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Transcript of Unit 3 – The Electron Chapter 5 Test:. Review of Atomic Theory Bohr determined that the e - travel...
![Page 1: Unit 3 – The Electron Chapter 5 Test:. Review of Atomic Theory Bohr determined that the e - travel around the nucleus according to energy Electrons must.](https://reader036.fdocuments.us/reader036/viewer/2022062519/5697bfd81a28abf838caeedc/html5/thumbnails/1.jpg)
Unit 3 – The Electron
Chapter 5
Test:
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Review of Atomic Theory
Bohr determined that the e- travel around the nucleus according to energyElectrons must have energy to keep
them away for the nucleus (opposites attract!)
Closer to nucleus the lower the energyEnergies are also observed with the
speed the electrons orbit the nucleus
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Atomic Orbitals
Orbitals - describes the as clouds where the e- will be found 90% of the time (Quantum Mechanical Model)3-D region of probabilityElectron cloud gives volume of the
atom Each atomic orbital has its own
general size – not defined as Bohr suggested
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Principal Quantum Numbers
Model assigns principal quantum numbers (n) that indicate relative size & energies of the orbitals On periodic table each row is an energy leveln specifies the energy levels (principal
energy levels)• Lowest energy level assigned n =1• When electron is in that level it is at its’ ground
state
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Energy Levels
Principal energy levels contain energy sublevels There are 4 energy sublevels – s, p, d, and f and
are given by the shapes the orbitals make. Each sublevel contains a different number of
orbitals Each orbital can only contain 2 electrons
Sublevel Description of Spectrum
# orbitals Max # e-
s Sharp 1 2
p Principal 3 6
d Diffuse 5 10
f Fundamental 7 14
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Organization of Energy Levels
PrincipalQuantum #n (energy level)
Sublevel sSpherical shaped
Sublevel pdumbbell shaped
Sublevel dSublevel f
7 orbitals5 orbitals3 orbitals1 orbital
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Orbital Shapes
Figures 5-15 & 5-16 on page 133 of Text Book
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Orbital Shapes
Image from http://www.chemcomp.com/journal/molorbs.htm
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Locating Electrons
We have two ways to show where the electrons are found in the atom Orbital filling diagrams Electron configurations
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Orbital Filling Diagrams
Show how the electrons fill into the orbitals Each box or circle represents an orbital
which can hold a max of 2 electrons Electrons must fill all of one energy sub-
level before starting into another Electrons are notated with an arrow (up or
down) Up arrows must fill the boxes first then
double up with the down arrows Arrows represent the spin of the electrons
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Orbital Filling Diagrams
Figure 5-17 on page 135 of Text Book
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Orbital Filling Diagrams
The three p orbitals fill in the order shown below:
The number of arrows must match the number of electrons contained in the atom Example: Carbon has six electrons
Page 136 in Text Book
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Electron Configuration
Shorthand method for describing the arrangement of electrons
Composed of the principal energy level followed by the energy sublevel and includes a superscript with the # of electrons in the sublevel
He 1s2
Energy Level Sublevel
# electrons in sublevel
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Putting it all together
Neon Atom Electron Configuration: 1s22s22p6
Orbital Filling Diagram Orbital image:
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Electron Configuration Shorthand
Give the symbol of the noble gas in the previous energy level in brackets
Give the configuration for the remaining energy level
Example:
Sulfur 1s22s22p63s23p4
[Ne]3s23p4
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Valence Electrons
Valence electrons: found in the outermost energy level
Electrons used for bonding Represented visually in Lewis-Dot
Structures Example: Carbon 1s2 2s2 2p2
Add up the number of electrons (superscripts) in the highest energy level
So, carbon has 4 valence electrons
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Lewis-Dot Structures
Element’s symbol represents the nucleus and inner-level electrons
Dots represent the valence electrons Dots are placed one at a time on the four
sides of the symbol then paired until all valence electrons are used.
Maximum of 8 electrons will be around the symbol
d sublevel electrons are not valence electrons – they are in a lower energy level!
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Lewis-Dot Structures
Page 140 in Text Book
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Ions
Atoms that have gained or lost electrons
The word atom implies that it is neutral! Denoted by a superscript charge (sign
and number) to the right of the element symbolExamples: Cl- and Mg2+
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Ions
Cation – positive ion (Cat Ions are Pawsitive) Loses electrons to become more positive Example: Be 1s22s2 → Be2+ 1s2
Anion – negative ion Gains electrons to become more negative Example: F 1s22s22p5 → F- 1s22s22p6
What do you notice about the ions’ electron configurations?
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Ions- Practice
Determine the # of protons, neutrons, and electrons and name as a cation or anionK1+
Cl1-
O2-
Mg2+
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Lewis Dot Diagrams - Ions
Metals lose electrons to form cations Nonmetals gain electrons to form
anions Set-up diagram using the ions
electron configuration, then place brackets around the diagram with a superscript of the charge
Example: O2- 1s22s22p6 [ O ]2-
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Light
Electromagnetic Radiation (light) is a form of energy with a wavelike nature
Visible light is only a small portion of the electromagnetic spectrum
Wave model does not explain all of light’s behavior Use a dual wave-particle model Explains why chemicals give off certain
colors of light when heated in a flame
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How Light is Emitted
As energy is absorbed (heat gained) the electrons move from their ground state to an excited state (higher energy level)
) ) ) ) ) ) ) 1 2 3 4 5 6 7
nucleus
initialposition
energy levels
ground state
Absorbs energy
finalposition
excited state
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How Light is Emitted
Electrons’ unstable in excited state Returns to ground state by releasing energy (light
quantum) Color of light determined by the # of energy
levels moved & amount of energy electron had
) ) ) ) ) ) ) 1 2 3 4 5 6 7
nucleus
finalposition
energy levels
ground state
Releases energy – gives of light
initialposition
excited state
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Visible Light
(high energy) V iolet I ndigo
B lue G reen
Y ellow O rangeR ed ← (low energy)
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Electromagnetic Spectrum
Figure 5-5 on Page 120 in Text Book
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Emission Of Light
Max Planck described the emission spectrum of objects that were heated, from this we get the following terms: Quantum – minimum amount of
energy that can be gained or lost by an atom (can be referred to as a packet of energy)
Photon – packet of light energy (light quantum), has wave & particle properties
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Spectra
Emission Spectra - Series of colored lines used to identify an element (each element has different spectrum) Shows all the wavelengths of light that are emitted
Spectroscope – instrument used to see the emission spectra
Absorption Spectra – Opposite of emission spectra Shows all the wavelengths of light that are
absorbed
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Spectra
Emission Spectra
Absorption Spectrum
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Photoelectric Effect Phenomenon where electrons are emitted from
a metal’s surface when light of a certain frequency shines on the surface
Solar panels use this to generate electricity (solar calculators too!)
Page 123 in Text Book
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End of Unit 3 Notes
Study for Test on