Modern Atomic Theory 1 Bohrs Model Nucleus Electron

Modern Atomic Theory 1

Bohr’s Model Nucleus Electron Orbit Energy Levels

3 ELECTROMAGNETIC RADIATION

Electromagnetic radiation. 4

5 Electromagnetic Radiation • Most subatomic particles behave as PARTICLES and obey the physics of waves.

Electromagnetic Spectrum Long wavelength --> small frequency Short wavelength --> high frequency increasing wavelength 6

7 Electromagnetic Spectrum In increasing energy, ROY G BIV

Excited Gases & Atomic Structure 8

Atomic Line Emission Spectra and Niels Bohr 9 Bohr’s greatest contribution to science was in building a simple model of the atom. It was based on an understanding of the LINE EMISSION SPECTRA of excited atoms. Niels Bohr (1885 -1962)

10 Electrons • Ground State – electrons are located in the lowest energy levels possible. • Excited State – electrons have absorbed energy and jump to a higher energy level. • When electrons go from a high energy level to a lower energy level they emit a photon of light.

An excited lithium atom emitting a photon of red light to drop to a lower energy state. 11

An excited H atom returns to a lower energy level. 12

13 Spectrum of White Light

Line Emission Spectra of Excited Atoms • Excited atoms emit light of only certain wavelengths • The wavelengths of emitted light depend on the element. 14

15 Spectrum of Excited Hydrogen Gas

16 Line Spectra of Other Elements

17 Bohr Model Bohr’s view of atomic structure was that an electron (e-) traveled about the nucleus in a specific orbit.

18 • The first energy level can hold a maximum of 2 electrons. • The second energy level can hold up to 8. • The third energy level can hold up to 18. • The fourth energy level can hold up to 32.

19 } Increasing energy Fifth Fourth Third Second First • Further away from the nucleus means more energy. • There is no “in between” energy • Energy Levels

Lewis Structures Valence shell: the outermost electron shell Valence electrons: electrons in valence shell electrons used in bonds Lewis structure: atom symbol = nucleus + inner edots represent valence electrons 20

Lewis Structures • Table 1. 4 Lewis Structures 21

22 Atomic Spectra and Bohr’s classical view is wrong. Need a new theory — now called QUANTUM or WAVE MECHANICS. e- can only exist in certain discrete orbits found within an energy level. Each principle energy level is made up of sublevels. Sublevels are made up of orbitals.

Quantum or Wave Mechanics 23 Schrodinger applied idea of ebehaving as a wave to the problem of electrons in atoms. He developed the WAVE EQUATION Solution gives set of math expressions called WAVE E. Schrodinger FUNCTIONS, 1887 -1961 Each describes an allowed energy state of an e-

Heisenberg Uncertainty Principle W. Heisenberg 1901 -1976 Problem of defining nature of electrons in atoms solved by W. Heisenberg. Cannot simultaneously define the position and momentum (= m • v) of an electron. We define e- energy exactly but accept limitation that we do not know exact position. 24

Arrangement of Electrons in Atoms Electrons in atoms are arranged as Principle Energy Levels (1, 2, 3…) Sublevels (s, p, d, f) Orbitals (can contain 2 electrons) 25

26 Energy Levels • Each energy level has a number called the PRINCIPAL ENERGY LEVEL, n • Currently n can be 1 thru 7, because there are 7 periods on the periodic table

27 Energy Levels n=1 n=2 n=3 n=4

28 Energy Sublevels • Each Principle Energy Level is composed of sublevels (s, p, d and f) • Each sublevel is composed of orbitals, each orbital can contain two electrons. # of Sublevel Orbitals s 1 p 3 d 5 f 7 Total # Electrons 2 6 10 14 PEL 1 st Found 1 st 2 nd 3 rd 4 th

Relative sizes of the spherical 1 s, 2 s, and 3 s orbitals of hydrogen. 29

p Orbitals • The three p orbitals lie 90 o apart in space 30

31 The shapes and labels of the five 3 d orbitals.

32 f Orbitals

Orbitals and the Periodic Table • Orbitals grouped in s, p, d, and f orbitals (sharp, proximal, diffuse, and fundamental) s orbitals f orbitals d orbitals p orbitals 33

Diagonal Rule • The diagonal rule is a memory device that helps you remember the order of the filling of the orbitals from lowest energy to highest energy • Electrons fill from the lowest possible energy to the highest energy 34

35 Diagonal Rule Steps: 1 s 2 s 3 s 1. Write the energy levels top to bottom. 2. Write the orbitals in s, p, d, f order. Write the same number of orbitals as the energy level. 3. Draw diagonal lines from the top right to the bottom left. 4. To get the correct order, 2 p 3 p 3 d follow the arrows! 4 s 4 p 4 d 4 f 5 s 5 p 5 d 5 f 5 g? 6 s 6 p 6 d 6 f 6 g? 6 h? 7 s 7 p 7 d 7 f 7 g? 7 h? By this point, we are past the current periodic table so we can stop. 7 i?

