not circular.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”An area where there is a chance of finding an electron.Think of fan blades
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.Atomic orbitals - regions where there is a high probability of finding an electron.Sublevels
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 only s orbital only 2 electrons 1s2 Second Energy Levels and p orbitals are available2 in s, 6 in p2s22p68 total electrons
14 By Energy Level Third energy level s, p, and d orbitals 2 in s, 6 in p, and 10 in d3s23p63d1018 total electronsFourth energy levels,p,d, and f orbitals2 in s, 6 in p, 10 in d, ahd 14 in f4s24p64d104f1432 total electrons
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 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.Phosphorus (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 The easy way to remember 2s 2p3s 3p 3d4s 4p 4d 4f5s 5p 5d 5f6s 6p 6d 6f7s 7p 7d 7f1s22 electrons
26 Fill from the bottom up following the arrows 2s 2p3s 3p 3d4s 4p 4d 4f5s 5p 5d 5f6s 6p 6d 6f7s 7p 7d 7f1s2 2s24 electrons
27 Fill from the bottom up following the arrows 2s 2p3s 3p 3d4s 4p 4d 4f5s 5p 5d 5f6s 6p 6d 6f7s 7p 7d 7f1s2 2s2 2p6 3s212 electrons
28 Fill from the bottom up following the arrows 2s 2p3s 3p 3d4s 4p 4d 4f5s 5p 5d 5f6s 6p 6d 6f7s 7p 7d 7f1s2 2s2 2p6 3s2 3p6 4s220 electrons
29 Fill from the bottom up following the arrows 2s 2p3s 3p 3d4s 4p 4d 4f5s 5p 5d 5f6s 6p 6d 6f7s 7p 7d 7f1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s238 electrons
30 Fill from the bottom up following the arrows 2s 2p3s 3p 3d4s 4p 4d 4f5s 5p 5d 5f6s 6p 6d 6f7s 7p 7d 7f1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s256 electrons
31 Fill from the bottom up following the arrows 2s 2p3s 3p 3d4s 4p 4d 4f5s 5p 5d 5f6s 6p 6d 6f7s 7p 7d 7f1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s288 electrons
32 Fill from the bottom up following the arrows 2s 2p3s 3p 3d4s 4p 4d 4f5s 5p 5d 5f6s 6p 6d 6f7s 7p 7d 7f1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d10 7p6108 electrons
34 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
35 Write these electron configurations Titanium - 22 electrons1s22s22p63s23p64s23d2Vanadium - 23 electrons1s22s22p63s23p64s23d3Chromium - 24 electrons1s22s22p63s23p64s23d4 expectedBut this is wrong!!
36 Chromium is actually: 1s22s22p63s23p64s13d5 Why? This gives us two half filled orbitals.Slightly lower in energy.The same principal applies to copper.
37 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
38 Section 13.3 Physics and the Quantum Mechanical Model OBJECTIVES:Calculate the wavelength, frequency, or energy of light, given two of these values.
39 Section 13.3 Physics and the Quantum Mechanical Model OBJECTIVES:Explain the origin of the atomic emission spectrum of an element.
40 LightThe study of light led to the development of the quantum mechanical model.Light is a kind of electromagnetic radiation.Electromagnetic radiation includes many kinds of wavesAll move at 3.00 x 108 m/s = c
41 Parts of a waveCrestWavelengthAmplitudeOriginTrough
42 Parts of Wave - p.372 Origin - the base line of the energy. Crest - high point on a waveTrough - Low point on a waveAmplitude - distance from origin to crestWavelength - distance from crest to crestWavelength is abbreviated by the Greek letter lambda = l
43 FrequencyThe number of waves that pass a given point per second.Units: cycles/sec or hertz (hz or sec-1)Abbreviated by Greek letter nu = nc = ln
44 Frequency and wavelength Are inversely relatedAs one goes up the other goes down.Different frequencies of light are different colors of light.There is a wide variety of frequenciesThe whole range is called a spectrum, Fig , page 373
46 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.
47 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
48 Light is a Particle Energy is quantized. Light is energy Light must be quantizedThese smallest pieces of light are called photons.Photoelectric effect?Energy & frequency: directly related.
49 Energy and frequency E = h x E is the energy of the photon is the frequencyh is Planck’s constanth = x Joules x sec.joule is the metric unit of Energy
50 The Math in Chapter 132 equations so far:c = E = hKnow these!
51 ExamplesWhat is the wavelength of blue light with a frequency of 8.3 x 1015 hz?What is the frequency of red light with a wavelength of 4.2 x 10-5 m?What is the energy of a photon of each of the above?
52 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.
53 Changing the energyLet’s look at a hydrogen atom
54 Changing the energyHeat or electricity or light can move the electron up energy levels (“excited”)
55 Changing the energyAs the electron falls back to ground state, it gives the energy back as light
56 Changing the energy May fall down in steps Each with a different energy
58 UltravioletVisibleInfraredFurther they fall, more energy, higher frequency.This is simplifiedthe orbitals also have different energies inside energy levelsAll the electrons can move around.
59 What is light? Light is a particle - it comes in chunks. Light is a wave- we can measure its wavelength and it behaves as a waveIf we combine E=mc2 , c=, E = 1/2 mv2 and E = hWe can get: = h/mvcalled de Broglie’s equationCalculates the wavelength of a particle.
60 Sample problemWhat is the approximate mass of a particle having a wavelength of 10-7 meters, and a speed of 1 m/s?Use = h/mv= 6.6 x 10-27(Note: 1 J = N x m; 1 N = 1 kg x m/s2
61 Matter is a Wave Does not apply to large objects Things bigger than an atomA baseball has a wavelength of about m when moving 30 m/sAn electron at the same speed has a wavelength of 10-3 cmBig enough to measure.