The document discusses the development and key features of the periodic table. It describes how early scientists like Newlands, Meyer, and Mendeleev organized the elements based on properties like atomic mass and predicted new elements. Moseley later showed atomic number was fundamental. The modern periodic table is organized into rows (periods) and columns (groups) based on electron configuration. Elements in the same group have similar properties due to their valence electrons. The document outlines trends in properties like atomic radius, ionization energy, and electronegativity across periods and down groups.

1. The Periodic Table and
Periodic Law
Chemistry Chapter 6

2. Main IdeasMain Ideas
 Periodic trends in the properties of atoms allow us to
predict physical and chemical properties.
 The periodic table evolved over time as scientists
discovered more useful ways to compare and organize
elements.
 Elements are organized into different blocks in the
periodic table according to their electron
configurations.
 Trends among elements in the periodic table include
their size and their ability to lose or attract electrons

3. Development of theDevelopment of the
Modern Periodic TableModern Periodic Table
Objectives:
 Trace the development of the periodic table
 Identify key features of the periodic table

4. Development of theDevelopment of the
Periodic TablePeriodic Table
In the 1700’s, Lavoisier compiled a list of all
the known elements of the time.
 33 elements

5. Development of theDevelopment of the
Periodic TablePeriodic Table
The 1800s brought large amounts of
information and scientists needed a way to
organize knowledge about elements.
 Advent of electricity – break down
compounds
 Development of the spectrometer –
identify newly isolated elements

6. Development of theDevelopment of the
Periodic TablePeriodic Table
The 1800s brought large amounts of
information and scientists needed a way to
organize knowledge about elements.
 Industrial revolution – new chemistry
based ingredients and compounds.
 70 known elements by 1870

7. Development of theDevelopment of the
Periodic TablePeriodic Table
The 1800s brought large amounts of
information and scientists needed a way to
organize knowledge about elements.
 John Newlands proposed an arrangement
where elements were ordered by
increasing atomic mass.

8. Law of OctavesLaw of Octaves
Newlands (1864)
noticed when the
elements were
arranged by
increasing atomic
mass, their properties
repeated every eighth
element.

9. Law of OctavesLaw of Octaves
Octaves was used due
to the musical
analogy, but was
widely dismissed.
Some elements didn’t
follow the pattern

10. The Periodic TableThe Periodic Table
 Meyer and Mendeleev both demonstrated a
connection between atomic mass and
elemental properties.

11. The Periodic TableThe Periodic Table
 Mendeleev’s Table – A Russian scientist –
gets the most credit because he published
first.
 Arranged elements by increasing mass and
columns with similar properties.
 Predicted the existence and properties of
undiscovered elements.
 Still some inconsistencies.

12. The Periodic TableThe Periodic Table
 Moseley discovered that each element had
a distinct number of protons.
 Once rearranged by increasing atomic
number, the table resulted in a clear
periodic pattern.

13. The Periodic TableThe Periodic Table
Periodic repetition of chemical and physical
properties of the elements when they are
arranged by increasing atomic number is
called periodic law.

14. Development of theDevelopment of the
Periodic TablePeriodic Table

15. The Modern PeriodicThe Modern Periodic
TableTable
 The modern periodic table contains boxes which
contain the element's name, symbol, atomic
number, and atomic mass.

16. The Modern PeriodicThe Modern Periodic
TableTable
 Rows of elements are called periods. (total of 7)
 Columns of elements are called groups. (total of 18)
 Elements in groups 1,2, and 13-18 possess a wide
variety of chemical and physical properties and are
called the representative elements.
 Elements in groups 3-12 are known as the transition
elements .

17. Types of ElementsTypes of Elements
Elements are classified as metals, non-metals, and
metalloids.
 Metals are made up of most of the representative
elements and all of the transition elements.
 They are generally shiny when smooth and clean,
solid at room temperature, and good conductors of
heat and electricity.
 Most are Ductile and Malleable –
 Ductile – the ability to be drawn into wire.
 Malleable – the ability to be pounded into sheets

18. Types of ElementsTypes of Elements
Elements are classified as metals,
non-metals, and metalloids.
 Alkali metals are all the elements in group 1 except
hydrogen, and are very reactive.
 Alkaline earth metals are in group 2, and are also
highly reactive.

19. Types of ElementsTypes of Elements
The transition elements (groups 3 - 12) are divided
into transition metals and inner transition
metals.
 The two sets of inner transition metals are called
the lanthanide series and actinide series and are
located at the bottom of the periodic table.
 Lanthanides are phosphors – elements that
emit light when struck by electrons.

