Chapter 6 - Periodic Table
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Chapter 6 - Periodic Table OBJECTIVES
1. List the early attempts of classification of the elements. 2. Match scientists and their contributions to the development of the P.T. 3. State the modern periodic law. 4. Distinguish groups and periods in the P.T. 5. Define chemical stability using the octet rule. 6. Use the periodic table to predict electron configurations. 7. List properties of metals and nonmetals. 8. Distinguish metals from nonmetals in terms of electronic configuration. 9. Use the P.T to locate the : alkali metals, alkaline earth metals, chalcogens, halogens, noble gases, transition elements, lanthanides, and actinides. 10. State the relationship between the properties of elements and their electronic configuration. 11. Define periodic, and list reasons why a group has properties related to position in the periodic table. 12. State the general relationship between size of the atoms with their positions in the P.T. 13. Given a pair of atoms, ions, or an atom and an ion, select the particle with the larger or smaller radius. 14. Predict and explain the oxidation number of an element given the element's location in the P.T. 15. Define first ionization energy. 16. Predict the relative first ionization energies of two elements given their positions in the P.T. 17. Distinguish metals from nonmetals in terms of ionization energy. 18. Define the effective nuclear charge, and list the factors that affect ionization energy, radius, and electron affinity. 19. Explain the variations in ionization energies of elements as their atomic numbers increase in a period or family of the P.T. 20. Define electron affinity and electronegativity. 21. Predict trends in electron affinities and electronegativities for elements, based on their location in the P.T. 22. Define family, or group, and explain what members of a chemical family have in common. 23. Describe some properties of the elements of the representative groups in the P.T. 24. State the relationship between the activities of metallic and nonmetallic elements in their relation to their locations in the P.T. 25. Given two elements, determine which is more chemically active within the metals and nonmetals.
READING:
Chapter 2, Section 2.3, Introduction to Periodic Table pages 35-38, Chapter 6, Section 6.7: Periodic Trends and the properties of Aroms, pages 166-172 Notes.
ASSIGNMENTs:
Dittos in the packet. Textbook, page 176, questions # 54-60 evens; page 177 questions # 64, 74
ch6_pt1.doc 1 04/05/18 Chapter 6 - Periodic Table Introduction 1. How many groups are in the periodic table? ______How many periods? ______2. Define atomic number ______What information does it convey? ______3. What is mass number ? ______4. What number, other than atomic number is represented for every element in the periodic table?____ 5. Why is this number a decimal number, rather than a whole number? ______6. What is an isotope? ______7. How can you use the periodic table to determine the number of electrons in an atom? ______8. How many elements are in period 1______? 2? ______3? ______4? _____ 9. In which group are the noble gases found? ______10. Why are the noble gases special? ______11. In how many groups are the s-sublevels filled? ______p-sublevels filled? ______d-sublevels filled? ______f-sublevels filled? ______12. How does the number of electrons in each sublevel compare with the number of groups in each block? ______13. What block in the periodic table contains the transition elements? ______The rare earth elements (lanthanides and actinides)? ______14. What do the elements in a group have in common? ______15. What is an ion? ______16. How does an atom form an ion?______17. How does an atom form a cation?______18. How does and atom form an anion?______19. Which groups form cations? ______20. Which groups form anions? ______21. Draw the electron diagram for elements in groups 1, 2, 13, 17, and 18.
