Electron Configuration & Periodicity Unit 3

Transcription

1 Name: Electron Configuration & Periodicity Unit 3 (seven class periods) Unit 3.1: First Ionization Energy & Photoelectron Spectroscopy 1) Coulombs Law a) The force of attraction between two charged objects is proportional to the product of their charges divided by the distance between them. A positive value is demonstrative of a repulsive force, while a negative value results in attraction F a is the force of attraction q 1 is the charge of one object q 1 is the charge of one object d is the distance between them b) First ionization energy i) The energy needed to remove 1 electron from an atom in its gaseous state ii) The neutral atom s electron requires energy to remove, therefor we discuss this energy as a reactant that must be added and is symbolized as positive. The ion & electron still hold this energy therefore as a product it could be given off and is symbolized as negative iii) As nuclear charge (Z) increases the attraction (F a) of electrons for the nucleus increases thus drawing the electrons closer to the nucleus. This decreases the distance between the nucleus and the electrons thus more energy is required to remove an electron iv) However as electrons are drawn to the nucleus they repel each other and develop shells based on this repulsion. This increases the distance between the nucleus and the electrons thus decreasing ionization energy (1) The first shell, the one closest to the nucleus, contains only 2 electrons. After that repulsion is too great for any additional electrons to be that close, so a second shell develops further from the nucleus (2) The second shell can contain the repulsion of up to 8 electrons before the repulsion is too great and another shell must develop. (3) If you look at the periodic table this information is easily visible. In the first row there are two elements, these two elements only have 1 & 2 electrons respectively. In the second row there are 8 elements which correspond to elements with 3 to 10 electrons, each of which would have 2 electrons in their first shell, and up to 8 in their second. v) In summary: Ionization energy increases as your move across the period due to increased nuclear charge (Z), and decreases as you move down a family due to increased number of shells (also known as principle energy levels)

2 c) Bohr Models i) A graphical representation of the layout of electrons around the nucleus of an atom based on the location of electrons at various energy levels ii) Steps Chlorine-52 (1) In the center of your model symbolize the nucleus by writing the number of protons and neutrons in the atom (2) in concentric circles around your nucleus place your electrons. Remember: (a) the row number allows you to predict the number of energy levels (circles) needed (b) the number of electrons in that energy level corresponds to the number of element s in that row, INCLUDING the element which you are drawing d) Photoelectron Spectroscopy i) High intensity lazars provide MASSIVE amounts of energy to a small sample. A photoelectric plate registers electrons that the sample emits and records the energy that the sample absorbed at that point (1) Peaks are representative of the number of electrons released at that energy. The larger the peak the more electrons that were removed ii) Electrons in the same shell require the same amount of energy to escape = 8 electrons (represent the second shell broken into two subshells)

3 Unit 3.2: Electron Configuration 1) Sublevels are named based on their probability distributions a) s sublevel/orbital i) can hold up to 2 electrons ii) spherical iii) elements in the first two families (alkaline metals & alkaline earth metals) have their highest energy level electrons in this orbital b) p sublevel/orbital i) can hold up to 6 electrons ii) named partial, and have a dumbbell shape iii) elements in groups 13 to 18 have their highest energy level electrons in this orbital c) d sublevel/orbital i) can hold up to 10 electrons ii) named dispersed and have several complex shapes iii) elements in groups 3 to 12 (transition metals) have their highest energy level electrons in this orbital iv) Penetrate 1 energy level closer to the nucleus than s or p sublevels. Therefore start at 3d NOT 4d d) f sublevel/orbital i) can hold up to 14 electrons ii) named fractal and have several complex shapes iii) elements in in the lanthanide and actinide series have their highest energy level electrons in this orbital iv) Penetrate 1 energy level closer to the nucleus than d sublevels and 2 energy levels more than s or p sublevels. Therefore start at 4f NOT 5 (based on the d sublevel) or 6 f (based on the s sublevel) e) Fill based on the Aufbau Principle, which provides which orbitals fill when.

