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# Unit 10 - PowerPoint PPT Presentation

Unit 10. Electrons in Atoms. Light as Energy. Wave Theory of Light : James Clerk Maxwell (mid-1800s) Electromagnetic radiation – a form of energy that exhibits wavelike behavior as it travels through space

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### Unit 10

Electrons in Atoms

Wave Theory of Light:

• James Clerk Maxwell (mid-1800s)

• Electromagnetic radiation – a form of energy that exhibits wavelike behavior as it travels through space

• Visible light – a form of electromagnetic radiation that is perceivable to human beings and is seen in the colors of the rainbow (ROY G. BIV)

• Crest – the top of a wave

• Trough – the bottom of a wave

• Wavelength – the distance from crest to crest or trough to trough in a wave

• Frequency – the number of wavelengths that pass a given point in a set amount of time

• Origin – the center/start or midpoint of a wave

• Amplitude – the distance from the origin to the crest or the trough of a wave

• Speed of light – c – the rate at which all forms of electromagnetic radiation travel through a vacuum = 3.00 x108 m/s

Light as a Wave

A. Wavelength ( “lambda”)

• Distance between corresponding parts on a wave (m)

B. Frequency( “nu”)

• # of peaks that pass a given point each second (cycles/sec  Hz)

B. Frequency( “nu”)

• # of peaks that pass a given point each second (cycles/sec  Hz)

• High frequency = Short Wavelength

• Low frequency = Long Wavelength

C. Wave Velocity (speed) = c

• Distance that a peak travels in a unit of time (generally one second)

• Units: m/s

• Speed of light = c = 3.00 x 108 m/s

D. Amplitude

• The height of a wave (measured in meters)

• Wavelength = speed of light / frequency

• Wavelength and frequency are inversely proportional

c = 

Example #1: The wavelength of the radiation which produces the yellow color of sodium vapor light is 5.89 x 10-7 m. What is the frequency of this radiation?

c = 

Example #2: What is the wavelength of a microwave having a frequency of 3.44 x 109 Hz?

ROY G. BIV

Can be separated by a prism

Visible Light

• All forms of light both visible and non-visible

• Ranges from radio waves to gamma rays

• First studied by Max Planck around 1900

• Came about to try and explain the parts of light behavior that didn’t fit Maxwell’s wave theory

Photoelectric Effect:

- Light is emitted in small discrete amounts called “photons”

• Photon: a single “particle” of light

• When an electron absorbs a photon of energy, the electron jumps from the ground state to its excited state

• Ground state – lowest energy level an electron occupies

• Excited state – temporary state when an e- is at a higher energy level

• quantum: the minimum amount of energy that can be gained or lost by an atom

• Photon: a single “particle” of light

• Planck’s Constant = h = 6.626 x 10-34 J.s

• Energy and frequency are directly proportional

• They both increase or decrease together

• SO…. Energy and wavelength are inversely proportional

Example #1: What is the energy of a photon from the violet portion of the rainbow if it has a frequency of 7.23 x 1014 s-1?

Example #2: What is the energy of a photon from the green portion of the rainbow if it has a wavelength of 4.90 x 10-7 m?

• Albert Einstein – 1905

Light is both a wave and a particle!!!!

• Demonstrations:

• Light bends

• The set of frequencies of light that are emitted when atoms of a gas are energized with electricity

• Each atom has its own distinct AES

• Intensive Property

• Hydrogen Atom Spectrum

• Neils Bohr extended his “Planetary Model” of the atom to address light

• Studied the Hydrogen atom and how different light was emitted when the atoms electrons moved between energy levels

• Bohr stated that electrons exist in set pathways called energy levels

• Energy levels are represented by the letter n

• So, the first energy level is n = 1,etc.

• The first energy level is closest to the nucleus, etc.

• Bohr concluded that the color of light depends on the fall of the electron from the excited state to the ground state

• Balmer, Lyman and Paschen each studied specific energy levels and the light emitted when those energy levels are involved

• Both the velocity and position of an electron cannot be determined simultaneously

• You can determine either the speed OR the location at a given time but not both

• Since an electron is so small and the methods for determining position and velocity can’t be done simultaneously, doing one creates error in the other

• Erwin Schrodinger – 1926

• Explained electron movement as wavelike rather than set as in the Bohr model (remember from Chapter 4)

• Gave electrons regional locations called orbitals and based them on probability calculations

• Electron locations are designated by a system of letters and numbers called “quantum numbers”

• These numbers give you the location of an electron ranging from the most general location to the most specific location

• The most general of the information

• Whole numbers from 1,2,3,4, etc.

