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 have energy to keep them away for the nucleus (opposites attract!) Closer to nucleus the lower the energy Energies 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 probability Electron 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 level n 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

Description of Spectrum 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

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 # electrons in sublevel He 1s2 Energy Level 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 symbol Examples: 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 anion K1+ 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 ) ) ) ) ) ) ) 1 2 3 4 5 6 7 energy levels As energy is absorbed (heat gained) the electrons move from their ground state to an excited state (higher energy level) final position initial position nucleus energy levels ) ) ) ) ) ) ) 1 2 3 4 5 6 7 Absorbs energy ground state excited state

How Light is Emitted ) ) ) ) ) ) ) 1 2 3 4 5 6 7 energy levels 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 initial position final position nucleus energy levels ) ) ) ) ) ) ) 1 2 3 4 5 6 7 excited state ground state Releases energy – gives of light

Visible Light B lue G reen Y ellow R ed ← (low energy) (high energy)  V iolet I ndigo B lue G reen Y ellow O range R 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 Study for Test on