1. Do Now Define an element.
What relationship exists between atomic number, protons and electrons?
Do you notice any pattern or relationships between elements of the same row or column?

2. The Periodic Table

3. Mendeleev’s Periodic Table
Dmitri Mendeleev

4. History of the Periodic Table
By 1860 more than 60 elements had been discovered.
Chemists learned about the new elements by reacting them with other elements to form new compounds.
In 1865, John Newlands, an English chemist, arranged the known elements according to their properties in order of increasing atomic mass.
The properties of elements seemed to repeat every eight elements. This pattern is known as the law of octaves.

5. History of the Periodic Table
In 1869, Dmitri Mendeleev, a Russian chemist, used Newland’s observations and other observations to produce the first PERIODIC Table.
He placed each element on a card with the element’s atomic mass, chemical and physical properties. He arranged the elements in many different ways according to various properties and looked for trends.
When he arranged the elements according to increasing atomic mass, he noticed that certain similarities occurred at regular intervals. This was considered a “periodic pattern”.
He arranged all the elements in a table according to increasing atomic mass, starting a new row every time the pattern repeated, and the elements with similar properties fell in the same vertical columns (with a few exceptions).
He had some blank spaces in his table where he thought new elements would be discovered. These blank spaces were used to predict the existence and discovery of new elements.

6. History of the Periodic Table
By 1886 three new elements were discovered that followed Mendeleev’s predication.
Most chemists were persuaded to accept his table, labeling Mendeleev the founder of the periodic law.

7. History of the Periodic Table
In order to keep the pattern of properties, Mendeleev had to switch some elements out of order based on atomic mass.
In 1911, Henry Moseley, an English chemist, was examining the spectra of 38 different metals. He noticed that the wavelengths of spectra lines correlated to atomic numbers, not atomic mass.
Moseley discovered a new pattern and organized the elements according to their increasing nuclear charge, or atomic number, i.e. increasing number of protons.
When the elements were arranged by increasing atomic number, the discrepancies in Mendeleev’s table disappeared.

8. Modern Russian Table

9. Chinese Periodic Table

10. Stowe Periodic Table

11. A Spiral Periodic Table

12. Triangular Periodic Table

13. “Mayan” Periodic Table

14. Meet the Elements Music Video

15. Do Now
Take out your Reference Tables and open to the Periodic Table of Elements
What do you think of when you hear the word metal?
What is a nonmetal?
Do you know of any examples of each?

16. The Periodic Table consists of:
METALS
NONMETALS
METALLOIDS

17. Properties of Metals
Metals are good conductors of heat and electricity
Metals are malleable (can be hammered or rolled into sheets)
Metals are ductile (can be made into wire)
Metals have high tensile strength
Metals have luster

18. Examples of Metals
Potassium, K reacts with water and must be stored in mineral oil
Copper, Cu, is a relatively soft metal, and a very good electrical conductor.

19. Examples of Metals Zinc, Zn, is more stable than potassium
Mercury, Hg, is the only metal that exists as a liquid at room temperature

20. Mixing metals
Alloy = mixture of a metal with another element, usually another metal
Alloys have properties different from the individual elements, usually eliminating some disadvantages.
Alloys are usually harder and more resistant to corrosion
Ex: brass = copper + zinc
Ex: sterling sliver = copper + silver
Ex: steel = iron + carbon + manganese + nickel
Ex: Stainless steel = iron + chromium

21. Properties of Nonmetals
Carbon, the graphite in “pencil lead” is a great example of a nonmetallic element.
Nonmetals are poor conductors of heat and electricity
Nonmetals tend to be brittle
Many nonmetals are gases at room temperature

22. Examples of Nonmetals Sulfur, S, was once known as “brimstone”
Microspheres of phosphorus, P, a reactive nonmetal

23. Examples of Nonmetals
Graphite is not the only pure form of carbon, C. Diamond is also carbon; the color comes from impurities caught within the crystal structure
Bromine is a nonmetal that exists as a liquid at room temperature.

24. Properties of Metalloids
Metalloids straddle the border between metals and nonmetals on the periodic table.
They have properties of both metals and nonmetals.
Metalloids are more brittle than metals, less brittle than most nonmetallic solids
Metalloids are semiconductors of electricity
Some metalloids possess metallic luster
*note them on your Periodic Table in your Reference Table!

