CH 5 ELECTRONS IN ATOMS5-1 Light and Quantized Energy

Some elements emit visible light when heated with a flame. This chemical behavior is due to the arrangement of e- in atoms.

ELECTROMAGNETIC RADIATION

Form of energy that exhibits wave-like behavior as it travels through space.

There are many types of electromagnetic radiation and all are represented in the ELECTROMAGNETIC SPECTRUM

ELECTROMAGNETIC SPECTRUM

PARTS OF A WAVE

Frequency (v, nu –The number of complete wavelengths that pass a given point each second. Units: wave/second = 1/s = s-1 = Hertz (Hz)

Wavelength (lambda) – The distance between identical points on successive waves. (crest to crest or trough to trough) Units: meters (m)

c = v c = speed of light, 3.00 x 108 m/s

WAVE NATURE OF LIGHT Max Planck theorized that all matter can

gain/ lose energy in small “chunks” of light (quanta). Quantum- minimum amt of energy that

can be gained or lost by an atom.o Ex: Iron when hot appears red or blue, emits

nrg that is quantized has a specific frequency.o Heating water – temp increases by molecules

absorbing a specific amt or quanta. Calculated as follows:

Equantum= hv

o E = Energy (J)o h = Planck’s constant 6.626 x 10-34 (J s)o v = frequency ( Hz or s-1)

PARTICLE NATURE OF LIGHT

Photoelectric effect – electrons are emitted from a metal’s surface when light of a specific frequency shines on the surface.

Albert Einstein (1905) assumed that light travelled as a stream of tiny particles or packets of energy called photons.

Photons- EM radiation w/ no mass that carries a quantum of energy. EM radiation has both wave- like and particle- like nature. Ephoton= hv Photon = quantum of energy

ATOMIC EMISSION SPECTRA

Set of frequencies of light waves emitted by an atom of an element.

Line spectrum – consists of several individual lines of color from light energy emitted by excited unstable atoms Only certain colors (frequencies) appear in an

element’s AES & it can be used to identify the element.

5-2 QUANTUM THEORY OF THE ATOM

Bohr Model of the Atom Used to explain why AES was set of

discontinuous lines of specific frequencies (color).

Proposed that Hydrogen atoms have only certain allowable energy states based on Planck’s and Einstein’s quantized energy. Ground state- lowest allowable energy states of an

atom. Excited state- atom gains energy; H atoms can have

many different excited states although it contains 1 e-. Electrons move around a H atom in circular orbit Orbits equal to a principal quantum number n,

where n=1 is lowest nrg level, closest to nucleus.

BOHR MODEL OF THE ATOM

Orbits/ levels are like rungs in step ladder Cannot stand b/w rungs, e-

can’t exist b/w levels (orbits).

E- move from 1 orbit to the next emitting or absorbing certain amts of nrg (quanta). The smaller the e- orbit, the

lower the energy state/level The larger the e- orbit, the

higher the energy state/level

n =1

n =2

n =3

n =4n =5n =6

nucleus

BOHR MODEL OF THE ATOM Hydrogen’s Line Spectrum (AES)

At n= 1 H atom is in ground state When nrg is added, e- moves to higher energy level,

n=2 (excited state). E- drop back to lower energy level n=1 and emitts a

photon equal to the difference b/w levels.

A photon is absorbed

A photon is emitted with E= hυ

HYDROGEN’S LINE SPECTRUM Lines which show up have specific energies which

correspond to a frequency of a color of light.

En

ergy

of

Hyd

roge

n A

tom

1

2

3

456n

A photon is emitted with E= hυ for each frequency

E= 4.85 x 10-19 J

E= 3.03 x 10-19 J

5-2 QUANTUM THEORY AND THE ATOM

Quantum mechanical model is the modern atomic model and comes fromA. Louis De Broglie: radiation (energy) behaves

like particles and vice versa.1. All particles w/ a mass have wave characteristics2. E- move around nucleus in a wave-like manner

B. Heisenberg uncertainty principle- impossible to know both the velocity and position of an e- at the same time.

C. Shrodinger: e-’s energy are limited to certain values (quantum) but does not predict path

1. Treated e-’s as waves2. Created wave function = predicts probability of

finding e- in a volume of space (location)

Moving Electron

Photon

Before

Electron velocity changes

Photon wavelengthchanges

After

Shrodinger’s wave eqn predicts atomic orbitals Atomic orbital - 3D regions around the

nucleus that describes the e-’s probable location.a. atomic orbital = fuzzy cloudb. Do not have a defined sizec. Shape = volume that contains 90% of the

probable location of e-’s inside that region.

