Chemistry Crosby High School Chapter 5 Atomic Structure T1 Valence Electrons T2 Models of Atoms T3 Electron configuration Chemistry Crosby High School

Chapter 5 Homework Read and outline 5.2 (exclude The Heisenberg uncertainty principle and The Schrodinger wave equation) and 5.3

Chapter 5 Plan Day Class work 1 Valence electrons and electron dot structures 2 Intro to models; orbital shape 3 Electron configuration & electron lab activity 4 Electron configuration 5 Worksheet 6 Review HW due 7 Chapter 5 test

Valence Electrons The electrons furthest away from the nucleus Participate in chemical bonds. Look at the group number to determine an element's number of V.E. Group 1 -1 Group 15 - 5 Group 2 - 2 Group 16 - 6 Group 13 - 3 Group 17 - 7 Group 14 - 4 Group 18 - 8

Practice How many valence electrons do the following elements contain? B Be K Na C F N

Electron Dot Diagrams (aka Lewis Dot diagrams) Show an atoms valence electrons

How to draw a Lewis Dot diagram 1. Draw the symbol 2. Determine valence electrons 3. Starting at 12 o’clock and moving clockwise, draw one dot on each side of the symbol, then draw more dots. X

Practice as a class Lewis Dot Diagram for Carbon Nitrogen Potassium

Your turn Draw the electron dot structure for Oxygen, Lithium, and Argon

More practice Page 162, #26-28, 30, 31

Atomic Models & Orbital Shape Key Topics: Bohr’s model Quantum mechanical model Principal energy levels

The Atomic Model as of the early 1900s It was Rutherford’s model Positively charged core (nucleus) in the middle of empty space. Electrons surrounded the nucleus Neutrons not discovered until 1932 by Chadwick

Revising the Atomic Model Limitations of Rutherford's Model: It could not explain the chemical properties of elements. For example, his model could not explain why metals give off characteristic colors when heated in a flame

Neils Bohr (1913) Physicist Developed a new atomic model: Electrons are found in specific paths, or orbits, around the nucleus Student of Rutherford’s.

Bohr’s Model Each orbit has a fixed energy Energy increases as you move away from nucleus Think of it like a ladder: lowest rung has lowest energy; climb the rungs to gain energy. Remember: valence electrons participate in chemical bonds….it is because they have the most energy Electrons cannot exist between energy levels

Energy Levels Fixed energies an electron can have Quantum = Energy needed to move an electron from one energy level to another

Flame Test Ground state vs. Excited state

Schrodinger (1926) Developed the Quantum Mechanical Model Describes the probability that electrons will be found in certain locations Most dense where probability is high and least dense where probability is low Unlike the Bohr model, this model does not specify an exact path the electron takes

Concept Check How can electrons in an atom move from one energy level to another? By losing or gaining an amount of energy – a quantum

Youtube Video https://www.youtube.com/watch?v=rcKilE9CdaA&list=PL8dPuuaLjXtPHzzYuWy6fYEaX9mQQ8oGr&index=5 (6 minutes of video)

Review Thomson Rutherford Bohr Schrodinger Subatomic Particles in the model Draw a picture of the model Write a 1 sentence summary of the model Thomson Rutherford Bohr Schrodinger

Quantum Mechanical Model and Energy Levels Instead of restricting electrons to limited circular orbits, Schrodinger developed a new idea of orbitals Orbitals describe the probability of finding an electron at various locations around the nucleus. Schrodinger’s numbers: size, shape and orientation. We will only cover size and shape

Probability of Finding an electron The region of space in which the electron can probably be found is called an orbital. In this lab, you will use a felt-tip maker and a target to investigate the probability distribution of marks about a central point.

What you will do: Place the target on a notebook or a stack of papers to create a cushioned landing. 2. Drop the marker 100 times from hip-height 3. Create a chart and record the number of marks in each region of the target. 4. Create a graph of your marks. Draw a smooth curve that shows the general trend. DO NOT simply connect the dots. 5. Write a conclusion that explains how the probability of finding an electron changes as you move outward from the nucleus.

Electron Arrangement in Atoms (In the Q.M.M.) Main energy levels (aka: quantum numbers) - n = 1, 2, 3, 4 Sublevels - s, p, d, f The sublevels contain the orbitals

Sublevels Each sublevel contains orbitals Orbitals have specific shapes s: 1 orbital – sphere p: 3 orbitals – dumbells d: 5 orbitals – pears Look on page 131

Summary of Principal Energy Levels and Sublevels (p. 155) Number of sublevels Type of sublevel Maximum # of Electrons n = 1 1 1s (1 orbital) 2 n = 2 2s (1 orbital), 2p (3 orbitals) 8 n = 3 3 3s (1 orbital), 3p (3 orbitals), 3d (5 orbitals) 18 n = 4 4 4s (1 orbital), 4p (3 orbitals), 4d (5 orbitals), 4f (7 orbitals) 32

Review Page 167, #59-61, 63, 64, 68, 70, A4 finished #1

Electron Configuration Key Question: What are the rules for writing the electron configurations of elements?

Electron Configurations Definition: the ways electrons are arranged in various orbitals around the nuclei of atoms. Three rules tell you how to find the electron configuration of an atom: Aufbau Principle Pauli Exclusion Principle Hund’s rule

The Rules 1. Aufbau Principle electrons occupy orbitals of the lowest energy first 2. Pauli Exclusion Principle Each orbital holds only 2 electrons remember s has 1 orbital, p has 3 orbitals and d has 5 orbitals So, how many electrons will you find in s sublevel? p sublevel? d sublevel?

What does all that mean? When writing electron configurations 1.) Follow this chart (Aufbau Principle) (5-C) Chart is in lab manual

2) Remember: ‘s’ orbital holds a max of 2 electrons, ‘p’ orbital holds a max of 6 electrons and ‘d’ orbital holds a max of 10 electrons 3) Remember, the atomic number tells you how many electrons are in each atom

Let’s Practice Phosphorous has an atomic number of 15, what is the electron configuration? Follow the chart 1s2 2s2 2p6 3s2 3p3 = 1s22s22p63s23p3

More Practice Electron configuration for Carbon (atomic number = 6): 1s22s22p2 Electron configuration for Argon (atomic number = 18) 1s22s22p63s23p6

Individual Practice On page 136, answer #8c and all of 9. - For # 9, just write the electron configuration, do not worry about unpaired electrons.

More Practice In your lab manual, complete pages 5-C (front and back) and the front of 5-D

Drawing atomic orbitals: Use the Aufbau chart to figure out electron configuration 2) Draw arrows to represent electrons **Pauli exclusion: Electrons spin in opposite directions** 3) Draw all ‘up’ arrows to represent electrons before drawing any ‘down’ arrows **Hund’s rule: Every orbital in a sublevel is occupied by one electron before any one sublevel is occupied by two electrons**

Look at Table 5.2 Remember: s has 1 orbital p has 3 orbitals d has 5 orbitals ***You must memorize the number of orbitals in each sublevel.

Practice Here is the electron configuration for Phosphorus: 1s22s22p63s23p3 Draw the electron orbital configuration 1s 2s 2p 3s 3p __ __ __ __ __ __ __ __ __

More Practice On the back of 5-D in your lab manual try # 1 and 2 Do 1 and 2 as a group, if they get those right, they may continue. A4 stopped here

Practice 1. Page 136, #9 Draw the electron orbitals for the each atom to determine the number of unpaired electrons 2. On page 137, answer questions 11 and 14

Chapter Review Page 152, # 35-44, 56, 59, 60, 69, 73, 77 Page 157, # 1-3, 5 - 10