4 Example 1 Write the complete electron configuration for each element a.) b.) Cl I 2) Noble Gas Configuration a) A short hand for the electron configuration b) Based off the core electrons of the atom i) Electrons found on the interior of the outer shell (valance shell) ii) These electrons are nearer the nucleus and thus VERY difficult to remove, due to their strong attraction and limited repulsion iii) Occurs when an energy level fills c) Find the PREVIOUS noble gas and place its symbol in brackets. Then build the valence shell from that element Example 2 Write the noble gas configuration for each element a.) b.) Cl I 3) Ions a) The universe tends towards the lowest energy possible (water travels down hill, unstable nuclei undergo decay, students would rather finger paint than learn chemistry etc.) b) Noble gasses are atoms with very little energy i) the amounts of electron repulsion is overcome by a strong effective nuclear charge resulting in atoms that are extremely inert. c) Atoms will gain, or lose, electrons to reach this lower energy state d) Metals will lose electrons, making them cations (positively charged atoms). e) Non-metals (generally) will gain electrons, making them anions (negatively charged atoms) f) Electrons are lost from their valance shells!

5 4) Orbital Notation a) Pictorial representation of energy levels and the electrons found with in b) 2 electrons can share an energy level provide they are spinning in opposite directions c) Two electrons cannot spin in the same direction in the same orbital (Pauli s exclusion principle) d) Due to sharing each subshell is broken into orbitals i) s orbital 2 electrons therefore one orbital ii) p orbital 6 electrons, therefore 3 orbitals (p x, p y, & p z) iii) d orbital 10 electrons, therefore 5 orbitals iv) f orbital 14 electrons therefore 7 orbitals e) Hund s Rule -- Electrons only pair when there is no open space in an orbital Example 3 Draw the orbital diagram for each element a.) Cl b.) I

6 5) Magnetism a) Diamagnetic i) magnetism dies ii) Elements with few/no unfilled subshells, therefore not magnetic iii) Elements at the ends of blocks b) Paramagnetic i) partially magnetic ii) Elements with several unfilled subshell, therefore able to be influenced by a magnetic field iii) Elements in the middle of blocks c) Ferromagnetic i) Elements with LOTS of unfilled subshells ii) Elements that are extremely magnetic iii) Found in the middle of the d-block and f-block Unit 3.3: Periodic Trends Increases Increases Increases 1) YOU MUST KNOW 4 TRENDS! A trend is NOT an explanation, in order to get FULL credit for the test you must also be able to explain a trend a) Explaining a trend from left to right, remember nuclear charge i) As atomic number increases there is an increase in the number of protons in the nucleus, thus increasing nuclear charge (Z) ii) Thus electrons experience a greater attration for the nucleus and are concequently pulled closer to it

7 b) Explaining a trend from top to bottom, remember principle energy level (number of shells) i) As atoms grain electrons those electrons repel one another and force additional elections into higher energy levels ii) At higher energy levels electrons are further from the nucleus and thus experience less attraction and further away 2) Atomic Radius a) The size of an atom, from the nucleus to where there is less than a 95% chance of finding an electron b) Decreases as you move from right to left (across the period) i) Increased nuclear charge (Z) causes electrons to experience a greater attraction for the nucleus, thus are pull in closer, shrinking the atom c) Increases as you move from top to bottom (down a family) i) Increased number of electrons cause electrons to repel each other thus producing more shells around the nucleus, these shells which are further from the nucleus contain electrons that have a lower attraction for the nucleus and are thus further away 3) First Ionization Energy a) The amount of energy needed to remove an electron from an atom in its gaseous state b) increases as you move from right to left (across the period) i) Increased nuclear charge (Z) causes electrons to experience a greater attraction for the nucleus, thus are pull in closer, shrinking the atom. Electrons nearer the nucleus have less initial energy and thus require more to be removed c) decreases as you move from top to bottom (down a family) i) Increased number of electrons cause electrons to repel each other thus producing more shells around the nucleus, these shells which are further from the nucleus contain electrons that have a lower attraction for the nucleus and are thus further away. Electrons further from the nucleus have more initial energy and thus require less energy to remove. 4) Electronegativity a) How attractive electrons from neighboring atoms are for the nucleus of an atom b) A scaled, developed by Linus Pauling, of values form 0 (noble gases) to 4 (fluorine) assigned to each element based on this attraction. c) increases as you move from right to left (across the period) i) Increased nuclear charge (Z) causes electrons to experience a greater attraction for the nucleus, thus are pull in closer, shrinking the atom. Since the atom is smaller, neighboring electrons are able to get closer to the nucleus thus causing a higher electronegativity value

