Honors Chemistry Unit 3 ( )

Transcription

1 Honors Chemistry Unit 3 ( ) Quantum numbers Electron orbital shapes Rules: o Aufbau principle o Hund s Rule o Pauli Exclusion principle Orbital notations Electron configuration Noble gas notation Families (research and present) Metals/nonmetals Trends o Atomic radius o Electronegativity o Ionization energy o Metallic and nonmetallic character Review Ions Oxidation # s 1

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3 We are learning to: 1. Describe the quantum mechanical model. 2. Apply the quantum mechanical model. 3. Describe the history/arrangement of the periodic table. 4. Describe periodicity as a result of electron configurations. We are looking for: 1a. Electrons can move from ground state to an excited state by gaining specific amounts of energy (quantum) and return to the ground state by releasing a specific amount of energy (quantum) in the form of photons. 1b. Define each of the 4 quantum numbers (principal, angular momentum, magnetic, spin). 1c. Describe the shape of the various orbital shapes (s, p, d, f). 2a. Use the Aufbau principle, Hund s rule, and Pauli exclusion to assign electron/orbital/noble gas configurations for a given element. 2b. Identify the exceptions to the Aufbau principle and Hund s rule for electron configurations. 2c.Give the 4 digit quantum number for a specific electron in an orbital notation. 2d. Indicate the specific electron in an orbital notation that is described by the given 4 digit quantum number. 2e. Identify the s, p, d, and f blocks on the periodic table. 3a. Doberiener arranged the elements into triads. 3b. Newlands arranged the elements by the law of octaves. 3c. Mendeleev arranged the elements by atomic mass (periodic law). 3d. Moseley arranged the elements by atomic number (modern periodic law). 3e. Identify and describe the unique properties of the families on the periodic table (alkali metals, alkaline earth metals, transition metals, inner transition metals, post transition/other metals, metalloids, halogens, noble gases, other nonmetals, hydrogen). 4a. Identify the number of valence electrons using the periodic table or given an electron configuration. 4b. Describe the periodic trends (metallic character, electronegativity, ionization energy, atomic radius, ionic radius). 4a. Given the name of a polyatomic ion, write the corresponding formula and charge. 4b. Given the formula and charge of a polyatomic ion, write the corresponding name. 3

4 Name Class Flame Test: Nichrome Wire Prelab: Complete the following on a separate sheet of paper: 1. How much of each solution containing a metal ion should you transfer to your test tube? 2. What should you do to the nichrome wire prior to dipping it in a solution containing the metal ion? 3. Which metals will be tested? 4. Create a neatly drawn data table with the following 2 column headings: metal ion and flame color. 5. Create a second neatly drawn data table with the following 3 column headings: unknown, flame color, and metal ion(s) present. Purpose: 1. To observe the different colors emitted by ions in a flame test. 2. Use observations to identify the metal ion(s) in an unknown sample. Materials: Solutions containing the following metal ions: Ba 2+ Ca 2+ Li + K + Na + Sr 2+ Cu 2+ Test Tubes Test Tube Rack Matches/Lighter Nichrome Wire Bunsen Burner Nitric Acid **Caution! Handle very carefully! Procedure: 1. Obtain a sample of your assigned solutions from the stock solutions in the flasks up front. You only need enough of the solution to cover the loop at the tip of the nichrome wire when it is dipped into the solution. 2. Light the Bunsen burner. 3. Place tip of nichrome wire in nitric acid. 4. Place tip of nichrome wire in flame until there is a constant orange/yellow color. 5. Dip the tip of the nichrome wire into one of the solutions 6. Place the tip of the nichrome wire into the flame. 7. Observe the color of the flame as the salt/solution burns. 8. Repeat steps 1-5 for each salt/solution including the unknowns. Questions: (answer these on your sheet of paper with your prelab & data table) 1. Can a flame be used to identify a metal ion? Why, or why not? 2. What is happening to the electrons when we see color? 4

