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**Recap of Last Class Recall:**

Werner Heisenberg formulated the Uncertanity Principle that states it is impossible for us to know an electron’s exact position (where it is) and momentum (where it is going) As a result, we cannot identify specific orbits that electrons travel in We can only identify regions of space within an atom where an electron is most likely to be found ORBITALS! Schrodinger’s complex math equation allows us to: Calculate the shape of the electron cloud Probability of finding the electron at distinct locations within those clouds

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**How do the Orbitals Fill Up with Electrons?**

An Introduction to Electron Configurations

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**Assigning an Electron’s Address Explore**

Complete the activity “Welcome to Atomos Apartments!” on page 208

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**Predicting Electron Locations**

We use electron configurations The way electrons are arranged in atoms There are rules to follow! Aufbau principle Electrons are added one at a time to the lowest energy orbitals available until all the electrons of the atom have been accounted for “aufbau” German for ‘build up or construct’

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aufbau chart 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f

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**Predicting Electron Locations**

Pauli’s Exclusion Principle An orbital can hold only two electrons

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**Predicting Electron Locations**

Hund’s Rule “Electrons must fill a sub-level such that each orbital has a spin up electron before they are paired with spin down electrons” A bus analogy: If you enter a bus and don’t know anyone on it, you will pick a seat that is completely empty rather than one that already has a person in it

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**Orbital Diagrams and Electron Configurations**

Electrons fill in order from lowest to highest energy The Pauli exclusion principle holds. An orbital can hold only two electrons Two electrons in the same orbital must have opposite signs (spins) You must know how many electrons can be held by each orbital 2 for s 6 for p 10 for d 14 for f Hund’s rule applies. The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons for a set of orbitals By convention, all unpaired electrons are represented as having parallel spins with the spin “up”

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**Electron Configuration Practice**

Just a thought… How do you determine the number of electrons in an element? Examples: Oxygen Magnesium Argon Scandium

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Orbital Notation

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Short-Hand Notation Use the Noble Gas symbol to abbreviate or shorten the electron configuration Krypton Rubidium Zirconium

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**How Can We “Locate” an Electron?**

Use Quantum Numbers!

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Quantum Numbers Each electron has a specific ‘address’ in the space around a nucleus An electrons ‘address’ is given as a set of four quantum numbers Each quantum number provides specific information on the electrons location

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**Electron Configuration**

state town house number street

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**Quantum Numbers state (energy level) - quantum number n**

town (sub-level) - quantum number l street (orbital) - quantum number ml house number (electron spin) - quantum number ms

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**Principal Quantum Number (n)**

Same as Bohr’s n Integral values: 1, 2, 3, …. Indicates probable distance from the nucleus Higher numbers = greater distance from nucleus Greater distance = less tightly bound = higher energy

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**Angular Momentum Quantum Number (l)**

Integral values from 0 to n - 1 for each principal quantum number n Indicates the shape of the atomic orbitals Table 7.1 Angular momentum quantum numbers and corresponding atomic orbital numbers Value of l 1 2 3 4 Letter used s p d f g

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**Magnetic Quantum Number (ml)**

Integral values from l to -l, including zero Relates to the orientation of the orbital in space relative to the other orbitals 3-D orientation of each orbital

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**Magnetic Quantum Number**

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**Electron Spin Quantum Number (ms)**

An orbital can hold only two electrons, and they must have opposite spins Spin can have two values, +1/2 and -1/2 Pauli Exclusion Principle (Wolfgang Pauli) "In a given atom no two electrons can have the same set of four quantum numbers"

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Closer! Complete the Closer on Page 206

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Homework Begin homework on page 209 – FRONT AND BACK!

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