22. What is the charge on the ions formed by elements in groups 1, 2, 17, and 18. ( Indicate + or - as well as magnitude of the charge)
ch6_pt1.doc 2 04/05/18 Chapter 6 - Periodic Table
Periodic Properties: ch6_pt1.doc 3 04/05/18 Chapter 6 - Periodic Table 1. For each of the following pairs, use the periodic table to select the atom that is larger in radius. a. Rb, Sr b. Cl, I c. Na, Rb d. Mg, Be e. S, P f. Ac, U g. B, Al h. Au, Ba
2. For each of the following pairs of particles, select the particle that is larger in radius. a. Ca, Ca+2 b. Cl, Cl- c. As-3, P-3 d. Pb4+ , Pb e. Mg2+ , Be2+ f. Te2- , Te g. C, C4- h. Ag, Ag1+
3. Predict the oxidation numbers for the following elements: a. Al b. N c. Cl d. Zn e. Mg f. S g. Na h. Mn
4. Which electrons were gained or lost to complete the outer octet and produce the following ions? a. K+1 b. O-2 c. Ga+3 d. P3- e. Sn4+ f. Br1- g. Ca2+ h. Sc3+
5. Which atom in each of the following pairs would have the lower first ionization energy? a. N, O b. Te, Sn c. Ne, F d. C, Ge e. Br, I f. Mg, Ca g. I, Sb h. Al, N I. F, S
6. Within a group, does the radii of atoms increase or decrease as the atomic number increases?
7. Does the radii of atoms within a period increase or decrease as the atomic number increases?
8. In each of the following pairs of atoms, pick the one that is larger. a. Mg, Na b. K, Ca c. Al, B d. Br, Cl e. F, N f. Ne, Ar
9. In each of the following pairs of particles, pick the one that is smaller. a. Fe, Fe3+ b. S2- , S c. Ac3+ , U3+ d. Br1- , Se2- e. Mo6+ , Mo f. As3- , As
10. Predict the oxidation number of the following elements: a. Li b. Be c. B d. C e. P f. O g. F h. Ar i. K
11. Predict the oxidation number of the following elements: a. Rb b. Co c. Pu d. Bi e. V f. Ba
12. In a group, will the ionization energy tend to increase or decrease with increasing atomic number?
13. In a period, will the ionization energy tend to increase or decrease with increasing atomic number?
14. Do metals generally have lower ionization energies? Explain?
15. Carbon has a first ionization energy of 1086.5 kJ/mol. Predict whether the first ionization energies of the following elements will be more or less than that of carbon. a. helium b. lithium c. fluorine d. silicon
ch6_pt1.doc 4 04/05/18 Chapter 6 - Periodic Table 16. As the distance between the nucleus and the outer electrons of an atom increases, will the ionization energy increase or decrease?
17. As the shielding effect increases, will the ionization energy increase or decrease?
18. As the positive charge on an ion increases, will the ionization energy increase or decrease?
Oxidation Numbers:
1. Predict the oxidation numbers for the following elements:
a. Al ______b. N ______ch6_pt1.doc 5 04/05/18 Chapter 6 - Periodic Table c. Cl ______d. Zn ______
e. Mg ______f. S ______
g. Na ______h. Mn ______
2. Which electrons were gained or lost to complete the outer octet and produce the following ions?
a. K+1 ______b. O-2 ______
c. Ga+3 ______d. P3- ______
e. Sn4+ ______f. Br1- ______
g. Ca2+ ______h. Sc3+ ______
First ionization Energy:
1. Which atom in each of the following pairs would have the lower first ionization energy?
a. N, O b. Te, Sn c. Ne, F d. C, Ge e. Br, I f. Mg, Ca g. I, Sb h. Al, N I. F, S
2. Within a group, does the radii of atoms increase or decrease as the atomic number increases?
3. Does the radii of atoms within a period increase or decrease as the atomic number increases?
4. In each of the following pairs of atoms, pick the one that is larger.
a. Mg, Na b. K, Ca c. Al, B d. Br, Cl e. F, N f. Ne, Ar
General Trends in the Periodic Table:
1. Arrange the elements Rb, Te, and I in order of a. increasing atomic radius ______
b. increasing ionization energy ______ch6_pt1.doc 6 04/05/18 Chapter 6 - Periodic Table c. increasing electronaffinity (electonegativity) ______
2. List the following species in order of decreasing radius: a. K, Ca, Ca+2, Rb ______
b. S, Te-2, Se, Te ______
3. Name and give the symbol for the element with the characteristics given below. a. Electron configuration 1s22s22p63s23p3 . ______
b. Lowest ionization energy in Group 17. ______
c. Alkali metal with the largest atomic radius. ______
d. Largest ionization energy in the third period. ______
Review(1)
1. In a column , the ionization energy tends to ____ with increasing atomic number. a. increase b. decrease c. remain the same
2. In a period, the ionization energy tends to ___ with increasing atomic number. a. increase b. decrease c. remain the same ch6_pt1.doc 7 04/05/18 Chapter 6 - Periodic Table 3. The _____ the electron affinity, the greater the ionization energy. a. greater b. lesser c. more constant
4. In an experiment designed to measure the first four ionization energies of aluminum, the I.E. would be the greatest. a. first b. second c. third d. fourth
5. ______generally have the lowest ionization energies. a. Noble gases b. Metalloids c. Nonmetals d. Metals