HYDROGEN’S ATOMIC ORBITALS

QUANTUM MECHANICAL MODEL

Like Bohr, electrons occupy space surrounding nucleus and exist in several principal energy levels = principal quantum number (n) Relative size and energies of atomic orbital n = 1,2, 3, etc. = period

Principal nrg levels consist of energy sublevels with different nrg values. Energy sublevels – shape of the atoms’ orbitals

s = sphericalp = dumbbelld, f= different shapes

QUANTUM MECHANICAL MODEL Principal energy levels have specific allowed

sublevels - shapes.

s sublevel is lower in energy and f has higher energy

1

2

3

4

n =s

s p

s p d

s p d f

QUANTUM MECHANICAL MODEL

Sublevels consist of orbitals of different orientation. Orbitals in same sublevel are = in energy (no matter

orientation) Orbitals only hold 2e- maximum with opposite spins (+

or – spins).Sublevel Orientations/ Orbitals Max #

e-s 1 2p 3 6d 5 10f 7 14

ORIENTATIONS/ ORBITALS PER SUBLEVEL s- spherical only 1 orbital orientation

p- dumbbell has 3 orbital orientations

d- 2dumbbells with 5 orbital orientations

f- 3dumbbells with 7 orbital orientations

http://winter.group.shef.ac.uk/orbitron/AOs/1s/index.html

5-3 ELECTRON CONFIGURATIONS

Electron configuration – arrangement of e- in atoms; lower nrg arrangements

Arrangements defined by:1. Aufbau principle – e- occupy lowest nrg

orbital availablea. All orbitals in a sublevel are = in nrg (px py pz )

b. Sublevels within an energy level have different energies

Ex: 2s lower in nrg than 2pc. Order of energy = s, p, d, fd. Sublevels in one energy level can overlap with

sublevels in another principal energy level.a. Ex: 4s lower in nrg than 3d

AUFBAU DIAGRAM

ELECTRON CONFIGURATIONS

2. Pauli exclusion principle – a max of 2 e- may occupy a single orbital only if they have opposite spins.

3. Hund’s rule – energy charged e- repel each other.

All same nrg orbitals are filled first with e- containing same spin before extra e- can occupy the same orbital with opposite spins. Ex: 3 orbitals of 2p

2px 2py 2pz

FILLING SUBLEVELS WITH ELECTRONS

Energy sublevels are filled from lower energy to higher energy following the diagram. ALWAYS start at the beginning of each level and

follow it until all e- in an element have been placed.

1s2s 2p3s 3p 3d4s 4p 4d 4f5s 5p 5d 5f6s 6p 6d7s 7p

Orbital diagram for Fe: Iron has how many e- ? 26 e-

1s 2s 2p 3s 3p 4s 3d Electron configuration for Fe:

Iron has 26 e- 1s2 2s2 2p6 3s2 3p6 4s2 3d6

Shortcut to the E- config. for Fe is Noble gas notation Group 18 or 8A are the Nobel Gases Argon has 18 e- Iron has 26 e- Noble gas notation:

1s2 2s2 2p6 3s2 3p6

ORBITAL DIAGRAM AND E- CONFIGURATIONS

][1s2 2s2 2p6 3s2 3p6 4s2 3d6[Ar] 4s2

3d6

VALENCE ELECTRONS AND ELECTRON DOT STRUCTURES

Valence electrons – outer energy level/orbital electrons which are involved in bonding. Valence electrons = groups 1A to 8A B GROUPS DO NOT COUNT

E- dot structures- consists of the element’s:a. Symbol - represents the atomic nucleus & inner-

level electrons b. Surrounded by dots- represent the valence

electrons.c. Ex: O = 1s2 2s2 2p4 or [He]2s2 2p4 ve- =6 in grp

6A O

PERIODIC TABLE SHORTCUT

Peri

od

s =

En

erg

y L

evel

Groups (A only) = Valence e- 1A

2A 3A 4A 5A 6A 7A

8A

Energy level = n-1 for d sublevel

Energy level = n-2 for f sublevel