1. Atoms & the Periodic Table

2. Chapter Outline What is Atom?
Chemical properties of Atoms: the Periodicity
Isotopes
Electrons in Atom: Quantum physics’ view
Valence electrons and the Periodic Table
Periodic trend: Atomic Radius, Metallic Character, Ionization Energy

3. Experiencing Atoms Atoms: incredibly small, yet compose everything
atoms are the pieces of Elements
properties of the atoms determine the properties of the elements

4. Within an Atom Atoms = (Protons + Neutrons) + Electrons
The nucleus (Protons + Neutrons) is only about cm in diameter yet with most of the mass of the atom
The electrons move outside the nucleus with an average distance of about 10-8 cm
The atom is electrically neutral : #proton (#p) = #electron (#e)

5. Comparison among Proton, Electron, Neutron

6. Elements each Element has a unique number of protons in its nucleus
Atomic number: the number of Protons in the nucleus of an atom
the elements are arranged on the Periodic Table in order of their atomic numbers
Name and Symbol of an Element
symbol either one or two letters
one capital letter or one capital letter + one lower case

7. The Periodic Table of Elements

8. The Size of Atoms Atomic Mass Unit (amu): 1 amu = 1.66  10-24 g
Hydrogen the smallest atom
mass of H atom= 1.67 x 10-24g ~ 1 amu
volume of H atom = 2.1 x 10-25cm3

9. Isotopes The same element could have atoms with different masses
Examples:
2 isotopes of chlorine atoms in nature: one weighs about 35 amu (Cl-35); another weighs about 37 amu (Cl-37)
Carbon-12 (C-12) is much more abundant than C-13.

10. Isotopes all isotopes of an element: chemically identical
undergo the exact same chemical reactions
the same number of protons
different masses due to different numbers of neutrons. Example: C-14 atom has eight neutrons; C-12 atom has six neutrons.
identified by their mass numbers
protons + neutrons

11. Isotopes Atomic Number (Z) Mass Number (A)
Number of protons
Mass Number (A)
Protons + Neutrons
Abundance = relative amount found in a sample
Example: Cl-35 (75%) vs. Cl-37 (25%)

12. Isotopic Symbol
Cl-35 has a mass number = 35, 17 protons and 18 neutrons ( ). The symbol for this isotope would be Cl 35 17
Atomic Symbol
A = mass number
Z = atomic number
#neutrons = A - Z AX Z

13. Example: How many protons, neutrons, and electrons in an atom of
Isotopic symbol  element
atomic number
#p  #e #n
Mass number = Atomic number (# protons, or #p) + #neutrons
U = uranium
Atomic Number = 92
#p = atomic number = 92
#e = #p = 92
Mass Number = #p + #n
238 = 92 + #n
146 = #n
#proton = 92
#neutron = 146
#electron = 92

14. Mass Number is Not the Same as Atomic Mass
Atomic mass (or Atomic Weight) is an experimental number determined from all naturally occurring isotopes
Mass number refers to the number of protons + neutrons in one isotope
natural or man-made
When given the relative abundance of all isotopes, we can find the Atomic mass

15. The Modern Periodic Table
Elements with similar chemical and physical properties are in the same column
columns are called Groups or Families
designated by a number and letter at top
rows are called Periods
each period shows the pattern of properties repeated in the next period

16. The Modern Periodic Table
Main Group = Representative Elements = ‘A’ groups
Transition Metals = ‘B’ groups
Aka Transition elements
Inner Transition Elements = Bottom rows = Rare Earth Elements
metals
really belong in Period 6 & 7

17. Main group vs. Transition metals, Inner transition metals
= Metalloid
= Nonmetal IA
VIIIA
IIA
IIIA
IIIB
VIIB
VIIIB IB
IIB

18. Metals: Physical vs. Chemical Properties
solids at room temperature, except Hg
reflective surface
shiny
conduct heat, electricity
Malleable (can be shaped)
Tend to Lose electrons and form Cations in reactions. Na  Na+ + e -
about 75% of the elements are metals
lower left on the table

19. Nonmetals: Physical vs. Chemical Properties
Elements found in all 3 states
poor conductors of heat or electricity
solids are brittle
Tend to gain electrons in reactions to become anions: Cl + e -  Cl-
upper right on the table
except H

20. Metalloids: between Metals and Nonmetals
show some properties of metals and some of nonmetals
also known as semiconductors
Properties of Silicon
shiny
conducts electricity
does not conduct heat well
brittle

