Mendeleev's Periodic Table and the History of the PT

1. Unit 6 Periodicity

2. History of the PT • Dimitri Mendeleev (1836-1907) • Russian chemistry professor • In writing a textbook of general chemistry, he compiled vast quantities of data on the elements. He realized as he examined the data that the chemical properties of the elements repeated themselves when the elements were placed in order of atomic mass. Based on his realization, Mendeleev arranged the elements in order of increasing atomic mass, and in such a way that elements with similar chemical properties fell in the same column. He published a primitive version of today’s periodic table in 1869. It contained the 62 elements that had been discovered at that time.

3. Mendeleev’s Periodic Table Dmitri Mendeleev

4. Genius of Mendeleev’s Work • Left spaces for elements not yet discovered. • He predicted that some still-unknown elements must exist to fit in the holes.

5. Inductive vs. Deductive Reasoning • Inductive reasoning is the use of detailed facts to form a general principle or model (going from specific to general). How did Mendeleev use inductive reasoning? • Deductive reasoning is the use of a general principle or model to draw specific inferences (going from general to specific). How did Mendeleev use deductive reasoning?

6. Mendeleev’s Periodic Table • Contained 1 inconsistency. • He placed the elements in order of atomic mass • Forced to break pattern a couple times to preserve the patterns he had discovered.

7. Henry Moseley • Shortly after Rutherford’s discovery of the proton in 1911, Henry Moseley (1887-1915) did experiments with x-rays to determine the number of protons in various elements. When Moseley arranged the elements of Mendeleev’s periodic table according to increasing atomic number and not atomic mass, the inconsistencies associated with Mendeleev's table were eliminated. The modern periodic table is based on Moseley's arrangement by atomic number. At age 28, Moseley was killed in action during World War I. As a direct result, Britain adopted the policy of exempting scientists from fighting in wars.

8. Periodic Law • From Mendeleev’s and Moseley’s work comes the Periodic Law: The properties of the elements are periodic functions of their atomic numbers. • What this means is that if we arrange the elements in order of increasing atomic number, we will periodically encounter elements that have similar chemical and physical properties. These elements appear in the same vertical column (group).

9. Modern Russian Table

10. Stowe Periodic Table

11. A Spiral Periodic Table

12. “Mayan” Periodic Table

13. The Periodic Table Period Group or Family Group or family Period

15. Irregular conformations of Cr and Cu Chromium steals a 4s electron to half fill its 3d sublevel Copper steals a 4s electron to FILL its 3d sublevel

16. Nonmetals Metalloids Metals

17. Hydrogen • The hydrogen square sits atop Family 1, but it is not a member of that family. Hydrogen is in a class of its own. • It’s a gas at room temperature. • It has one proton and one electron in its one and only energy level. • Hydrogen only needs 2 electrons to fill up its valence shell.

18. The Properties of a Group: the Alkali Metals • Group 1 (s-block) • Easily lose valence electron • (Reducing agents) • React violently with water • ns1 configuration • silvery apperance • soft (cut w/ a knife) • Most Reactive! • Not found as free elements. • React with halogens to form • salts

19. Alkali Metals • The alkali family is found in the first column of the periodic table. • Atoms of the alkali metals have a single electron in their outermost level, in other words, 1 valence electron. • They are shiny, have the consistency of clay, and are easily cut with a knife.

20. Alkali Metals • They are the most reactive metals. • They react violently with water. • Alkali metals are never found as free elements in nature. They are always bonded with another element.

21. Alkaline Earth Metals • They are never found uncombined in nature. • They have two valence electrons. • Alkaline earth metals include magnesium and calcium, among others.

22. Transition Metals • Transition Elements include those elements in the B families. • These are the metals you are probably most familiar: copper, tin, zinc, iron, nickel, gold, and silver. • They are good conductors of heat and electricity.

23. Transition Metals • The compounds of transition metals are usually brightly colored and are often used to color paints. • Transition elements have 1 or 2 valence electrons, which they lose when they form bonds with other atoms. Some transition elements can lose electrons in their next-to-outermost level.

24. Boron Family • The Boron Family is named after the first element in the family. • Atoms in this family have 3 valence electrons. • This family includes a metalloid (boron), and the rest are metals. • This family includes the most abundant metal in the earth’s crust (aluminum).

25. Carbon Family • Atoms of this family have 4 valence electrons. • This family includes a non-metal (carbon), metalloids, and metals. • The element carbon is called the “basis of life.” There is an entire branch of chemistry devoted to carbon compounds called organic chemistry.

26. Nitrogen Family • The nitrogen family is named after the element that makes up 78% of our atmosphere. • This family includes non-metals, metalloids, and metals. • Atoms in the nitrogen family have 5 valence electrons. They tend to share electrons when they bond. • Other elements in this family are phosphorus, arsenic, antimony, and bismuth.