5 Heisenberg s Uncertainty Principle: It is impossible to determine both the position and momentum of an electron at the same time, but we can give a probable location of an electron around the nucleus called an orbital. Orbital = 3D region around the nucleus that indicates the probable location of an electron To further define the orbitals (locations of electrons), quantum numbers are used. Quantum Numbers Your home address is your most probable location, even though you are not there all the time. Your address is made up of four parts: state, city, street, and house number. The four quantum numbers (n, l, m, and s) can be thought of as the address for an electron in an atom. The principle quantum number, n, describes the energy level or average distance the electron is from the nucleus therefore signifying the size of the electron cloud. Although n can have numerical values from 1 to, seven is the highest energy level we will use. In other words, there can be seven states, if we use our address analogy. The second quantum number (angular momemtum), l, describes the sublevels within the energy level. These sublevels have different shapes and are represented by the letters s, p, d, f. The s sublevel is spherical in shape. The p sublevel resembles a dumbbell shape or sometimes called flower petal shape or peanut shape. The d shaped sublevel is more complicated and is described as double lobed and the f sublevel is even more complicated. There are pictures of each of these sublevels on the cover of this packet. Using our address analogy, a city is a sublevel of the state. Just as there are small states having fewer cities than large states, energy levels have different numbers of sublevels. The first energy level, n=1, is the closest to the nucleus and therefore it is the smallest electron cloud. It is so small that there is only one sublevel. The fourth energy level, n=4, is much farther away from the nucleus and therefore has more room for more sublevels. The fourth energy level has four different sublevels. The sublevels also correlate to a number as follows: s=0, p=1, d=2, f=3. These will be used when giving the four digit quantum number for a specific electron within an atom. The third quantum number, m, is the magnetic quantum number. Each sublevel contains one or more orbitals. The space occupied by one pair of electrons is called an orbital. In our address analogy, the orbitals are like the streets within a city. Just as streets are oriented north-south or east-west, orbitals are oriented in a magnetic field along the X, Y, and Z axes surrounding the nucleus of the atom. The s (spherical)sublevel contains only one orbital. The p sublevel contains three orbitals arranged along the X, Y, and Z axes. The d sublevel is made up of five orbitals, while the f sublevel contains seven orbitals. Using the address analogy, there would be seven streets in the f city. The numerical values for m range 5

6 from l to +l. Again, these will be used when giving the four digit quantum number for a specific electron within an atom. The fourth quantum number is the s or spin quantum number. It describes the clockwise or counterclockwise rotation of the electron. The numerical vales for s are +½ and -½. n l m s 1. Principle Quantum # 2. (n) Energy Level

7 Rules/Principles to Following When Assigning Electron Locations Aufbau Principle: An electron occupies the lowest energy level available which is given by the Aufbau principle filling order. Pauli Exclusion Principle: No 2 electrons in the same atom can have the same 4 quantum numbers. Hund s Rule: Orbitals of equal energy (and sublevel) are each occupied by one electron before any orbital is occupied by a second electron. All electrons in singularly occupied orbitals must have the same spin. Aufbau Principle Filling Order 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 4f 14 5s 2 5p 6 5d 10 5f 14 6s 2 6p 6 6d 10 6f 14 7s 2 7p 6 7d 10 7f 14 7

8 Aufbau Principle Filling Order Applied to the Periodic Table:

9 Orbital Notation Orbital notation uses arrows up/down (for electrons) on labeled lines (orbitals). Underneath the lines (orbitals) numbers are used to represent the energy level and letters are used to show the sublevel (s, p, d, f) for the electrons. An up versus down arrow indicates the spin of the electron. When 2 electrons occupy the same orbital, they must spin in opposite directions. The orbitals are filled by following the Aufbau principle filling order and Hund s rule with the exception of copper and chromium. Ex) orbital notation for cobalt _ _ _ 1s 2s 2p 3s 3p 4s 3d **The number of arrows should correspond to the number of electrons in the given atom of the element or electrons in the given ion of the element. Using the Aufbau filling order and Hund s rule, write the Orbital Notation for each of the following elements on a separate sheet of paper: 1. Carbon 5. Antimony 9. Molybdenum 2. Argon 6. Tungsten 10. Tin 3. Vanadium 7. Calcium 11. Copper** 4. Chromium** 8. Germanium 12. Ba 2+ Electron Configuration Electron configuration follows the same filling order as orbital notation but instead of using a line for each orbital and arrows for each electron, the number of electrons in the sublevel (s, p, d, f) is indicated by a superscript after the sublevel letter. Ex) electron configuration for cobalt 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 7 Electron Configuration practice: On a separate sheet of paper, write the full electron configuration for each of these elements. 1. Beryllium 5. Iodine 9. Uranium 2. Sulfur 6. Europium 10.Seaborgium 3. Scandium 7. Osmium 11. O 2-4. Copper 8. Actinium 12. Mg 2+ 9

