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Chapter 13 Electrons in Atoms

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1 Chapter 13 Electrons in Atoms

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 particles John Dalton - one type of atom for each element

4 J. J. Thomson’s Model Discovered electrons
Atoms were made of positive stuff Negative electron floating around “Plum-Pudding” model

5 Ernest Rutherford’s Model
Discovered dense positive piece at the center of the atom- nucleus Electrons would surround it Mostly empty space “Nuclear model”

6 Niels Bohr’s Model He 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”

7 Bohr’s Planetary Model

8 The Quantum Mechanical Model
Erwin Schrodinger derived an equation that described the energy and position of the electrons in an atom Things that are very small behave differently from things big enough to see. The quantum mechanical model is a mathematical solution It 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 Orbitals Principal 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
d 5 10 3 7 14 4 f

13 By Energy Level First Energy Level Second Energy Level
only s orbital only 2 electrons Second Energy Level s and p orbitals are available 8 total electrons Third energy level s, p, and d orbitals 18 total electrons Fourth energy level s,p,d, and f orbitals 32 total electrons

14 By Energy Level Any more than the fourth and not all the orbitals will fill up. You simply run out of electrons The orbitals do not fill up in a neat order. The energy levels overlap Lowest 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.

17 Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s 2p
Aufbau diagram - page 367

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 Phosphorus Need 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 orbital Notice the opposite spins only 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 orbital only 11 more...

22 Increasing energy The next electrons go into the 2p orbital
3d 4d 5d 7p 6d 4f 5f The next electrons go into the 2p orbital only 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 orbital only 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 shapes 3 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

27 Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s 2p

28 Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s 2p

29 Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p
Now Bromine 1s22s22p63s23p64s23d104p5

30 Exceptional Electron Configurations

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 electrons 1s22s22p63s23p64s23d2 Vanadium - 23 electrons 1s22s22p63s23p64s23d3 Chromium - 24 electrons 1s22s22p63s23p64s23d4 expected But 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: 1s22s22p63s23p64s23d9 But the actual configuration is: 1s22s22p63s23p64s13d10 This 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 spectra The light is emitted given off Sample 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 energy Let’s look at a hydrogen atom

42 Changing the energy Heat or electricity or light can move the electron up energy levels (“excited”)

43 Changing the energy As 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

45 { { {

46 Ultraviolet Visible Infrared Further they fall, more energy, higher frequency. This is simplified the orbitals also have different energies inside energy levels All 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 electron And hitting the electron changes the frequency of the light.

50 Before After Photon changes wavelength Photon Moving Electron
Electron Changes Velocity Moving Electron Fig , p. 382

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