1. ATOMIC STRUCTURE
Electrons in Atoms

2. In the early 1900’s scientists studying the behavior of atoms observed that certain elements emitted visible light when heated by a flame or if subjected to a high voltage spark
Much of our understanding of how electrons behave in atoms comes from the study of light
In order to understand this, you must know about the nature of light…

3. LIGHT
Visible light (the light we see with our eyes) is a type of electromagnetic radiation
All other electromagnetic radiation is invisible

4. ELECTROMAGNETIC RADIATION
Electromagnetic (EM) radiation is the transmission and emission of energy in the form of electromagnetic waves
Visible light, X-rays, microwaves , infrared waves (IR), ultraviolet waves (UV), radio waves, etc.

5. ELECTROMAGNETIC SPECTRUM
The electromagnetic spectrum encompasses all forms of electromagnetic radiations

6. CHARACTERISTICS OF WAVES
A wave can be thought of as a vibrating disturbance by which energy is transmitted
Waves can be characterized by their:
Wavelength  (lambda): distance between identical points on a successive wave
nanometers (nm)
Frequency  (nu): the number of waves that pass through a particular point in 1 second
hertz (Hz); 1 Hz = 1 cycle/s
Amplitude: the vertical distance from the midline to the peak or trough

8. High Frequency
Low Frequency

9. The speed of a wave is a product of its wavelength and frequency:
SPEED OF LIGHT
Another important property of waves is their speed
dependent on type of wave and medium it’s traveling through
The speed of a wave is a product of its wavelength and frequency:
c = 
Notice  and  are inversely proportional
In a vacuum, All electromagnetic waves travel at the speed of light (3.00 x 108 m/s)

10. PLANCK’S QUANTUM THEORY
According to the theory atoms and molecules can emit or absorb energy only in discreet quantities, which Planck called “quantum”
quantum (meaning “fixed amount”) is the smallest quantity of energy that can be gained or lost by an atom
The energy E of a single quantum of energy is given by
E =h
where h is called Planck’s constant (6.63 x J•s) and  is the frequency of radiation
According to Planck’s theory energy is always released or absorbed by matter in whole-numbers of h, 2h, 3h …

11. THE PHOTOELECTRIC EFFECT
Albert Einstein used Planck’s quantum theory to explain the photoelectric effect
When light shines on a clean metal surface the surface emits electrons (can be converted to electrical energy)
For each metal there is a minimum frequency of light needed to emit electrons
red light in incapable of releasing electrons from sodium metal even if intense
faint violet light releases electrons easily

13. DUAL NATURE OF LIGHT Ephoton = h
In explaining the photoelectric effect Einstein proposed that light (electromagnetic radiation) has dual nature:
Can behave like a wave
Can behave like a stream of particles (photons)
PHOTON - a particle of electromagnetic radiation
A photon has no mass
A photon carries a quantum of energy
Einstein calculated that a photon’s energy depends on its frequency
Ephoton = h

14. ATOMIC EMISSION SPECTRA
Light of a neon sign is produced by passing electricity through a tube of neon gas
The atoms in the tube absorb energy and become excited and unstable
They become stable by releasing the energy as light

15. ATOMIC EMISSION SPECTRA
If the light emitted from neon is passed through a prism neon’s atomic emission spectrum is produced

16. ATOMIC EMISSION SPECTRA
The atomic emission spectrum of an element is its “fingerprint”
It’s the set of frequencies EM waves the element’s atoms emit (give off)
Each element’s atomic emission spectrum is unique and can be used to identify the element
Hydrogen’s emission spectrum:
violet
410 nm
blue-violet
434 nm
blue-green
486.1 nm
red
656.2 nm

18. Light Bulb
(white light)
Hydrogen Bulb

19. THE BOHR MODEL OF THE ATOM
Bohr proposed:
Electrons move around the nucleus in circular orbits (“rings”) with distinct energy levels
smaller orbits have lower energy, larger orbits higher energy
In other words, electrons found closer to the nucleus has less energy than electrons found at greater distances from the nucleus
Bohr assigned a quantum number (n) to each level

20. BOHR’S MODEL OF THE ATOM+
Energy
n = 7
n = 6
n = 5 -
n = 4
n = 3
n = 2
n = 1
This model is often called the planetary model

21. BOHR’S ATOM CONTINUED
The lowest energy state of an atom is its ground state
When an atom gains energy (through heating for example) it is in an excited state
in an excited state the electron absorbs the energy & jumps to higher energy level when it jumps back down to its ground state it releases excess energy in the form of light
Even though hydrogen contains only one electron, it can have many excited states

22. BOHR MODEL CONTINUED
Because electrons jump between orbitals that have specific energy levels only certain frequencies of electromagnetic radiation can be given off (only certain colors can be emitted):
If an excited electron drops from n=3 to n=2 red light is emitted
If an excited electron drops from n=4 to n=2 blue-green light is emitted
5-2: blue light
6-2: violet light
This is how Bohr explained hydrogen’s emission spectrum

23. E = h
E = h

25. Wait!
Bohr’s model explained the emission spectrum of Hydrogen, but it did not explain the emissions of any other element!
It was eventually found that Bohr was incorrect:
Electrons do not travel in circular orbits around the nucleus
In 1924, a young French scientist Louis de Broglie ( ) proposed an idea that eventually accounted for the fixed energy levels of Bohr’s model and better explained the behavior of electrons

26. THE QUANTUM MECHANICAL MODEL OF THE ATOM
If waves can act like particles, particles can act as waves!
Electrons behave like waves
Can’t know electrons position/path around the nucleus:
Electrons move about in a cloud around the nucleus in what appears to be a random pattern
The Quantum Model only predicts where an electron is likely to be found

27. HEISENBERG UNCERTAINTY PRINCIPLE
According to this principle it is fundamentally impossible to know the exact position and velocity of a particle at the same time
locating an electron produces uncertainty in the position and motion of the electron
It’s like trying to locate a helium filled balloon in a completely darkened room:
If you locate it by touch, you change it’s position-Once you find it, it’s already somewhere else!

