Ch 5 electrons in atoms notes

CH 5 ELECTRONS IN ATOMS
5-1 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 10 8 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.
Ex: Iron when hot appears red or blue, emits nrg that is quantized has a specific frequency.
Heating water – temp increases by molecules absorbing a specific amt or quanta.
Calculated as follows:
E quantum = h v
E = Energy (J)
h = Planck’s constant 6.626 x 10 -34 (J s)
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.
E photon = h v
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 =4 n =5 n =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= h υ

HYDROGEN’S LINE SPECTRUM
Lines which show up have specific energies which correspond to a frequency of a color of light.
A photon is emitted with E= h υ for each frequency E= 4.85 x 10 -19 J E= 3.03 x 10 -19 J Energy of Hydrogen Atom 1 2 3 4 5 6 n

5-2 QUANTUM THEORY AND THE ATOM
Quantum mechanical model is the modern atomic model and comes from
Louis De Broglie: radiation (energy) behaves like particles and vice versa.
All particles w/ a mass have wave characteristics
E- move around nucleus in a wave-like manner
Heisenberg uncertainty principle- impossible to know both the velocity and position of an e- at the same time.
Shrodinger: e-’s energy are limited to certain values (quantum) but does not predict path
Treated e-’s as waves
Created wave function = predicts probability of finding e- in a volume of space (location)
Moving Electron Photon Before Electron velocity changes Photon wavelength changes After

Shrodinger’s wave eqn predicts atomic orbitals
Atomic orbital - 3D regions around the nucleus that describes the e-’s probable location.
atomic orbital = fuzzy cloud
Do not have a defined size
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 = spherical
p = dumbbell
d, f= different shapes

Principal energy levels have specific allowed sublevels - shapes.
s sublevel is lower in energy and f has higher energy
s 1 2 3 4 n = s p s p d s p d f

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 2
p 3 6
d 5 10
f 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:
Aufbau principle – e- occupy lowest nrg orbital available
All orbitals in a sublevel are = in nrg (p x p y p z )
Sublevels within an energy level have different energies
Ex: 2s lower in nrg than 2p
Order of energy = s, p, d, f
Sublevels in one energy level can overlap with sublevels in another principal energy level.
Ex: 4s lower in nrg than 3d

AUFBAU DIAGRAM

ELECTRON CONFIGURATIONS
Pauli exclusion principle – a max of 2 e- may occupy a single orbital only if they have opposite spins.
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
2p x 2p y 2p z

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.
1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 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-
1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6
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:
ORBITAL DIAGRAM AND E- CONFIGURATIONS 1s 2 2s 2 2p 6 3s 2 3p 6 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 [ Ar ] 4s 2 3d 6 ] [

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:
Symbol - represents the atomic nucleus & inner-level electrons
Surrounded by dots- represent the valence electrons.
Ex: O = 1s 2 2s 2 2p 4 or [He]2s 2 2p 4 ve- = 6 in grp 6A
O

PERIODIC TABLE SHORTCUT
Periods = Energy Level
Groups (A only) = Valence e- Energy level = n-1 for d sublevel Energy level = n-2 for f sublevel 1A 2A 3A 4A 5A 6A 7A 8A

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Ch 5 electrons in atoms notes

by Esther Herrera

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