6. The most active ____ have the highest electronegativities.
7. Carbon has a first ionization energy of 1121.6 kilojoules per mole. Predict whether the first ionization energies of the following elements will be more or less that that of carbon. a. helium ______c. fluorine ______
b. lithium ______d. silicon ______
8. Fill in the following blanks with the word "increases" or "decreases." a. As the distance between the nucleus and the outer electrons of an atom increases, the I.E. ______.
b. As the shielding effect increases, the I.E. ______.
c. As the positive charge on an ion increases, the I.E. ______.
Explain the following statements:
9. Ionization energy tends to increase with increasing atomic number along any horizontal row.
10. Ionization energy decreases with increasing atomic number down any vertical column.
11. Explain why there is a tremendous increase between the fourth and fifth ionization energies of the element carbon.
12. Underline the atom in each of the following pairs that has the lower first ionization energy. a. Li, Na b. Kr, Rb c. Cs, Ba d. Cl, Br e. F, Ne f. S, Cl
Properties of Metals and Nonmetals
1. Compare the following in terms of :
Metals Nonmetals
Ionization Energy ch6_pt1.doc 8 04/05/18 Chapter 6 - Periodic Table Metals Nonmetals
Electronegativity
Luster
Deformability
Conductivity of Heat
Conductivity of Electricity
Phase at Room Temperature
Ion Formation
Number of Electrons in outermost Energy Level
2. How do metals and nonmetals differ in terms of electronic configurations?
3. The most active metals are found in what part of the Periodic Table? ______
What are the properties that determine the activity of a metal? ______
4. The alkali metals react with water when cold producing ______
5. The alkaline metals react only with hot water or steam producing ______
6. The most active nonmetals are found in what part of the Periodic Table? What are the factors that determine the activity of a nonmetal? ______
7. The most active nonmetal is ______
8. Draw an arrow showing the increased metallic properties in a period and in a group.
Review(2)
1. a. If n= 4, there may be ______sublevels. b. The p sublevel mat contain ____ pairs of electrons. c. Two electrons occupying the same orbital must have opposite _____. d. The possible quantum numbers for the 4p1 electron are : n = ______; l = ______; m = ______; s = ______
2. How many sublevels are possible for a. n = 6 ______; b. n = 1 ______ch6_pt1.doc 9 04/05/18 Chapter 6 - Periodic Table
3. Match the quantum number with their proper descriptions:
a. number of energy level and describes electron cloud size ______(1) l b. shape of electron cloud ______(2) m c. direction in space of each orbital ______(3) n d. spin of the electron ______(4) s
4. Since the activity of a nonmetal depends upon the ease with which the atom gains electrons, using the periodic table arrange the following elements in order of increasing activity.
iodine, fluorine, bromine, and chlorine ______
5. Metallic ions are ______their corresponding atoms. a. smaller than b. the same size as c. bigger than
6. The tendency to lose electrons ____ as we move down a column. a. increases b. remains the same c. bigger than
7. ______first proposed the law of octaves. a. Mendeleev b. Dobereiner c. Meyer d. Newlands
8. The most stable atoms are those of a. metals b. metalloids c. noble gases d. nonmetals
9. The elements in a(n) ______are all in the same horizontal line.
10. A(n) ______usually has five or more electrons in its outer energy level.
11. In the ______series electrons are added to the 5f sublevel.
12. The ______is formed when elements with similar electron configurations are placed in columns in order of their increasing principal quantum numbers.
