1. The Atom
Chapter 2 Chem2A

4. Elements An element is a pure substance that cannot be
broken down into simpler substances by a
chemical reaction.
Consists of a dense core, surrounded by a negatively charged cloud
Contains three types of subatomic particles
Proton
Neutron
Electron

5. The Proton Charge = +1.602x10-19 C Located in the nucleus
Mass = x g
~1 atomic mass unit (amu)
1 amu = 1.66 x g
1/12 the weight of a 12C atom

6. The Neutron No charge (O C) Located in the nucleus
Mass = x g
~1 amu

7. The Electron Charge = -1.602 x 10-19 C
Located outside the nucleus in an e- “cloud”
Mass = x g
~0 amu

8. Plum Pudding Model
J.J. Thomson
1904

9. Gold Foil Experiment
Ernest Rutherford
1909

10. Thomson vs. Rutherford
Rutherford used the angle of deflection to calculate size of
the nucleus
Nucleus approximated at 10,000 times smaller than the
radius of the atom
Atom is actually mostly empty space

11. Atoms

14. Periods

15. Groups

16. Periodic Table of Elements

17. Classification of Periodic Table

18. Elements and Symbols C = Carbon N = Nitrogen O = Oxygen Cl = Chlorine
Ba = Barium
U = ?
Cf = ?
Bk = ?
Md = ?
Es = ?

19. K =
Pb =
W =
Sb =
Co =
Cu =
Cr =
Mg =
Mn =

20. Metals Shiny Conduct Electricity Ductile Malleable (Shapeable)
Can be drawn into a thin wire
Malleable (Shapeable)
High M.P. and B.P. RT
Except Hg

21. Non-Metals Don’t conduct well Not usually ductile Brittle
Low M.P. and B.P.
Many are gasses at RT

22. Metalloids Chemical characteristics in between metals and non-metals
Includes all atoms falling on the dividing line between metals and non metals

23. Classification of Periodic Table

24. Alkali Metals Group 1 (1A) Li, Na, K etc. Soft, shiny metals
Conduct heat and electricity
React violently with H2O
Form H2(g) and alkaline (basic) solutions

25. Akali(ne) Earth Metals
Group 2 (2A)
Be, Mg, Ca etc.
Not as reactive as Alkali Metals, but still quite reactive
Tend to make basic solutions when placed in water

26. Transition Metals Groups 3-12 Tend to have high densities and B.P.
All are metals
Often used for electrical conduction
Often have vivid colors when in solution
Used for pigments

27. Colors of Transition Metal Compounds
Nickel
Cobalt
Copper
Zinc
Iron

28. Lanthanides Elements 57-71 Lanthanum (La) to Lutetium (Lu)
Commonly used in lasers
Can deflect UV and infrared rays

29. Actinides/Actinoids Elements 89-103 Actinium (Ac) to Lawrencium (Lr)
Only Actinium, Thorium (Th), and Uranium (U) occur naturally
Others created by neutron bombardment
Radioactive

30. Groups 13(3A) – 16(6A) No common name
Boundary between metals and non-metals occurs here
Contain elements abundant in earth’s crust, atmosphere, and living things
Contains the metalloids

31. Halogens Group 17 (7A) Love to form salts with metals
NaCl, KBr, CaCl2
Like to form diatomic molecules
F2, Cl2, Br2

32. Noble Gases Group 18 (8A) Very unreactive
Don’t like to bond to other molecules
Generally not abundant

33. Diatomic Molecules
Dinitrogen (N2)
Molecules consisting of only two atoms of either the same or different elements O2 CO
Homonuclear Diatomic Molecule: a molecule made of two atoms of the same element H2
Heteronuclear Diatomic Molecule: a molecule made of two atoms that are different elements NO

34. Allotropism
The existence of multiple pure forms of an element, in the same phase (solid, liquid, or gas), that differ in structure
Different forms are called allotropes
Can exhibit varied physical properties and chemical behaviors
Don’t confuse allotropes with isotopes!!! O2 O3

36. 2.2 Structure of the Atom From the periodic table: Atomic Number
= # of protons
Atomic Symbol
Atomic Mass
Any given element is neutral
# of protons (+) = # of electrons (-)
# of neutrons = Atomic Mass (rounded) - # Protons
7 – 3 = 4 neutrons

