1. Electrons in Atoms A chapter dedicated just for you electrons
Electrons in Atoms A chapter dedicated just for you electrons! Chapter 5

2. Light Waves Amplitude = wave’s height Wavelength = wave’s distance Frequency = wave cycles (hertz)

3. Electromagnetic radiation = A big word for Light

4. Electromagnetic Radiation
Waves have a frequency
Use the Greek letter “nu”, , for frequency, and units are “cycles per sec” or Hertz (Hz)
All radiation:  •  = c
where c = velocity of light = x 108 m/sec

5. Electromagnetic Spectrum
Long wavelength --> small frequency INVERSE
Short wavelength --> high frequency RELATIONSHIP

6. Electromagnetic Spectrum
In increasing energy, ROY G BIV

7. White light is actually all the colors together
A prism bends the different colors to where we can see them

8. Atomic Emission Spectra and Niels Bohr
Bohr’s greatest contribution to science was in building a simple model of the atom.
It was based on an understanding of the ATOMIC EMISSION SPECTRA of excited atoms
- Remember the Bohr-rings!
Niels Bohr
( )

9. Atomic Emission Spectra Excited Atoms
Excited atoms emit light of only certain wavelengths
The wavelengths of emitted light depend on the element
Each element has its own Atomic Emission Spectra… like a FINGERPRINT
ELECTRICITY or HEAT EXCITED ELECTRONS!

11. 5.2 = More about electrons Heisenberg Uncertainty Principle
Problem of defining nature of electrons in atoms solved by W. Heisenberg
Cannot simultaneously define the position of an electron
We define e- energy exactly but do not know exact position
W. Heisenberg

12. We know area the propeller is, but we cannot locate its exact position

13. Arrangement of Electrons in Atoms
Electrons in atoms are arranged as
LEVELS
ORBITALS
SUBLEVELS

14. QUANTUM NUMBERS n (principal) ---> energy level = FROM BOHR!
The shape, size, spin, and energy of each orbital is a function of 4 quantum numbers which describe the approximate location of an electron
n (principal) ---> energy level = FROM BOHR!
l (orbital) ---> shape of orbital
ml (magnetic) ---> designates a particular suborbital
s (spin) ---> spin of the electron (clockwise or counterclockwise)
Think of the 4 quantum numbers as the address of an electron… Country > State > City > Street

15. QUANTUM NUMBERS
So… if two electrons are in the same place at the same time, they must be repelling
(like charges repel!)
The Pauli Exclusion Principle says that an atomic orbital may describe at most 2 electrons. If two electrons are in the same energy level, the same sublevel, and the same orbital, they must repel = spin (S)
Pauli = spinning electrons (repel)

16. Energy Levels
Each energy level has a number called the PRINCIPAL QUANTUM NUMBER,
n (energy level)
Currently n can be 1 thru 7, because there are 7 periods on the periodic table
The period #s = energy levels

17. Energy Levels
n = 1
n = 2
n = 3
n = 4

18. Types of Orbitals
The most probable area to find these electrons takes on a shape
So far, we have 4 shapes
They are named s, p, d, and f
No more than 2 e- assigned to an orbital – one spins clockwise, one spins counterclockwise

19. Types of Orbitals (l)
s orbital
p orbital
d orbital

20. Relative sizes of the spherical 1s, 2s, and 3s orbitals of hydrogen.

21. Aufbau’s Principal
Electrons occupy the orbitals of lowest energy first!
Aufbau = electron order….
Least to most energy!

22. Orbitals and the Periodic Table
Orbitals grouped in s, p, d, and f orbitals (sharp, proximal, diffuse, and fundamental)
s orbitals
d orbitals
p orbitals
f orbitals

24. Electron Configurations
A list of all the electrons in an atom
Must go in order (Aufbau principle)
2 electrons per orbital, maximum (Pauli Ex.)
We need electron configurations so that we can determine the number of electrons in the outermost energy level
These are called valence electrons
The number of valence electrons determines how many and what this atom can bond to in order to make a molecule
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14… etc.