13. Which group in the Periodic Table has the outer configuration: a. ns2 b. ns2np2 c ns2np5
14. Give the total number of outermost level electrons (valence electrons) of an element in Group a. 1 b. 13 c. 14 d. 16 e. 17
15. In what group of the Periodic Table do all the elements have: a. 2 valence electrons b. 5 valence electrons c. 6 valence electrons?
NOTES Objective: Describe the origin of the periodic table. State the periodic law.
Origin:
1. Dobereiner, Johann (1780-1849) Arranged elements in triads:
Ca 40 Cl 35.5 Sr 87.6 Av: 88.5 Br 79.9 Av 81.3 Ba 137 I 127 ch6_pt1.doc 10 04/05/18 Chapter 6 - Periodic Table 2. John Newland ( 1837-1898) 1863: arranged elements in order of increasing atomic masses. Noticed: properties repeat every 8th element. Law of Octaves: Properties of elements repeat every 8th element. He arranged his elements in 7 groups of 7 elements each.
> Li Be B C N O F Na > Na Mg Al Si P S Cl K etc.
Concluded: atomic masses are related to chemical properties.
3. Dimitri Mendeleev ( 1834 -1907) 1869 Periodic Law: Properties of elements are a function of their atomic masses.
Strengths of Mendeleev’s Table: Carefully planned. Contained details. Suggested periods of different lengths. Elements with similar properties were arranged in horizontal rows. Left empty boxes, if no element fitted in the spot.
Predicted properties of the undiscovered elements - for the elements that were unknown in his time.
He predicted the existence of the following elements: ekaboron (Sc) eka-aluminum (Ga); ekasilicon (Ge); ekamanganeese (Tc); dvi-manganese (Rh); ekatantalum (Po).
Noticed irregularities in masses, but used properties to arrange the elements: I 126.9 Te 127.6; switched their position in periodic table.
Stated the periodic law.
4. Moseley (1887-1915)
1914: discovered the relationship between the atomic number (# of protons in nucleus) and the place of the element in the Period Table using x-ray techniques.
MODERN PERIODIC LAW: Properties of elements are a periodic function of their atomic numbers.
Objective: Describe the nature of periods and groups in the Periodic table.
Period: Horizontal row. Begins on the left with an active metal, and ends with a noble gas (except 7th period) Numbered 1 through 7 Period 1 is a short period, contains only two elements.
Group: Family, vertical columns numbered 1 - 18. Elements in a group have similar physical and chemical properties. ch6_pt1.doc 11 04/05/18 Chapter 6 - Periodic Table
Every element is a member of both a group and a period.
Periodic Table and Electron Structure:
1. The electron configuration of the outer energy levels (valence level) determines the chemical properties of the elements.
2. Every member of a group has the same electron arrangement in its valence energy level. Helium is an exception: 2 electrons.
3. The number of the period in which an element is found is the same as the number of the energy level of is valence electrons.
4. It is also the same as the number of the occupied energy levels in atoms of the element.
Short Periods: period 1 ( 2 elements), period 2, period 3 (8 elements each)
Long Periods: periods 4 and 5 ( 18 elements) periods 6, 7 (32 elements each)
Transition elements: any element with an atom that has an incomplete d-sublevel, or that gives rise to a cation or cations with incomplete d-sublevels.
Rare Earth Elements: elements with incomplete f-sublevels, or cations with incomplete f sublevels. Lanthanides and Actinoids.
