2. Why do all elements want to obtain a noble gas electron configuration?

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1 AP Chemistry Ms. Ye Name Date Block Do Now: 1. Complete the table based on the example given Location Element Electron Configuration Metal, Nonmetal or Semi-metal Metalloid)? Group 1, Period 1 Group 11, period 4 Group 14, Period 3 Group 17, Period 4 Hydrogen (H) Group/Family Name 1s 1 Nonmetal (none) 2. Why do all elements want to obtain a noble gas electron configuration? 3. Which of the following elements has the most similar properties to Ca? (elements in the same have the most similar properties because they have the same ) a. K b. Sc c. Sr d. Ar Review-Properties of Elements & Periodic Trends: Effective Nuclear Charge-a measure of the positive attractive force of the nucleus towards negatively charged electrons due to the number of ; how much attractive force an electron feels can be affected by the number of shielding electrons (defined below) Electron Shielding Effect-electrons in the energy levels closer to the nucleus protects the electrons in the outer shells and lessens the effect of the positive, attractive force of the nucleus.

2 1. Atomic Radius: size of an atom When looking at elements going down a GROUP, atomic radius increases o As you go down a group, more are being added, therefore increasing the size. When looking at elements going across a PERIOD, atomic radius decreases o As you go across a period, the therefore the nucleus the electrons of the atom, and the radius decreases Examples: For each pair of elements below, circle the one with the larger atomic radius. a. Na and Cl c. C and B e. K and Se g. Br and Ca b. Mg and Sr d. Ar and Ne f. Sb and B h. Ge and C **Use nuclear charge to explain/show why Be has a smaller atomic radius than Li. Include a Bohr diagram for both Li and Be. 2. Electronegativity: tendency to attract or gain an electron When looking at elements going across a PERIOD, electronegativity increases o As you go across a period, the therefore the nucleus (gain) o Elements towards the right side of the periodic table are closer to becoming a noble gas they want to gain electrons o Exception: When looking at elements going down a GROUP, electronegativity decreases o As you go down a group, the atomic radius increases. The, therefore and the attraction for electrons Examples: For each pair of elements below, circle the one with the greater Electronegativity. a. Na and Cl c. C and B e. K and Se g. Br and Ca b. Mg and Sr d. Ar and Ne f. Sb and B h. Ge and C

3 3. Ionization Energy: energy required to remove an electron When looking at elements going down a GROUP, ionization energy decreases o As you go down a group, the atomic radius. With more shells being added, the, therefore and attractive force on the electrons, making them easier to remove. When looking at elements going across a PERIOD, ionization energy increases o As you go across a period, the therefore the making it harder to remove an electron o Elements towards the right side of the periodic table don t want to lose electrons (they want to gain electrons) to become like a noble gas. Therefore, it is difficult (requires more energy) to remove an electron Examples: For each pair of elements below, circle the one with the greater Ionization Energy. a. Na and Cl c. C and B e. K and Se g. Br and Ca b. Mg and Sr d. Ar and Ne f. Sb and B h. Ge and C **Use electron shielding to explain why Mg has a lower ionization energy and a lower electronegativity than Be. Include a Bohr diagram for both Mg and Be. More about Ionization Energy. First Ionization Energy energy required to remove 1 electron Second Ionization Energy energy required to remove 2 nd electron Third Ionization Energy energy required to remove 3 rd electron In general 1 st I.E. 2 nd I.E. 3 rd I.E. Because as the # of electrons decreases, the nucleus (# protons doesn t change) has a stronger pull on the electrons that are remaining. Based on the relative jump between ionization energies, you can tell how many the element has

4 1. a. Between which 2 ionization energies do you see the biggest jump? b. How many valence electrons would this element have? c. What group would this element be found in? 2. What group would this element be found in? Ions and Ionic Radius When an atom loses electrons and becomes a cation, its radius becomes than that of the neutral atom o # protons # electrons, therefore increasing the effective nuclear charge, meaning that there is a stronger pull of the electrons towards the nucleus. When an atom gains electrons and becomes an anion, its radius becomes than that of a neutral atom o When electrons get added to the same energy level, they repel each other *Note: the term isoelectronic refers to

5 Properties of Metals vs. Nonmetals Metals Malleable (can be hammered/molded into sheets) Ductile (can be drawn/pulled into a wire) Have luster (are shiny when polished) Good conductors (allow heat & electricity to flow throw them) Nonmetals Not malleable or ductile; instead, they are brittle (shatter easily) Lack luster; instead, they are dull They are either poor or nonconductors Reactivity of Metals vs. Nonmetals Reactivity of a metal is related to its o The, the the metal o Trend (within the metals on the periodic table): Going down a group: Going across a period: o Most reactive metal: Reactivity of a nonmetal is related to its o The, the the nonmetal o Trend (within the nonmetals on the periodic table): Going down a group: Going across a period: o Most reactive nonmetal:

