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Home , Period 5 element

Accelerated Chemistry Study Guide The , Chapter 5

Terms, definitions, and people

 Dobereiner

 Newlands

 Mendeleev

 Moseley

 Periodic table

 Periodic Law

 family

Page 1 of 38  alkali

 metal

 semimetal,

 s-, p-, d-, and f-blocks

electrons (s and p)

 isoelectronic

 abbreviated

Page 2 of 38  periodic trend

 ionic radius

 ionization energy, successive ionization energies

Specific topics to know

 periodic trends within a period

 periodic trends within a family

Page 3 of 38 The World of Chemistry Episode 7 - The Periodic Table

1. What two types of properties are described in the video?

2. What are some examples of physical properties?

3. How many elements are on the modern periodic table? How many of these can be found in nature?

4. Why do the symbols for some elements (such as ) seem to have no relationship to their name?

5. What is meant by the atomic and mass number of an element?

6. Elements in the periodic table are arranged by increasing ______number.

7. What is a group of elements? a period?

8. What are the ingredients used in the making of glass? What determines the color of glass?

9. What are alkali ? Describe their reaction with water.

10. How does the size of an change - a. as you go down a group of elements? b. as you go from left to right in a period of elements?

11. Who developed the periodic table?

12. What did Mendeleev do for elements that had not yet been discovered?

13. How did Glenn Seaborg change the periodic table?

14. Why are the electrons in the outer shell of an atom important?

Page 4 of 38 The Periodic Table

 Elements in the periodic table are arranged in order of increasing .

 Atomic number indicates the number of protons, and if the element is neutral the number of protons is equal to the number of electrons.

 The periodic table is arranged in four blocks: s, p, d, and f that group together elements with similar electron configurations.

 The horizontal rows are called periods.

 The vertical columns are called groups or families. Elements in the same column are chemically similar to each other because each element in the family has the same number of valence electrons.

 The number of valence electrons (outer shell electrons) is equal to the number of s and p electrons in the highest principal .

 The Periodic Table consists of , and metals, and elements become more “metallic” as you move from the right to the left.

 Periodic Trends. Make sure that in each case you can describe the trend (up/down and left/right) and also explain why the trend occurs.

Atomic Size—the distance from the nucleus of an atom to the outermost electrons. It depends entirely on the outer shell electrons. Notice that though the nucleus makes up most of the mass of the atom, it takes up very little space. In a typical atom, the nucleus would be about the size of a marble on the 50-yard line of a football stadium with the electrons orbiting like crazed fans in the upper stands.  Vertical trend: Atomic size increases moving down the periodic table  Explanation: Atomic size increases going down because the principal energy level increases (2p, 3p, 4p for example). This means that the outer shell electrons have more energy and can move farther from the nucleus.

 Horizontal trend: Atomic size decreases when moving left to right.  Explanation: As you move left to right along a period, the atomic size decreases because outer shell electrons (valence electrons) are added to the same principal energy level or lower (ex: 4s then 3d then 4p in period 4) while the number of protons increases. In other words, as the atomic number increases, the number of protons increases. Remember that protons attract electrons because protons are positive and electrons are negative. The fact that there are more protons causes the electrons to be drawn in towards the nucleus.

1 of 3 Page 5 of 38 Ionic size -- the distance from the nucleus of an ion to the outermost electrons. The trends in atomic radius are different from the trends in ionic radius. It is important to keep in mind that the number of electrons is different for ions than it is for neutral elements. It is possible to predict how much bigger or smaller an ion will be by comparing its relative number of protons and electrons and comparing that number to a neutral element with the same number of protons or electrons.

Ionization energy—the energy required to remove an electron. In equation form, this looks like: X  X+ + e- Notice that this equation looks reasonable for metals: Na  Na+ + e- Mg  Mg+ + e- Sn  Sn+ + e- And should look unreasonable for nonmetals: Cl  Cl+ + e- O  O+ + e-  Vertical trend: ionization energy decreases moving down the periodic table  Explanation: as with ionic radius, electrons are in higher principal energy levels moving down the PT. This means that they are of higher energy and farther from the nucleus. This makes them easier to remove. There is another factor at work here: inner shell electrons (the ones that aren’t the outer shell or valence electrons) get in the way of the protons’ ability to hang on to the electrons. This is called the shielding effect.

 Horizontal trend: ionization energy increases going left to right () on the PT  Explanation: this occurs for the same reason as the horizontal trend in atomic size. More protons, same principal energy level—electrons are held more strongly and harder to remove.