8 d) decreases as you move from top to bottom (down a family) i) Increased number of electrons cause electrons to repel each other thus producing more shells around the nucleus, these shells which are further from the nucleus contain electrons that have a lower attraction for the nucleus and are thus further away. Since the atom is larger the nucleus is shielded from neighboring electrons by its many layer of electrons which repel those electrons 5) Ionic Radius a) Atoms will gain or lose electrons to reach a lower energy state i) Metals will lose electrons, forming positive ions (cations) ii) Nonmetals will gain electrons, forming negative ions (anions) b) When an atom gains electrons those electrons are added to its valence shell thus expanding the atoms radius due to repulsive forces c) When an atom loses electrons those electrons are removed from the atoms valence shell thus reducing repulsive forces and reducing the atoms radius.

9 Developing the Concept of Shells, Subshells, Electron Configurations, and More 3. Consider PART I: Discovering how electrons are ÔarrangedÕ in an atom 1. Describe the nature of the interaction between protons and electrons in an atom? Consider using some or all of the following terms in your description: attraction, repulsion, neutral, positive, negative, charge, distance, nucleus, force, energy, CoulombÕs Law. 2. For each situation below, compare the relative energy necessary to separate positive and negative electrical charges. Compare A to B How many electrons do you see in the picture? How many protons? Which of these electrons is the easiest (requires the least amount of energy) to remove (ionize)? Justify your answer. Compare the energy required to remove the electron from 3 with the energy in 2a 2c Compare A to C The first ionization energy is defined as the minimum energy that must be added to a neutral atom, in the gas phase, to remove an electron from that atom. This definition can be represented by the following chemical equation: energy + A(g)! A + (g) + e Ð 4. In the ionization equation above identify which species is at lower energy, A(g) or A + (g) + e Ð? Justify your answer. 5. Explain why energy is required (an endothermic process) to remove the electron in a neutral atom. 6. The value of the first ionization energy for hydrogen is 1312 kj mol -1. energy + H(g)! H + (g) + e Ð! On the graph on the next page use a short horizontal line to indicate the energy of H(g) and a short horizontal line to indicate the energy of H + (g) + e Ð. Be sure to consider your responses to Q4 and Q5 above.

10 H(g) H + (g) + e Ð In the energy diagram below locate (draw a horizontal line) the first ionization energy for hydrogen and the first ionization energy for helium. 7. What does the difference in energy in the lines in your diagram above represent? The values for the first ionization energy for a hydrogen and helium atom are provided in the table below.! Atom 1H 2He 3Li Ionization Energy (kj mol Ð1 ) ! 8. Based on comparisons you made in Question 2 how would you explain the difference in the values for the first ionization energy for hydrogen and helium? 10. How does the diagram illustrate the relative ease with which an electron can be removed from each atom? 11. Predict a value for the first ionization energy for lithium. Do not add your prediction to the figure just yet. Justify your prediction (look back at Question 2 if you need guidance). 9. How does your explanation account for the relative charge on hydrogen and helium and the distance of the electron(s) from the nucleus?