25. Silicon, Si – A Metalloid
Silicon has metallic luster
Silicon is brittle like a nonmetal
Silicon is a semiconductor of electricity
Other metalloids include:
Boron, B
Germanium, Ge
Arsenic, As
Antimony, Sb
Tellurium, Te

26. Elements classified as metals and nonmetals

27. Lab
Alien Nation

28. Do Now
Take out your Reference Table and open to the Periodic Table of Elements
Do you notice any patterns as you look at the properties of elements?
Hints: scan left to right, and scan from the top down

29. The Periodic Table PERIODIC LAW THE MODERN PERIODIC TABLE
The physical and chemical properties of the elements are periodic functions of their atomic numbers (elements with similar properties appear at regular intervals).
THE MODERN PERIODIC TABLE
An arrangement of the elements in order of increasing atomic number so that the elements with similar properties fall in the same column or group.

30. Periodic Table with Group Names
PERIODS= horizontal rows
GROUPS= vertical columns

31. VALENCE ELECTRONS
Elements within the same group have the same number of valence electrons.

32. VALENCE ELECTRONS
Most loosely bound electron is called the VALENCE ELECTRON!
Valence electrons participate in chemical reactions, so elements with similar valence electrons react in similar ways.
Noble gases are stable with _8_ valence electrons.
All elements with the exception of Hydrogen and Helium want _8_ valence electrons to be stable.
Hydrogen and Helium want _2_ valence electrons to be stable.

33. Orbital filling table

34. Orbital filling table

35. s-BLOCK and p-BLOCK are considered “MAIN GROUP ELEMENTS”
s-Block: Includes GROUP 1 and GROUP 2
GROUP 1 - Alkali Metals
GROUP 2 – Alkaline Earth Metals
p-Block: Elements vary greatly in properties, includes GROUP 13 through GROUP 18
GROUP 17 – Halogens
GROUP 18 – Noble Gases

36. GROUP 1 Alkali Metals Properties: Metal Easily lose valence electron
Chemically reactive – do not occur as free elements in nature
Soft, silvery
Good conductor of electricity
React violently with water
React with halogens to form salts

37. EXCEPTION: Hydrogen is in group 1, but has different properties than the alkali metals
*note that Hydrogen is a nonmetal on your Periodic Table in your Reference Table!

38. Hydrogen Most common element in the universe.
It consists of one proton and one electron.
It can react with many other elements:
With oxygen to make water
With carbon to make organic compounds
With nitrogen to make ammonia

39. GROUP 2 Alkaline Earth Metals
Properties:
They are metals, but have less metallic characteristics than Alkali Metals
2 valence electrons
Chemically reactive – do not occur as free elements in nature
Harder, denser, stronger, higher boiling points, and slightly less reactive than alkali metals

40. Group 13 Properties: 3 valence electrons
Lose 3 electrons to form 3+ ions
Boron is a metalloid and does not form a 3+ ion readily

41. Group 14 Properties: 4 valence electrons
Carbon is a nonmetal with two distinct crystalline forms known as allotropes (graphite and diamond)
Silicon and Germanium are metalloids used in the computer industry
Tin and Lead are metals

42. Group 15 Properties: 5 valence electrons
Nitrogen is a nonmetal that exists naturally as N2, a diatomic molecule with a triple bond.
Phosphorus is a nonmetal that exists in two allotropic forms (red and white)

43. Group 16 Properties: 6 valence electrons
Oxygen is a nonmetal that exists naturally in two allotropic forms (O2 and O3)
Sulfur is a nonmetal that exists naturally in several allotropic forms
Selenium is a nonmetal
Tellurium is a metalloid
Polonium is a metal

44. GROUP 17 Halogens (“Salt makers”)
Properties:
Most reactive nonmetals.
7 valence electrons – need 1 electron to become stable
They vigorously react with metals to form salts.
F2 - most reactive element
Cl2
Br2 I2
Gases at STP
Liquid at STP
Solid at STP

45. GROUP 18 NOBLE GASES (or INERT GASES)
Properties:
Full set of valence electrons: most elements have 8 valence electrons, except Helium with 2 valence electrons, but it is still associated with this group because its properties match these elements.
Extremely stable and occur as monoatomic gases in nature
Although they do not readily combine with other elements, but compounds of Krypton and Xenon have been prepared.