• Remember 1 is closest to the nucleus, 2 next, etc.

• The principal energy level is the same as the row number of the periodic table for a given element

• Example:

Valence electrons in calcium reside in the n = 4 energy level

• One step more specific than the energy level

• Letters s, p, d, and f

• Energy level 1 has only sublevel s

• Energy level 2 has s and p

• Energy level 3 has s, p, and d

• Energy level 4-7 have s, p, d, and f

• The most specific piece of information is about the number and location of the electrons within the sublevel

• Can be represented by an exponent or arrows depending on the method of notation

• The s sublevel has 1 orbital

• The p sublevel has 3 orbitals

• The d sublevel has 5 orbitals

• The f sublevel has 7 orbitals

• Orbital - region within a sublevel where an e- can be found (homes for e-)

• Aufbau Principle: electrons fill from the lowest energy level to the highest (they don’t skip around)

• Pauli ExclusionPrinciple: each orbital can hold a maximum of 2 electrons at a time (they must have opposite spins)

• Hund’s Rule: orbitals of equal energy in a sublevel must all have 1 electron before the electrons start pairing up

• The system of numbers and letters that designates the location of the electrons

• 3 major methods:

• Full electron configurations

• Abbreviated/Noble Gas configurations

• Orbital diagram configurations

Example Notation:

• 1s2 2s1 (Pronounced “one-s-two, two-s-one”)

A. What does the coefficient mean?

Principle energy level

B. What does the letter mean?

Type of orbital (sublevel)

C. What does the exponent mean?

# of electrons in that orbital

1. Determine the total number of electrons the atom has (for neutral atoms it is equal to the atomic number for the element).

Example: F

atomic # = # of p+ = # of e- =

2. Fill orbitals in order of increasing energy.

3. Make sure the total number of electrons in the electron configuration equals the atomic number.

When writing electron configurations:

• d sublevels are n – 1 from the row they appear in

• f sublevels are n – 2 from the row they appear in

Write electron configurations for the following elements:

He:

P:

Rh:

Br:

Ca:

Ce:

i. Where are the noble gases on the periodic table?

ii. Why are the noble gases special?

iii. How can we use noble gases to shorten regular electron configurations?

Example: Barium

1. Look at the periodic table and find the noble gas in the row above where the element is.

2. Start the configuration with the symbol for that noble gas in brackets, followed by the rest of the electron configuration.

Practice! Write Noble Gas Configurations for the following elements:

Rubidium:

Bismuth:

Arsenic:

Zirconium:

• Recall there were three values we talked about which indicated the location of an electron:

• Principal energy level: n

• (row of the periodic table)

• Sublevel: s, p, d, f

• Orbital: each orbital can hold up to 2 e-

• These two electrons form an “electron pair” which must have opposite spin (point in opposite directions)

• Principal energy level: n (= row of periodic table)

• Sublevel: s, p, d, f

• Orbital: each orbital can hold 2 e-

Think of these in terms of a hotel:

• Floor – most general; hotels have many floors

• (like the energy level in an atom)

• Wing – each floor has a few wings or corridors

• (like the sublevel in an atom)

• Room – most specific location

• (like the orbital in an atom)

Orbital diagrams use boxes (sometimes circles) to represent energy levels and orbitals. Arrows are used to represent the electrons.

= orbital

sublevels

Don’t forget - orbitals have a capacity of two electrons!! Two electrons in the same orbital must have opposite spin so draw the arrows pointing in opposite directions.

Example: oxygen 1s22s22p4

2p

Increasing Energy 

2s

1s

• First, determine the electron configuration for the element.

• Next draw boxes for each of the orbitals present in the electron configuration.

• Boxes should be drawn in order of increasing energy (see the Aufbau chart).

• Arrows are drawn in the boxes starting from the lowest energy sublevel and working up. This is known as the Aufbau principle.

• Add electrons one at a time to each orbital in a sublevel before pairing them up (Hund’s rule)

• The first arrow in an orbital should point up; the second arrow should point down (Pauli exclusion principle)

• Double check your work to make sure the number of arrows in your diagram is equal to the total number of electrons in the atom.

• # of electrons = atomic number for an atom

Practice! Draw orbital diagrams for the following elements:

Sulfur:

Nickel: