1. Atoms and the Periodic Table CHE 101 Sleevi
Chapter 2
Atoms and the Periodic Table
CHE 101
Sleevi

2. Atoms and the Periodic Table
Element
Substance that cannot be broken down into simpler substances by a chemical reaction
Identified by a one or two letter symbol
1st letter capitalized
2nd letter, if present, lowercase
Position on the periodic table based on chemical properties

3. Atoms and the Periodic Table3 3

4. The Periodic Table

5. Elements and the Periodic Table
Metals
Left of the stair-step line on the periodic table
Shiny (lustrous)
Malleable
Solids at room temperature except Hg which is a liquid
Good conductors of heat and electricity

6. Elements and the Periodic Table
Nonmetals
Right of the stair-step line on the periodic table
Usually not shiny
Can be solids, liquids or gases at room temperature (Br is the only liquid nonmetal)
Poor conductors of heat and electricity

7. Elements and the Periodic Table
Metalloids
Elements that abut the stair-step line on the periodic table
Have properties that are intermediate between metals and nonmetals
Silicon is brittle, but conducts electricity, is a shiny blackish silver color

8. Elements and the Periodic Table
Compounds
Substance formed by chemically combining two or more elements together
Use chemical symbols to describe the compound
Symbols identify the types of elements (atoms) in the compound
Subscripts show the ratio of the number of each type of atom in the smallest representative particle of the element

9. Chemical Formulas Examples H2O one molecule contains two H atoms
one O atom
C3H6 one molecule contains
three C atoms
six H atoms

10. Representations of Chemical Compounds
Typical colors used for atoms:

11. Elements of Biological Interest
Building block elements
Comprise 96% of mass of human body
C, O, H, N
Major nutrients (> 100 mg/day)
K, Na, Cl, Mg, S, Ca, P
Trace elements (< 15 mg/day)
As, B, Cr, Co, Cu F, I, Fe, Mn, Mo, Ni, Se, Si, Zn

12. Other Elements of Interest
Most abundant elements on Earth
O, Si, Al, Fe, Ca, Na, K, Mg, H, Ti
Make up 98.9% of mass of the crust, ocean and atmosphere

13. Atomic Structure
Since the late 19th Century, experimental evidence has determined the atom is composed of particles
Particle
Discovery
Proton
Goldstein
Electron
Thomson
Neutron
Chadwick

14. Characteristics of Fundamental Subatomic Particles
Symbol
Charge
Mass (g)
Relative Mass
(amu)
Location
Proton p +1
1.67 x 10-24 1
Nucleus
Neutron n
Electron e- -1
9.07 x 10-28
1/1836
OutsideNucleus 14

15. Structure of the Atom Nucleus: Electron cloud: location of protons
and neutrons
dense core of the atom
location of most of the
atom’s mass
location of electrons
comprises most of the atom’s volume 15 15

16. Atomic Structure Atoms are neutral (#p = #e-)
Nucleus is central, dense core of the atom
Most of the mass of the atom is in the nucleus
Mass of electrons negligible compared to mass of nucleus
Chemical properties dependent upon the electrons

17. Identification of Atoms
Atomic number (Z)
Number of protons
Identifies the element
Mass number (A)
= # of protons + # of neutrons
In neutral atom:
# of protons = # of electrons

18. Structure of the Atom Atomic Number
From the periodic table:
Atomic number (Z) is
the number of protons
in the nucleus. 3 Li
Every atom of a given element has the same
atomic number.
Every atom of a given element has the same
number of protons in the nucleus.
Different elements have different atomic numbers. 18 18

19. Quick Review Particle Symbol Charge Relative Mass Location in Atom
Purpose
Proton
Neutron
Electron

20. Isotopes 6C Chemically alike
Atoms with same atomic number but different mass number (different number of neutrons
Chemically alike 6C 12
Z = atomic number
A = mass number
Z = # p = 6
A = # p + # n = 12 20

21. element name – mass number
Isotopes
Isotopes often listed as:
element name – mass number
carbon-14
nitrogen-14
uranium-238
C-12
Pb-207
Write the chemical symbols for these isotopes, including both atomic number and mass number 21

22. Data on the Periodic Table
atomic number (Z) 6 C
12.01
element symbol
atomic weight (amu)
Note:
atomic weight also called atomic mass or average atomic mass
Mass number does NOT appear on the periodic table