20. The Modern PeriodicThe Modern Periodic
TableTable
 Non-metals are elements that are generally gases
or brittle, dull-looking solids, and poor conductors
of heat and electricity.
 Group 17 is composed of highly reactive elements
called halogens.
 Group 18 gases are extremely unreactive and
commonly called noble gases.

21. The Modern PeriodicThe Modern Periodic
TableTable
 Metalloids have physical and chemical properties
of both metals and non-metals, such as silicon and
germanium. They are found along the stair step
of the table starting with Boron

23. QuestionsQuestions
What is a row of elements on the periodic
table called?
A. octave
B. period
C. group
D. transition

24. QuestionsQuestions
What is silicon an example of?
A. metal
B. non-metal
C. inner transition metal
D. metalloid

25. Practice ProblemsPractice Problems
 Page 181 #1-7

26. Classification of theClassification of the
ElementsElements
Objectives:
 Explain why elements in the same group
have similar properties.
 Identify the four blocks of the periodic table
on their electron configuration.

27. Organizing the ElementsOrganizing the Elements
by Electron Configurationby Electron Configuration
Electron configuration determines the
chemical properties of an element.
 Recall electrons in the highest principal
energy level are called valence electrons.

28. Organizing the ElementsOrganizing the Elements
by Electron Configurationby Electron Configuration
 All group 1 elements have one valence
electron.
 All group 2 elements have two valence
electrons.

29. Organizing the ElementsOrganizing the Elements
by Electron Configurationby Electron Configuration

30. Organizing the ElementsOrganizing the Elements
by Electron Configurationby Electron Configuration
 The energy level of an element’s valence electrons
indicates the period on the periodic table in which it is
found.
 The number of valence electrons for elements in groups
13-18 is ten less than their group number.
 After the s-orbital is filled, valence electrons occupy the p-
orbital.
 Groups 13-18 contain elements with completely or
partially filled p orbitals.

31. Organizing the ElementsOrganizing the Elements
by Electron Configurationby Electron Configuration

32. Organizing the ElementsOrganizing the Elements
by Electron Configurationby Electron Configuration

33. Organizing the ElementsOrganizing the Elements
by Electron Configurationby Electron Configuration
 The d-block contains the transition metals and is the
largest block.
 There are exceptions, but d-block elements usually
have filled outermost s orbital, and filled or partially
filled d orbital.
 The five d orbitals can hold 10 electrons, so the d-block
spans ten groups on the periodic table.

34. Organizing the ElementsOrganizing the Elements
by Electron Configurationby Electron Configuration
 The f-block contains the inner transition metals.
 f-block elements have filled or partially filled outermost
s orbitals and filled or partially filled 4f and 5f orbitals.
 The 7 f orbitals hold 14 electrons, and the inner
transition metals span 14 groups.

35. Practice ProblemsPractice Problems
 Page 186 #8-15

36. Periodic TrendsPeriodic Trends
Objectives:
 Compare period and group trends of several
properties.
 Relate period and group trends in atomic
radii to electron configuration

37. Atomic RadiusAtomic Radius
Atomic radius – is determined by the amount of
positive charge in the nucleus and the number of
valence electrons of an atom. It is usually measured
in picometers (10-12
).
 For metals, atomic radius is half the distance between
adjacent nuclei in a crystal of the element.
 For nonmetals, the atomic radius is the distance
between nuclei of identical atoms.

38. Atomic RadiusAtomic Radius

39. Atomic RadiusAtomic Radius
The periodic trend: decreases from left to right
(periods) and increases top to bottom (groups)
due to the increasing positive charge in the
nucleus.

40. Atomic RadiusAtomic Radius

41. Atomic RadiusAtomic Radius
 Atomic radius generally increases as you move
down a group.
 The outermost orbital size increases down a group,
making the atom larger.
 Valence electrons are not shielded from the
increasing nuclear charge because no additional
electrons come between the nucleus and the valence
electrons.

42. Ionic RadiusIonic Radius
Ions – atom(s) that gain or lose one or more electrons to
form a net charge.
Ionic radius is the radius of a charged atom.
 When atoms lose electrons and form positively charged
ions, they always become smaller.
 Lost electrons are usually valence electrons and could
leave the outer orbital empty and therefore smaller.
 Electrostatic repulsion between remaining electrons
decreases and pulls closer to nucleus.

43. Ionic RadiusIonic Radius
 When atoms gain electrons and forms a
negatively charged ion, they become larger.
 Increased electrostatic repulsion increases
distance of outer electrons.

44. Ionic RadiusIonic Radius
Periodic Trend: radius of an ion decreases from left
to right (periods) until charge changes and then
the radii increases dramatically. After the change,
the radius continues to decrease. Ionic radii
increases top to bottom (groups) until change in
charge.