25. Electron Configurations
2p4
Number of electrons in the sublevel
Energy Level (n)
Sublevel
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14… etc.

26. Diagonal Rule 1 2 3 4 5 6 7 s s 2p s 3p 3d s 4p 4d 4f s 5p 5d 5f 5g?
Steps:
Write the energy levels top to bottom.
Write the orbitals in s, p, d, f order. Write the same number of orbitals as the energy level.
Draw diagonal lines from the top right to the bottom left.
To get the correct order, follow the arrows! 1 2 3 4 5 6 7 s
s 2p
s 3p 3d
s 4p 4d 4f
By this point, we are past the current periodic table so we can stop.
s 5p 5d 5f 5g?
s 6p 6d 6f 6g? 6h?
s 7p 7d 7f 7g? 7h? 7i?

27. Let’s Try It!
Write the electron configuration for the following elements:
H 1s1
Li 1s2 2s1
N 1s2 2s2 2p3
Ne 1s2 2s2 2p6
K 1s2 2s2 2p6 3s2 3p6 4s1
Zn 1s2 2s2 2p6 3s2 3p6 4s2 3d10

29. Orbital Diagrams Graphical representation of an electron configuration
One arrow represents one electron
Shows spin and which orbital within a sublevel

30. Orbital Diagrams Hund’s Rule
In orbitals of EQUAL ENERGY (p, d, and f), place one electron in each orbital before making any pairs
All single electrons must spin the same way
I nickname this rule the “Monopoly Rule”
In Monopoly, you have to build houses EVENLY. You can not put 2 houses on a property until all the properties have at least 1 house.

31. Lithium
Group 1A
Atomic number = 3
1s22s1 ---> 3 total electrons

32. Carbon Group 4A Atomic number = 6 1s2 2s2 2p2 --->
6 total electrons
HUND’S RULE!
Same energy level
= 1 electron until
All filled

33. Draw these orbital diagrams!
Lithium (Li)
Beryllium (Be) it helps to do the
Boron (B) e- configuration
Carbon (C) first
Nitrogen (N)
Neon (Ne)
Sodium (Na)

36. Ch. 6.1 & 6.2 Dmitri Mendeleev (1869)
Mendeleev published classification schemes for elements known to date.
The periodic table has organized the similar properties and reactivities by certain elements
Later, Henri Moseley established that each elements has a unique atomic number

37. The Periodic Table

38. Periodic Law
When elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties

39. Periodic Table Expanded View
The Periodic Table can be arrange by subshells
The s-block is Group IA and & IIA
The p-block is Group IIIA - VIIIA
The d-block is the transition metals
The f-block are the Lanthanides and Actinide metals

40. Periodic Table: Metallic arrangement
Layout of the Periodic Table:
Metals vs. nonmetals

41. METALS 80 percent of all elements Are GOOD conductors of
heat and electric current
(Give up electrons)
Have a high luster or sheen
Reflects light
All metals are solids at room temperature (except mercury Hg)
They are ductile (can be drawn into wires)
They are malleable (can be hammered down)

42. NONMETALS Upper-right corner Variation in physical properties:
Accept electrons
Most are gas at room temperature (N and O)
Few are solids (sulfur and phosphorus)
One is a liquid (bromine—dark red)
Most nonmetals have properties opposite of metals:
Poor conductors (except carbon)
Brittle when solid

43. METALLIODS Hug the stair-step line
Have properties similar to both metals and nonmetals, depending on certain conditions
Give up or accept electrons
EXAMPLE:
Pure silicon is a poor conductor of electrical current
But… if silicon is mixed with boron, it is a great conductor of electrical current
Used to make computer chips

44. Across the Periodic Table
Periods: Are arranged horizontally across the periodic table These elements have the same number of valence shells.
2nd Period
6th Period

45. Down the Periodic Table
Family: Are arranged vertically down the periodic table = These elements have the same number electrons in the the valence shell