28. SO WHERE CAN ELECTRONS BE FOUND?
In the quantum model, the nucleus is not surrounded by orbits, but by atomic orbitals
Atomic Orbital: a three-dimensional pocket of space around the nucleus that the electron is most likely to be found
An electron has a 90% chance of being found in the atomic orbital
That is the best we can do!
Electron probability density for hydrogen

29. Where 90% of the
e- density is found
for the 1s orbital
e- density (1s orbital) falls off rapidly
as distance from nucleus increases

31. ATOMIC ORBITAL ORGANIZATION
Principal energy level (n: 1-7)
(n) indicates relative size and energy of orbital. As n increases so do energy and size
Energy sublevels (s, p, d, f)
sublevels labeled according to shape
s: spherical; p: dumbbell; d/f: varied
Orbitals: Each sublevel has a specific number of orbitals:
s: 1orbital
p: 3 orbitals
d: 5 orbitals
f: 7 orbitals Each orbital can hold two electrons!!!

33. increasing
energy

34. Each orbital can hold two electrons4p
Energy 3d 4s 3p 3s H
Hydrogen 1
1.008 He
Helium 2
4.003 Li
Lithium 3
6.941 C
Carbon 6
12.01 B
Boron 5
10.81 Be
Beryllium 4
9.012 O
Oxygen 8
16.00 Ne
Neon 10
20.18 N
Nitrogen 7
14.01 F
Fluorine 9
19.00 2p       2s   1s  

35. ORDER OF ORBITALS (FILLING) IN MULTI-ELECTRON ATOM

36. ELECTRON CONFIGURATION
An atoms electron configuration is the way an atom’s electrons are distributed among the orbitals of an atom
The most state stable electron configuration is an atom’s ground state
Ground state: all electrons are in the lowest possible energy state
Electron configuration represented by writing symbol for the orbital and a superscript to indicate the number of electrons in the orbital Li: 1s2 2s1

37. The Pauli Exclusion Principle
The two electrons in an orbital must spin in opposite directions   1s  2s  2p       4s 3d 3s 3p
paramagnetic
unpaired electrons
diamagnetic
paired electrons

38. HUND’S RULE Negatively charged electrons repel each other, so:
Electrons won’t pair up unless they have to
Once there is one electron in every orbital…the pairing will begin! 2s 1s  2p  1.
Add an electron: 2s 1s  2p  2.
Add an electron: 2s 1s  2p  3.
Add an electron: 2s 1s  2p  4.

39. ELECTRON CONFIGURATION
The periodic table can be divided into four distinct blocks based on valence electron configuration
electron configuration explain the recurrence of physical and chemical properties

41. SHORTHAND (NOBLE GAS) NOTATION
Shows electron filling starting from previous noble gas:
Na: 1s22s22p63s1
Noble gas configuration: [Ne]3s1

42. Electron Configuration for Cation and Anions
What is the electron configuration for a sodium atom?
Na: 1s22s22p63s1
What is the electron configuration for a sodium ion (Na+)?
Na+: 1s22s22p6 or [Ne]

43. TRANSITION CATIONS
When transition metals form cations electrons are removed from the s-orbital before the d-orbital
Mn: [Ar]4s2 3d3
Mn2+: [Ar]3d3
This happens because d-orbitals are more stable than s-orbitals in transition metals
The s-orbitals are always in a higher energy level

44. ISOELECTRONIC
What do you notice about the following electron configurations:
F- : 1s22s22p6 or [Ne]
O2- : 1s22s22p6 or [Ne]
N3- : 1s22s22p6 or [Ne]
They have the same number of electrons and therefore identical ground-state electron configurations
These three ions are isoelectronic

45. VALENCE ELECTRONS
Valence electrons are the electrons found in the outermost orbitals of the atom (highest energy level)
They are the electrons involved in bonding and are responsible for the chemical properties of an element
Carbon has 4 valence electrons: 1s2 2s2 2p2
How many does magnesium have?

46. VALENCE ELECTRONS & GROUP NUMBER
One of the most important relationships in chemistry:
Atoms in the same group have similar chemical properties because they have the same number of valence electrons!
For the representative elements (group A elements)
group # = number of valence electrons
Exceptions: He is in group 8 but only has 2 valence electrons

47. The octet/duet rule
Atoms will gain, lose or share electrons to achieve noble gas configuration, meaning all atoms want a full outer orbital:
2 valence electrons for He
8 valence electrons for all other noble gases

48. ELECTRON-DOT STRUCTURE
Chemists often represent valence electrons in electron-dot structures
Electron-Dot Structure consists of the element’s symbol surrounded by dots representing the atom’s valence electrons
valence electrons are placed one to each of the four sides first,
when each side has one dot, you may begin doubling up S