# Periodic Table & Periodic Trends I. Importance of Classification II. History & Development law of octaves III. Periodic Law

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1 Periodic Table & Periodic Trends I. Importance of Classification A. Makes large sums of information manageable. B. In chemistry, it reduces the number of reactions that need to be studied. II. History & Development A. In the 1700s Antoine Lavoisier made a list of the 33 known elements and placed them in 4 categories. B. In 1860 Stanislo Cannizzaro presented a standard method for measuring atomic masses. C. In 1864 John Newlands arranged the elements in groups of 8 and noticed that elements grouped this way had similar properties. The law of octaves related the pattern of properties based on a musical analogy. D. In 1870 Lothar Meyer and Dimitri Mendeleev observed a connection between atomic mass and element properties. Mendeleev worked out the first Periodic Table similar to the Modern Periodic Table. a. He arranged it in order of increasing atomic mass. b. Left spaces for undiscovered elements and allowed for the prediction of their properties. E. In 1911 Henry Moseley discovered that the elements differed by an increase in positive charge. When the table was organized by an increase one in the positive charge many of the exceptions to Mendeleev s arrangement were removed. The Periodic Table is now arranged by increasing atomic number. III. Periodic Law: Physical & chemical properties of elements are periodic functions of their atomic numbers. A. Periodic- any regular repeating pattern B. Occurs because the valence electrons determine how an element will react. IV. Regions of the Periodic Table There are three main regions of the periodic table: A. Metals largest region of the table and left of the stair step. Basic properties: good conductors (poor insulators) of heat and electricity, ductile, malleable, most are solids at room temperature, shiny. B. Nonmetals to the right of the stair step. Basic properties: good insulators (poor conductors) of heat and electricity, most are either brittle solids or gasses at room temperature, Br is the only liquid. C. Metalloids the elements bordering the stair step and located between the metals and nonmetals. Metalloids have the properties of both metals and nonmetals and are considered semiconductors or semimetals. The elements that are metalloids are B, Si, Ge, As, Sb, Te.

2 V. Period or Series- the horizontal rows of the Periodic Table. There are seven periods and they correspond to the seven energy levels (n), range from 1-7. VI. Group or Family- Vertical columns of the Periodic Table. All the elements in a column have similar properties because they have the same number of valence electrons. Range from A. Group 1 or IA- Alkali Metal or Sodium Family 1. Francium is the most active member (Fr) 2. All have one valence electron 3. They have a +1 oxidation number in a compound 4. Highly reactive B. Group 2 or IIA- Alkaline Earth Metals or Calcium Family 1. Radium is the most active member 2. They have two valence electrons 3. They have a +2 oxidation number in 4. Highly reactive C. Group 13 or IIIA- Boron Family 1. Boron is the most active member 2. They have three valence electrons 3. They have a +3 oxidation number in 4. Boron is a metalloid D. Group 14 or IVA- Carbon Family 1. Carbon is the most active member 2. They have four valence electrons 3. They have a +2 or +4 oxidation number in 4. Silicon and Germanium are metalloids E. Group 15 or VA- Nitrogen Family 1. Nitrogen is the most active member 2. Nitrogen & Phosphorus are nonmetals 3. As & Sb are metalloids 4. They have five valence electrons 5. They have a -3 oxidation number in F. Group 16 or VIA- Oxygen Family 1. Oxygen is the most active member 2. O, S, & Se are nonmetals 3. Tellurium (Te) is a metalloid 4. They have 6 valence electrons 5. They have a -2 oxidation number in G. Group 17 or VIIA- Halogens or Fluorine Family 1. Fluorine is the most active member 2. All are nonmetals 3. They have 7 valence electrons 4. They have a -1 oxidation number in 5. Highly reactive H. Group 18 or VIIIA- Noble Gas or Inert Gas Family 1. These are inactive; do not react, inert 2. Have 8 valence electrons except for He which has 2 valence electrons VII. Transition Metals Are found in the center of the periodic table and make up the majority of the elements. The Lanthandide and Actinide Series are specifically called the Inner Transition Metals. VIII. Periodic Trends/Properties A. Atomic Radius (AR)- Radius-half the distance between the nuclei in a molecule consisting of identical atoms; influenced by nuclear pull and number of energy levels. Zeff- the effective nuclear charge of an atom; the degree of attraction of the positive nucleus to the negative electron cloud. Going across a period: atomic number = atomic radii -Zeff increases attraction of the nucleus pulls the electron cloud closer smaller radius Going down a group: atomic number = atomic radii -Increased number of energy levels (n) increases the distance over which the nucleus must pull and thus reduces the attraction for electrons -Full energy levels provide some shielding between the nucleus and valence electrons

