Unit 3 – The Electron

Chapter 5

Test:

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

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

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

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

Organization of Energy Levels

PrincipalQuantum #n (energy level)

Sublevel sSpherical shaped

Sublevel pdumbbell shaped

Sublevel dSublevel f

7 orbitals5 orbitals3 orbitals1 orbital

Orbital Shapes

Figures 5-15 & 5-16 on page 133 of Text Book

Orbital Shapes

Image from http://www.chemcomp.com/journal/molorbs.htm

Locating Electrons

We have two ways to show where the electrons are found in the atom Orbital filling diagrams Electron configurations

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

Orbital Filling Diagrams

Figure 5-17 on page 135 of Text Book

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

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

Putting it all together

Neon Atom Electron Configuration: 1s22s22p6

Orbital Filling Diagram Orbital image:

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

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

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!

Lewis-Dot Structures

Page 140 in Text Book

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+

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?

Ions- Practice

Determine the # of protons, neutrons, and electrons and name as a cation or anionK1+

Cl1-

O2-

Mg2+

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-

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

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

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

Visible Light

(high energy) V iolet I ndigo

B lue G reen

Y ellow O rangeR ed ← (low energy)

Electromagnetic Spectrum

Figure 5-5 on Page 120 in Text Book

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

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

Spectra

Emission Spectra

Absorption Spectrum

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

End of Unit 3 Notes

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