10 Noble Gas Notation Noble gas notation is a shorthand version of electron configuration. A noble gas symbol in brackets is used to represent all the inner electrons for an atom of the given element and the configuration continues from that point on showing only the outermost electrons. Ex) Using the electron configuration for cobalt 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 7 These are the inner electrons and they would correspond to the electron configuration of argon(ar). The Nobel gas configuration for cobalt would be [Ar] 4s 2 3d 7 Noble Gas Configuration practice: On a separate sheet of paper, write the full Noble gas configuration for each of these elements. 1. Aluminum 3. Beryllium 5. Gold 2. Iron 4. Copper 6. Plutonium Mixed Notation Practice Complete the following on a separate sheet of paper: 1. Write the orbital notation for the following elements: a. Nitrogen b. Chlorine c. Gallium d. Cesium 2. Write the full electron configuration for the following elements: a. Titanium b. Xenon c. Polonium 3. Write the Noble gas notation for the following elements: a. Strontium b. Iodine 4. How many electrons would sulfur need to gain to achieve a Noble gas notation? 5. Would sodium tend to gain or lose electrons to achieve a Noble gas notation? How many? More Practice!!!! Complete the following on a separate sheet of paper: 1. Name each of the 4 quantum numbers along with its symbol and the information it gives for the electron. 2. Write the orbital notation for arsenic. 3. Find the 27 th electron (based on filling order, including Hund s rule) from #2 and give the four digit code for it. 4. Write the full electron configuration for Cesium. 5. Draw the electron that is represented by /2 10

11 Name 1. Write the orbital notation for: a. 14Si b. 29Cu 2. Write the electron configuration for: a. 16S b. 39Y 3. Write the Noble gas notation for: a. 22Ti b. 49In 4. Given the following orbital notation; predict the 4 digit quantum number that describes the circle electron: a. 4 digit quantum # 1s 2s 2p b. 1s 2s 2p 3s 3p 4s 3d 4digit quantum # 5. Given the 4 digit quantum number; draw /label the electron that it describes: a ½ b ½ 6. Would calcium tend to gain or lose electrons to achieve a Noble gas notation? How many? 11

12 Quantum Numbers Review: 1. What is the maximum number of electrons that can be in the a. second energy level b. third energy level c. fourth energy level 2. Which quantum number signifies the size of the electron cloud? 3. The sublevel or shape of the electron cloud is designated by which quantum number? 4. Which quantum number is used to represent the orbital? 5. The s quantum number is used to describe the clockwise and counterclockwise rotation of the electron in an orbital. What two numerical values can s have? 6. When n has the numerical value 4, what values can l have? 7. When l =3, what values can m have? 8. How many orbitals are contained in each p sublevel? 9. How many orbitals are contained in each d sublevel? 10. How many electrons can be in one orbital? 11. What is the maximum number of electrons that can be in a p sublevel? 12. What is the maximum number of electrons that can be in an f sublevel? 13. What is the maximum number of electrons that can be in a d sublevel? 14. What is the name of the scientist who stated that no two electrons in the same atom can have the same set of four quantum numbers? 15. What is the name of the scientist who pointed out that it is impossible to know both the exact position and momentum of an electron at the same time? 16. What is the name of the scientist who treated the electron mathematically as a wave? 12