Chapter 6 – Periodic Table Packet
Page 2: 11. a. 2 b. 6 c. 10 d. 1 21. 12. s1-1 s2-2 s2p1 – 13 22. group 1 - +1 group2=+2 1. a. 18 b. 7 13. a. d b. f group17=-1 group 18=0 2. a. # of protons 14. same electron configuration in 3. #p +#n outermost energy level Page 3: 4. atomic mass 15. charged atom or group of 1. Periodic Table 5. mixture of isotopes atoms 2. Atomic mass 6. same # of p; different # of n 16. gains or loses an electron 3. Periodic law 7. atomic number 17. loses electrons 4. Atomic mass 8. a. 2 b. 8 c. 8 d. 18 18. gains electrons 5. Period 9. 18 19. groups 1-14 6. Group 10. inert 20. 15-17 7. Family ch6_pt1.doc 12 04/05/18 Chapter 6 - Periodic Table 8. Transition element 18. increase 9. Lanthanoid series 1. a. high, low b.low, high c. 10. Actinoid series Page 6: Oxidation numbers high, low d. high, low e. 11. Metal 1.a +3 b. –3 c. –1 high, low f. high, low g. 12. Nonmetal d. +2 e. +2 f. –2 solid, gas h. high, high i. 13. Semimetal g. +1 h. +2 Low, high 14. Alkali metal 2. see Page 4 #4 2. Metals never have an valence 15. alkali earth metal electron in the p orbital. 16. halogen First Ionization Energy 3. Group 1; the # of valence 17. noble gas 1. a. N b. Sn c. F electrons d. Ge e. I f. Ca 4. Hydroxide bases Page 4/5: g. Sb h. Al 5. Hydroxide bases 1. a. Rb b. I c. Rb d. Mg e. P 6. Group 18; the octet # of f. Ac g. Al h. Ba Page 7: General Trends in the electrons 2. a. Ca b. Cl- c. As-3 d. Pb e. Periodic Table 7. Depends which PT you look at Mg+2 f. Te-2 g. C-4 h. Ag 1 a. Rb>Te>I b. Rb
f.Ca g. Sb h. Al i. S 1. b 2. 6;1 6. increase 2. a 3. a. n b. l c. ml d. ms 7. decrease 3. a 4. I, Br, Cl, F 8. a. Na b. K c. Al d. Br e.N f. 4. d 5. A Ar 5. d 6. A 9. +3 +3 - a. Fe b. S c. U d. Br e. 6. nonmetal 7. D +6 Mo f. As 7. a. more b. less c. more d. 8. C 10. a. +1 b. +2 c. +3 d. +4, -4 e. less 9. Period –3, +3, +5 f. –2 g. –1 h. 0 8. a. decreases b. decreases c. 10. Nonmetal i.+1 increases 11. Actinoid 11. 1. +1 b.+2,+3 c. +3, +4, +6 9. answer unreadable 12. Group d. +3, +5 e. +2,+4,+5 f. +2 10. no answer here 13. A. 2 b. 14 c. 17 12. decrease 11. no answer here 14. A. 1 b. 3 c. 4 d. 6 e.7 13. increase 12. a. Na b. Rb c. Cs D. Br e. F 15. A. 2 b. 15 c. 16 14. yes f. S 15. a. more b. less c. more d. less 16. decrease Page 9: Properties of Metals and 17. decrease Nonmetals
Atomic mass = 1 Atomic mass = 7 Atomic mass = 9 Atomic mass = 11 clear gas soft silvery-white solid silvery-white solid, yellow brown solid barely soluble in air reacts moderately with water reacts with hydrochloric relatively unreactive with burns in air when ignited to release H2 gas and form a acid to release H2 gas air, or acid d = 0.0001 g/cm3 basic solution, density = 1.8 g/cm3 density = 2.3 g/cm3 melting point = -259 oC tarnishes in air melting point = 1278 oC melting point = 2300 oC boiling point = -253 oC density = 0.53 g/cm3 boiling point = 2970 oC boiling point = 2550 oC does not conduct electricity melting point = 186 oC conducts electricity poor conductor of electricity o forms XCl and X2O; boiling point = 1336 C forms XCl2 and XO; forms XCl3 and X2O3 XCl is acidic in water conducts electricity XO is basic in water
forms XCl and X2O; X2O is basic in water Atomic mass = 12 Atomic mass = 14 Atomic mass = 16 Atomic mass = 19
ch6_pt1.doc 13 04/05/18 Chapter 6 - Periodic Table black solid, relatively clear gas, barely soluble in clear gas, barely soluble in pale yellow gas, unreactive with water, or water water reacts violently with water to acid density = 0.0013 g/cm3 density = 0.0014 g/cm3 release oxygen gas and form density = 2.2 g/cm3 melting point = -210 oC melting point = -218 oC an acidic solution melting point = 3550 oC boiling point = -196 oC boiling point = -183 oC density = 0.0017 g/cm3 boiling point = 4200 oC does not conduct electricity does not conduct electricity melting point = -223 oC o poor conductor of electricity forms XCl3 and X2O3; forms XCl2 and XO; boiling point = -188 C forms XCl4 and XO2; X2O3 is acidic in water does not conduct electricity XO2 is acidic in water forms XCl and X2O