Electronegativity: the tendency for an atom to attract electrons to itself when forming a bond with another element. The vertical and horizontal trends for electronegativity (and their explanations) are the SAME as those for ionization energy. There is an exception: , not , has the highest electronegativity on the periodic table and similarly the are the group with the highest electronegativity, not the noble gases. This is because the noble gases have a full outer shell of electrons and do not tend to attract additional electron density. The halogens, on the other hand strongly attract electrons in order to fill their outer shell (remember that they have seven valence electrons, but would be most stable with eight)

2 of 3 Page 6 of 38 Electron affinity: the energy change that occurs when an atom gains an electron. This is how electron affinity can be expressed as an equation: Ne (g) + e-  Ne- electron affinity = 29 kJ/mol F (g) + e-  F- electron affinity = -328 kJ/mol

Notice that the reaction is highly exothermic (releases energy, which you know because the heat change is negative) in the case of fluorine and endothermic in the case of . This makes sense because fluorine is made much more stable when it gains an electron, so it releases heat as it becomes more stable. The gain of an electron is favored for fluorine.

In some ways, electron affinity can be viewed as the opposite of ionization energy, but there are some ways in which they are not completely opposite. For instance, the noble gases have a high ionization energy but low and positive electron affinity

3 of 3 Page 7 of 38

Page 8 of 38 Regions of the Periodic Table: families, states of matter

Regions of the Periodic Table: s-, p-, d-, f-blocks

Page 9 of 38 Regions of the Periodic Table: metals, nonmetals, metalloids

Valence electrons

Page 10 of 38 Periodic trends: atomic size, ionization energy

Periodic trends: electron affinity

Page 11 of 38 Periodic trends: electronegativity

Periodic trends: summary

Page 12 of 38 Page 13 of 38 Page 14 of 38 Page 15 of 38

Page 16 of 38 Page 17 of 38 Page 18 of 38 ChemQuest 16

Information: Shielding

FIGURE 1: “Bohr Diagrams” of , and

nucleus charge = +5 nucleus charge = +6 nucleus charge = +7 Boron Carbon Nitrogen

Because the nucleus is positively charged, it exerts an attractive force on the electrons. However, the three electrons in boron’s outer energy level do not feel the full +5 attraction from the 5 protons in boron’s nucleus. Before the +5 attraction gets to the outer energy level it gets partially cancelled (or “shielded”) by the two electrons in the first energy level. The two electrons in the first energy level weaken the attractive force by two. Therefore to the outer energy level it only “feels” like a +3 charge rather than a +5 charge from the nucleus.

Consider the diagram of carbon. An electron in the outer energy level only “feels” a charge of +4 coming from the nucleus because the two electrons in the first energy level shield two of the positive charges from the nucleus.

Critical Thinking Questions

1. How large is the charge that the second energy level of nitrogen “feels” from the nucleus?

2. Why does the first energy level in each of the three above diagrams only contain two electrons?

3. How many electrons can fit in the second energy level of any atom?

4. How many electrons can fit in the third energy level?

Page 19 of 38 5. How many energy levels does aluminum have? How many electrons should be in each energy level?

6. Draw a Bohr diagram for aluminum similar to those above.

7. Explain why the second energy level of aluminum only feels a +11 attraction instead of a +13 attraction from aluminum’s electrons.

8. How large is the charge that the third energy level of an aluminum atom “feels” from the nucleus?

Information: Charge and Distance

As you know, opposite charges attract. Examine the following diagrams of charged metal spheres.

FIGURE 2: Diagram A Diagram B Diagram C

+3 -3 +4 -4 +5 -5

The attraction between the two charged metal spheres in each diagram is represented by an arrow. The metal spheres are pulled closer together in diagram C because of the +5 to -5 attraction is stronger than the +4 to -4 attraction in Diagram B and the +3 to -3 attraction in Diagram A.

Electrons behave the same way as the metal spheres are ddepicted in Figure 2. Consider the Bohr diagrams of boron, carbon and nitrogen in Figure 1. Recall that Boron’s outer electrons feel a +3 attraction from the nucleus. Carbon’s outer electrons feel a +4 attraction. In question 1, you found out that nitrogen’s outer electrons feel a +5 attracttion.

This attraction between the nucleus and outer energy levvel determines the size of the atom. If the attraction is strong the atom is small; if the attraction is weak, the atom spreads out and is larger.