37. Problems H C Fe Pd W
Atomic Weight
# Protons
# Electrons
# Neutrons

38. Calculator Boot Camp

39. 2.3 Isotopes A. Isotopes, Atomic Number, and Mass Number
Isotopes are atoms of the same element that have
a different number of neutrons.
the number of protons (Z)
Mass number (A) = +
the number of neutrons
Mass number (A) 35 37
Mass number (A) Cl Cl
Atomic number (Z) 17 17
Atomic number (Z)
# of protons = 17
# of protons = 17
# of electrons = 17
# of electrons = 17
# of neutrons = 35 – 17 = 18
# of neutrons = 37 – 17 = 20

40. Identify the atomic mass, number of protons, and number of neutrons for the following elements.

41. 2.3 Isotopes B. Atomic Weight
The atomic weight is the weighted average of the
masses of the naturally occurring isotopes of a
particular element reported in atomic mass units.
From the periodic table: 6 C
12.01
atomic number
element symbol
atomic weight (amu)

42. 2.3 Isotopes B. Atomic Weight
HOW TO Determine the Atomic Weight of an Element
What is the average atomic weight of
chlorine?
Example
List each isotope, its mass in atomic
mass units, and its abundance in nature.
Step [1]
Isotope
Mass (amu)
Isotopic Abundance
Cl-35
34.97
75.78% =
Cl-37
36.97
24.22% =

43. 2.3 Isotopes B. Atomic Weight
HOW TO Determine the Atomic Weight of an Element
Multiply the isotopic abundance by the mass
of each isotope, and add up the products.
Step [2]
The sum is the atomic weight of the element.

44. Using the percent abundances below, calculate the average atomic mass for Carbon
Antimony (Sb) has two stable isotopes, 121Sb and 123Sb with masses of u and u, respectively. Calculate the percent abundances of these two isotopes

46. Atomic Orbitals and Electron Configurations

47. The shells are numbered n = 1, 2, 3, etc.
Electronic Structure
Electrons are confined to discrete regions
called shells
The shells are numbered n = 1, 2, 3, etc.
Electrons in lower numbered shells are closer to
the nucleus and are lower in energy.
Electrons in higher numbered shells are further
from the nucleus and are higher in energy.

48. Electronic Structure
Shells with larger numbers (n) are farther from the nucleus and can hold more electrons.
# of electrons in each shell is calculated using 2(n2), where n = shell number.
Number of Electrons
in a Shell
Shell (n) 4 32 3 18
increasing
number of
electrons
increasing
energy 2 8 1 2

49. Electronic Structure
Shells are divided into subshells, identified by the letters s, p, d, and f.
The subshells consist of orbitals.
An orbital is a region of space where the probability of finding an electron is high.
Each orbital can hold two electrons.
Subshell
Number of Orbitals s 1
increasing
energy p 3 d 5 f 7

50. S Orbitals Spherical shape Lowest energy of the orbitals
Contains 1 orbital: s
1 orbital can hold 2 electrons MAX
= 2 electrons max

51. P Orbitals The p orbital has a dumbbell shape
Higher energy than S Orbitals
3 separate orbitals px, py, pz
Each orbitals can hold 2 electrons MAX
= 6 electrons max

52. D Orbitals Double dumb bell shape or single dumb bell with a donut
Higher energy than s and p
5 orbitals
dxy, dxz, dyz, dx2-y2, dz2
= 10 electrons max

53. F Orbitals Complex Shapes Highest Energy Orbital 7 Orbitals
=14 electrons max

55. Electronic Structure

56. Electron Configuration
The electron configuration shows how the electrons
are arranged in an atom’s orbitals.
The ground state is the lowest energy arrangement.
Rules to Determine the Ground State Electronic
Configuration of an Atom
Rule [1]
Electrons are placed in the lowest energy
orbital beginning with the 1s orbital.
Orbitals are then filled in order of
increasing energy.

57. Electron Configuration
Rules to Determine the Ground State Electronic
Configuration of an Atom

58. Electron Configuration
Rules to Determine the Ground State Electronic
Configuration of an Atom
Rule [2]
Each orbital holds a maximum of 2 electrons.
Rule [3]
When orbitals are equal in energy:
1 electron is added to each orbital until all
of the orbitals are half-filled.
Then, the orbitals can be completely filled.

