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The Periodic Table and Periodic Law
Section 6.1 Development of the Modern Periodic Table
Section 6.2 Classification of the Elements
Section 6.3 Periodic Trends
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3. Section 6.1 Development of the Modern Periodic Table
Trace the development of the periodic table.
Identify key features of the periodic table.
atomic number: the number of protons in an atom
The periodic table evolved over time as scientists discovered more useful ways to compare and organize the elements.
Section 6-1

4. Section 6.1 Development of the Modern Periodic Table (cont.)
periodic law
group
period
representative elements
transition elements
metal
alkali metals
alkaline earth metals
transition metal
inner transition metal
lanthanide series
actinide series
nonmetals
halogen
noble gas
metalloid
Section 6-1

5. CHEMICAL SYMBOLS must have one, and only one, capital letter
represent elements (there are 118 currently)
some come from Latin or Greek [which is why the symbol is not from 1st letter(s)]
Au (gold) “aurum” (means “shining dawn”)
Ag (silver) “argentum”
Fe (iron) “ferrum”
K (potassium) “kalium”
Hg (mercury) “hydrargyrum” (means liquid silver)
Pb (lead) “plumbum”
Sn (tin) “stannum”
Na (sodium) “natrium”
Cu (copper) “cuprum”
W (tungsten) “wolfram”
Sb (antimony) “stibium”

6. The temporary naming system Any element symbol made of three letters is temporary and comes from the atomic number of the element.
nil = 0
un = 1
bi = 2
tri = 3
quad = 4
pent =5
hex = 6
sept = 7
oct = 8
enn = 9

7. Development of the Periodic Table
In the 1700s, Lavoisier compiled a list of all the known elements of the time. It contained 33 elements organized in four categories
Section 6-1

8. Development of the Periodic Table (cont.)
The 1800s brought large amounts of information and scientists needed a way to organize knowledge about elements.
Electricity was used to break down compounds
Spectrometers were developed to identify new elements. Helium was 1st discovered this way on the sun and is named for the Greek God of the sun Helios.
The industrial revolution led to many chemical industries
Section 6-1

9. John Newlands (1864) proposed an arrangement where elements were ordered by increasing atomic mass.

10. Development of the Periodic Table (cont.)
Newlands noticed when the elements were arranged by increasing atomic mass, their properties repeated every eighth element. He called this arrangement the “Law of Octaves” because properties repeated every 8th element
Section 6-1

11. Development of the Periodic Table (cont.)
Meyer and Mendeleev both demonstrated a connection between atomic mass and elemental properties. (Mendeleev given more credit and is called the “father” of the periodic table. Several places of his table were flawed, though. [see table today..}
Mendeleev was also able to predict the discovery of new elements. He left blank spaces after noting trends of known elements. He predicted Ga, Sc, and Ge.
Section 6-1

12. Moseley rearranged the table by increasing atomic number, and resulted in a clear periodic pattern with no exceptions.
Periodic repetition of chemical and physical properties of the elements when they are arranged by increasing atomic number is called periodic law.
“periodic” means repeating.

13. Development of the Periodic Table (cont.)
Section 6-1

14. The Modern Periodic Table
The modern periodic table contains boxes which contain the element's name, symbol, atomic number, and atomic mass.
Section 6-1

16. The Modern Periodic Table (cont.)
Columns of elements are called groups or families. These are labeled 1 to 18 or with Roman numerals with A/B.
Rows of elements are called periods or series.
Elements in groups 1,2, and possess a wide variety of chemical and physical properties and are called the representative elements.[the A groups]
Elements in groups 3-12 are known as the transition metals. [the B groups]
Section 6-1

17. The Modern Periodic Table (cont.)
Elements are classified as metals (left of stair-like line),non-metals (right of stair-like line), and metalloids (on stair-like line)
Metals are elements that are generally shiny when smooth and clean, solid at room temperature, and good conductors of heat and electricity.
Alkali metals are all the elements in group 1 except hydrogen, and are very reactive. Francium would be the most reactive metal on the chart.
Alkaline earth metals are in group 2, & are also highly reactive (but not as reactive as group 1)
Section 6-1

18. The Modern Periodic Table (cont.)
The transition elements are divided into transition metals and inner transition metals (or rare earth elements).
The rare earth elements contain the two sets of inner transition metals; these are called the lanthanide series and actinide series and are located at the bottom of the periodic table.
Section 6-1

19. The Modern Periodic Table (cont.)
Non-metals are elements that are generally gases or brittle, dull-looking solids, and poor conductors of heat and electricity.
Group 16 is called the chalcogens. [chalcogen means “acid forming”]
Group 17 is composed of highly reactive elements called halogens [halogen means “salt forming”] Fluorine is the most reactive.
Group 18 gases are extremely unreactive and commonly called noble gases.[used to be called “inert” gases. Inert means unreactive]
Section 6-1

20. The Modern Periodic Table (cont.)
Metalloids have physical and chemical properties of both metals and non-metals, such as silicon and germanium. Found on the stair-like line.
Section 6-1

21. The Modern Periodic Table (cont.)
Section 6-1

22. A B C D Section 6.1 Assessment
What is a row of elements on the periodic table called?
A. octave
B. period
C. group
D. transition A B C D
Section 6-1

23. A B C D Section 6.1 Assessment What is silicon an example of? A. metal
B. non-metal
C. inner transition metal
D. metalloid A B C D
Section 6-1

24. Section 6.2 Classification of the Elements
Explain why elements in the same group have similar properties.
valence electron: electron in an atom's outermost orbitals; determines the chemical properties of an atom
Identify the four blocks of the periodic table based on their electron configuration.
Elements are organized into different blocks in the periodic table according to their electron configurations.
Section 6-2