Page 20 of 38

Critical Thinking Questions

9. Which atom is larger: nitrogen or carbon? Why?

10. In a atom, the force of attraction from the nucleus to the outer energy level is +4. Using a Bohr diagram of a silicon atom as an illustration, explain in detail why this is true.

11. Draw Bohr diagrams for and .

12. a) Find the size of the charge attraction between the nucleus and outer energy level for sulfur and for chlorine.

b) Which atom do you predict to be larger: sulfur or chlorine?

c) Explain, in detail, your reasoning to part b.

13. Notice and compare the locations of boron, carbon and nitrogen on the periodic table. Now compare their sizes. Do the same with sulfur and chlorine. There is a general trend in size as you proceed from left to right across the periodic table. What is this trend? In other words, how do in the same row of the periodic table compare to each other in size?

Page 21 of 38 Information: Bohr Diagrams and the Size of Atoms

Examine the following “Bohr Diagrams” of three atoms from the periodic table. FIGURE 3 Atom A Atom B Atom C

Critical Thinking Questions

14. Give the name and atomic number of each atom from Figure 3. Atom A Atom B Atom C Name of the element

Atomic number

15. a) Concerning atoms A, B and C, what is similar about their location in the periodic table?

b) In atoms A, B, and C compare the force of attraction from the nucleus to the outer level of electrons.

16. Draw Bohr diagrams for neon and .

17. a) What is similar about the location of neon and argon in the periodic table

b) Compare the force of attraction between the outer level electrons and the nucleus for neon and argon.

c) Using your answers to 15b and 17b, what can be said about elements in the same column and the force of attraction between their outer energy level and nucleus?

Page 22 of 38 18. Atom A is larger than atom B. The attraction between the nucleus and outer level electrons is equal in atoms A and B, so what other reason could there be for Atom A’s larger size? Propose an explanation based on that structure of atoms A and B.

19. Which is larger—neon or argon? Why is this atom the largest?

20. In general, there is a trend in the sizes of atoms as you move down a column of the periodic table. a) What is this trend?

b) Why does this trend exist? (Explain the basis for the trend.)

21. Order the following lists of elements in order from smallest to largest. a) K, As, Br

b) P, Sb, N

c) S, Ca, Mg, Cl

Page 23 of 38 Skill Practice 16

1. What force of attraction does the second energy level of a atom “feel” from the nucleus? Draw a Bohr diagram and use it to explain your answer.

2. Using the concepts of shielding and attraction, explain why sulfur is smaller in radius than silicon.

3. Why can’t you tell by looking at the periodic table whether chlorine or is larger?

4. Order the following elements from smallest to largest.

A) Al, Na, S, Mg B) C, Sn, Pb, Si

C) K, Se, Ca, Br D) Be, Ca, C, B, Mg

E) Ga, Al, Cl, P F) O, Se, S, Ne

Page 24 of 38 ChemQuest 17

Information: Separating Charges

Examine Figure 1 below where there are three pairs of metal spheres that have different amounts of charge on them. The spheres in diagram C are closer than the others because they have the strongest attraction.

FIGURE 1: Attraction between metal spheres. Diagram A Diagram B Diagram C

+3 -3 +4 -4 +7 -7

Critical Thinking Questions

1. Which would be harder to separate: the two spheres in Diagram A or the two spheres in Diagram B? Why?

2. Draw Bohr diagrams of the following atoms: lithium, nitrogen and fluorine.

3. For each of the atoms in question two compare the attraction between the nucleus and thhe outer level of electrons. Which atom has the strongest attraction between the nucleus and outer electrons?

4. Which would be more difficult: if you wanted to remove an electron from an atom of fluorine or from an atom of nitrogen?

5. Would it be easier to remove an electron from lithium or nitrogen? 6.

Page 25 of 38 Information: Ionization Energy

The amount of energy that it takes to completely remove an electron from an atom is called ionization energy. The first ionization energy is the energy required to remove one electron from an atom’s outer energy level. The second ionization energy is the energy needed to remove a second electron from the energy level. The second ionization energy is always higher than the first ionization energy. Because noble gases have eight electrons in their outer energy level, they are very stable and therefore it takes a very high amount of energy to remove an electron from a noble gas.

Critical Thinking Questions

7. Which would have a higher first ionization energy: phosphorus or aluminum? (You may want to draw a Bohr diagram to help you determine the answer.)

8. a) Phosphorus and aluminum are in the same row of the periodic table. Lithium, nitrogen and fluorine are also in the same row. Using your answers to questions 3, 4 and 6 what do you notice about the ionization energy of elements proceeding from left to right across a row of the periodic table? Does the ionization energy increase or decrease as you go across a period?

b) Explain why you believe this trend exists.

