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**Chapter 13 Electrons in Atoms**

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**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”.

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**Greek Idea Democritus and Leucippus**

Matter is made up of solid indivisible particles John Dalton - one type of atom for each element

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**J. J. Thomson’s Model Discovered electrons**

Atoms were made of positive stuff Negative electron floating around “Plum-Pudding” model

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**Ernest Rutherford’s Model**

Discovered dense positive piece at the center of the atom- nucleus Electrons would surround it Mostly empty space “Nuclear model”

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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”

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**Bohr’s Planetary Model**

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**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.

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**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.

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**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.

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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

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**Summary # of shapes Max electrons Starts at energy level s 1 2 1 p 3 6**

d 5 10 3 7 14 4 f

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**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

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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.

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**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.

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**Section 13.2 Electron Arrangement in Atoms**

OBJECTIVES: Explain why the electron configurations for some elements differ from those assigned using the aufbau principle.

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

Aufbau diagram - page 367

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**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

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**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

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**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...

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**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...

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**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...

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**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...

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**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

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**Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p**

Now do Oxygen.

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

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

1s22s2

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

1s22s22p4

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**Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p**

Now Bromine 1s22s22p63s23p64s23d104p5

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**Exceptional Electron Configurations**

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**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

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**Write these electron configurations**

Titanium - 22 electrons 1s22s22p63s23p64s23d2 Vanadium - 23 electrons 1s22s22p63s23p64s23d3 Chromium - 24 electrons 1s22s22p63s23p64s23d4 expected But this is wrong!!

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**Chromium is actually: 1s22s22p63s23p64s13d5 Why?**

This gives us two half filled orbitals. Slightly lower in energy. The same principal applies to copper.

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**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

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**Section 13.3 Physics and the Quantum Mechanical Model**

OBJECTIVES: Calculate the wavelength, frequency, or energy of light, given two of these values.

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**Section 13.3 Physics and the Quantum Mechanical Model**

OBJECTIVES: Explain the origin of the atomic emission spectrum of an element.

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**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.

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**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.

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**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

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**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.

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Changing the energy Let’s look at a hydrogen atom

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

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Changing the energy As the electron falls back to ground state, it gives the energy back as light

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**Changing the energy May fall down in steps**

Each with a different energy

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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.

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**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

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**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

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**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.

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**Before After Photon changes wavelength Photon Moving Electron**

Electron Changes Velocity Moving Electron Fig , p. 382

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