Chem Ch 5 Notes.pdf

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Chemistry

Chapter 5 Section 1

Revising the Atomic Model

Big Idea

"Electrons and the Structure of the Atom"

Energy Levels In Atoms

Limitations of Rutherford's Atomic Model

*Rutherford thought that electrons moved about the nucleus like planets about the sun

*His model could not explain the chemical properties of atoms

The Bohr

Model

*Proposed

that electrons

are only

found in

specific paths

(orbits)

around the

nucleus

Energy Levels In Atoms

*Theory worked for Hydrogen (1 e-) but not for atoms with more than one electron

The Quantum Mechanical Model

*Schrodinger devised and solved a

mathematical equation describing the behavior of electrons

*The quantum mechanical model = electrons in energy levels, but does not assign them a specific path, but a probable location instead

Atomic Orbitals

*atomic orbital = describes the probability of finding an electron at various locations around the nucleus *energy levels are labeled by principal quantum numbers (n)

*n can be 1, 2, 3, 4, 5 etc...

*Far all principal energy levels greater than 1, there are several orbitals with different shapes

*These constitute energy sub levels *The different orbitals are designated by letters

s orbital n = 1 n = 2

p orbitals

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Summary of Principal Energy Levels and Sublevels

Principal Energy Level

Number of Sublevels

Type of Sublevel

Max. Number of Electrons

n = 1

n = 2

n = 3

n = 4

1

2

3

4

1s (1 orbital)

2s (1 orbital) 2p (3 orbitals) 3s (1 orbital), 3p (3 orbitals) 3d (5 orbitals)

4s (1 orbital), 4p (3 orbitals) 4d (5 orbitals), 4f (7 orbitals)

2

8

18

32

*n = energy level

*energy level also equals the number of sub levels

*n2 = the number of orbitals in that energy level *2 elections is the max in each orbital

*Max number of e- that can occupy a principal energy level = 2n2

Useful Info!!

Classwork

*5.1 Lesson Check

*#1-7

*Development of Atomic Models

*Pg 133 #1 & 2

*Homework

*Read 5.2

*SG 5.1

Chemistry

Chapter 5

Section 2

Electron Configuration

Big Idea

Electrons and the Structure of Atoms

Electron Configurations

*Atoms, ions, and molecules will go toward the lowest possible energy state

*High energy systems are unstable

*In atoms, the electrons and the nucleus work together to form the most stable arrangement *The way in which electrons are arranged around the nucleus is called electron configurations

Electron Configurations

*There are three rules that we use to determine electron configurations

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Aufbau Principle

*Electrons enter orbitals in the lowest energy level first *This means they enter the s level first

Aufbau Principle

3d

In

cr

ea

si

ng

E

ne

rg

y

1s 2s 3s 4s 5s 6s 7s

2p 3p 4p 5p 6p 7p

4d 5d 6d

5f 6f

Notice the number alone does not determine how high the energy level is

Pauli Exclusion Principle

*An atomic orbital may describe at most two

electrons

*This means that there are only two electrons in

each orbital at most; there can be just one

however

*If there are two electrons in one orbital they

must have opposite spins (clockwise and

counterclockwise)

1s 2s 2px_{2py 2pz} 3s

Electron Configurations

*When electrons occupy orbitals of equal energy, one electron enters each orbital until all the orbitals contain one electron with parallel (the same direction) spin

Hunds's Rule

*There is short hand for showing electron configurations *Oxygen has 8 electrons

1s2 2s2 2p4

*Notice that the sum of the superscripts equals the number of electrons

1s 2s 2px _{2py 2pz} 3s

Exceptional Electron Configurations

*There are exceptions to the aufbau principal *Only happens after atomic number 23

Cu electron configuration using aufbau

Correct Cu

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To help remember the order

3d

In

cr

ea

si

ng

E

ne

rg

y

1s 2s 3s 4s 5s 6s 7s

2p 3p 4p 5p 6p 7p

4d 5d 6d

5f 6f

Chromium has 24 electrons

chp898843_700k.asf

Video!

Classwork

*Practice Problems # 8 & 9

*5.2 Lesson Check

Homework

*Read 5.3

*Do SG through 5.2

Chemistry

Chapter 5 Section 3

Atomic Emission Spectra and the Quantum Mechanical Model

Big Idea

Electrons and Atomic Structure

Light and Atomic Spectra

Wave Model

Light Consists of

Electromagnetic

Waves

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Light and Atomic Spectra

*All electromagnetic waves, in a vacuum move at the speed of 3.0 x 1010 cm/s

*Each wave cycle starts at the origin

*Amplitude ( ) - wave's hight from the origin to the crest (or the trough) (units of amplitude = cm)

*Frequency (v) - is the number of wave cycles to pass a given point per unit time (units of frequency = 1/s)

Light and Atomic Spectra

C = v

C = the speed of electromagnetic waves C = 3.0 x 108 m/s

Lambda ( ) = wavelength in cm v = frequency in 1 per second

The SI unit of cycles per second is called a Hertz (Hz) Hz is also equal to s-1 or 1/s

Light and Atomic Spectra

_{*Each color of visible light has its own characteristic}

wavelength and frequency

*If you pass light through a prism, you can separate the light into each color

*The rainbow that it shows is called a spectrum *Red light has the longest wavelength and the shortest frequency

*As wavelength increases, frequency decreases = inverse relationship

Light and Atomic Spectra

*If you have an element in its gaseous state and pass electricity through it, it will give you a specific color

*Then if you pass that light through a prism, you will get an atomic emission spectrum *You don't get a rainbow like with white light *Each element has a unique emission spectrum

Practice Problem

What is the wavelength of radiation with a frequency of 2.0 x 1013Hz?

Practice Problem

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The Quantum Concept and the Photoelectric Effect

*A man named Planck was trying to figure

out why iron changes colors as it is heated

*He figured out the following equation

*E = h x v

**v = frequency of the radiation (in 1/s)

**Plank's constant = h = 6.6262 x 10-34Js

**Radiant energy = E (in Joules)

Practice Problem

Calculate the energy of a quantum of radiant energy with a frequency of 4.25 x 1011/s.

Practice Problem

What is the frequency of a quantum of radiant energy that has an energy of 2.12 x 10-22J.

The Quantum Concept and the Photoelectric Effect

The photoelectric effect

*Metals eject electrons called

photoelectrons when light shine on them

*Specific frequencies of light must be

shone on them to get them to project

photoelectrons back at you

An Explanation of Atomic Spectra

*You need a specific amount of energy

to make the electrons in an atom

increase by one energy level

*Once you do this, you will see a

specific spectrum emission

Drop to n =1 Ultraviolet

Drop to n =3 Infrafed

Drop to n =2 Visiable

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The Heisenberg Uncertainty Principle

*States that it is impossible to know both the velocity and the position of a particle at the same time

*Very important when dealing with very small particles like electrons

*Schrodinger's quantum mechanical

description of atoms takes this uncertainty into account

Classwork

*Sample Problems #15 - 18 *Lesson Check 5.3

Homework *Finish SG