46. Monoatomic Atoms All of the noble/inert gases are monatomic.
Monatomic means one atom.

47. Magic seven
*note them on your Periodic Table in your Reference Table!
Diatomic molecules= always exist as 2 of the same atoms bonded to each other H2 N2 O2 F2
Cl2
Br2 I2

48. Transition Metals (d-Block)
They have “typical metallic properties”
Luster, ductile, malleable, good conductors of heat and electricity
Less reactive than Group 1 and 2
Many are unreactive (for example, palladium, platinum and gold are found as pure elements in nature)
These elements begin in Period 4 and include Groups 3-12.

49. Transition Metals (d-Block)
As ions, the transition elements form colorful solutions.

50. Lanthanides and Actinides (f-Block)
Lanthanide Series- shiny metals similar in reactivity to Alkaline Earth Metals
Actinides Series- all radioactive (mostly synthetic)

51. Synthetic Elements Synthetic elements do not occur in nature
Ex: Technetium (Tc): atomic number 43
Ex: All elements with atomic numbers > 92 (transuranium elements), which are all radioactive
These elements are produced artificially by bombardment or as by-products of fission reactors.

52. What did you learn today?

53. What did you learn today?
The placement or location of elements on the Periodic Table gives an indication of physical and chemical properties of that element. The elements on the Periodic Table are arranged in order of increasing atomic number.
The number of protons in an atom (atomic number) identifies the element. The sum of the protons and neutrons in an atom (mass number) identifies an isotope.
Elements can be classified by their properties and located on the Periodic Table as metals, nonmetals, metalloids (B, Si, Ge, As, Sb, Te), and noble gases.
Elements can be differentiated by their physical properties. Physical properties of substances, such as density, conductivity, malleability, solubility, and hardness, differ among elements.
Elements can be differentiated by chemical properties. Chemical properties describe how an element behaves during a chemical reaction.
For Groups 1, 2, and on the Periodic Table, elements within the same group have the same number of valence electrons (helium is an exception) and therefore similar chemical properties.

54. Quiz
Quiz on Intro to the Periodic Table

55. Lab
Periodic Trends lab

56. Do Now
What is a trend?
Have you noticed any trends in history? Have some trends become more abundant, while others have become more scarce?
List some examples.

57. Periodic Trends ELECTRONEGATIVITY ATOMIC RADIUS IONIC RADIUS
IONIZATION ENERGY
ELECTRON AFFINITY
METALLIC CHARACTER

58. The trend depends on 3 factors:
NUCLEAR CHARGE= atomic number= # of protons- the higher the nuclear charge the more the electrons are pulled toward the nucleus.
PRINCIPAL ENERGY LEVEL- Principal energy level is determined by the period an atom is located; the higher the principal energy level, the higher the potential energy, the larger the atom.
ELECTRON CLOUD SHIELDING EFFECT- inner electrons shield the outer electrons so they are less attracted to the nucleus so the atom is larger

59. Electronegativity
A measure of the ability of an atom in a chemical compound to attract electrons.
FLUORINE is the most electronegative atom on the Periodic Table
PERIOD TREND: Electronegativity increases across a period
(the closer to 8 valence electrons the more they will pull on the electrons)
GROUP TREND: Electronegativity decreases down a group
(the larger the atomic radius the less ability of that atom to attract electrons due to electron shielding)

60. Periodic Table of Electronegativities

62. Electronegativity Review
Element with most Electronegativity:
Element with least Electronegativity:

63. Atomic Radius PERIOD TREND: Radius decreases across a period
Atomic radius = half of the distance between nuclei of identical atoms bonded together (bond radius).
This is determined by how strongly an atom’s nucleus (positive charge) is attracted to its outermost electrons.
PERIOD TREND: Radius decreases across a period
Nuclear charge increases across a period so electrons are more closely attracted to the nucleus
GROUP TREND: Radius increases down a group
Addition of principal energy levels, so inner electrons shield outer electrons and the electrons are farther away from nucleus, so atom is larger.

64. Table of Atomic Radii

67. Review Ions Ions are formed when atoms gain or lose electrons
Ions have a positive or negative charge
Ionic Radii: The loss or gain of electrons by an atom causes a corresponding change in size.

