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The Periodic Table

The Periodic Table

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The Periodic Table

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  1. The Periodic Table

  2. Periodic Table of Elements • There are 117 elements (January, 2007) • Your table contains 113 • 94 of the elements are naturally occurring, the rest are man-made • Most of the elements were discovered between 1735-1843

  3. History/Development • Development of the table has occurred over 300 years and continues today • Dmitri Mendeleev is commonly credited with creating the periodic table in 1869

  4. Original Table (Mendeleev) • Classified elements in horizontal rows based on their atomic mass • When there was a repeat of properties the elements were placed in the next row • Concluded that similar properties appear at regular intervals when the elements are listed in order of increasing atomic mass

  5. Original Table (Mendeleev) • Elements with similar properties were located in the same vertical columns, if no known element had the expected properties to fit the particular space he left that space empty • Assumed that elements not yet discovered would fit into the empty spaces • Predicted some properties of these unknown elements (gallium, scandium, and germanium)

  6. Modern Periodic Law - 1914 • Henry Moselely bombarded elements with high speed electrons and they emitted X-rays with a certain wavelength • He found that each element differed by one proton • Concluded that the regularity, or periodicity, of the properties is a function of the atomic number (Modern Periodic Law)

  7. Questions • The modern periodic table is arranged by increasing _______________________. • The original periodic table was arranged by increasing _____________________.

  8. Organization of the Modern Periodic Table • The periodic table is arranged in order of increasing ATOMIC NUMBER • Horizontal rows are called PERIODS • There are 7 periods • Vertical columns are called GROUPS • There are 18 groups

  9. Periods • Horizontal Rows • The number of each period indicates the principle energy level in which the valence electrons are located • The number of valence electrons increases as you go from left to right • The properties of the elements change systematically through a period

  10. Groups (Families) • Vertical Columns • The outermost shell of an atom contains the same number and arrangement of valence electrons • Elements that have similar chemical properties are located in the same group

  11. Group/Period Examples • Are elements in the same period or group more similar? Explain why. • Which elements have the most similar chemical properties? • K and Na • K and Cl • K and Ca • K and S 3. Why do elements of a given group on the periodic table show similar chemical properties?

  12. States of Matter • Solid • The majority of the elements are solids • Liquid • The only liquids are Hg and Br, which are found on the right • Gas • H, O, N, F, Cl and the Noble Gases (Group 18), located on the right

  13. Nonmetals Metalloids Metals

  14. Metals/Nonmetals/Metalloids • Where are the metals located? • Where are the nonmetals located? • Where are the metalloids located? • List all of the Metalloids

  15. Lose electrons to form positive ions Solids at STP Except Hg High melting and boiling points Good thermal (heat) and electrical conductors Luster (shine) Malleable (bendable) Ductile – can be made into wires Gain electrons to form negative ions Solids, 1 liquid (Br), gases at STP Low melting and boiling points Poor conductors of heat and electricity Dull Brittle Not ductile Properties of Metals and Nonmetals

  16. Metalloids • Properties of both metals and nonmetals • B, Si, As, Te, Ge, and Sb (front + middle back of staircase)

  17. Metal/Nonmetal Questions 1. Atoms of metals tend to a. lose electrons to form positive ions b. lose electrons to form negative ions c. gain electrons to form positive ions d. gain electrons to form negative ions 2. The majority of elements on the table are classified as a. metals b. nonmetals c. metalloids 3. Which property is generally characteristic of metallic elements? a. low electrical conductivity b. high heat conductivity c. existence as brittle solids d. low melting points

  18. Metal/Nonmetal Examples 4. At room temperature, which substance is the best conductor of electricity? a. nitrogen b. neon c. sulfur d. silver 5. Which element is brittle in the solid phase and a poor conductor of electricity? a. calcium b. strontium c. sulfur d. copper 6. The majority of elements on the table are in what physical state at STP? a. solid b. liquid c. gas

  19. Atomic Radius • ½ the distance between any two nuclei • Given on Reference Table S in picometers • 1pm = 1x10-12m

  20. Trend within a Period (Left to Right) • Atomic Radius decreases • As you move across a period the number of protons increases, resulting in a stronger nuclear charge therefore, electrons are pulled closer to the nucleus

  21. Trend within a Group (Top to Bottom) • Atomic Radius increases • Remember period number = PEL • As you move down a group there are additional rings, therefore the valence electrons are further away from the nucleus, resulting in a larger radius

  22. Atomic Radius Examples • Which sequence of elements is arranged in order of decreasing atomic radii? • Al, Si, P • Li, Na, K • Cl, Br, I • N, C, B 2. What is the radius of Ca? 3. What is the radius of Sr? 4. Explain why Sr has a larger atomic radius than Ca.

  23. Ionic Radius (IR) • Radius that results from the loss or gain of electrons

  24. Metals (Left Side) • Tend to lose 1 or more electrons when forming positive ions • Radius will decrease • Ionic Radius < Atomic Radius Ex: Na+ is smaller than Na

  25. Non-metals (Right Side) • Tend to gain 1 or more electrons when forming negative ions • Radius will increase • Ionic Radius > Atomic Radius Ex: Cl- is larger than Cl

  26. Ionic Radius Examples • Which ion has the largest radius? a. Na+ b. Mg2+ c. K+ d. Ca2+ 2. Which of the following elements has an ionic radius smaller than its atomic radius? • Neon b. Nitrogen c. Sodium d. Sulfur 3. The Na+ ion has a smaller radius than the Ne atom, even though they both contain 10 electrons. Explain why this is so.