27. Oxygen Family • Atoms of this family have 6 valence electrons. • Most elements in this family share electrons when forming compounds. • Oxygen is the most abundant element in the earth’s crust. It is extremely active and combines with almost all elements.

28. Halogen Family • The elements in this family are fluorine, chlorine, bromine, iodine, and astatine. • Halogens have 7 valence electrons, which explains why they are the most active non-metals. They are never found free in nature. • Halogen atoms only need to gain 1 electron to fill their outermost energy level. • They react with alkali metals to form salts.

29. Noble Gases • Noble Gases are colorless gases that are extremely un-reactive. • One important property of the noble gases is their inactivity. They are inactive because their outermost energy level is full. • Because they do not readily combine with other elements to form compounds, the noble gases are called inert. • The family of noble gases includes helium, neon, argon, krypton, xenon, and radon. • All the noble gases are found in small amounts in the earth's atmosphere.

30. Rare Earth Elements • The thirty rare earth elements are composed of the lanthanide and actinide series. • One element of the lanthanide series and most of the elements in the actinide series are called trans-uranium, which means synthetic or man-made.

32. Properties of Metals • Metals are good conductors of heat and electricity • Metals are malleable • Metals are ductile • Metals have high tensile strength • Metals have luster

33. Examples of Metals Potassium, K reacts with water and must be stored in kerosene Copper, Cu, is a relatively soft metal, and a very good electrical conductor. Zinc, Zn, is more stable than potassium Mercury, Hg, is the only metal that exists as a liquid at room temperature

34. Propertiesof 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

35. Examples of Nonmetals Microspheres of phosphorus, P, a reactive nonmetal Sulfur, S, was once known as “brimstone” Graphite is not the only pure form of carbon, C. Diamond is also carbon; the color comes from impurities caught within the crystal structure

36. 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

37. 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

38. Electron Shielding Shielding electrons: electrons in the energy levels between the nucleus and the valence electrons. They are called "shielding" electrons because they "shield" the valence electrons from the force of attraction exerted by the positive charge in the nucleus.

39. Atomic Radius Definition: Half of the distance between nuclei in covalently bonded diatomic molecule • Radius decreases across a period • Increased effective nuclear charge due to decreased shielding • Radius increases down a group • Each row on the periodic table adds a “shell” or energy level to the atom

40. Table of Atomic Radii

41. Period Trend:Atomic Radius

42. Ionic Radii • Positively charged ions formed when • an atom of a metal loses one or • more electrons Cations • Smaller than the corresponding • atom • Negatively charged ions formed • when nonmetallic atoms gain one • or more electrons Anions • Larger than the corresponding • atom

43. Ionic Radii • Positively charged ions formed when • an atom of a metal loses one or • more electrons Cations • Smaller than the corresponding • atom • Negatively charged ions formed • when nonmetallic atoms gain one • or more electrons Anions • Larger than the corresponding • atom

44. Graphic courtesy Wikimedia Commons user Popnose

45. Table of Ion Sizes

46. Ionization Energy Definition: the energy required to remove an electron from an atom • Tends to increase across a period • As radius decreases across a period, the electron you are removing is closer to the nucleus and harder to remove • Tends to decrease down a group • Outer electrons are farther from the nucleus and easier to remove

47. Ionization of Magnesium Mg + 738 kJ  Mg+ + e- Mg+ + 1451 kJ  Mg2+ + e- Mg2+ + 7733 kJ  Mg3+ + e-

48. Periodic Trend:Ionization Energy

49. Trends in Electron Affinity • electron affinity (electron-liking): the energy change that accompanies the addition of an electron to an atom. • *If energy is released in the process of adding an electron, EA is negative. A negative EA means that the atom wants to gain the electron. The larger the negative number, the more it wants to gain the electron. It is a favorable process. Nonmetal atoms have large negative EAs. • *If energy is absorbed in the process of adding an electron, EA is positive. A positive EA means that the atom does not want to gain the electron. The larger the positive number, the more it does not want to gain the electron. It is an unfavorable process. Metal atoms have small negative or positive EAs. • group trend: EA decreases(less favorable/value is more +) going down a group. • It is harder to add the electron when it is farther from the nucleus. The nucleus can’t “grab onto it” as well • periodic trend: EA increases (more favorable/value is more - ) going across a period. • This is due to nuclear charge - across a period, nuclear charge increases, so it becomes easier to add an electron. • Note that this periodic trend supports the idea that nonmetals have a much greater tendency to gain electrons than metals do.

50. Electron Affinity • Electron Affinity • Increases UP and to the RIGHT