68. Ionic Radii Cations Anions Positively charged ions
Smaller than the corresponding atom
Metal tend to form CATIONS by LOSING an electron.
Metal cations are smaller than their corresponding neutral atoms (they lost electrons).
Negatively charged ions
Larger than the corresponding atom
Nonmetals tend to for ANIONS by GAINING an electron.
Nonmetal anions are larger than neutral atoms because they gained electrons.

69. Here is a simple way to remember a cation and an anion:
This is a cat-ion.
He’s a “plussy” cat!
This is Ann Ion.
She’s unhappy and negative.

71. Ionic Radius Review
The loss or gain of electrons by an atom causing a corresponding change in size.
Metals
Tend to ____ their valence electrons.
What happens to their valence electrons?__________
Before: After:
Therefore the atomic radius of metals is ___________ than their ionic radius.

74. Ionic Radius Review
The loss or gain of electrons by an atom causing a corresponding change in size.
Nonmetals
Tend to _________ electrons.
What happens to their radius?__________
Before: After:
Therefore the atomic radius of nonmetals is ________________ than their ionic radius.

76. Table of Atomic Radii
and
Ionic
Radii

77. Ionic Radius Metals form CATIONS Loss of electron
Get a Positive Charge
Nonmetals
Form
ANIONS
Gain of electron
Get a Negative charge
Radius Decreases when it becomes an ion
Radius Increases when it becomes an ion

78. Do Now Period 2: Li Be C N O F Ne
Which elements do not want to lose their electrons? ____________

79. A + ionization energy  A+ + e-
Ionization energy = the energy required to remove an electron from an atom
A + ionization energy  A+ + e-
PERIOD TREND: Tends to increase across a period
Within same principle energy level: due to increased nuclear charge the electrons are more closely attracted to nucleus
GROUP TREND: Tends to decrease down a group
Addition of principle energy levels, so outer electrons are farther from the nucleus due to electron shielding, so less attracted to the nucleus

82. Another Way to Look at Ionization Energy

83. Ionization of Magnesium
Ionization energy increases for successive electrons taken from the same atom
Mg kJ  Mg+ + e-
Mg kJ  Mg e-
Mg kJ  Mg e-

84. Ionization Energy Review
Element with the most Ionization Energy:
Element with the least Ionization Energy:
Why do noble gases have such high ionization energy but no electronegativity?

85. A + e- + energy  A- (unstable)
Electron Affinity
The energy change that occurs when an electron is acquired by a neutral atom is called the atom’s electron affinity.
A + e-  A- + energy
A + e- + energy  A- (unstable)
PERIOD TREND: Tends to increases across a period
due to increasing nuclear charge (remember: the Halogens gain electrons most readily)
GROUP TREND: Tends to decrease down a group
due to increasing electron shielding from addition of principle energy levels

86. REACTIVITY METALS- most reactive in lower left corner
NONMETALS- most reactive in top right corner, but it does not include noble gases

93. Reactivity Review:
Place the following elements in increasing order of reactivity: (least reactive first)
Sodium, potassium, calcium
Group 1, Group 1, Group 2, so…
Calcium, sodium, potassium

94. Metallic Character
As you go down a group, the metallic character increases.
As you go across a period, the metallic character decreases.

95. Melting Point
Melting points are the amount of energy required to break a bond(s) to change the solid phase of a substance to a liquid. 
The stronger the bond between the atoms of an element, the higher the energy requirement in breaking that bond.
Metals generally possess a high melting point.
Most nonmetals possess low melting points.

97. Boiling Point
Boiling points are the amount of energy required to break a bond(s) to change the liquid phase of a substance to a gas.
The stronger the bond between the atoms of an element, the higher the energy requirement in breaking that bond.
Metals generally possess a high boiling point.
Most nonmetals possess low boiling points.

99. Summation of Periodic Trends
Electronegativity
Electronegativity

100. Periodic Trends Definition Trend in Period Trend in Group
Electronegativity
The measure of the ability of an atom in a chemical compound to attract electrons
Increases across a period
Decrease down a group
Atomic Radius
Half the distance between nuclei of identical atoms bonded together
Decreases across a period
Increases down a group
Ionization Energy
The energy required to remove an electron from an atom
Decreases down a group
*Use Table S in your Reference Tables!

101. Lab Reactivity lab Videos: http://www.youtube.com/watch?v=m55kgyApYrY

102. Quiz
Quiz on Periodic Table Trends

103. Test
Test on the Periodic Table