4.1 Refinements of the atomic model
Models of the atom so far: Dalton – atoms are like little “bb’s” - then the
electron gets discovered Thomson – atom is like a charged “bb” Rutherford - Gold foil experiment – hollow charged
“bb” Bohr model of the atom (1913) – Neils Bohr – Danish
Physicist The Bohr model of the atom comes from the idea that
light is waves of energy
http://web.visionlearning.com/custom/chemistry/animations/CHE1.2-an-atoms.shtml
The Bohr Atom (1913)
All the positive charge was in the nucleus Electrons orbited the nucleus much like planets orbit the
sun (at fixed distances) The closer the electrons to the nucleus, the less energy it
has. The farther the electron is from the nucleus, the more
energy it has.
The Electromagnetic Spectrum
Visible light, x-rays, ultraviolet radiation, infrared radiation, microwaves and radio waves are all part of the electromagnetic spectrum
The Electromagnetic Spectrum
The spectrum consists of electromagnetic radiation – energy that travels like a wave
Waves can be described by the wave equation which includes velocity (c = speed of light), wavelength (λ) and frequency (ν).
Wavelength (definition) = the distance between peaks of a wave
Light through prism leads to high energy (violet) low energy (red)
The Electromagnetic Spectrum
ROYGBIV - colors of the visible spectrum Bright Line Spectrum (BLS) – caused by e- emitting
energy as they return to lower energy levels energy level.
heat sodium - yellow light 2 c heat lithium - red light elements can appear to give off the same color light, but
each will have its own BLS BLS - used to determine identity of an element BLS - validates Bohr’s idea that electrons jump to different
energy levels and give off different wavelengths of light
The Electromagnetic Spectrum
Light from the sun (white light) appears as a continuous spectrum of light.
Continuous Spectrum of Light (definition) = There are no discrete, individual wavelengths of light but rather all wavelengths appear, one after the other in a continuous fashion
Spectroscopy (definition) = the study of substances from the light they emit.
We will use spectroscopes (An instrument that splits light into its component colors) and flame tests to study elements because each element emits a different spectrum of light when exited .
Birght Line Spectrum
Bohr proposed that the energy possessed by an e- in a H- atom and the radius of the orbit are quantized (bls) Quantized (definition): a specific value (of energy)
The ramp is an example of a continuous situation in which any energy state is possible up the ramp
Like a set of stairs, the energy states of an electron is quantized – i.e. electrons are only found on a specific step
Bohr’s Energy Absorption Process Light or energy excites an e- from a lower
energy level (e- shell) to a higher energy level
These energy levels are “ quantized “ (the e- cannot be in between levels), the e- disappears from one shell and reappears in another
This absorption or excitation process is called a quantum leap or quantum jump
Bohr’s Energy Absorption Process Ground State Analogy = a spring and two balls
This is an energy emission process and
what we observe in the hydrogen line spectrum
Both the atom and e- now have higher
energyThe e- absorbs energy in the ground state and is
excited to a higher level
Bohr’s Energy Absorption Process When energy is added, the electron is found in the
“excited state.” The Excited State (definition) = an unstable, higher
energy state of an atom An illustration of Bohr’s Hydrogen atom (from ground to
excited state):
Bohr’s Energy Absorption Process The atomic line spectral lines - when an e- in an excited
state decays back to the ground state
The electron loses energy, light (colors) is emitted and the e- returns to the ground state
This is another illustration of bls.
The Bohr Model - Summary
1. When an atom absorbs energy, its electrons are promoted to a higher energy level. When the electron drops back down, energy is given off in the form of light.
2. Each distance fallen back is a specific energy, and therefore, a specific color.
3. Since electrons can fall from level 5 to 4, 5 to 3, etc., many colors are produced.
The Bohr Model - Summary
Bohr also predicted that since electrons would occupy specific energy levels and each level holds a specific number of electrons
The maximum capacity of the first (or innermost) electron shell is two e-.
Any element with more than two e-, the extra e- reside in additional electron shells.
