catalyst 1. 2. 3.. ways to organize shapes organize a sphere on xyz coordinate plane organize dumb...

CatalystCatalyst1.2.3.

Ways to Organize ShapesWays to Organize Shapes• Organize a sphere on xyz coordinate plane• Organize dumb bells on xyz coordinate plane

Today’s Learning TargetsToday’s Learning Targets• 1.6 – I can characterize an electron based on its 4

quantum numbers (n, l, ml, and ms). I can explain what each of these numbers indicate and discuss the importance of these numbers.

• 1.7 – I can describe the shape, number, and energy level of the s, p, d, and f orbitals. Furthermore, I can draw the s and p orbitals.

• 1.8 – I can write the electron configuration and orbital diagram for any element on the Periodic Table using the Pauli Exclusion Principle and Hund’s Rule.

OrbitalsOrbitals• The solution to the Schrödinger equation gives us

orbitals for a particular element.• Every orbital has a special shape and energy

level.• We keep the idea of energy level (n = 1,2,3, etc.)

from the Bohr model, but remove idea of fixed orbits.

• There are 4 characteristic shapes that an orbital can take.

s Orbitalss Orbitals• Spherical shape• Only one possible orientation around the nucleus.• Seen whenever n (energy level) is 1 or greater

All s orbitals have only one possible

orientation.

p Orbitalsp Orbitals• All p-orbitals are dumb-bell shaped.• Because of the unique shape there are 3 possible

orientations around the nucleus (one along each of the 3 axis).

• Seen whenever n is 2 or greater

d Orbitalsd Orbitals• All d orbitals have a four-leaf clover shape.• 4 possible orientations within the plane and one

odd shape (dz2)

• Appear when n is 3 or higher.

f Orbitalsf Orbitals• These orbitals appear whenever n is 4 or higher• Only seen in elements that have many electrons.

Quick TalkQuick Talk• 1 partner explain s and d orbitals. 1 partner

explain p and f orbitals. Be able to draw s and p orbital

Quantum NumbersQuantum Numbers• We can describe any orbital that an electron

exists in using 4 quantum numbers.• No 2 orbitals can have the same 4 quantum

numbers.• Quantum Numbers:

1. Principal Quantum Number (n)2. Angular Momentum Quantum Number (l)3. Magnetic Quantum Number (ml)

4. Spin Magnetic Quantum Number (ms)

Principal Quantum Number Principal Quantum Number (n)(n)

• Describes the energy level of the orbital• Can have any integer value of 1 or greater (1, 2,

3, etc.)• The bigger n becomes, the higher the energy

level

Angular Momentum ( )Angular Momentum ( )• Describes the shape of our orbital (s, p, d, or f)• Can have any value from 0 to (n – 1)• Each orbital has an assigned angular momentum:

Value of 0 1 2 3

Letter used

s p d f

Quick WriteQuick Write• If I have an element at the n = 3 energy level,

then what type(s) of orbitals do I have?

Magnetic Quantum Number Magnetic Quantum Number (m(m ) )

• ml is any number – to +

• Every orbital can take on a certain number of “allowed” orientations.

• Tells you the number of total electrons an orbital can hold

• The number of allowed orientations is the sum of – to +

Quick WriteQuick Write• How many allowed orientations are there for the s

orbital ( =0)? How many electrons can it hold?• How many allowed orientations are there for the

d orbital ( =2)? How many electrons can it hold?

Spin Magnetic Quantum Spin Magnetic Quantum Number (mNumber (mss))

• Any orbital can contain, at most, 2 electrons• If electrons are in the same orbital, then they

must have opposite spins or ms values

• An electron can either have a spin of +½ or -½

SummarizeSummarize

Around the WorldAround the World

• Around the room there are 10 problems. Cycle through the problems to practice writing electron configurations.

• Complete all problems

5 Minute Break5 Minute Break

Who would perform at Who would perform at the ultimate concert?the ultimate concert?

Concert of a LifetimeConcert of a Lifetime

Concert of a LifetimeConcert of a Lifetime

Stage

Orbitals and their EnergyOrbitals and their Energy• Orbitals, based on many

different reasons, have varying energies.