3 B. Ionization Energy (I.E.)- I.E.-Measures how tightly an atom holds on to its electrons. Equal to the amount of energy needed to remove an e - from a gaseous atom or ion. Always endothermic. Going across a period: atomic number = ionization energy - Zeff increases attraction of nucleus and holds e - more tightly - Some exceptions occur in groups II-III & V-VI Drops in I.E. between groups II-III occur because p-orbital e - do not penetrate the nuclear region like s-orbital e - so they are not as tightly held Drops in I.E. between groups V-VI occur because increased repulsion created by first pairing of electrons outweighs Zeff less energy is required to remove electrons Going down a group: atomic number = ionization energy - An increased number of energy levels (n) increases the distance over which the nucleus must pull the attraction for electrons is reduced - Full energy levels provide some shielding between nucleus and valence electrons Generally Speaking: - Metals- low I.E. easy to remove electrons - Nonmetals- high I.E. hard to remove electrons - Transition elements- intermediate I.E. - Noble gases- very high I.E. Atomic Radius First Ionization Energy C. Electron Affinity (E.A.)- E.A.- Measures an atoms ability to attract electrons Involves the addition of an e - to a gaseous atom or ion (NOTE: E.A. is NOT opposite of I.E.) Affinity can be exothermic or endothermic; a negative sign (negative E.A.) indicates the direction of energy flow out of the system Across a period: atomic number = electron affinity Down a group: atomic number = electron affinity Generally Speaking: - Nonmetals- high E.A. - Metals- low E.A. D. Electronegativity- Property that measures the attraction of an atom for the pair of outer shell electrons in a covalent bond with another atom

4 Patterns are the same as electron affinity Focus is on the attraction that the nucleus has for electrons Across a period: atomic number = electronegativity - Zeff increases attraction of nucleus and strengthens attraction for electrons Down a group: atomic number = electronegativity - Increased number of energy levels (n) increases the distance over which the nucleus must pull the attraction for electrons is reduced - Full energy levels provide some shielding between nucleus and valence electrons E. Ionic Radius (I.R.)- I.R.- The distance from the nucleus to the outer edge of the electron cloud in a charged atom Positive Ions (cations)- - Are smaller than their respective neutral atoms - Result from the loss of valence electrons - The ratio of p to e increases electrons are held closer and with more strength Negative Ions (anion)- - Are larger than their respective neutral atoms - Result from the addition of valence electrons - Ratio of p to e decreases electrons are not held as closely - Ratio of e to e increases repulsions; this also plays a role in expanding e cloud F. Reactivity- Depends on whether the element reacts by losing electrons (metals), or gaining electrons (nonmetals) Metals: - More reactive as you move down a group, directly tied to I.E. - Metals react by losing electrons and loosely bound electrons result in a more active metal - With an increase in the number of energy levels (n) comes an increase in distance from nuclear attraction and thus more loosely held electrons available for reacting. Nonmetals: - More reactive as you move up a group - Nonmetals tend to gain electrons - A strong nuclear attraction will result in a more reactive nonmetal atoms with high Zeff and least number of energy levels (n) should be the most reactive nonmetal because its nucleus exerts the strongest pull. Electronegativity

5 Name: Date: Pd.: Part A: Use the words in the box to fill in the blanks. Worksheet 1-Atomic Structure & the Periodic Table metals isotopes average atomic mass electron cloud groups metalloids transition elements atomic number electrons nucleus mass number periods chemical symbol quarks periodic table 1. A capital letter or a combination of a capital letter and a small letter that is used to represent an element is called a(n). 2. The horizontal rows of elements are called. 3. An average of masses of all the isotopes that occur in nature for an element is the. 4. Vertical columns of elements are called. 5. Elements in the middle of the periodic table, periods 4 through 7, are called the. 6. The number of protons in an atom is the. 7. Protons and neutrons can be subdivided into by colliding them. 8. The center of an atom where protons and neutrons are located is the. 9. A total count of the neutrons and protons in an atom is the. 10. Atoms of the same element but with different numbers of neutrons are. 11. Elements that are found on the left side of the periodic table are. 12. Elements that have some properties of both metals and nonmetals are. 13. The particles that move about the nucleus and have a negative charge are. 14. The region around the nucleus occupied by electrons is a(n). 15. A chart that shows the classification of elements is called the. Part B: Answer the following questions in the space provided. 1. In the modern periodic table, elements are ordered a. According to decreasing atomic mass c. According to increasing atomic number b. According to Mendeleev s original design d. Based on when they were discovered 2. The modern periodic law states that a. No two electrons with the same spin can be found in the same place in an atom b. The physical and chemical properties of the elements are functions of their atomic number c. Electrons exhibit properties of both particles and waves d. Chemical properties of elements can be grouped according to periodicity, but physical properties can t 3. Mendeleev noticed that properties of elements appeared at regular intervals when the elements were arranged in order of increasing. a. density b. reactivity c. atomic number d. atomic mass