The Bohr Model - Summary
Group IA
Lithium
VIA
Oxygen
VIIA
Fluorine
VIIIA
Neon
IA
Sodium
Electron Configurations for Selected Elements
The number of e- per shell = 2n2 (where n is the shell number)
Short Hand Bohr Model
Write the symbol of the element Use a ) to represent each shell Write the # of e- in each shell Ex. Element Short-Hand e-
Configuration
Hydrogen H )1e-
Lithium Li )2e- )1e-
Fluorine F )2e- )7e- Sodium Na )2e- )8e- )1e-
The Truth About Bohr Models At atomic # 19 (z = 19), there is a a break in the pattern. One
would expect that energy level #3 would continue to fill up. However, the next two electrons go into the next energy level. Look at K and Ca:
IA VIIIA
H+ ) 1 IIA IIIA IVA VA VIA VIIA
He+ ) 2
Li+ ) ) 2 1
Be+ ) ) 2 2
B+ ) ) 2 3
C+ ) ) 2 4
N+ ) ) 2 5
O+ ) ) 2 6
F+ ) ) 2 7
Ne+ ) ) 2 8
Na+ ) ) ) 2 8 1
Mg+ ) ) ) 2 8 2
Al+ ) ) ) 2 8 3
Si+ ) ) ) 2 8 4
P+ ) ) ) 2 8 5
S+ ) ) ) 2 8 6
Cl+ ) ) ) 2 8 7
Ar+ ) ) ) 2 8 8
K+ ) ) ) ) 2 8 8 1
Ca+ ) ) ) ) 2 8 8 2
The Truth Continued……
So, there is a relationship between the main column # and the number of outershell electrons.
Column # = the number of valence electrons And, there is a relationship between the row # and
the number of energy levels. Row # = the number of shells The Bohr model truly works well for the H atom only
for elements larger than H the model does not work.
Bohr Summed Up
Bohr made 2 huge contributions to the development of modern atom theory He explained the atomic line spectra in terms of electron
energies He introduced the idea of quantized electron energy
levels in the atom
The Bohr atom lasted for about 13 years and was quickly replaced by the quantum mechanical model of the atom. The Bohr model is a good starting point for understanding the quantum mechanical model of the atom
Do ws# 1, question 1- Use short-hand configuration
Quantum Numbers & Atomic Orbitals
The Bohr model describes the atom as having definite orbitals occupied by electron particles.
Schrödinger (1926) introduced wave mechanics to describe electrons – proved Bohr’s Model to be a lie Based his idea that electrons behaved like light
(photons). Electrons show diffraction (interference)
properties like light. Treats electrons as waves that are found in
orbitals. Orbitals (definition) = clouds that show region of
probable location of a particular electron.
Wave Mechanical Model
The Bohr model really is the wave mechanical model
There are many types of orbitals – we can see them on the periodic table
Subatomic Orbitals
Type # sublevels Total # e Shape
s 1 2 sphere
p 3 6 peanut
d 5 10 dumbbell
f 7 14 flower
S P D
Quantum Numbers
An electron’s address principle (n): what shell, level, the e- is in n =
1,2,3...7
azimuthal (l): energy sub level - s, p, d, f
magnetic – orientation of orbital about the nucleus (s has only 1, p has 3, etc.)
spin - clockwise or counterclockwise (+1/2 or -1/2)
Label Your Periodic Tabel
On your periodic table, shade azimuthal s,p,d,f blocks different colors
Label the principal quantum numbers…1-7
Label the valence electrons across the top
Electron Configuration
Electron Configuration - a representation of the arrangement of electrons in an atom
Electron Configuration
Examples of electron Configuration 1. Li 1s22s1
2. C 1s22s22p6
principle azimuthal
# of e- in that shell
Electron Configuration
Take note that after 4s is filled, 3d is than filled before 4p.
…… 6s than 4f than 5d than 6p When writing out the electron
configuration, always write your numbers in numerical order Y 1s22s22p63s23p64s23d104p65s24d1 – NO!
Y 1s22s22p63s23p63d104s24p64d15s2
Electron Configuration
Short Hand Write the name of the last noble gas Write the electron config. that follows
Ex. Fe [Ar]3d64s2
Exceptions Cr [Ar] 3d54s1
Cu, Ag, Au- all s’s donate 1 e- to make the d orbital full
Cu [Ar] 4s13d10
Orbital Notation Electrons enter orbitals in a set pattern. For the most
part, they follow these rules: 1) The Aufbau Principle - electrons must fill lower
energy levels before entering higher levels.
Orbital Notation
Orbitals are like "rooms" within which electrons "reside".
The s subshell has one s-orbital. The p subshell has
three p-orbitals. The d subshell has 5 and f has 7.
Each orbital can hold at most 2 electrons
Orbital Notation
2. Hund’s Rule (better known as the Bus Rule) Before any second electron can be placed in a sub level,
all the orbitals of that sub level must contain at least one electron – spread out the e- before pairing them up.
3. Pauli Exclusion Principle - electrons occupying the same orbital must have opposite spin.
See a good online illustration at http://www.avogadro.co.uk/light/aufbau/aufbau.htm
Orbital Notation
We can also do shorthand orbital notation (outer shell only)
Ca N
Fe
Ag [Kr] 4d105s1
Ag [Kr]
4d 5s
Significance of Electron Configurations Valence shell electrons - outermost electrons involved
with bonding no atom has more than 8 valence electrons Noble gases - 8 valence electrons – least reactive of all
elements Lewis Dot structures: NSEW (cheating) also show
correct way, count to 8
Lewis Dot Structures