11 The actual value of the first ionization energy of lithium is 520 kj mol -1. Add this value for to the figure on the previous page. 12. How would you explain the ionization energy for lithium compared to the ionization energy for helium? Compared to hydrogen? 13. Predict the relative value of the energy necessary to remove a second electron (called the second ionization energy) from lithium. Support your prediction with an explanation. 14. Based on the first ionization energies for hydrogen, helium and lithium that you represented in the figure on the previous page, what can you infer about the distance of the electrons from their respective nuclei.!!!! The first ionization energies for selected elements from the second period of the periodic table are provided in the table below. Atom 3Li 4Be 6C 7N 9F 10Ne Ionization Energy (kj mol Ð1 ) Explain the trend in ionization energies in terms of the charge of the nucleus and the relative location of the electrons. The first ionization energy for the element sodium is given in the following table. Below is a table containing the electron energies for each of the 18 electrons in an argon atom. The graph of this data is shown. Electron Electron Energy Removed (kj mol Ð1 ) 1!1521 2!2666! 3!3931! 4!5751! 5!7238! 6!8781! 7!11995! 8!13842! 9!40760! 10!46186! 11!52002! 12!59653! 13!66198! 14!72918! 15!82472! 16!88576! 17!397604! 18!427065! 17. Make observations about the graph in terms of the relative energies of the electrons and their relationship to each other. Atom 11Na 12Mg 14Si 15P 17Cl 18Ar Ionization Energy (kj mol Ð1 ) Predict the values for the first ionization energy for the other selected third period elements. Explain how you arrived at your predictions. 18. Based on your responses from the previous questions how many ÔgroupsÕ (levels or shells) of electrons are shown for Argon? 19. Indicate the number of electrons in each group/level that you identified?

12 20. On the graph below draw a horizontal line (to the right of the y-axis) that represents an average energy level for each of the groups of electrons that you identified. Label the levels 1, 2, etc.é beginning from the lowest energy level. What do these lines represent? Describe the electron structure (location of the electron) of the atom. Consider using some or all of the following terms in your description; nucleus, electron, energy, distance, level, proton, shell, arrangement, attraction, repulsion, positive, negative, charge, location. PART II: Do all electrons in the same level have the same energy? 21. How would you describe the relative energy separation of these energy levels? 22. An electron from which level requires the least amount of energy to remove? The largest amount of energy to remove? One important conclusion based on the first ionization energy experimental data is that electrons in higher shells require less energy to remove. We have examined experimental data that relates the energy required to remove an electron to the shell the electron occupies. In which shell does an electron require more energy to remove, an electron in the second shell or the fourth shell? An interesting question that cannot be answered from the experimental data of the first ionization energy isé Do all electrons in the same shell require the same amount of energy to remove? We CAN answer this question if we look at photoelectron spectroscopy (PES) data for the atoms. In a photoelectron spectroscopy experiment any electron can be ionized when the atom is excited. Like with the first ionization, only one electron is removed from the atom. However in a PES experiment it can be ANY electron, not just the electron that requires the least amount of energy to remove.

13 Examine the PES spectrum for hydrogen shown in the figure. The label on the y-axis is energy and the units are in megajoules(m J mol Ð1 ) 2. What is the relationship between the photoelectron spectrum and the first ionization energy for hydrogen? Helium is next, but before looking at its photoelectron spectrum answer the following questions: 3. How many electrons does helium have in its first shell? 4. Refer back to Part I of this activity, and obtain the first ionization energy for a helium atom. Can you predict what the PES would look like if a. the same amount of energy is required to remove each of the electrons? b. different amounts of energy are required to remove each electron? Go to back to the previous figure and sketch both scenarios. 1. What does the x-axis depict? Explain.

14 Examine the PES for helium and compare it to your prediction from the previous question. Look at this PES and compare it to the prediction you made in the previous question. 5. Explain the relative energy of the peak(s) and the number of electrons represented by each peak in the PES for helium and for hydrogen. 8. For each peak in the PES of lithium, identify the shell the electrons represented by that peak occupy. Be sure to comment about the relative energy of the peak(s) and the number of electrons for each peak for Li.) 6. For lithium a. How many electrons does lithium have? b. What shells (levels) do those electrons occupy? 7. Predict what you expect the PES for lithium to look like. (Note: you do not have to predict the exact energies of each electron, you can make a reasonable estimate based on the first ionization energies for lithium and helium - refer back to Part I of this activity. The next element in the Periodic Table is beryllium. 9. How many electrons does beryllium have and what shells do those electrons occupy? 10. For the PES for beryllium predict a. how many peaks b. the number of electron for each peak c. estimate the relative energies.