23. 1 amu = 1/12 the mass of a C-12 atom
Average Atomic Mass
Weighted average of the mass of the stable, naturally occurring isotopes of an element
Ave at mass = Σ (rel. abund. * isotope mass)
Mass units for individual atoms in atomic mass units (1 amu)
1 amu = 1/12 the mass of a C-12 atom

24. Calculating the Average Atomic Mass
Data needed
Number of isotopes
Relative abundance of each isotope
Isotope mass
Calculation
Weighted average
Sum of products of rel abundance and isotope mass

25. Calculating the Average Atomic Mass
Example

26. The Modern Periodic Table
Based on work of Mendeleev, 1869
about 65 known elements
arranged elements based on trends in physical and chemical properties
left spaces for other elements, not yet discovered
generally organized according to increasing atomic weight 26

27. 27

28. Modern Periodic Table
Numerous people worked on periodic table at this time
Henry Moseley, British Physicist, 1913
determined nuclear charge of atoms (Z)
arranged elements in order of increasing atomic number 28

29. Periodic Law
When elements are arranged in order of increasing atomic number, there is a periodic pattern in their physical and chemical properties 29

30. Structure of the Periodic Table
Group
Vertical columns
Number appears above the column
Period
Row
May or not have written number
1st period contains H and He

31. Group Names Main group elements Transition Metal elements
Also called representative elements
Designated by Group numbers 1A – 8A
Transition Metal elements
Short columns in the middle (“short stack”)
Designated by Group numbers ending in B
Inner transition elements
Lanthanide and actinide series
No group numbers

32. Periodic Table of the Elements
Noble
Gases
Periodic Table of the Elements
Alkaline Earth
Metals 1A 2A 3A 4A 5A 6A 7A 8A
Halogens
Transition Metals
Alkali
Metals 3B 4B 5B 6B 7B 8B 1B 2B
Metalloids
Inner
Transition
Metals 32

33. Group Names of the Main Group Elements
Group Number
Group Name
Properties of Both Groups 1A
Alkali Metals
Soft, shiny metals
Low melting points
Good conductors (heat, electricity)
React with water to form basic solutions 2A
Alkaline Earth Metals

34. Group Names of the Main Group Elements
Group Number
Group Name
Properties 7A
Halogens
Exist as diatomic molecules
Very reactive 8A
Noble gases
Exist as monatomic species
Very stable
Rarely combine with other elements

35. Location of the Elements
Since periodic table is a grid, elements can be located by “coordinates” of their group number and period number
name the element in Group 2A, period 3
name the element in Group 5B, period 4
Element properties can be identified by knowing its group
name the group for Mg
name the group for Ne 35

36. Modern Atomic Theory and Electron Configurations
1911 – Rutherford’s Gold Foil Expt
1913 – Bohr proposed fixed orbits for electrons based on energy
studied hydrogen emission spectrum
Bohr’s model did not resolve all observations
Early 1900’s – Birth of Quantum Mechanics
1926 – Schrödinger proposed quantum mechanical model of the atom 36

37. Quantum Mechanical Model of the Atom
Principal energy levels (shells)
correspond to the discrete energy levels in the atom that can be occupied by electrons
designated by quantum number n (n = 1, 2, 3, 4, …)
contain sublevels
quantum number n designates number of sublevels at that energy level
Sublevels (subshells)
correspond to orbital type within a principal energy level (s, p, d, f) 37

38. Orbitals Solution to the Schrödinger equation
Shape describes the probability of where electron will be found in space, at that energy
Overlay the nucleus
Increase in size as the principle energy level increases
Contain no more than 2 electrons 38

39. Electronic Structure Orbital Shapes
The s orbital has a spherical shape.
The p orbital has a dumbbell shape. 39 39

40. Sublevels: Orbital Types & Shapes
spherical
1 possible s orbital at a given PEL p
dumbbell
3 possible p orbitals at a given PEL
px, py, pz d
more complex, multiple lobes
5 possible d orbitals at a given PEL 40

41. Energy Levels in an Atom
Principal Energy Level
(shells)
Sublevels
(subshells) 1 s 2
s, p 3
s, p, d 4
s, p, d, f 41

42. Orbitals and Electrons by Sublevel
Max # Electrons s 1 2 p 3 6 d 5 10 f 7 14 42

43. Max Number of Electrons by Principal Energy Level
Sublevels
Max # Electrons 1 s 2
s, p
2 + 6 = 8 3
s, p, d
= 18 4
s, p, d, f
= 32 43