6 4. The discovery of the noble gases changed Mendeleev s periodic table by adding a new. a. period b. series c. group d. sublevel block 5. The most distinctive property of the noble gases is that they are a. metallic b. radioactive c. metalloids d. largely unreactive 6. Moseley's work led to the realization that elements with similar properties occurred at regular intervals when the elements were arranged in order of increasing a. atomic mass b. density c. radioactivity d. atomic number 7. What are the elements whose discovery added an entirely new row to Mendeleev's periodic table? a. noble gases b. alkali metals c. transition elements d. metalloids 8. What are the radioactive elements with atomic numbers from 90 to 103 in the periodic table called? a. the noble gases b. the lanthanides c. the actinides d. rare-earth elements 9. What are the elements with atomic numbers from 58 to 71 in the periodic table called? a. the lanthanide elements b. the noble gases c. the actinide elements d. the alkali metals 10. Argon, krypton, and xenon are a. alkaline earth metals b. noble gases c. actinides. d. lanthanides. 11. Elements in a group or column in the periodic table can be expected to have similar a. atomic masses. b. atomic numbers c. numbers of neutrons d. properties. 12. To which group do lithium and potassium belong? a. alkali metals b. transition metals c. halogens d. noble gases 13. To which group do fluorine and chlorine belong? a. alkaline-earth metals b. transition elements c. halogens d. actinides 14. The electron configuration of aluminum, atomic number 13, is [Ne] 3s2 3p1. Aluminum is in Period a. 2 b. 3 c. 6 d Elements to the right side of the periodic table (p-block elements) have properties most associated with a. gases b. nonmetals c. metals d. metalloids. 16. Within the p-block elements, the elements at the top of the table, compared with those at the bottom, a. have larger radii c. have lower ionization energies. b. are more metallic d. are less metallic. 17. The most reactive group of the nonmetals are the a. lanthanides b. transition elements c. halogens d. rare-earth elements 18. The element that has the greatest electronegativity is a. oxygen. b. sodium c. chlorine d. fluorine 19. In a row in the periodic table, as the atomic number increases, the atomic radius generally a. decreases c. increases b. remains constant d. becomes unmeasurable

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9 Name: Date: Pd.: Practice Problems- Periodic Table 1. Within a group of elements, as the atomic number increases, the atomic radius a. increases c. decreases regularly b. remains approximately constant d. decreases, but not regularly 2. In the alkaline-earth group, atoms with the smallest radii a. are the most reactive c. are all gases. b. have the largest volume. d. have the highest ionization energies. 3. As the atomic number of the metals of Group 1 increases, the ionic radius a. increases. c. remains the same. b. decreases. d. cannot be determined. 4. Across a period in the periodic table, atomic radii a. gradually decrease. b. gradually decrease, then sharply increase. c. gradually increase. d. gradually increase, then sharply decrease. 5. Which is the best reason that the atomic radius generally increases with atomic number in each group of elements? a. The nuclear charge increases. c. The number of energy levels increases. b. The number of neutrons increases. d. A new octet forms. 6. The energy required to remove an electron from an atom as you move left to right from potassium through iron. a. generally increases c. does not change b. generally decreases d. varies unpredictably 7. The force of attraction by Group 1 metals for their valence electrons is a. weak b. zero c. strong d. greater than inner shell electrons 8. Across a period, the atomic radii of d-block elements generally a. increase. b. decrease c. remain constant d. increase and then decrease. 9. As with main-group elements, ionization energies of d-block elements generally across a period. a. increase b. decrease c. remain constant d. drop to zero 10. In, d-block elements most often form ions with charge a. 2 b. 1 c. 1+ d. 2+

10 Chemistry: Ionization Energies Directions: Below is a table of the 1 st, 2 nd, and 3 rd ionization energies for the first 20 elements. On the graph, plot the 1 st ionization energy vs. atomic number. (The atomic number should be along the x-axis.) Then, on the same graph, plot the 2 nd ionization energy vs. atomic number, and similarly for the third ionization energy. Use all of the elements of a good graph. After completing your graph, answer the questions at the bottom of this page. Atomic Number Chemical Symbol 1 st Ionization Energy (kj/mol x 10 3 ) 2 nd Ionization Energy (kj/mol x 10 3 ) 3 rd Ionization Energy (kj/mol x 10 3 ) 1 H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca In general, what happens to the 1 st ionization energy as you go across a period? 2. In general, what happens to the 1 st ionization energy as you go down a group / family? 3. List the elements for which the 2 nd ionization energy is significantly higher than the 1 st (say, more than four times higher). 4. Explain why the elements you listed in your answer to question three have such large 2 nd ionization energies. 5. List the elements for which the 3 rd ionization energy is significantly higher than the 2 nd (say, more than four times higher). 6. Explain why the elements you listed in your answer to question five have such large 3 rd ionization energies.

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