Chapter 5

Chapter 5

Quantum Theory and Electron Configurations

It’s all about color… In terms of atomic models, so far: Dalton (1803) = Tiny, solid particle Thomson (1897) = “Plum Pudding” model – Electrons stuck on the outside of a big positive charge Rutherford (1911) = Positively-charged nucleus with electrons moving around it Rutherford’s model of the atom not quite right Could not explain chemical properties of elements Could not explain color changes when metal is heated

Bohr Model of the Atom Niels Bohr’s model of the atom Electron found only on specific, circular paths around nucleus Each orbit has fixed energy level Hypothesis: When electrons are excited (added energy), jump into higher energy levels. When they moved back into lower energy levels - gave off light. Electrons do not exist between levels (think of rungs on a ladder) Electrons absorb and emit only certain quanta (amounts) of energy Quantum of energy = fixed amount of energy required to move from one energy level to another energy level

Bohr’s Model Nucleus Electron Orbit Energy Levels Chapter 5

Bohr’s Planetary Model of the Atom Electrons must have enough energy to keep moving around the nucleus Electrons orbit nucleus in defined energy levels, just like planets orbit the sun Each energy level assigned a principal quantum number n. Lowest energy level called ground state (n=1) Higher energy levels (n=2, 3, 4...) excited states Model worked OK for hydrogen but not so good for other elements Nucleus n = 1 n = 2

} Bohr’s Model Fifth Fourth Increasing energy Third Second First Further away from the nucleus means more energy. There is no “in between” energy Energy Levels Fifth Fourth Third Increasing energy Second First Nucleus Chapter 5

Electron starts on lowest energy level (ground state) Higher energy levels = excited states Add energy to electron – moves to excited state Energy levels are not evenly spaced Energy Level 2 Energy Level 3 Energy Level 1 Nucleus

Electron starts on lowest energy level (ground state) Higher energy levels = excited states Add energy to electron – moves to excited state Energy Energy levels are not evenly spaced Electron returns to lower state – emits/gives off quantum of energy Energy Level 2 Energy Level 3 Energy Level 1 Nucleus

Bohr used this theory to explain the lines in the atomic emission spectra for hydrogen Chapter 5

Each of these lines corresponds to different energy changes 434 nm 656 nm 486 nm 410 nm Chapter 5

Chapter 5

Chem I - Mon, 9/22/15 Do Now Get to work on the PEN worksheet from last class Homework MEAL paragraph if not finished Agenda History Intro to quantum Electron Config

Quantum-Mechanical Model of the Atom Since the Bohr model had a very limited use, a new and very different model of the atom exists The Quantum Mechanical Model (1926) contains: Quantum energy levels Dual wave/particle nature of electrons Electron clouds In the new model, don’t know exactly where electrons are - only know probabilities of where they could be

Heisenberg Uncertainty Principle Heisenberg Uncertainty Principle = impossible to know both the velocity (or momentum) and position of an electron at the same time

Quantum-Mechanical Model of the Atom Orbital = region around nucleus where an electron with a given energy level will probably (90%) be found Four kinds of orbitals s - spherical in shape, lowest orbital for every energy level p - dumbbell shaped, second orbital d - complex “flower” shape, third orbital f - very complex shape, highest orbital

s-orbitals All s-orbitals are spherical. As n increases, the s-orbitals get larger.

p- orbitals Three p-orbitals: px, py, and pz Lie along the x-, y- and z- axes of a Cartesian system. Dumbbell shaped, gets larger as n increases

d and f - orbitals There are five d and seven f-orbitals.

Quantum Mechanical Model Principle Energy Levels (n) Labeled from 1-7 First energy level is n=1 Contains sublevels (s, p, d and f) Each energy level contains the number of sublevels equal to its value for n If n=3, there are three sublevels

Chapter 5

Quantum Mechanical Model In each sublevel there are atomic orbitals Atomic orbitals – describe a space where an electron is likely to be found Type of subshell Shape of orbitals Number of orbitals Orbital ‘names’ s Spherical 1 p Dumbbell 3 px, py, pz d Cloverleaf (and one donut) 5 f Multi-lobed 7

Quantum Mechanical Model Each orbital can contain two electrons. Since negative-negative repel, these electrons occupy the orbital with opposite spins.