15 Below is the PES for boron. 13. Briefly describe how to interpret the PES for boron. The next element in the Periodic Table is boron. 11. How many electrons does boron have and what shells do those electrons occupy? 12. For the PES for boron predict a. how many peaks b. the number of electron(s) for each peak c. estimate the relative energies 14. Predict what changes in the PES you would expect to see going across period 2 of the periodic table, from carbon to neon? Look at the PES for these second period elements.

16 Below is the PES for the period 2 elements from boron to neon. 15. Answer the following questions after looking at the PES for hydrogen through neon. a. Would you agree or disagree with the following statement? Explain your answer. ÔThe electrons in the second shell all have the same energy.õ b. How many ÔsubshellsÕ are found in the second shell? c. How many ÔsubshellsÕ are found in the first shell? d. How many electrons are in each subshell in the second shell? In the first shell? e. Moving systematically from lithium to neon; i. How many electrons are in the first shell? ii. What happens to the energy required to remove an electron in the first shell moving from left to right in the second period? Support your observation with an explanation. iii. What happens to the energy of the electrons in the outer most shell? 16. Look at the PES for the elements in the third period (sodium Ð argon) and describe your observations. Any surprises? Explain. A notation has been agreed upon for writing an electron configuration to identify the location of the shell and subshell of each electron in an atom. Shells are labeled with a number; 1, 2, 3, etc. and subshell are labeled with letters; s, p, d, and f. Every shell contains an s subshell. 17. Write the complete electron configuration for the first ten elements in the periodic table? Look at the PES for potassium, calcium and scandium. 18. Explain what happens in the PES for scandium that has not occurred in any element prior. 19. If one electron is removed from scandium, which electron (identify the shell and subshell) requires the least amount of energy to remove?

17 Name: Date: Period: Name: Date: Period: Bohr Model Worksheet 3. Neon is often found in lasers. Draw the Bohr model for neon. Directions Draw the Bohr Models showing all the electrons in each energy level. 1. Magnesium compounds are used in the production of uranium for nuclear reactors. Draw the Bohr model for magnesium. Niels Bohr 4. Argon gas can be found in Geiger counters which are used to detect radiation. Draw the Bohr model for argon. 2. Sodium is found in salts that can be used to seed clouds to increase rainfall. Draw the Bohr model for sodium. Page 1 of High School Technology Initiative (HSTI) Educational Materials: The ATOM: Structure Page 2 of High School Technology Initiative (HSTI) Educational Materials: The ATOM: Structure

18 Name: Date: Period: Name: Date: Period: 5. Aluminum alloys are used in airplane construction due to their low density. Draw the Bohr model for aluminum. 7. Lithium can be found in Mount PalomarÕs 200-inch telescopic mirror. Draw the Bohr model for lithium. 6. Oxygen is often added to rocket fuel as an oxidizer. Draw the Bohr model for oxygen. 8. Sulfur dioxide is often used at water treatment facilities to dechlorinate water. Draw the Bohr model for sulfur. Page 3 of High School Technology Initiative (HSTI) Educational Materials: The ATOM: Structure Page 4 of High School Technology Initiative (HSTI) Educational Materials: The ATOM: Structure

19 Electron Configurations Worksheet Write the complete ground state electron configurations and orbital notations for the following: # of e Element (atom) e - configuration Orbital Notations/ diagrams 1) lithium 2) oxygen 3) calcium 4) nitrogen 5) potassium 6) chlorine 7) hydrogen 8) copper 9) neon 10) phosphorous Write the abbreviated ground state electron configurations, noble gas configuration, for the following: # of electrons Element Electron Configuration 11) helium 12) nitrogen 13) chlorine 14) iron 15) zinc 16) barium 17) bromine 18) magnesium 19) fluorine 20) aluminum Page 1 of 8