• n does not determine energy levels when comparing orbitals

Electron Electron

ConfigurationsConfigurations• Electron configurations

describe the distribution of each electron among the various orbitals in the atom

Pauli Exclusion PrinciplePauli Exclusion Principle• Within an atom, no two electrons can have the

same set of 4 quantum numbers.

Orbital Orbital

energy energy

diagramsdiagramsAufbau principle: Build up each atom from the preceding atom by “filling” electrons in from the bottom.

Hund’s RuleHund’s Rule• When placing electrons into a group of similar

orbitals, electrons enter empty orbitals first before they form pairs

Class ExampleClass Example• Draw the electron configuration for fluorine.

Table TalkTable Talk• Draw the electron configuration for magnesium

3 Essential Principles3 Essential Principles1. Pauli Exclusion Principle – Every electron gets its own unique quantum number per element.2. Hund’s Rule – Electrons spread out within the same energy level.3. Aufbau Principle – Fill the lowest energy orbital first

How to Remember Order of How to Remember Order of Orbital FillingOrbital Filling

1s2s3s4s5s6s

2p3p4p5p6p

3d4d5d6d

4f5f6f

Periodic Table Quantum Periodic Table Quantum Shortcut! Shortcut!

• The position of the element determines the last orbital and energy level filled.

Condensed Electron Condensed Electron ConfigurationsConfigurations

• Abbreviate configurations using Noble Gases.• Choose the Noble Gas in the row just above the

element of interest.• Substitute this Noble Gas for all electrons prior to

the row the element is in.

1s22s22p63s23p4 [Ne]3s23p4

10 electrons = Neon!

Class ExampleClass Example• Write the electron configuration for zinc using

condensed notation

Table TalkTable Talk• Write the electron configuration for Tin.

JTPS: Analyze the JTPS: Analyze the ConfigurationConfiguration

• Which configuration for chromium is more stable. WHY?:

1s22s22p63s23p64s23d4 1s22s22p63s23p64s13d5or

Strange Electron Strange Electron ConfigurationsConfigurations

• Due to the closeness in energy of the 3d and 4s orbitals, electrons will spread out in order to satisfy Hund’s Rule.

• Seen most often in chromium and copper.

SummarizeSummarize

Coach and CorrectCoach and Correct

Question 1Question 1• In 3 words or less for each:

o Describe the Pauli Exclusion Principleo Describe Hund’s Ruleo Describe the Aufbau Principle

Question 2Question 2• What is the principal quantum number, number of

electrons, and angular momentum of 5p6

Question 3Question 3• Write the long-form electron configuration for

sulfur

Question 4Question 4• Write the condensed electron configuration for

titanium

Question 5Question 5• Why do electrons fill the lowest energy level first?

Question 6Question 6• Looking at the following electron configurations,

which represents a chemically unreactive element? Justify your answer!

a) 1s____2s ↑ _b) 1s ↑↓ 2s ↑↓_c) 1s ↑↓ 2s ↑↓ 2p ↑ ↑ ____d) 1s ↑↓ 2s ↑↓ 2p ↑↓_ ↑↓ _↑↓_ e) [Ar]4s ↑↓ 3d ↑↓ ↑ ↑ ↑ ↑_

Question 7Question 7• The contour representation of one of the orbitals

for the n = 3 shell of hydrogen atom is shown belowa) What is the quantum number l for this orbital?b) What is the notation for this orbital

Question 8Question 8• An element has a valence shell configuration of

ns2np5, what element(s) is it?

Question 9Question 9• You friend writes the electron configuration for

phosphorus as: [He]3s23p3. What is wrong with the way that they wrote the configuration?

Question 10Question 10• Write the condensed electron configuration for

uranium

Closing TimeClosing Time• You should have all Chapter 6 reading and

homework problems done by Monday/Tuesday to be on track.

• Also, read corresponding chapter in Cracking the AP Chemistry Exam

• Pre-lab for Lab 2 due at start of class Monday/Tuesday.

• No pre-lab = you cannot start lab on time = staying after school to finish the lab.