21. = Alkaline Earth Metals
= Alkali Metals
= Alkaline Earth Metals
= Noble Gases
= Halogens
= Lanthanides
= Actinides
= Transition Metals
add pictures of elements from text

22. = Transuranium element
= Transition Metals
= Rare Earth Metals
= Transuranium element
add pictures of elements from text U

23. Important Element - Hydrogen
nonmetal
colorless, diatomic gas H2
very low melting point & density
reacts with Nonmetals to form molecular compounds
HCl is acidic gas
H2O is a liquid
reacts with Metals to form hydrides
metal hydrides react with water to form H2
Nickel-metal hydride (NiMH) used in rechargeable battery
HX dissolves in water to form acids

24. Important Element - Carbon
Three forms of pure carbon:
Diamond: hardest substance in nature
Graphite: soft and slippery solid
Buckminsterfullerene: a molecule made of 60
(Images from public domain wikipedia.com)
carbon atoms in a sphere

25. Carbon as backbone for Organic/Biochemical Molecules
Carbon atoms capable of forming robust bonds with many other elements and themselves.
Examples:
Small molecules: Butane, Sugar, Fatty acid, Vitamins
Big molecules (Polymers): Starch, Kevlar, Teflon, Protein, and DNA

26. Group IA: Alkali Metals
Usually Hydrogen is included
All metals: soft, low melting points
Flame tests ® Li = red, Na = yellow, K = violet
Chemical Property:
Very reactive. React with water to form basic (alkaline) solutions and H2.
releases a lot of heat
Tend to form water soluble compounds, such as table salt and baking soda.
colorless solutions
lithium
sodium
potassium
rubidium
cesium

27. Group IIA: Alkali Earth Metals
Physical properties: harder, higher melting, and denser than alkali metals
flame tests ® Ca = red, Sr = red, Ba = yellow-green
Chemical properties:
reactive, but less than corresponding alkali metal
form stable, insoluble oxides. oxides are basic = alkaline earth
reactivity with water to form H2, ® Be = none; Mg = steam; Ca, Sr, Ba = cold water
beryllium
magnesium
calcium
strontium
barium

28. Group VIIA: Halogens nonmetals F2 & Cl2 gases; Br2 liquid; I2 solid
all diatomic
very reactive
Cl2, Br2 react slowly with water
Cl2 + H2O ® HCl + HOCl
(chlorine)
react with metals to form ionic compounds
HX all acids
HF weak < HCl < HBr < HI
fluorine
chlorine
bromine
iodine

29. Group VIIIA: Noble Gases
all gases at room temperature,
very low melting and boiling points
very unreactive, practically inert
very hard to remove electron from or give an electron to

30. Atomic Orbitals Quantum Physicists including Schrödinger:
Electrons move very fast around the nucleus
Electrons show up with a particular probability at certain location of the atom
Orbital: A region where the electrons show up a very high probability when it has a particular amount of energy
generally set at 90 or 95%

31. Electron Shells Cl (17p+ & 17e-)
Electron Shell: the main energy level for the orbital. Principal quantum number n = 1, 2, …
For a chlorine atom, three shells of electrons:
The innermost shell (n = 1, 2 electrons) has the lowest energy
The outmost shell (n = 3, 7 electrons) has the highest energy
17 p+
2e-
8e-
7e-
Cl (17p+ & 17e-)

32. Each Shell have Subshells
Each Electron Shell has one or more Subshells (s, p, d, f)
the number of subshells = the Principal quantum number n
n = 1, one subshell (1s);
n = 2, two subshells (2s, 2p)
n = 3, three subshells (3s, 3p, 3d)
n = 4, three subshells (4s, 4p, 4d, 4f)
each Subshell has orbitals with a particular shape
the shape represents the probability map
90% probability of finding electron in that region

33. Shapes of Subshells
s Orbital
p Orbitals: px , py , pz
d Orbitals

34. f orbitals
Tro: Chemistry: A Molecular Approach, 2/e 34

35. Shapes of f orbitals: 4f orbitals (downloaded from public domain) The coloration corresponds to the sign of function.