20 Electron Configuration Elements (atoms) and Ions Write the electron configuration and orbital notations for the following Atoms and ions: Element / Ions F Atomic number # of e - Electron Configuration F 1- O O -2 Na Na 1+ Ca Ca +2 Page 2 of 8

21 Al 3+ Al N N 3- S 2- Cl 1- K 1+ S Br 1- Mg 2+ Page 3 of 8

22 Electron Configuration Practice Directions: Write and draw the electron configurations of each of the following atoms. Example: Co : 27 e - 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 7 Co 1s 2s 2p 2p 2p 3s 3p 3p 3p 3d 3d 3d 3d 3d 4s 1. Scandium: 2. Gallium: 3. Silver: 4. Argon: 5. Nitrogen: 6. Lithium: 7. Sulfur: Page 4 of 8

23 Electron Position and Configuration Position: Draw the Electron Position of each of the following atoms. Example: He: 1. Li 3. O 2. C 4. Ar Directions: Draw the electron configurations of each of the following atoms. Example: F 1s 2s 2p 2p 2p 1. Chlorine: 5. Sodium: 2. Nitrogen: 6. Potassium: 3. Aluminum: 7. Sulfur: 4. Oxygen: 8. Calcium Page 5 of 8

24 Electron Configuration Practice In the space below, write the expanded electron configurations (ex. = 1s 2 2s 1 ) of the following elements: 1) Sodium 2) potassium 3) chlorine 4) bromine 5) oxygen In the space below, write the abbreviated electron configurations (ex. Li= [He]2s 1 ) of the following elements: 6) manganese 7) silver 8) nitrogen 9) sulfur 10) argon In the space below, write the orbital notation (arrows) of the following elements: 11) manganese 12) silver 13) nitrogen 14) sulfur 15) argon Determine what elements are denoted by the following electron configurations: 16) 1s 2 2s 2 2p 6 3s 2 3p 4 17) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 1 18) [Kr] 5s 2 4d 10 5p 3 19) [Xe] 6s 2 4f 14 5d 6 20) [Rn] 7s 2 5f 11 Determine which of the following electron configurations are not valid: 21) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 4d 10 4p 5 22) 1s 2 2s 2 2p 6 3s 3 3d 5 23) [Ra] 7s 2 5f 8 24) [Kr] 5s 2 4d 10 5p 5 25) [Xe] Page 6 of 8

25 Electrons, Valence, and Lewis Dot Structures Chem 544/545 Dr. Brielmann Name Period 1. How many electrons are present in: Helium (He) Carbon (C) Neon (Ne) Sodium (Na) Zinc (Zn) 2. How many valence electrons are present in: Helium (He) Carbon (C) Neon (Ne) Sodium (Na) Potassium (K) Fluorine (F) Chlorine Bromine 3. Draw Lewis Dot Structures for the following elements: Helium (He) Carbon (C) Neon (Ne) Sodium (Na) Ne 4. Correct the following Lewis Dot Structures: Oxygen Nitrogen Beryllium Fluorine O N Be F 5. Fill in the following table: number of electrons: Carbon Carbon anion Carbon cation C - + C number of valence electrons Lewis structure Page 7 of 8

26 S-C-5-3_Periodic Trends Worksheet and KEY 10. For each of the following, circle or highlight the correct element that best matches the statement on the right. Li Si S metal N P As smallest ionization energy K Ca Sc largest atomic mass S Cl Ar member of the halogen family Al Si P greatest electron affinity Ga Al Si largest atomic radius V Nb Ta largest atomic number Te I Xe member of noble gases Si Ge Sn 4 energy levels Li Be B member of alkali metals As Se Br 6 valence electrons H Li Na nonmetal Hg Tl Pb member of transition metals Na Mg Al electron distribution ending in s 2 p 1 Pb Bi Po metalloid B C N gas at room temperature Ca Sc Ti electron distribution ending in s 2 d 2 Source:

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