13 17. What is the name of the scientist that stated all orbitals within a sublevel must contain one electron each, all with the same spin, before two electrons will occupy the same orbital within a sublevel? Orbital Notation and 4-Digit Quantum Number Practice Show the electron for the following 4 digit codes: 1) /2 2) /2 3) /2 4) /2 5) /2 6) /2 7) /2 8) /2 Give the orbital notation and the four digit code for the last electron for the following problems. 9) P 3- (#15) 10) Rb 1+ (#37) How many electrons does phosphorus have to gain/lose to be like the closest noble gas? How many electrons does rubidium have to gain/lose to be like the closest noble gas? 13

14 Periodic Table History ~ B.C.: elements:,,, By, elements were known. ~ : Wrote the first extensive list of elements. : (German Chemist) Noticed that,, and had similar properties and that Br s atomic mass was between that of Cl & I. He found three other groups with similarities and he called these groups. : (English Chemist) Arranged the now known 62 elements from to. He noticed that every element had chemical and physical properties; their properties were repeating. This became known as o Li Be B C N O F o Na Mg Al Si P S Cl : (Russian Chemist) Organized the elements by atomic also but also made it into form to help his students. Elements with properties were put into the same. Original Periodic Law : the properties of the elements are a function of their atomic masses. Considered the of the modern periodic table. He left where elements seemed to be. There were places where elements were put before elements because of their : o Te-I Co-Ni - (He felt the mistake was in measuring the mass and this would be corrected with further research.) 14

15 : (English Chemist) Arranged elements by atomic. Gave rise to the new Periodic Law, Properties of the elements are a periodic function of their : Glenn Discovered new elements. Only living person for whom an element was named. elements have been identified (naturally and artificially). o elements occur naturally on Earth. Examples: gold, aluminum, lead, oxygen, carbon o elements have been created by scientists. Examples: technetium, americium, seaborgium Periodic Table Terminology /Series = /Family = 15

16 Metals Located to the of the staircase line. Characteristics: o to form positive (+) ions (cations) o o o o o React with Most metallic element (most reactive metal) = Nonmetals Located to the of the staircase line. Characteristics: o Tend to to form negative (-) ions (anions) o o o Most nonmetallic element (most reactive nonmetal) = Metalloids (Semimetals) Located and of the staircase line, except,, and. Characteristics: o o o o 16

17 Element Directions: ( is our assigned element) Due Date: You need an unopened following: On one side:. Decorate it with the -assigned element symbol -element name - atomic # (above the symbol) -average atomic mass (below the symbol) Another side: -who discovered & when -where did the element name come from? -Family your element belongs to. -at least 3 uses for the element. Another side: -3 physical properties -3 chemical properties -orbital notation -electron configuration -Nuclear symbol for the most common isotope of the element If you can t fit all the information for a side on that side, it can be placed on another side. You will be presenting your creation to the class on the due date. 17

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19 Periodic Table Family Webquest Background: The periodic table is one of the most important instruments that a chemist can use. It is organized in a way so that elements with similar properties are grouped together. There are many patterns that you can find within the periodic table. You are going to use the Internet to discover what these patterns are. Directions: You can use any website to find information about families of the periodic table. Some helpful websites are and but others can be used. Part I: Families 1. What is the only element in a family by itself? The s block 2. Li, Na, K, Rb, Cs, Fr a. What is the name of this family? b. How reactive are they (High, Medium, Low)? c. What is the most reactive element in this family? d. How many valence electrons do they have? What is the oxidation # (charge) that they form? e. Are they metals, nonmetals, or metalloids? 3. Be, Mg, Ca, Sr, Ba, Ra a. What is the name of this family? b. How reactive are they (High, Medium, Low)? c. How many valence electrons do they have? What is the oxidation # (charge) that they form? d. Are they metals, nonmetals, or metalloids? The p block 4. F, Cl, Br, I, At a. What is the name of this family? b. How reactive are they (High, Medium, Low)? c. What is the most reactive element in this family? d. How many valence electrons do they have? What is the oxidation # (charge) that they form? e. Are they metals, nonmetals, or metalloids? 5. He, Ne, Ar, Kr, Xe, Rn a. What is the name of this family? b. How reactive are they (High, Medium, Low)? c. How many valence electrons do they have? What is the oxidation # (charge) that they form? d. Are they metals, nonmetals, or metalloids? 19