44. Electronic Structure Shells
Shells with larger numbers (n) are farther from the
nucleus, have a larger volume, and can therefore
hold more electrons.
The distribution of electrons in the first four shells:
Number of Electrons
in a Shell
Shell (n) 4 32 3 18
increasing
number of
electrons
increasing
energy 2 8 1 2 44 44

45. Electron Configurations
Arrangement of electrons around the nucleus
Based on set of rules that comes from Quantum Mechanical model of the atom 45

46. Electron Configurations
Rules determine how orbitals are filled, based on number of electrons in the atom
Aufbau Principle
Pauli Exclusion Principle
Hund’s Rule 46

47. Aufbau Principle Electrons enter orbitals of lowest energy first
Filling order:
1s 2s 2p 3s 3p 4s 3d 4p 5s …..
Number in front = quantum number (PEL)
Letter designates sublevel (orbital type)
Maximum number of electrons depends upon the orbital type 47

48. Aufbau Diagram 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f48

49. Pauli Exclusion Principle
An orbital can contain 0, 1, or 2 electrons
Electrons have a property called spin
When there are two electrons in one orbital they have opposite spin 49

50. Hund’s Rule
When electrons occupy orbitals of equal energy, one electron enters each orbital with parallel spin a. b. c. 50

51. Following the rules… Elements through Vanadium follow these rules
Most elements after that follow the rules
There are exceptions 51

52. Writing Configurations
Diagram
Use horizontal lines for orbitals,
Arrows for electrons,
Direction of arrows for spin
Label rows with PEL and orbital designation (1s, 2s, 2p, etc.)
Follow Aufbau Rule, Pauli Exclusion Principle, Hund’s rule 52

53. Examples
Lecture Problems 53

54. Electron Configuration Shorthand
Write PEL and orbital type as number, letter
Use superscript for number of electrons in the orbital
1s22s22p3
Element can be identified by adding up the superscripts (# electrons) 54

55. Mapping Aufbau Diagram to the Periodic Table
Electrons enter s orbitals (1A, 2A)
Electrons enter p orbitals (3A – 8A)
Electrons enter d orbitals (Transition metals)
Electrons enter f orbitals (Inner transition metals)
Principal energy level for s and p given by the row number 55

56. Periodic Table of the Elements56

57. Electron Configurations and the Periodic Table
FIGURE 2.9 The Blocks of Elements in the Periodic Table

58. Electron Configurations
Element Symbol
Shorthand Notation
Noble Gas Notation Li
1s22s1
[He] 2s1 Na
1s22s22p63s1
[Ne] 3s1 K
1s22s22p63s23p64s1
[Ar] 4s1 Rb
1s22s22p63s23p64s24p65s1
[Kr]5s1 58

59. Unpaired Electrons

60. Practice!!
Lecture Problem 60

61. Valence Electrons Outermost electrons
Highest occupied principle energy level (largest quantum number)
Number of valence electrons = group number for the A groups
Determine the chemical properties of the elements
Determine number and type of bonds 61

62. Valence Configurations

63. More Practice!
Lecture Problems 63

64. Valence Electrons Electron-Dot Symbols
Dots representing valence electrons are placed
on the four sides of an element symbol.
Each dot represents one valence electron.
For 1–4 valence electrons, single dots are used.
With > 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 64 64

65. Practice! Draw Lewis dot structures for reach of the following:
S P Sr C
Br Kr Na Al

66. Periodic Trends
Properties of the elements are dependent upon their location on the periodic table (periodic law)
Properties have trends with respect to location of elements on periodic table
Atomic Size
Ionization Energy 66

67. Periodic Trends Atomic Size
The size of atoms
increases down a
column, as the
valence e− are
farther from the
nucleus.
Increases
Decreases
The size of atoms decreases across a row, as the number of protons in the nucleus increases.
The increasing # of protons pulls the e− closer to the nucleus, making the atoms smaller. 67 67

68. Periodic Trends Ionization Energy
The ionization energy is the energy needed to remove
an electron from a neutral atom.
Na energy  Na e–
Decreases
Ionization energies
decrease down a
column as the
valence e− get
farther away from
the positively charged nucleus.
Increases
Ionization energies increase across a row as the
number of protons in the nucleus increases. 68 68

69. Examples! Arrange the atoms in order of increasing size B, C, Ne
F, S, Al
Kr, Ne, Xe
Ca, Mg, Be

70. Examples! Arrange the atoms in order of increasing ionization energy
P, Si, S
Ne, Kr, Ar
C, F, Be
Ca, Al, N