6 Periodic Table Trends MC Questions: 1. Which general trend is found in a period as the elements are considered in order of increasing atomic number? A. Increasing atomic radius B. Increasing electronegativity C. Decreasing atomic mass D. Decreasing first ionization energy 2. As the elements of a group are considered in order from top to bottom, the first ionization energy of each successive element will A. Decrease B. Increase C. Remain the same 3. As atoms of elements a group are considered in order from top to bottom, the electronegativity of each successive element A. Decreases B. Increases C. Remains the same 4. As the elements of a group are considered in order from top to bottom, the covalent radius of each successive element increases. This increase is primarily due to an increase in A. Atomic number C. The number of protons occupying the nucleus B. Mass number D. The number of occupied energy levels 5. Which element in Period 3 has the largest covalent atomic radius? A. Cl B. Al C. Na D. P *Justify your answer: 6. Which of these elements in Period 3 has the least tendency to attract electrons? A. Mg B. Al C. S D. Cl *Justify your answer: 7. Which element in group 18 of the Periodic Table has the highest first ionization energy? A. Kr B. Ar C. Ne D. He *Justify your answer: 8. Atoms of which of the following elements have the smallest covalent radius? A. Si B. P C. S D. Cl *Justify your answer: 9. The atoms of which element in Group 16 have the greatest tendency to gain electrons? A. O B. S C. Se D. Te *Justify your answer: 10. Which list of elements from Group 2 on the Periodic Table is arranged in order of increasing atomic radius? A. Be, Mg, Ca B. Ca, Mg, Be C. Ba, Ra, Sr D. Sr, Ra, Ba *Justify your answer

7 Model 1: Hydrogen and Lithium Hydrogen Atom Lithium Atom 1. In the Hydrogen and Lithium atoms, what force of attraction holds the electron(s) in the atom? 2. The amount of energy necessary to remove an electron from an atom is called the ionization energy of that electron. What is the relationship between the ionization energy of an electron and the net attractive force that holds an electron in an atom? 3. Consider the electrons in an atom of lithium as diagrammed in the model above. Which electron, 1 or 3, will require more energy to be removed? Support your answer by discussing the attractive forces in the atom and how they might be different for electrons 1 and 3

8 Model 2: Photoelectron Spectra of Lithium Refer to the PES graph for Lithium above: 4. What are the units of the x-axis? What is unusual about the way the x-axis values are graphed? 5. Which of the peaks in the graph represents electrons that are more tightly held by the nucleus? Explain your reasoning 6. The number of peaks in a PES spectrum reveal the number of energy sublevels occupied by electrons in an atom. Based on the energy values of the peaks, label each peak with the electrons in a lithium atom to which they correspond. 7. Why is the higher energy peak about twice as high as the lower energy peak?

9 8. Using the lithium PES spectrum as a starting point, sketch how the spectrum of the next larger element (Beryllium) would look like. Recall that Be will have one more proton in its nucleus and on more electron in its sublevels. Model 3: Bohr Model Modified Photoelectric Spectra of Neon Bohr Model of Shell model of Ne Ne PES of Ne C A B 9. Based on the PES of Ne, which is a better electron model Bohr or Shell model? Why? 10. In the table below, label the peaks for the PES of Ne with the correct sublevels and number of electrons. Peak A B C Sublevel # of electrons 11. Write the electron configuration for Neon. 12. What do you notice about your answers to 10 and 11?

10 SUMMARY: 13. What does the PES tell us about the electrons in an atom? How can you use the PES of an atom to determine the identity of the atom? Practice #1 14. Label the peaks with the correct sublevels and number of electrons for each atom. 15. Consider the attractive and repulsive forces in the atoms of sulfur and phosphorus. a. Explain why most of the peaks in the sulfur spectrum are shifted to the left relative to the peaks in the phosphorus spectrum. b. Explain why peak E in the sulfur spectrum is shifted slightly right compared to peak E in the phosphorus spectrum. 16. Sketch the PES spectra for chlorine using the spectra for P and S as a guide

11 Practice #2 17. Answer the following questions based on the principles of atomic and electronic structure. Element #1 C Relative Number Of Electrons A B D 10, Binding Energy (ev) Element #2 C Relative Number Of Electrons A B D E 10, Binding Energy (ev) a. The diagrams above represent the photoelectron spectra (PES) for two different elements located in Period 3. Identify Element #1 and Element #2. Label each peak in each PES with the name of the orbital in which the electrons are located. Identify the orbital with its principal quantum number (n) and orbital type, such as 1s or 3p. Identity of Element #1 A B C D Identity of Element #2 A B C D E b. Explain why peak E is twice as high as peak D in the PES for Element #2. c. Explain why peak A in Element #2 has a higher binding energy than peak A in Element #1.

12 Practice #3 18. Look at the PES below. Identify the sublevel & number for each peak; write the electron configuration; and identify the element Electron configuration: Element identity is 19. Look at the PES below. Identify the sublevel & number for each peak; write the electron configuration; and identify the element Electron configuration: Element identity is