59. Electron Configuration Orbital Diagrams
An orbital diagram uses a box to represent each
orbital and arrows to represent electrons.
an orbital
a single,
unpaired
electron
an electron
pair
Two electrons must have paired spins (opposite directions) to fit into the same orbital.

60. Electron Configuration A. First-Row Elements (Period 1)
Orbital
Notation
Electron
Configuration
Element
H (Z = 1)
1 electron
He (Z = 2)
2 electrons

61. Electron Configuration B. Second-Row Elements (Period 2)
Orbital
Notation
Electron
Configuration
Element
Li (Z = 3)
3 electrons
C (Z = 6)
6 electrons
Ne (Z = 10)
10 electrons

62. Problems
Determine the electron configuration and orbital diagrams for the following atoms N O Na Mg Fe W

63. Electron Configuration
The electron configuration can be shortened by using Noble Gas Notation.
Write the Symbol of the previous Noble Gas, then add the electronic configuration of the additional electrons.
element: C
previous
noble gas:

64. Electron Configuration C. Other ElementsCa
20 electrons

65. Electron Configurations and the Periodic Table

66. 2.7 Valence Electrons
The chemical properties of an element depend on the number of electrons in the valence shell.
The valence shell is the outermost shell (the highest value of n).
The electrons in the valence shell are called valence electrons. Be Cl
1s22s2
1s22s22p63s23p5
valence shell: n = 2
valence shell: n = 3
# of
valence electrons = 2
# of
valence electrons = 7

67. 2.7 Valence Electrons A. Relating Valence Electrons to Group Number
Elements in the same group have similar electron configurations.
Elements in the same group have the same number of valence electrons.
The group number, 1A–8A, equals the number of valence electrons for the main group elements.
The exception is He, which has only 2 valence electrons.
The chemical properties of a group are therefore very similar.

68. 2.7 Valence Electrons A. Relating Valence Electrons to Group Number
Period 1: H
1s1 He
1s2
Period 2: Li
2s1 Be
2s2 B
2s22p1 C
2s22p2 N
2s22p3 O
2s22p4 F
2s22p5 Ne
2s22p6
Period 3: Na
3s1 Mg
3s2 Al
3s23p1 Si
3s23p2 P
3s23p3 S
3s23p4 Cl
3s23p5 Ar
3s23p6

69. Valence Electrons B. Electron-Dot Symbols
Dots representing valence electrons are placed on the four sides of an element symbol.
Each dot represents one valence electron.
For 1 to 4 valence electrons, single dots are used. With more than 4 valence electrons, the dots are paired.
Element: H C O Cl
# of Valence electrons: 1 4 6 7
Electron-dot symbol: H C O Cl

70. Periodic Trends A. Atomic Size
The size of atoms
increases down a
column
Increases
Decreases
The size of atoms decreases across a row

71. Chapter 2 Review
The periodic table will be folded up for quiz and exam
Know the names and elemental symbols of s block, p block, and select d block atoms
W, Cu, Ag, Fe, Hg, Ti

72. Chapter 2 Review
Know how to classify atoms as metals, non-metals, and metalloids.
Know the period classifications
Alkali, Alkaline, Transition Metals, main block, halogens, noble gas

73. Classification of Periodic Table13 Al 84 Po

74. Chapter 2 Review
Location, and charge of the 3 particles that make up an atom
Proton (nucleus), Neutron (nucleus), Electron (diffuse cloud surrounding nucleus)
Determining how many protons, neutrons and electrons an atom has

75. Chapter 2 Review Orbitals s, p, d, f Know everything about s and p
Know how many electrons all can have
Know how they relate in energy

76. Chapter 2 Review What are isotopes
How do isotopes relate in terms of weight
How to calculate the average atomic weight

77. Chapter 2 Review
Know how to fill in (yes ill give you boxes) molecular orbitals
Know how to write noble gas abbreviations for any atom

78. Chapter 2 Review
Know what valence electrons are and how to ID them in an electron configuration
Know the relationship between atoms in the same period
Know how to write electron dot symbols