20 6. Al, Ga, In, Sn, Tl, Pb, Bi, Po a. What is the name of this family? 7. C, N, O, P, S, Se a. What is the name of this family? The d block a. What is the name of this family? a. How reactive are they (High, Medium, Low)? b. How many valence electrons do they have? What is the oxidation # (charge) that they form? c. Are they metals, nonmetals, or metalloids? The f block a. What are the two family names for the f block? Part II: General Questions about families 1. What family has all three states of matter at room temperature? 2. Which elements are: a. Gases at room temperature? b. Liquids at room temperature? 3. Which elements are radioactive and manmade? 4. Which family does not form compounds (does not bond with other elements)? Part III: Categories of the periodic table- Metals, Nonmetals, and Metalloids 1. Metals a. Where are they located? b. Give four main characteristics. 2. Nonmetals a. Where are they located? b. Give four main characteristics. 3. Metalloids a. Where are they located? b. Give four main characteristics. 20

21 Oxidation Numbers: The charge an atom acquires by gaining or losing electrons. Label your periodic table with oxidation numbers as instructed by your teacher. Isoelectronic Configurations Elements with similar electronic configurations tend to have similar chemical and physical properties. It is possible for elemental ions to have exactly the same electronic configuration as other elements or ions. When two elements and/or ions have the same electronic configuration it is said that they are "isoelectronic" with one another. When two chemical species are isolectronic they again tend to have similar chemical properties. Examples of Isoelectronic Elements and/or Ions Element or ion pair Electronic configuration Li +, He 1s 2 Be 2+, He 1s 2 F -, Ne 1s 2 2s 2 2p 6 S 2-, Ar 1s 2 2s 2 2s 6 3s 2 3p 6 Mg 2+, Na + 1s 2 2s 2 2p 6 Ca +, K [Ne]4s 1 P 3-, S 2-1s 2 2s 2 2p 6 3s 2 3p 6 Which of the atoms/ions are isoelectronic with each other? C, Cl -, Mn 2+, B -, Ar, Zn, Fe 3+, Ge 2+ 21

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23 Properties of Metals, Nonmetals, & Metalloids Define the following terms/physical properties & determine how you can test a sample for this property. 1) Malleability 2) Luster 3) Conductivity 4) For reactivity with acid, how will you determine if a chemical reaction has occurred? Purpose: To classify unknown samples as metal, nonmetal, or metalloid based on the observed characteristics. Use the procedures you described above to test each physical property of the samples. To test the chemical property (reaction w/ acid) place a small piece of the sample into a test tube and then add 5-10 drops of acid. Record your observations in the table on the back. 23

24 Properties of Metals, Nonmetals, & Metalloids Table Sample # Color Luster (lustrous) Malleability (malleable) Conductivity Reaction w/ acid Classification See your notes for the basic properties of metals, nonmetals, and metalloids. Then classify each of your samples as a metal, nonmetal, or metalloid. 24

25 Periodic Trends 25

26 "Periodic Table Properties (Trends) Atomic Radius Definition: Left to right across a period radius gets Why? 1. Electron electron Makes radius larger Very weak interactions 2. Proton electron Makes radius smaller Very strong interactions (proton to electron attraction is a million times stronger than electron to electron repulsion) Down a group/family radius gets Why? Energy level, n, from top to bottom Radius gets larger by adding electrons to higher (farther away) energy levels. Think Bohr model (adding rings) Ionization Energy Define: Which element has the highest value? Which element has the lowest value? Electronegativity Define: Which element has the highest value? Which element has the lowest value? 26