9. Draw a Bohr diagram of . Do you expect the first ionization energy to be high or low for sodium?

10. The second ionization energy of all elements is higher than the first ionization energy, but in sodium the second ionization energy is extra large. Considering the structure of the sodium atom explain why this might be. (HINT: after one electron is removed, what is sodium’s electron arrangement like?).

11. A certain atom in the third period has a third ionization energy that is unusually high. Using the same reasoning you used for question 9, name this atom.

Page 26 of 38 Information: Trend in Ionization Energy in Groups

Consider the following figure of diagrams of two magnets that you wish to separate.

FIGURE 2: Separating Magnets Diagram D Diagram E

Critical Thinking Questions

12. Would it be easier to separate the magnets in Diagram D or those in Diagram E in Figure 2? Assume that each magnet is attracted to the other and that the size of the attraction is the same. Take only the distance between magnets into consideration.

13. How does the distance between the outer level of electrons and the nucleus in phosphorus compare to the distance between the outer level of electrons and the nucleus in nitrogen?

14. Based on your answer to question 12, would it be easier to remove an electron from nitrogen or from phosphorus?

15. Nitrogen and phosphorus are in the same column on the periodic table. From your answers to questions 12 and 13 what happens to the ionization energy as you move down a column in the periodic table?

Page 27 of 38 Skill Practice 17

1. If an atom has a “high first ionization energy” does this mean that it is relatively easy or relatively hard to remove an electron from the atom?

2. Arrange the following atoms in order from lowest to highest 1st ionization energy.

A) Ca, Se, As, Br B) As, N, P, Bi

C) Ga, Al, S, Si D) Li, K, O, C

E) Te, O, S, Po F) In, Te, Sn, I

3. A certain atom in the 2nd period has an unusually high 3rd ionization energy. Name this element. Draw a Bohr diagram and use it to illustrate why you were able to identify this atom.

4. Compare the trends for size and for ionization energy. As the size of an atom increases, what happens to the ionization energy? Explain why the ionization energy seems to depend on the size.

Page 28 of 38 Page 29 of 38 Page 30 of 38

Selected Chapter 6 Notes

6-1 Alkali metals and alkaline earth metals  alkali metals form cations with + 1 charge  alkaline earth metals form cations with +2 charge  both are water- and -reactive  react with water to produce gas, sometimes explosively  become more reactive (gradually more terrifying) as we go down the group

6-2 Transition metals  used in a variety of materials  can form cations with different charges: Fe2+ and Fe3+ or Cu+ and Cu2+  are often highly colored in solution  (Cu, Ag, Au) are all highly inert and have electron configurations that are exceptions to the (4s13d10 for Cu) Coinage metals are used in currency and in coatings for substances to prevent

In general, metals usually lose electrons while nonmetals gain electrons. Typically a metal and a nonmetal form an ionic compound, while two or more nonmetals usually form a covalent compound.

Carbon  Component of all organic compounds  Forms long chains, so it can form large biomolecules and other molecules with repeating units called polymers  Usually forms covalent, not ionic compounds Halogens  Group 17 with 7 valence electrons, form ions with a -1 charge  Highly reactive, and react by “stealing” an electron  Become “gradually more terrifying” as you go UP the group Noble gases  Have 8 valence electrons (Helium has two)  Are inert, typically don’t have any reactions with any other elements

Page 31 of 38

Page 32 of 38 Name______Period Date______

Accelerated Chemistry

“Searching the Periodic Table”

Directions: Answer each of the following using your Periodic Table and the patterns and trends discussed in class.