Quantum Mechanical Model The total number of orbitals of an energy level is n2. For the third principle energy level, n=3, which means there are 9 orbitals These orbitals are 3s, 3px, 3py, 3pz and the 5 d orbitals Remember, we no longer think of orbitals as concentric circles, but we can say that n=4 extends farther from the nucleus than n=1.

Valence Electrons Only those electrons in the highest principle energy level

Electron Configuration and Orbital Notation Aufbau Principle – electrons fill lower energy orbitals first, “bottom-up” n=1 fills before n=3 Will an electron fill the 1s or the 2s orbital first? Energy 2px 2py 2pz 2s 1s

Electron Configuration &Orbital Notation Hund’s Rule – electrons enter same energy orbitals so that each orbital has one electron before doubling up Each of the first electrons to enter the equal energy orbitals must have the same spin If we have 7 electrons, how will they fill in the below orbitals? Energy 2px 2py 2pz 2s 1s

Electron Configuration and Orbital Notation Pauli Exclusion Principle – an orbital can contain no more than 2 electrons. Electrons in the same orbital must have different spins. If we have 8 electrons, how will they be arranged? Energy 2px 2py 2pz 2s 1s

Apartment Analogy Atom is the building Floors are energy levels Rooms are orbitals Only two people per room

Orbital Diagrams Draw each orbital as a box. Each electron is represented using an arrow. Up arrows – clockwise spin Down arrows – counter-clockwise spin Determine the total number of electrons involved. Start with the lowest energy level (1s) and start filling in the boxes according the rules we just learned.

Orbital Diagram 2s 1s 3s 4s 2p 3p 4p 3d Energy

Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s 2p Chapter 5

Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s 2p The first to electrons go into the 1s orbital Notice the opposite spins only 13 more Chapter 5

Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s 2p The next electrons go into the 2s orbital only 11 more Chapter 5

Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p The next electrons go into the 2p orbital only 5 more Chapter 5

Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p The next electrons go into the 3s orbital only 3 more Chapter 5

Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s The last three electrons go into the 3p orbitals. They each go into separate shapes 3 unpaired electrons 1s22s22p63s23p3 Chapter 5

Orbital Diagrams Orbital diagrams are used to show placement of electrons in orbitals. Need to follow three rules (Aufbau, Pauli, Hund’s) to complete diagrams Li Be B C N Ne Na

Orbitals and Energy Levels Principal Sublevels Orbitals Energy Level n = 1 1s 1s (one) n = 2 2s 2p , 2s (one) + 2p (three) n = 3 3s , 3p , 3d 3s (one) + 3p (three) + 3d (five) n = 4 4s , 4p , 4d , 4f 4s (one) + 4p (three) + 4d (five) + 4f (seven) Chapter 5

Summary shapes Max electrons Starts at energy level s p d f Chapter 5

Orbitals and Energy Levels 4p 4d 4f n = 4 and so on.... 3s 3p 3d n = 3 Increasing energy 2s 2p n = 2 1s n = 1

Electron Configuration Let’s determine the electron configuration for Phosphorus Need to account for 15 electrons Chapter 5

Writing Electron Configuration Determine the total number of electrons. Write the principle energy level number as a coefficient, the letter for the subshell, and an exponent to represent the number of electrons in the subshell. He: 1s2

The Kernel (Noble Gas) Notation Determine the total number of electrons Find the previous noble gas and put its symbol in brackets Write the configuration from that noble gas forward as usual

Writing electron configurations Examples O 1s2 2s2 2p4 Ti 1s2 2s2 2p6 3s2 3p6 3d2 4s2 Br 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p5 Core format O [He] 2s2 2p4 Ti [Ar] 3d2 4s2 Br [Ar] 3d10 4s2 4p5 Chapter 5