27 Ionic Radius:. Cations Anions The Anion radius is compared to the atom; Why? The Cation radius is compared to the atom; Why? When comparing ions of different elements, determine how many electrons each ion has and how many protons each ion has. If the electrons are the same, the one with more protons will have the smaller radius. Why? Which has a smaller atomic radius? K + or Cl - 27

28 Summary of Periodic Table Trends Top to Bottom Left to Right Atomic Radii Ionization Energy Electron Affinity Electronegativity Metallic Character 1. ionization energy (easy to remove electrons) 2. electronegativities (don t attract electrons) 3. Luster shiny 4. conductor of electricity and heat 5. and ductile Most metallic element = NonMetallic Character 1. ionization energy (hard to remove electrons) 2. electronegativities (attract electrons) 3. or no metallic luster 4. electrical and thermal conductors 5. solids Most nonmetallic element = Metalloid Character 1. Characteristics of both metals and nonmetals. 2. between those of metals and nonmetals. 3. energies between those of metals and nonmetals. Example: Silicon high luster, brittle, good conductor 28

29 Periodicity Review Worksheet 1. Which atom in the following pairs would have the larger atomic radii? P or Cl P or Ge Mg or Ca Na or Mg K or Mg Sn or Pb Sn or As Al or Cl As or Se Cl or Br Si or S Sr or Ra Si or O Ca or Al H or He 2. Which atom in the following pairs would have the greater first ionization energy? Sr or Ba Cs or Ba Cl or Ar In or Sn Xe or Kr Mg or Sr Ca or Cu 3. Which member in the following pairs would have the greater electronegativity? F or Cl F or O F or Any element S or Cl Ca or K Ca or F Al or Si S or O N or O Nonmetal or Metal C or S Mg or Na Na or K 4. What is the difference between a cation and an anion? 5. Identify three anions in Period 2 that can have the same number of electrons when they become ions. Do they gain or lose electrons? 6. Identify three cations in Period 3 that can have the same number of electrons when they become ions. Do they gain or lose electrons? 7. Identify the ions that are most likely to have an ion charge (oxidation state) of -2? (Hint: there are five of them) 8. Identify the ions that are most likely to have an ion charge (oxidation state) of +2? (Hint: there are six of them) 9. Why do elements in the same family generally have similar properties? 10. Which element has the lowest ionization energy? Why? 29

30 11. In a given period, are cations larger or are anions bigger? 12. Is the radius of a cation larger or smaller than the neutral atom? Why? 13. Is the radius of an anion larger or smaller than the neutral atom? Why? 14. Which member of the following pairs would have a larger radii? Br or Br - S or S 6+ O or O 2- P or P 3- Ca or Ca 2+ Al or Al 3+ Li or Li + N 3- or Al Which ion in the following pairs would have the larger radii? Li + or Be 2+ Mg 2+ or Be 2+ Cs + or Be 2+ Cu + or Cu 2+ Cr 3+ or Cr 6+ Na + or Al 3+ Zn 2+ or Fe 2+ N 3- or F - S 2- or O 2- I - or Cl - P 3- or S 2- Br - or S Consider all elements in period 3 for the following (Na through Ar) has the largest atomic radius has the greatest electron affinity has the highest first ionization energy is the most reactive metal is the most reactive non metal is the least reactive are metalloids are most likely to be cations are most likely to be anions loses 3 electrons to have noble gas configuration gains 3 electrons to have noble gas configuration 17. Consider all elements in group 16 (O through Po) is the least reactive is the most reactive has the greatest electron affinity has the greatest first ionization energy has the smallest atomic radius 30

31 Name Class Period Ion Practice 1. An isotope has 106 proton, 157 neutrons, and 106 electrons: a. Write the nuclear symbol for this isotope b. What is the name of this element? c. Is this an atom or an ion? d. What is the mass number of this isotope e. What is the atomic number? f. What is the net charge? 2. An isotope has 29 protons, 34 neutrons, and 28 electrons: a. Write the hyphen notation for this isotope b. What is the name of this element? c. Is this an atom or an ion? d. What is the mass number of this isotope e. What is the atomic number? f. What is the net charge? 3. a. The species 104 Rh 3+ has protons, neutrons and electrons b. The species 12 C has protons, neutrons and electrons c. The species 130 Te 2- has protons, neutrons and electrons 4. Which element will produce an ion with 15 protons, 16 neutrons and 18 electrons?. 5. Which element will produce an ion with 20 protons, 20 neutrons and 18 electrons?. 31