Atomic mass = 23 Atomic mass = 24 Atomic mass = 27 Atomic mass = 28 soft, silvery solid silvery-white solid, silvery-white solid, gray solid, relatively reacts vigorously with water reacts with hydrochloric reacts with hydrochloric unreactive with air, water , to release H2 gas and form a acid to releasw H2 gas acid to release H2 gas or acid basic solution, tarnishes in density = 1.7 g/cm3 density = 2.7 g/cm3 density = 2.4 g/cm3 air melting point = 651 oC melting point = 660 oC melting point = 1420 oC density = 0.97 g/cm3 boiling point = 1107 oC boiling point = 2057 oC boiling point = 2355 oC melting point = 98 oC conducts electricity conducts electricity poor conductor of electricity o boiling point = 880 C forms XCl2 and XO; forms XCl3 and X2O3; forms XCl4 and XO2 conducts electricity XO is basic in water XO2 is acidic in water forms XCl and X2O; X2O is basic in water
Atomic mass = 31 Atomic mass = 32 Atomic mass = 35.5 Atomic mass = 39 red or yellow solid, yellow solid, relatively green-yellow gas, soft-silvery solid, yellow ignites spontaneously unreactive with air, water, or reacts with water to form an reacts violently with water to in air acid acidic solution release hydrogen gas and density = 2.2 g/cm3(red) density = 2.0 g/cm3 density = 0.0032 g/cm3 form a basic solution, density 1.8 g/cm3 -(yellow) melting point = 116 oC melting point = -103 oC tarnishes in air melting point = 44 oC boiling point = 445 oC boiling point = -35 oC density = 0.87 g/cm3 boiling point = 280 oC does not conduct electricity does not conduct electricity melting point = 62 oC o poor conductor of electricity forms XCl2 and XO; forms XCl and X2O; boiling point = 760 C forms XCl3 and X2O3; XO2 and XO3 X2O is acidic conducts electricity X2O3 is acidic in water are acidic in water forms XCl and X2O, X2O is basic in water
A t o m P r o p e r t i e s E x e r c i s e
Different properties of the elements are related in systematic way to their atomic number. This is known as the Periodic Law.
In this exercise you will investigate the relationships between the atomic numbers of the first 20 elements (H through Ca) and the following properties:
1. First Ionization Energy 2. Atomic Radius (covalent radius)
You will do this by making full-page graphs of the two properties (Y-axis) versus atomic number (X-axis). The property values can be found in your Periodic Table or in the provided Table.
Procedures:
1. Prepare 2 graphs: ch6_pt1.doc 14 04/05/18 Chapter 6 - Periodic Table a) Ionization energy versus atomic number (ionization energy on the y-axis, atomic number on the x- axis). b) Atomic radius versus atomic number
2. Connect each point with solid lines.
3. Now, using a different color for each set of elements, connect all of the elements found in Group 1 of the Periodic Table. Repeat this procedure for the elements in Group 16 and Group 18.
Questions:
1. As you go from element 1 to element 20, what is the general overall pattern for? a. ionization energy? b. for atomic radius?
2. A horizontal row on the Periodic Table is called a “period” . a. List the elements in period 1; in period 2; in period 3; in period 4. b. Within a period, how do the values change for (1) Ionization energy (2) atomic radius?
3. Within the same column ( 1, 16, 18 ), how do the values for ionization energy and atomic radius change with increasing atomic number?
4. Explain why the radius and the first ionization energies of the elements are considered periodic properties.
5. State the modern version of the Periodic Law. How does it differ from Mendeleyev’s version?
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