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**Chapter 5 Electrons in Atoms**

Mr. Samaniego Lawndale High School

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**Summary of Atomic Theory**

Year Event 400BC Democritus proposes idea of atom 1808 Dalton develops Atomic Theory 1897 Thomson uses cathode ray to discover electron 1916 Millikan measures the mass of an e- 1919 Rutherford uses gold foil experiment to discover nucleus

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Structure Of An Atom So by this point, we know that protons and neutrons are located in the nucleus and electrons are around the outside of the nucleus

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**Section 5.1 – Models of the Atom**

In 1897 J. J. Thomson discovered the electron Observed that a magnet deflected the straight paths of the cathode rays

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Atoms were known to be electrically neutral which meant that there had to be some positively charged matter to balance the negative charges

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Ernest Rutherford’s experiment disproved the plum pudding model of the atom and suggested that there was a positively charged nucleus because most of the alpha particles went straight through the gold foil BUT, Rutherford’s atomic model could not explain the chemical properties of elements

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**The Bohr Model (Niels Bohr)**

In 1913, Niels Bohr came up with a new model (Bohr was a student of Rutherford) He noticed that light given out when atoms were heated always had specific amounts of energy, so Niels Bohr proposed a model that electrons in an atom must be orbiting the nucleus and can reside only in fixed energy levels

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Energy Levels Each of the electrons in Bohr’s model has a fixed amount of energy called energy levels This is similar to steps of a ladder (can climb up the ladder, but cannot step in between the steps) Quantum is the amount of energy required to move an electron from one energy level to another The further away from the nucleus, the more energy the electron has

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**The Quantum Mechanics Model (Erwin Schrodinger)**

While Rutherford’s model described the path the electron moves, Erwin Schrodinger solved mathematical equations to describe the behavior of electron Similar to Bohr’s model, Schrodinger describes the energy of electrons with certain values but does not involve an exact path the electron takes around the nucleus

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**The Quantum Mechanics View of the Atom (Schrodinger)**

The quantum mechanical model does not describe the exact path an electron takes around the nucleus, but determines the probability of finding an electron in a certain area

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**Quantum Mechanical Model**

In this model, electrons move similar to a rotating propeller blade You cannot tell its precise location at any instant because it’s a blurry region, but you have information regarding the probability of finding an electron within a certain volume of space Similar to a fuzzy cloud…the probability of finding an electron is higher where the cloud is more dense

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Atomic Orbitals An Atomic Orbital is a region of space where there is a high probability of finding an electron Each energy sublevel corresponds to an orbital of different shape describing where the electron is likely to be found (there are 4 different types of shapes)

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Shapes of Orbitals

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Shapes of Orbitals

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Shapes of Orbitals

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**Chapter 5.2 – Electron Arrangement in Atoms**

Each electron in an atom is assigned a set of four quantum numbers. These help to determine the highest probability of finding the electrons. Three of these numbers (n, l, m) give the location of the electron The fourth (s) describes the orientation of an electron in an orbital.

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**Quantum letters can be thought of like the numbers and letters on a concert ticket**

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**Labeling Electrons in Atoms**

Probable Location of electron Probability Probable location of Finding Beyonce or Taylor Energy level (n) High Probability Hotel Name Sublevel (l) Higher Probability Hotel Floor Number Orbital (m) Highest probability Hotel Room Number

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**Electron Configuration**

Sublevel # of Orbitals Available # of Electrons Available s 1 2 p 3 6 d 5 10 f 7 14

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**Electron Configurations**

Filling order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p Example He = 2 electrons 1s2 Example Li = 3 electrons 1s22s1 Example B = 5 electrons 1s22s22p1

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**Practice Problems Li 5. P N 6. Si Be 7. Mg C 8. Al**

Write electron configurations for the following atoms Li P N Si Be Mg C Al

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**Electron Configurations can be written in terms of noble gases**

To save space, configurations can be written in terms of noble gases Example 1: Ne = 1s22s22p6 S = 1s22s22p63s23p4 Or S = [Ne] 3s23p4 Example 2: Ar = 1s22s22p63s23p6 Mn = 1s22s22p63s23p64s23d5 Mn = [Ar] s23d5

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**Reading the Periodic Table**

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**Locating Electrons in Atoms**

So far we have discussed 3 quantum numbers n= principal quantum level (principal energy level) l= Sublevel m = magnetic quantum number (shape of orbitals) 1s2 n Number of electrons in sublevel l

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**s = spin When an electron moves, it generates a magnetic field.**

s describes the direction an electron spins They must spin in opposite directions Spin= up down There are two values of s: +1/2 and -1/2

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Orbital Diagrams The electron configuration gives the number of electrons in each sublevel but does not show how the orbital of a sublevel are occupied by the electrons

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Orbital Diagrams Used to show how electrons are distributed within sublevels Each orbital is represented by a box and each electron is represented by an arrow Notice that each box is drawn higher than the last set because it has more energy Example: Boron 1s22s22p1 2p 2s 1s

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**Orbital Diagrams 2p 2s 1s 2p 2s 1s**

Steps to writing orbital diagrams:ex F (Z=9) Write the electron configuration 1s22s22p5 2. Construct an orbital filling diagram using boxes for each orbital 3. Use arrows to represent the electrons in each orbital. 2p 2s 1s 2p 2s 1s

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**Rule #1 - Aufbau Principle**

Electrons must occupy the orbital with the lowest energy first Example: Oxygen 1s22s22p4 2p 2p 2s 2s 1s 1s

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**Rule #2 - Pauli Exclusion Principle**

Orbitals can only have two electrons max The 2 electrons must have opposite spins Example: Oxygen 1s22s22p4 2p 2p 2s 2s 1s 1s

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Rule #3 - Hund’s Rule Orbitals of equal energy are each occupied by one electron before any pairing occurs Example: Oxygen 1s22s22p4 2p 2p 2s 2s 1s 1s

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**Draw orbital diagrams for these elements**

Li P N Si Be Mg C Al

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**Section 5.3 - Atomic Spectra**

When atoms absorb energy, electrons move into higher energy levels (excited state) When these electrons return to their lower energy levels, they lose energy by emitting light Atomic Emission Spectrum – the discrete lines representing the frequencies of light emitted by an element

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**Calculating Wavelength of Light**

c = speed of light (3 x 108 m/s2) = wavelength (called lambda) = frequency (called nu)

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Practice 1. Calculate the wavelength of a yellow light if the frequency is 5.10 x 1014 sec-1 or Hz. Answer = 5.88 x 10-7m 2. What is the wavelength of 1.50 x 1013 sec-1? Answer = 2.00 x 10-5m 3. What frequency is radiation with a wavelength of 5.00 x 10-8m? Answer = 6.00 x 1015 sec-1 or Hertz

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Atomic Spectra Each discrete line in an emission spectrum corresponds to one exact frequency of light emitted by the atom Ground State – lowest possible energy of the electron in the Bohr model The light emitted by an electron moving from higher to a lower energy level has a frequency directly proportional to the energy change of the electron

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**Chapter 5 Assessment Page 148**

Homework Chapter 5 Assessment Page 148 #’s 22-24, 27, 29, 30-39, 50-53, 57, 60, 68, 70-72

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