32 6. a. A Calcium atom will (lose or gain) electrons. How many? Is the calcium atom bigger or smaller than the calcium ion? b. A Francium atom will (lose or gain) electrons. How many? Is the francium atom bigger or smaller than the francium ion? c. A Fluorine atom will (lose or gain) electrons. How many? Is the fluorine atom bigger or smaller than the fluorine ion? d. A Oxygen atom will (lose or gain) electrons. How many? Is the oxygen atom bigger or smaller than the oxygen ion? e. A Carbon atom will (lose or gain) electrons. How many? Is the carbon atom bigger or smaller than the carbon ion? 7. Considering a 26 Mg atom and and a 26 Mg 2+ ion, label the following true or false: They both have the same number of protons They both have the same number of electrons They both have the same number of neutrons The magnesium ion has 14 electrons and the magnesium atom as 10 electrons The net charge on the magnesium ion is 2+ The Bohr model of the magnesium ion has 0 electrons in the outer most shell. The Mg 2+ ion is larger than the Mg atom. 32

33 Martian PT Name Honors Chemistry Worksheet "Trends" "Periodic Chart For Mars And Its 33 Known Elements" Place the following elements in their proper place in the Martian periodic table. Remember, natural laws are the same for the whole universe. Note that Mars has no transition metals. a, b, c, d, e, f, g, h, i, j, k, l, m, n, o, p, q, r, s, t, u, v, w, x, y, z,!, #, $, %, +, =,? undiscovered 1. The most reactive metallic element is x. 2. The most reactive nonmetallic element is!. 3. Inert gases (like our noble gases) are $, %, a, and d. a has the highest ionization energy of this group and the least dense, $ has the lowest ionization energy for this group, and d has a smaller atomic radius than %.. 4. The lightest element of all is #. 5. All the following elements are in the 3rd energy level and have n is the largest atomic radius g is a metalloid = is the most reactive nonmetal? is a pretty reactive metal + is next to k but k has the higher ionization energy e is in this period as well 6. Element g has 14 protons. 7. f has a total of 7electrons. 8. c has an atomic mass of 5 and its final electron is in the 2 nd energy level. 9. r would correspond to our alkali metals and is in the 4th energy level. 10. The! family is made up of the elements!, =, s and p in order of increasing atomic radii. 11. j is the most dense of all Martian atoms and is radioactive and its electron configuration would end with 5p

34 12. q is in period 5 and has an oxidation number of m is in period 2 and will form a compound with! that has the formula m! 2. In other words, m has two electrons that it would like to give to!. 14. h is like our element carbon and is in same family as t, w and g. 15. t is bigger than w in atomic radius. 16. The Martian solvent is like our most important liquid and has the formula # 2y. 17. o is in the same family as e and has a lower electronegativity than e. 18. i is the only metalloid in the family of y. 19. l has a dot notation of 3 dots and an oxidation number of z is in the same family as q and has a slightly higher ionization energy than q. 21. b is slightly smaller that q. 22. u has a final electron configuration of 4p v is next to s and has a larger atomic radius than s.. Extension Problems: A. If you discovered the Martian element that is listed as undiscovered at this time on the Martian periodic table, what would you name it? B. What chemical symbol would you give this element and why? C. Name at least three characteristics that this element would have because of its location on the Martian periodic table? 34