36. Shells & Subshells

37. How does the 1s Subshell Differ from the 2s Subshell?

38. Subshells and Orbitals
Among the subshells of a principal shell, slightly different energies
for multielectron atoms, the subshells have different energies: s < p < d < f
each subshell contains one or more Orbitals
s : 1 orbital
p : 3 orbitals
d : 5 orbitals
f : 7 orbitals
within one subshell, different orbitals have the same energy. Example: 2px, 2py and 2pz

39. Electron Configurations
Definition: The distribution of electrons into the various energy shells (n = 1,2,3,…) and subshells (s, p, d, f) in an atom in its ground state
Each energy shell and subshell has a maximum number of electrons it can hold
Subshell s = 2, p = 6, d = 10, f = 14
Shell n: 1 = 2e, 2 = 8e, 3 = 18e, 4 = 32e
Electrons fill in the energy shells and subshells in order of energy, from low energy up
Aufbau Principal (“Construction” in German)

40. 6s 6p 6d 7s 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d
Energy 2s 2p 1s

41. Order of Subshell Filling in Ground State Electron Configurations
1. Diagram
putting each energy shell on
a row and listing the subshells,
(s, p, d, f), for that shell in
order of energy, (left-to-right) 1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 7s
2. draw arrows through
the diagonals, looping back
to the next diagonal
each time

42. Spinning Electron(s) in Orbital
Experiments showed Electrons spin on an axis
generating their own magnetic field
Pauli Exclusion Principle
each Orbital may have a maximum of 2 electrons, with opposite spin
Two electrons sharing the same orbital must have Opposite spins
so there magnetic fields will cancel
analogous to two bar magnets in parallel: only opposite alignment could stabilize each other.

43. Orbital Diagrams often an orbital as a square
the electrons in that orbital as arrows
the direction of the arrow represents the spin of the electron
unoccupied
orbital
orbital with
1 electron
orbital with
2 electrons

44. Filling the Orbitals in a Subshell with Electrons
Energy shells fill from lowest energy to high
1 → 2 → 3 → 4
Subshells fill from lowest energy to high
s → p → d → f
Orbitals of the same subshell have the same energy. Three 2p orbitals; Five 3d orbitals
 Electrons prefer “spreading out” in orbitals of same subshell before they pair up in orbitals.
Hund’s Rule
Example: 2p3 _ _ _ instead of  ____

45. Electron Configuration of Atoms in their Ground State
Electron configuration: a listing of the subshells in order of filling with the number of electrons in that subshell written as a superscript
Kr = 36 electrons = 1s22s22p63s23p64s23d104p6
a shorthand way : use the symbol of the previous noble gas in [] for the inner electrons, then just write the last set
Rb = 37 electrons = 1s22s22p63s23p64s23d104p65s1 = [Kr]5s1

46. Example: Ground State Orbital Diagram and Electron Configuration of Magnesium
1s22s22p63s2 = [Ne]3s2 1s 2s 2p 3s 3p 

47. Practice: Write Electron Configuration for the following atoms at the Ground state
Calcium
Sulfur
Sodium
Chlorine
Important: you are required to be able to write electron configuration for first three rows of elements

48. Valence Electron vs. Core Electron
Valence Electron: the electrons in all the subshells with the highest principal energy shell
Example: electrons in bold
Mg = [Ne]3s2 O = [He]2s22p4
Br = [Ar]4s23d104p5
Core electrons: electrons in lower energy shells
Chemists have observed that one of the most important factors in the way an atom behaves, both chemically and physically, is the Number of Valence electrons

49. Valence Electrons Rb = 37 electrons = 1s22s22p63s23p64s23d104p65s1
the highest principal energy shell that contains electrons is the 5th : 1 valence electron + 36 core electrons
Kr = 36 electrons = 1s22s22p63s23p64s23d104p6
the highest principal energy shell that contains electrons is the 4th : 8 valence electrons + 28 core electrons

50. Electrons Configurations and the Periodic Table

51. Electron Configurations from the Periodic Table
Example: Be 2s2 B 2s22p1 C 2s22p2 N 2s22p3 O 2s22p4
Elements in the same period (row) have Valence Electrons in the same principal energy shell.
#Valence electrons increases by one from left to right
Elements in the same group have the same #valence electron and they are same kind of subshell
Example: IIA: Be 2s2 Ca 3s2 Sr 4s2 Ba 5s2
VIIA: F 2s22p5 Cl 3s23p5 Br 4s24p5 I 5s25p5

52. Electron Configuration & the Periodic Table
Elements in the same Group have similar chemical and physical properties  their valence shell electron configuration is the same
No. Valence electrons for the main group elements is the same as the Group Number
Example:
Group IA: ns1 ;
Group IIIA: ns2np1
Group VIIA: ns2np5

53. Electron Configuration & the Periodic Tables1 s2
p1 p2 p3 p4 p5 s2 1 2 3 4 5 6 7 p6
d1 d2 d3 d4 d5 d6 d7 d8 d9 d10
f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14

54. Electron Configuration from the Periodic Table
Inner electron configuration = Noble gas of the preceding period
Outer electron configuration: from the preceding Noble gas the next period (Subshells)  Element
the valence energy shell = the period number
the d block is always one energy shell below the period number and the f is two energy shells below

