Electrons In AtomsWhere are they?

Development of Atomic ModelsPlum Pudding Model (1897)

J.J. ThomsonElectrons scattered in a “sea” of positive

charges

Development of Atomic ModelsRutherford’s Model (1911)

Discovered nucleus (disproves Plum Pudding)Electrons orbit nucleus like planets around the

sunCannot explain many of the properties of atoms

Development of Atomic ModelsBohr Model (1913)

Electrons move around nucleus in circular orbits at specific allowed distances

These distances relate to allowable energy levels

Energy levels – fixed energies an e- can haveQuantum of energy – energy needed to move

an electron from one E level to another

Development of Atomic ModelsMore Bohr Model

Electrons can gain or lose energyGround state – lowest energy level availableExcited state – higher energy levelAbsorb energy (gain E)

Go from lower to higher E levelsEmit energy (lose E)

drop from higher to lower E levels Give off E in the form of radiation (quanta of light)

Development of Atomic ModelsMore Bohr Model

Energy levels get closer together as they get farther from the nucleus

Problem: Works well with the hydrogen atom but not much else

Development of Atomic ModelsQuantum Mechanical Model (1926)

Modern description of electrons in atomsCloud model or Quantum TheorySchrodinger – developed mathematical

equation to predict atomic behaviorElectrons NOT in exact pathHeisenberg Uncertainty Principle

Impossible to know both location and energy of an electron

Can measure one or the other – NOT both Exact motion of electron unknown

Development of Atomic ModelsMore Quantum Mechanical Model

Determines allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus

These locations are called principal energy levels

Within these energy levels are sublevelsSublevels are subdivided into atomic orbitals

Development of Atomic ModelsMore Quantum Theory

Atomic Orbitals Region of space in which there is a high probability

of finding an electron 4 types of orbitals – s, p, d, and f Different orbitals have different shapes Each orbital can hold up to 2 electrons 2 electrons in same orbital must have opposite spins

(+1/2 and -1/2)

s and p orbitals

d sublevels

f orbitals

Quantum Numbers4 numbers used to describe electron location

1) Principal Energy Level (Principal Quantum Level)

n = 1, 2, 3…

2) Energy Sublevel Number specifies s, p, d, or f sublevel l = 0 to n-1

l = 0 s sublevel l = 1 p sublevel l = 2 d sublevel l = 3 f sublevel

Quantum Numbers3) Orbital quantum number (m)

m = -l to +l Specifies which orbital within a sublevel the

electron is located Within sublevels, orbitals differ only in spatial

orientation, not energy

4) Spin quantum number (ms) ms = +1/2 or -1/2 1st electron in orbital has + spin

Energy Levels, Sublevels, and Orbitals

Principal Energy Level

Number of Sublevels

Types of sublevels

1 1 1s (1 orbital)2 2 2s (1 orbital), 2p (3 orbitals)3 3 3s (1 orbital, 3p (3 orbitals)

3d (5 orbitals)4 4 4s (1 orbital), 4p (3 orbitals)

4d (5 orbitals), 4f (7 orbitals)

*** Remember: Each orbital can hold 2 electrons ***

Orbitals and ElectronsEnergy level, n Sublevel

(Orbitals)Max #

Electrons / sublevel

Max # Electrons /

energy level1 s 2 2

2 s p 2 6 8

3 s p d 2 6 10 18

4 s p d f 2 6 10 14

32

Maximum # electrons / energy level = 2n2

where n = energy level

s sublevel – 1 orbital, 2 electrons p sublevel – 3 orbitals, 6 electronsd sublevel – 5 orbitals, 10 electrons f sublevel – 7 orbitals, 14 electrons

Electron ConfigurationThe way in which electrons are arranged in

various orbitals around the nucleus of an atom

Aufbau PrincipleElectrons occupy the orbitals of lowest energy first

Electron ConfigurationPauli Exclusion Principle

An atomic orbital may describe at most two electrons

Opposite spins Boxes represent orbitals and arrows represent

electrons3s sublevel with 1 electron

4s sublevel with 2 electrons

Electron ConfigurationHund’s Rule

Electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin as large as possible

Orbitals of equal energy each get 1 electron before any pair up

2p sublevel

Electron ConfigurationDiagonal rule

s holds 2p holds 6d holds 10f holds 14

Atomic Structure Practice3a) 1s22s22p3 Nitrogen ( 7 electrons)

b) 1s22s2 Beryllium (4 electrons)

c) 1s22s22p63s23p3 Phosphorus (15 electrons)

d) 1s22s22p63s23p63d54s2 Manganese (25 electrons)

e) Potassium (19) f) Zirconium (40) g) Promethium (61) h) Selenium (34)

Atomic Structure Practice4a) Cu0 (29 e-) 1s22s22p63s23p64s23d9

Cu+ (28 e-) 1s22s22p63s23p64s23d8

Cu2+ (27 e-) 1s22s22p63s23p64s23d7

4b) Al0 (13 e-) 1s22s22p63s23p1

Al3+ (10 e-) 1s22s22p6

Electron ConfigurationShort Cut Method

Rare Gas Configuration, Noble Gas Configuration, or Inert Gas Configuration (Either name OK)

Relate back to the previous rare gasPut that element in [ ]Start at s sublevel using whatever period the

element is inNickel 1s22s22p63s23p64s23d8 or [Ar] 4s23d8