1. the most metallic element in the halogen family

2. valence electrons are 6s26p1

3. lowest ionization energy in Period 4

4. largest atom in period 3

5. most non-metallic element in Group 2

6. number of electrons that Ti4+ has

7. noble gas element which is isoelectronic with Ti4+

8. element having one more electron than Cu

9. Group to which Sr belongs

10. Group to which Br belongs

11. element in Group 14 with smallest atomic radius

12. valence electrons of the halogens

13. element in the second period with more ionization energy than F

14. Group 13 element in the 5th Period

15. three ions that are isoelectronic with He

16. two elements for which no compounds are known

17. physical state of all alkali metals

18. Group 16 element with the least atomic mass

19. Period 5 element having a slightly smaller volume than Sb

20. element with one less proton than Au

Page 1 of 2

Page 33 of 38 Name______Period Date______

21. halogen with lowest

22. valence electrons of the alkaline earth elements

23. element having exactly two electrons in 4p orbitals

24. element having the largest first ionization energy in the fourth period

25. Group 15 element in Period 5

26. atom having the largest diameter in Period 3

27. atom having the largest mass in Period 3

28. atom having the electron configuration [Kr]5s2

29. Period to which U belongs

30. number of half-filled orbitals in an atom of Hg

31. Period 3 element with the greatest metallic character

32. element with electron configuration [Xe]6s24f145d4

33. ion of Al which is isoelectronic with Ne

34. atom with smallest size

35. ion most commonly formed by Se (Z=34)

36. element with lowest melting point in group 1

37. number of electrons in an atom of , V

38. largest ion of elements in Period 2

39. smallest ion of elements in Group 16

40. atom with highest electron affinity in Period 4

Page 2 of 2

Page 34 of 38 Name: ______Period ____ Date: ______

Accelerated Chemistry 2017-2018 Practice Test Chapters 5&6, Periodic Table and the Elements 67 Points

Part I. Multiple Choice (2 points each, total 26 points)

Identify the letter of the choice that best completes the statement or answers the question.

____ 1. Who classified elements into triads? a. Newlands c. Mosley b. Mendeleev d. Dobereiner

____ 2. On the periodic table, the groups are the a. horizontal rows c. rare earth elements b. transition metals d. vertical columns

____ 3.

On the periodic table, the elements represnted by only section A are the a. metalloids c. p b. s block d. metals

____ 4. What are the elements in Group 2 of the periodic table called a. halogens c. alkali metals b. noble gases d. alkaline earth metals

____ 5. The 2nd period of the periodic table contains a. the s-block elements only c. the s- p- d- and f-block elements b. the s- p- and d-block elements only d. the s-and p-block elements only

Page 1 of 4 Page 35 of 38 Name: ______Period ____ Date: ______

Part II. Free Response Answer the question in the space provided. (57 total points)

6. Identify something these group 14 elements have in common: Carbon (C), (Ge), and (Sn).

What do they have in common? (2 points)

Also Classify each element as a metal, nonmetal, or semimetal. Then state one property that each element is likely to have. (Total 6 points)

(1 point each) Classification Property

C:

Ge:

Sn:

7. Write abbreviated electron configurations for each of the following: element 6, element 19, and element 33. To which chemical family group does each belong? How many valence electrons does each have? (3 points each, total 9 points)

element 6:

element 19:

element 33:

Page 2 of 4 Page 36 of 38 Name: ______Period ____ Date: ______

8. Rank the following elements from smallest to largest based on the designated property: (12 Total Points)

a. Atomic Radius: Mg,Al,S,Ca (3 points)

+1 -2 -1 +2 b. Atomic Radius: Rb , Se , Br , Sr (3 points)

c. Ionization Energy: Li,Be,Na,K (3 points)

d. Electronegativity: Ga, Ge, AS, In (3 points)

9. Identify the period 3 element that has an unusually high sixth ionization energy. (2 points)

10. Identify two elements in the 4d sublevel most likely to have positive electron affinities. Explain why each is a likely candidate. (8 points)

Page 3 of 4 Page 37 of 38 Name: ______Period ____ Date: ______

11. Fictitious symbols are used for the first 18 ellements in the periodic table. Use the clues below to write the fictitious symbol in the appropriate spot on the periodic table provided. Symbols for real elements do not represent those elements. HINT: You do not have to complete each clue in order. ((1 point for each correct answer, total 18 points)

1. has valence electrons are 3s23p1

2. is the most metallic element in Group 2

3. is a noble gas element which is isoelectronic withh 4-

4. is an element in Group 14 with smallest atomic radius

5. is an element that is in the alkaline earth family with more ionization energy than .

6. is a Group 16 element with the least atomic mass

7. is a Period 3 element having a slightly smaller atomic radius than

8. is the element with the largest electronegativity

9. is an element having exactly two electrons in 4p orbitals

10. is an element having the largest first ionization energy in the fourth period

11. is an Group 15 element in Period 4

12. does not readily form compounds

13. is an atom in the p-block having an unusually low ionization energy in Period 3

14. is atom having the electron configuration [ ]3s2

15. has 13 filled orbitals

16. is a period 3 element with the lowest electronegattivity

17. is the atom with smallest size

18. has a positive electron affinity and is a gas at room temperature.

Page 4 of 4 Page 38 of 38