35 Periodic Table Trends Review Worksheet For each of the following, circle the correct element 1. Li Si S Metal 2. N P As Smallest Ionization Energy 3. K Ca Sc Largest Atomic Mass 4. S Cl Ar Member of the Halogen Family 5. Al Si P Greatest Metallic Characteristics 6. Ga Al B Largest Atomic Radius 7. V Nb Ta Largest Atomic Number 8. Te I Xe Member of Noble Gases 9. Si Ge Sn Has electrons in 4 energy levels 10. Li Be B Member of Alkali Metals 11. As Se Br Largest Electronegativity 12. H Li Na Nonmetal 13. Hg Tl Pb Member of Transition Metals 14. Na Mg Al Electron distribution ending in s 2 p Sb Bi Pb Metalloid 16. B C N Greatest Nonmetal characteristics 17. Ca Sc Ti Electron distribution ending in s 2 d Be K Ga Member of the Alkaline Earth Metals 19. Si Al P Semiconductor 20. F - Cl - I - Smallest Radius 35

36 Honors Chemistry Review Unit 4 1. What contribution did the following people make to the creation of the periodic table: a. Newlands: b. Mendeleev: c. Moseley: d. Dobereiner: 2. What does the new periodic law state? 3. In the periodic table, the vertical columns are called and the horizontal rows are called. 4. Elements in the same group have the same number of. 5. What family of the periodic table is the most stable because their valence shell is full of electrons? 6. Give three properties of metals: 7. Give three properties of nonmetals: 8. What is ionization energy? 9. Why does the ionization energy decrease as you move down a group? 10. Why does the atomic radius increase as you move down a group? 11. As you move across a period, the decreases and the,, and increases. 12. As you move down a group, the atomic radius and the ionization energy, electron affinity, and electronegativity. For questions #13-36, use the following elements to answer the questions. Elements may be used more than once and for some questions, there will be more than one answer. 36

37 Na Si Br Ag Ne Sb H Ga S Ca Es Pa Cl Ba Li He Xe Fr F Al 13. The most reactive metal: 14. The most reactive nonmetal: 15. A metal: 16. A nonmetal: 17. A metalloid: 18. A halogen: 19. An alkali metal: 20. An alkaline earth metal: 21. A noble gas: 22. A transition metal: 23. A post transition metal: 24. An inner transition metal: 25. Other nonmetal: 26. In a family all by itself: 27. Radioactive: 28. Made in a laboratory: 29. A gas: 30. Found in a free-state (not part of a compound): 31. Has multiple oxidation states: 32. The only noble gas that does not contain eight valence electrons: 33. Has the same number of valence electrons as Nitrogen: 34. Has the same oxidation state as Potassium: 35. Has an oxidation state of -1: 36. Has 4 valence electrons: 37

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39 Trends in the Periodic Table 1. On the top graph, plot ionization energy (y-axis) vs, atomic number (x-axis). On the bottom graph, plot a separate graph of atomic radius (y-axis) vs, atomic number (x-axis). For each graph, connect successive dots with straight lines, label the axes, and include a title. 2. In order to make the graph easier to analyze, you should mark each group of elements with the same color circle/dot.! a. Group 1 Alkali Metals = Red (H, Li, Na) b. Group 2 Alkaline Earth Metals = Blue (Be, Mg) c. Group 3 Boron Family = Green (B, Al) d. Group 5 Nitrogen Family = Orange (N, P) e. Group 7 Halogens = Yellow (F, Cl) f. Group 8 Noble Gases = Purple (He, Ne, Ar) 3. Examine your graph of atomic radius vs. atomic number. a. Which elements are found at the main peaks on your graph? i. What do these elements have in common? b. Which elements are found at the main valleys on your graph? i. What do these elements have in common? 4. Generally, as you go from left to right across a period, what happens to atomic radius? Explain why. 5. Generally, as you go down a group in the periodic table, what happens to atomic radius? Explain why. 6. Examine your graph of ionization energy (IE) vs. atomic number. a. Which elements are found at the main peaks on your graph? i. What do these elements have in common? b. Which elements are found at the main valleys on your graph? i. What do these elements have in common? 7. Generally, as you go from left to right across a period, what happens to IE? Explain why. 39

40 8. Generally, as you go down a group in the periodic table, what happens to IE? Explain why. 9. Why do you think there is an inverse relationship between atomic radius and ionization energy? 40

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