  27. Ionization Energy • The energy required to remove the most loosely bound electron from an atom • Low IE = greater tendency to lose electrons and form positive ions • High IE = greater tendency to gain electrons and form negative ions • Given on Reference Table S in kilojoules per mole (kJ/mol)

  28. IE within a Period (Left to Right) • IE increases • Number of protons increases, resulting in a stronger nuclear charge • The nucleus has a better hold on the electrons, therefore more energy is required to remove an electron

  29. IE within a Group (Top to Bottom) • IE decreases • The principle energy levels increase, so the valence electrons are further away • Protons cannot hold onto the valence electrons as well, therefore less energy is required to remove an electron

  30. Electronegativity (e-neg) • Scale that measures the ability of an atom to attract electrons from another atom • Reference Table S • Scale ranges from 0.0-4.0 • Fluorine is the highest = 4.0 • Difference between electronegativity between two atoms can be used to determine the type of bond

  31. Trend within a Period (Left to Right) • Electronegativity Increases • More protons, resulting in a stronger nucleus, therefore the nucleus is better able to attract electrons

  32. Trend within a Group (Top to Bottom) • Electronegativity decreases • The atom is larger, so the nucleus is further away from the valence shell, therefore the nucleus is less able to attract electrons

  33. IE/e-neg Examples • Which element will lose electrons the easiest? a. Na b. Cl c. K d. F 2. Which element would be most likely to gain electrons? a. Na b.Cl c. K d. F

  34. Reactivity Metals • Most reactive = loses electrons the easiest (low ionization energy) • Lower left corner of the table (Fr) Nonmetals • Most reactive = gains electrons the easiest (high electronegativity) • Upper right corner of the table (F, Cl, O) – not group 18

  35. Periodic Trends Questions • What is the ionization energy of K? • What is the ionization energy of Ca? • Explain why K has a lower ionization energy than Ca. • According to the reference table, which of the following elements has the smallest radius? a. Ni b. Co c. Ca d. K

  36. Periodic Trends Questions 5. An element with high ionization energy would most likely be? a. A nonmetal with low electronegativity b. A nonmetal with high electronegativity c. A metal with low electronegativity d. A metal with high electronegativity 6. What happens to S when it becomes S2-? a. It loses two electrons and the radius increases b. It loses two electrons and the radius decreases c. It gains two electrons and the radius increases d. It gains two electrons and the radius decreases

  37. Periodic Trends Questions 7. Which of the following would have the largest radius? a. Na b. Na+1 c. Cl Cl-1 8. Which of the following has the greater ionization energy, Na or Na+? Explain your answer.

  38. Groups 1 and 2 • Properties: typical metallic characteristics • High reactivity (valence electrons are easily lost) • Only occur in nature as compounds • Reactivity increases as you move down the group • Group 1 elements are more reactive than group 2 elements • For metals – Low IE = high reactivity (electrons are easily lost) – Fr is the most reactive • Na and Water • K and Water

  39. Examples: • Which atom is the most reactive? • Na b. Mg c. K d. Ca • Which group 15 element has the least metallic character? • N b. P c. As d. Sb 3. Explain why reactivity increases as you move down group 1.

  40. Group 17 - Halogens • Typical Nonmetals • High electronegativity – F is the highest • High ionization energy • Are so reactive that they cannot exist as in the monoatomic form • Exist in nature as diatomics (HOFBrINCl) • F2 and Cl2 are gases • Br2 is a liquid • I2 and At2 are solids • F is the most reactive nonmetal • For nonmetals – high electronegativity = high reactivity

  41. Group 18 – Noble Gases • Exist as gases at STP • Exist as monatomic molecules (not combined with anything Example: He, Ne, Ar • The outermost ring is complete, therefore they are VERY STABLE (unreactive)

  42. Groups 3-12 (Transition Elements) • Hard solids • High melting points (except mercury) • Multiple Positive Oxidation States • They can react with electrons from both s and d sublevels • Different numbers of electrons can be lost • Colored Ions • Easily excited (since d and s sublevels are close)

  43. Group Examples • The presence of which ion usually produces a colored solution? a. K+ b. F- c. Fe2+ d. S2- • Which solution would be colored? a. KNO3 b. Ca(NO3)2 c. Cu(NO3)2 d. Al(NO3)3 • Which element at STP exists as monatomic molecules? a. Ne b. N c. Cl d. O • Which is a solid at STP? a. F2 b. Cl2 c. Br2 d. I2 • Which element in Period 4 is classified as an active nonmetal? a. Ga b. Ge c. Br Kr

  44. Group Examples • Which noble gas would most likely form a compound with fluorine? a. He b. Ne c. Ar d. Kr • Which element in Period 3 is the most reactive metal? a. Na b. Mg c. N d. Cl • Which element in Group 15 has the most metallic character? a. Bi b. As c. P d. N

  45. Group Examples • Why is hydrogen not considered to be a member of Group 1? • Why is hydrogen considered to be a member of Group 1? • Why is it unlikely for sodium to form the Na2+ ion?

  46. Allotropes • Some nonmetals can exist in 2 or more forms in the same phase • Allotropes have different physical and chemical properties because their atoms are arranged differently Examples: • oxygen and ozone (O2 and O3) • Graphite and Diamond (carbon) Diamond Graphite