45. Ionic RadiusIonic Radius

46. Ionization EnergyIonization Energy
Ionization energy is the energy needed to remove
an electron from the positive charge of the
nucleus of a gaseous atom. (how strongly a
nucleus holds on to an electron.)
 First ionization energy is the energy required to
remove the first electron.
 Removing the second electron requires more
energy, and is called the second ionization
energy.

47. Ionization EnergyIonization Energy
 Atoms with large ionization energies have a
strong hold of its electrons and are less likely to
form positive ions.
 Atoms with low ionization energies lose their
outer electrons easily and readily form positive
ions.
 The ionization at which the large increase in
energy occurs is related to the number of valence
electrons.

48. Ionization EnergyIonization Energy
Periodic Trend: First ionization energy increases from left
to right across a period. First ionization energy
decreases down a group because atomic size increases
and less energy is required to remove an electron
farther from the nucleus.

49. Ionization EnergyIonization Energy
 The octet rule states that atoms tend to gain, lose
or share electrons in order to acquire a full set of
eight valence electrons. The octet rule is useful for
predicting what types of ions an element is likely
to form.

50. Ionization EnergyIonization Energy

51. ElectronegativityElectronegativity
Electronegativity of an element indicates its
relative ability to attract electrons in a
chemical bond. Measured in Paulings:
numbers 4 and less.

52. ElectronegativityElectronegativity
Periodic Trend: electronegativity decreases down a group
and increases left to right across a period.

53. QuestionsQuestions
The lowest ionization energy is the ____.
A. first
B. second
C. third
D. fourth

54. QuestionsQuestions
The ionic radius of a negative ion becomes
larger when:
A. moving up a group
B. moving right to left across period
C. moving down a group
D. the ion loses electrons

55. Practice ProblemsPractice Problems
 Page 189 #16-19; Page 194 #20-23

56. Key ConceptsKey Concepts
 The elements were first organized by increasing
atomic mass, which led to inconsistencies. Later,
they were organized by increasing atomic number.
 The periodic law states that when the elements are
arranged by increasing atomic number, there is a
periodic repetition of their chemical and physical
properties.
 The periodic table organizes the elements into
periods (rows) and groups (columns); elements with
similar properties are in the same group.

57. Key ConceptsKey Concepts
 Elements are classified as either metals, nonmetals, or
metalloids.
 The periodic table has four blocks (s, p, d, f).
 Elements within a group have similar chemical properties.
 The group number for elements in groups 1 and 2 equals the
element’s number of valence electrons.
 The energy level of an atom’s valence electrons equals its
period number.

58. Key ConceptsKey Concepts
 Atomic and ionic radii decrease from left to right across
a period, and increase as you move down a group.
 Ionization energies generally increase from left to right
across a period, and decrease as you move down a
group.
 The octet rule states that atoms gain, lose, or share
electrons to acquire a full set of eight valence electrons.
 Electronegativity generally increases from left to right
across a period, and decreases as you move down a
group.

59. Chapter QuestionsChapter Questions
The actinide series is part of the
A. s-block elements.
B. inner transition metals.
C. non-metals.
D. alkali metals.

60. Chapter QuestionsChapter Questions
In their elemental state, which group has a
complete octet of valence electrons?
A. alkali metals
B. alkaline earth metals
C. halogens
D. noble gases

61. Chapter QuestionsChapter Questions
Which block contains the transition metals?
A. s-block
B. p-block
C. d-block
D. f-block

62. Chapter QuestionsChapter Questions
An element with a full octet has how
many valence electrons?
A. two
B. six
C. eight
D. ten

63. Chapter QuestionsChapter Questions
How many groups of elements are
there?
A. 8
B. 16
C. 18
D. 4

64. Chapter QuestionsChapter Questions
Which group of elements are the least
reactive?
A. alkali metals
B. inner transition metals
C. halogens
D. noble gases

65. Chapter QuestionsChapter Questions
On the modern periodic table, alkaline earth
metals are found only in ____.
A. group 1
B. s-block
C. p-block
D. groups 13–18

66. Chapter QuestionsChapter Questions
Bromine is a member of the
A. noble gases.
B. inner transition metals.
C. earth metals.
D. halogens.

67. Chapter QuestionsChapter Questions
How many groups does the d-block span?
A. two
B. six
C. ten
D. fourteen

68. THE END

69. Chapter QuestionsChapter Questions

70. Chapter QuestionsChapter Questions

71. Chapter QuestionsChapter Questions

72. Chapter QuestionsChapter Questions

73. Chapter QuestionsChapter Questions