25. Getting outer shell (valence) e- configurations and dot diagrams from the periodic table.

26. Organizing the Elements by Electron Configuration
Recall electrons in the highest principal energy level are called valence electrons.
All group 1 elements have one valence electron. [end with s1]
Section 6-2

27. Organizing the Elements by Electron Configuration (cont.)
The energy level of an element’s valence electrons indicates the period (row) on the periodic table in which it is found.
The number of valence electrons for elements in groups is ten less than their group number.
Section 6-2

28. Organizing the Elements by Electron Configuration (cont.)
Section 6-2

29. The s-, p-, d-, and f-Block Elements
The shape of the periodic table becomes clear if it is divided into blocks representing the atom’s energy sublevel being filled with valence electrons.
Section 6-2

30. The s-, p-, d-, and f-Block Elements (cont.)
s-block elements consist of group 1 and 2, and the element helium.
Group 1 elements have a partially filled s orbital with one electron.
Group 2 elements have a completely filled s orbital with two electrons.
Section 6-2

31. The s-, p-, d-, and f-Block Elements (cont.)
After the s-orbital is filled, valence electrons occupy the p-orbital.
Groups contain elements with completely or partially filled p orbitals.
Section 6-2

32. The s-, p-, d-, and f-Block Elements (cont.)
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.
Section 6-2

33. The s-, p-, d-, and f-Block Elements (cont.)
The f-block (rare earth elements) 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.
Section 6-2

34. STABILITY RULES
A full outer shell is most stable
atoms with a full “s” and full “p” [2+6=8]
atoms with 8 e- in outer shell have an octet
octet rule says 8 e- in outer shell make the atom very stable and unreactive [noble gases]
any full subshell is pretty stable
p6, d10, f14 atoms tend to be stable (note: s2 atoms are not stable
half full subshells are kind of stable
p3, d5, and f7 atoms are a little stable
some atoms deviate from predicted configurations because they rearrange their electrons to be more stable Ex: Copper predicted to be: [Ar]4s23d9
but really is: [Ar]4s13d10

35. A B C D Section 6.2 Assessment
Which of the following is NOT one of the elemental blocks of the periodic table?
A. s-block
B. d-block
C. g-block
D. f-block A B C D
Section 6-2

36. A B C D Section 6.2 Assessment Which block spans 14 elemental groups?
A. s-block
B. p-block
C. f-block
D. g-block A B C D
Section 6-2

37. Section 6.3 Periodic Trends
Compare period and group trends of several properties.
principal energy level: the major energy level of an atom
Relate period and group trends in atomic radii to electron configuration.
ion
ionization energy
octet rule
electronegativity
Trends among elements in the periodic table include their size and their ability to lose or attract electrons
Section 6-3

38. Atomic Radius
Atomic radius (size) is a periodic trend influenced by electron configuration.
For metals, atomic radius is half the distance between adjacent nuclei in a crystal of the element.
Section 6-3

39. Atomic Radius (cont.)
For elements that occur as molecules (nonmetals), the atomic radius is half the distance between nuclei of identical atoms.
Section 6-3

40. Atomic Radius (cont.)
There is a general decrease in atomic radius from left to right, caused by increasing positive charge in the nucleus.
Valence electrons are not shielded from the increasing nuclear charge because no additional electrons come between the nucleus and the valence electrons.
Section 6-3

41. Atomic Radius (cont.)
Section 6-3

42. Atomic radius generally increases as you move down a group.
Atomic Radius (cont.)
Atomic radius generally increases as you move down a group.
The outermost orbital size increases down a group, making the atom larger.
Section 6-3

43. Ionic Radius
An ion is an atom or bonded group of atoms with a positive or negative charge.
When atoms [metals] lose electrons and form positively charged ions, they always become smaller for two reasons:
The loss of a valence electron can leave an empty outer orbital resulting in a small radius.
Electrostatic repulsion decreases allowing the electrons to be pulled closer to the radius.
Section 6-3

44. Ionic Radius (cont.)
When atoms [nonmetals] gain electrons, they can become larger, because the addition of an electron increases electrostatic repulsion.
Section 6-3

45. Ionic Radius (cont.)
The ionic radii of positive ions generally decrease from left to right.
The ionic radii of negative ions generally decrease from left to right, beginning with group 15 or 16.
Section 6-3

46. Both positive and negative ions increase in size moving down a group.
Ionic Radius (cont.)
Both positive and negative ions increase in size moving down a group.
Section 6-3

47. Ionization Energy
Ionization energy is defined as the energy required to remove an electron from a gaseous atom.
The energy required to remove the first electron is called the first ionization energy.
Section 6-3

48. Ionization Energy (cont.)
Section 6-3

49. Ionization Energy (cont.)
Removing the second electron requires more energy, and is called the second ionization energy.
Each successive ionization requires more energy, but it is not a steady increase.
Section 6-3

50. Ionization Energy (cont.)
Section 6-3

51. Ionization Energy (cont.)
The ionization at which the large increase in energy occurs is related to the number of valence electrons.
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.
Section 6-3

52. Ionization Energy (cont.)
Section 6-3

53. Ionization Energy (cont.)
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.
Section 6-3

54. Ionization Energy (cont.)
The electronegativity of an element indicates its relative ability to attract electrons in a chemical bond.
Electronegativity decreases down a group and increases left to right across a period.
Section 6-3

55. Ionization Energy (cont.)
Section 6-3

56. A B C D Section 6.3 Assessment
The lowest ionization energy is the ____.
A. first
B. second
C. third
D. fourth A B C D
Section 6-3

57. A B C D Section 6.3 Assessment
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 A B C D
Section 6-3