36 Why are d and f orbitals always in lower energy levels? • d and f orbitals require LARGE amounts of energy • It’s better (lower in energy) to skip a sublevel that requires a large amount of energy (d and f orbtials) for one in a higher level but lower energy This is the reason for the diagonal rule! BE SURE TO FOLLOW THE ARROWS IN ORDER!

Electron Configurations A list of all the electrons in an atom (or ion) • Must go in order (Aufbau principle) • 2 electrons per orbital, maximum • We need electron configurations so that we can determine the number of electrons in the outermost energy level. These are called valence electrons. • The number of valence electrons determines how many and what this atom (or ion) can bond to in order to make a molecule 1 s 2 2 p 6 3 s 2 3 p 6 4 s 2 3 d 10 4 p 6 5 s 2 4 d 10 5 p 6 6 s 2 4 f 14… etc. 37

Electron Configurations 4 2 p Energy Level 38 Number of electrons in the sublevel Sublevel 1 s 2 2 p 6 3 s 2 3 p 6 4 s 2 3 d 10 4 p 6 5 s 2 4 d 10 5 p 6 6 s 2 4 f 14… etc.

Shorthand Notation • A way of abbreviating long electron configurations • Since we are only concerned about the outermost electrons, we can skip to places we know are completely full (noble gases), and then finish the configuration 39

Shorthand Notation 40 • Step 1: Find the closest noble gas to the atom (or ion), WITHOUT GOING OVER the number of electrons in the atom (or ion). Write the noble gas in brackets [ ]. • Step 2: Find where to resume by finding the next energy level. • Step 3: Resume the configuration until it’s finished.

Shorthand Notation • Chlorine – Longhand is 1 s 2 2 p 6 3 s 2 3 p 5 You can abbreviate the first 10 electrons with a noble gas, Neon. [Ne] replaces 1 s 2 2 p 6 The next energy level after Neon is 3 So you start at level 3 on the diagonal rule (all levels start with s) and finish the configuration by adding 7 more electrons to bring the total to 17 [Ne] 2 3 s 5 3 p 41

Valence Electrons are divided between core and valence electrons B 1 s 2 2 p 1 Core = [He] , valence = 2 s 2 2 p 1 Br [Ar] 3 d 10 4 s 2 4 p 5 Core = [Ar] 3 d 10 , valence = 4 s 2 4 p 5 42

Rules of the Game No. of valence electrons of a main group atom = Group number (for A groups) Atoms like to either empty or fill their outermost level. Since the outer level contains two s electrons and six p electrons (d & f are always in lower levels), the optimum number of electrons is eight. This is called the octet rule. 43

44 Keep an Eye On Those Ions! • Electrons are lost or gained like they always are with ions… negative ions have gained electrons, positive ions have lost electrons • The electrons that are lost or gained should be added/removed from the highest energy level (not the highest orbital in energy!)

45 Keep an Eye On Those Ions! • Tin Atom: [Kr] 5 s 2 4 d 10 5 p 2 Sn+4 ion: [Kr] 4 d 10 Sn+2 ion: [Kr] 5 s 2 4 d 10 Note that the electrons came out of the highest energy level, not the highest energy orbital!

46 Keep an Eye On Those Ions! • Bromine Atom: [Ar] 4 s 2 3 d 10 4 p 5 Br- ion: [Ar] 4 s 2 3 d 10 4 p 6 Note that the electrons went into the highest energy level, not the highest energy orbital!

Orbital Diagrams • Graphical representation of an electron configuration where each orbital is shown. • One arrow represents one electron • The direction of the arrow represents the electron spin, electrons in the same orbital have opposite spin. 47

Orbital Diagrams • One additional rule: Hund’s Rule – In orbitals of EQUAL ENERGY (p, d, and f), place one electron in each orbital before making any pairs – All single electrons must spin the same way. 48

Lithium 49 Group 1 A Atomic number = 3 1 s 22 s 1 ---> 3 total electrons

Carbon Group 4 A Atomic number = 6 1 s 2 2 p 2 ---> 6 total electrons Here we see for the first time HUND’S RULE. When placing electrons in a set of orbitals having the same energy, we place them singly as long as possible. 50