46. Groups
The specific colors of the squares distinguish the groups of elements
Group 1A = Alkali Metals
Group 2A = Alkaline Earth Metals
Alkali is the arabic word al aqali which means wood ashes
Wood ashes are rich in Na and K

48. Groups
The specific colors of the squares distinguish the groups of elements Group 7A = Halogens Halogen comes from the Greek word hals (salt) and the Latin word genesis (to be born) Chlorine, bromine, and iodine can be prepared from their salts

49. Electron Configurations
Elements can be sorted into noble gases, representative elements, transition metals, or inner transition metals based on their electron configurations The noble gases are group 8A… they rarely take place in reactions IF YOU LOOK AT THEIR CONFIGURATIONS… THERE VALENCE SHELL IS FULL!!!

50. Infamous Families of the Periodic Table
Notable families of the Periodic Table

51. Representative Elements
Groups 1A- 7A They display a wide range of physical and chemical properties For representative elements… The group number is equal to the number of electrons in the highest occupied energy level

52. Transition Elements
Transition metals are characterized by electrons in the d sublevel Inner transition metals are characterized by electrons in the f sublevel The periodic table is organized into ‘blocks’ Based on sublevels (except for helium)

54. General Periodic Trends 6.3
Atomic and ionic size
Ionization energy
(electronegativity– learn in future chapter)

55. Atomic Radius
Is one-half the distance between the nuclei of two atoms of the same element when atoms are joined

56. Atomic Size Size increases- going down a group
Each element has 1 more proton and 1 more electron than the preceding one
The increasing nuclear charge pulls the electrons in the highest occupied energy level closer to the nucleus and the atomic size decreases
Size decreases- going across a period

57. Which is Bigger?
Na or K ? K
Na or Mg ? Na
Al or I ? Al

58. What are IONS??
An atom or group of atoms that has a positive or negative charge
An atom is electrically neutral because it has equal numbers of protons and electrons…
EX: Na has 11 protons and 11 electrons = 0 charge
Positive (cation) and negative (anion) ions form when electrons are transferred between atoms
ELECTRONS determine a charge!

59. IONS Cations = positive charge Lose electrons
Usually in metallic elements
Anions = negative charge
Gain electrons
Usually nonmetallic elements
How do you think the trend is going go?

60. Ion Sizes Does the size go
up or down when losing an electron to form a cation?

61. Ion Sizes CATIONS are SMALLER than the atoms from which they comeLi +
, 78 pm
2e and 3 p
Forming a cation
Li,152 pm
3e and 3p
CATIONS are SMALLER than the atoms from which they come
The electron/proton attraction has gone UP and so size DECREASES
(forced closer to nucleus)

62. Ion Sizes
Does the size go up or down when gaining an electron to form an anion?

63. Ion Sizes Forming an anion-
, 133 pm
10 e and 9 p
F, 71 pm
9e and 9p
Forming an anion
ANIONS are LARGER than the atoms from which they come
The electron/proton attraction has gone DOWN and so size INCREASES

64. Cl or Cl2+ ? Cl K+ or K ? K F or F-? F- I- or Br- ? I- (b/c size)
Which is Bigger?
Cl or Cl2+ ? Cl
K+ or K ? K
F or F-? F-
I- or Br- ? I- (b/c size)

65. Ionization Energy
IE = energy required to remove an electron from an atom (in the gas phase)
This is called the FIRST ionization energy because we removed only the OUTERMOST electron

66. Trends in Ionization Energy
IE increases across a period because the positive charge increases.
Metals lose electrons more easily than nonmetals
Nonmetals lose electrons with difficulty (they like to GAIN electrons)

67. Trends in Ionization Energy
IE increases UP a group
Because size increases

69. Which has a higher 1st ionization energy?
Mg or Ca ? Mg
Al or S ? S
Cs or Ba ? Ba

70. Electronegativity
Is a measure of the ability of an atom in a molecule to attract electrons
Will talk about more later…