55. Electron configuration & Chemical Reactivity
Chemical properties of the elements are largely determined by No. Valence electrons
Why elements in groups? Since elements in the same column have the same #valence electrons, they show similar properties

56. Electron Configuration: Noble Gas
Noble gases have 8 valence electrons
except for He, which has only 2 electrons
Noble gases are especially nonreactive
He and Ne are practically inert
 The reason: the electron configuration of the noble gases is especially stable

57. Everyone Wants to Be Like a Noble Gas! Alkali Metals (Group 1A)
have one more electron than the previous noble gas, [NG]ns1
tend to lose their extra ONE electron, resulting in the same electron configuration as a noble gas
forming a cation with a 1+ charge
Na  Na+
Li  Li+

58. Everyone Wants to Be Like a Noble Gas! Halogens (Group 7A)
one fewer electron than the next noble gas: [NG]ns2np5
Reactions with Metals: tend to gain an electron and attain the electron configuration of the next noble gas: [NG]ns2np5 + 1e  [NG]ns2np6
forming an anion with charge 1-: Cl  Cl-
Reactions with Nonmetals: tend to share electrons so that each attains the electron configuration of a noble gas

59. Everyone Wants to Be Like a Noble Gas! Summary
Alkali Metals as a group are the most reactive metals
they react with many things and do so rapidly
Halogens are the most reactive group of nonmetals
one reason for their high reactivity: they are only ONE electron away from having a very stable electron configuration
the same as a noble gas

60. Stable Electron Configuration And Ion Charge
Metals:  Cations
by losing enough electrons to get the same electron configuration as the previous noble gas
Nonmetals:  Anions
by gaining enough electrons to get the same electron configuration as the next noble gas

61. Trends in Atomic Size Increases down a group
valence shell farther from nucleus
effective nuclear charge fairly close
Decreases across a period (left to right)
adding electrons to same valence shell
effective nuclear charge increases
valence shell held closer

62. Trends in Atomic Size

63. Metallic Character Metals Nonmetals malleable & ductile
shiny, lusterous, reflect light
conduct heat and electricity
most oxides basic and ionic
form cations in solution
lose electrons in reactions – oxidized
Nonmetals
brittle in solid state
dull
electrical and thermal insulators
most oxides are acidic and molecular
form anions and polyatomic anions
gain electrons in reactions - reduced

64. Trends in Metallic Character

65. Electron Configuration Affects the Size of Atoms and Metallic Character: Within a Group
Within the same Group, from top to bottom:
As quantum number n increases for the valence electron(s)
valence electron(s) further away from the nucleus
Larger Atomic Radius
weaker Coulombic force (electrostatic force) withholding valence electrons
electrons easier to be lost
Stronger metallic character

66. Be (4p+ & 4e-) Mg (12p+ & 12e-) Ca (20p+ & 20e-) Example: Group IIA

67. Electron Configuration Affects the Size of Atoms and Metallic Character: Over the Period
Within the same Period (row), from left to right:
Same quantum number n for the valence electron(s)
As Nucleus has increasing number of protons (p+)
Stronger Coulombic force (electrostatic force) withholding valence electrons
Valence Electrons closer the nuclues
Smaller Atomic Radius
Valence electrons harder to be lost
Weaker metallic character

68. Li (3p+ & 3e-) Be (4p+ & 4e-) B (5p+ & 5e-) O (8p+ & 8e-)
Example: Period 2
2e-
1e- 3+
2e- 4+
2e-
3e- 5+
Li (3p+ & 3e-)
Be (4p+ & 4e-)
B (5p+ & 5e-) 6+
2e-
4e- 8+
2e-
6e-
10+
2e-
8e-
O (8p+ & 8e-)
C (6p+ & 6e-)
Ne (10p+ & 10e-)

69. Ionization Energy (IE)
In an atom, electrons (“-” charge) are attracted to the nucleus (“+” charge). Energy is required to remove the electron from an atom.
Na + energy  Na+ + e-
Neutral atom IE Cation
Higher IE corresponds to lower Metallic property.

70. Trends in Ionization Energy
Decreases down a group
valence shell farther from nucleus
effective nuclear charge fairly close
Increases across a period (left to right)
adding electrons to same valence shell
effective nuclear charge increases
valence shell held closer

71. Practice: Rank elements K, Mg, S, F: A. increasing metallic character B. increasing atomic radii C. increasing ionization energy
F, S, Mg, K
K, Mg, S, F