CHEM 105 Problem Set 10 PDF

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Problem Set 10 – Periodic Trends Chem 105 1. (a) What is the difference between a group and period in the Periodic Table? Periods are horizontal rows. Elements in the same period share the same electron configuration. Identify electron shells groups are vertical columns elements in a group share similar chemical or physical properties and they have the same number of valence electrons Electronegativity decreases as you go down the group and increases as you go up due to an increased pull in the nucleus as a result of fewer energy levels. Electronegativity increases as you go from left to right across the period with addition of more protons as atomic numbers (b) The Periodic Table can be divided into roughly 4 blocks according to the type of subshell being filled. State the subshell being filled in each of these blocks. S orbital, p orbital, d orbital, and f orbital 2. (a) There are also different ‘families’ of elements in the Periodic Table. List a transition metal, an alkali metal, a halogen, a noble gas, an alkaline earth metal, a nonmetal, a lanthanide, an actinide, and a metalloid in that order. Iron (Fe), Lithium (Li), Fluorine (F), Helium (He), Beryllium (Be) , Carbon ©, Cerium (Ce), Thorium (Th), Boron (B) (b) In which set of elements on the periodic table are the p orbitals being filled? a. alkali metals b. halogens c. actinides d. lanthanides e. transition metals 3. Size, ionization energy, and electron affinity are all properties that show definite trends across the Periodic Table. (a) Describe the trends for each of these properties going down a group and across a period. Size: increases down a group decreases across the period from left to right Ionization energy: increase down a group, Decrease across the period from left to right Electron affinity: electron affinity increases down a group and across the period from left to right (b) Explain the reason for the change in atomic radii across a period as well as down a group. Atomic radii increases down a group as the valence electrons are further from the nucleus. There are more inner shells as you go down a group. Atomic radii decreases across a period as the number of protons increase as the electrons are on the same outer shell. Therefore the

negative electrons experience a stronger attraction with the positive nucleus and are pulled closer to it Identify each statement as true or false. If false, re-write the statement so it is true: (c) Ionization energies are always negative. Ionization energies are always positive (endothermic = remove electrons = require energy) (d) Oxygen has a larger first ionization energy than fluorine. Fluorine has a larger first ionization energy than oxygen (fluorine has the highest ionization energy; fluorine has more electrons (7) than oxygen (6) in its outer shell, hence it is more difficult to remove an electron from fluorine than from oxygen) (e) The second ionization energy of an atom is always larger than its first ionization energy. True: because it always takes more energy to remove an electron from a positively charged ion than from a neutral ion 4. (a) When an atom forms an anion, its radius increases. When an atom forms a cation, its radius decreases. Explain why this is so. To form an anion, the atom gains electrons, hence the radius increases because there’s more negative charge than positive charge (nucleus) so the pull is less strong. To form a cation, the atom loses electrons, hence the radius decreased because there’s more positive charge than negative charge so the pull is stronger (b) Arrange the following atoms and ions in order from smallest to largest radius. a. Sr, Kr, Br, Rb, Ar Rb, Sr, Br, Kr, Ar b. Ne, Ne+, Ne– Ne-, Ne, Ne+ c. Al, Cs, Ne, Rb, S Cs, Rb, Al, S, Ne d. Cl, Cl–, Cs, Cs+ Cs, Cs+, Cl-, Cl e. K+, Ca2+, S2-, Cl-, Ar S2- , Cl-, K+, Ca2+ 5. (a) The compound, NaCl, consists of sodium cations and chlorine anions. The sodium and chlorine ions are not the same size as they are in the elemental state. Explain how their relative sizes change (larger or smaller) in relationship to their size in the elemental state. Sodium is smaller than its elemental state because it loses electrons to chlorine to form a full shell. Chlorine is large than its elemental state because it gains electrons from sodium to form a full shell (b) In the depiction of the crystal structure of salt to the right, which ions are the purple ions, and which are the green? Sodium = purple Chlorine = green 6. Which of the following atoms has the smallest first ionization energy? Ionization energy increases moving from left to right in period decreased moving top to bottom down group a. Al b. P c. Sr d. Ga e. Rb 7. (a) The first ionization energies (IE) of hydrogen and helium are about 1300 kJ/mol and 2300 kJ/mol respectively. Yet, the electron removed from both of these originated in a 1s orbital. Explain in 1-3 sentences the large IE difference observed between these elements.

The valence electrons in both hydrogen and helium are on the same valence shell 1s helium has a higher atomic number, meaning that it has 1 more proton than hydrogen, so the electron experiences a stronger attraction to the positive nucleus (b) Another particularly interesting exception to the trend in 1st ionization energy is found by comparing nitrogen and oxygen. The 1st IE is higher for N than for O, even though O has 1 more proton in its nucleus. Offer an explanation. HINT: it may be helpful to look at the valence electron diagrams for N and O. The electron lost in O is from a filled shell 2p 4 and the electron lost in N is from a half filled shell 2p3 the electron from the filled shell is easier to remove due to the repulsion between the two electrons form the same sub-shell. Also, a half filled shell is more stable 8. An electron can be removed from Na or Rb by electromagnetic radiation (the photoelectric effect). Which element, Na or Rb, would require the shortest wavelength to remove the electron? Shorter wavelength = higher energy. Na requires more energy (shorter wavelength) to remove an electron because its valence electron is closer to the nucleus (Na = 3s1 Rb = 5s1) 9. Based on their positions in the periodic table, predict which atom of the following pairs have the most negative electron affinity: a. Cl, Ar b. K, Co c. S, Ge d. Sn, Te 10. “Photo-gray” lenses for eyeglasses darken in bright sunshine because the lenses contain tiny, transparent AgCl crystals. Exposure to light removes electrons from Cl– ions, forming a chlorine atom in an excited state (indicated below by the asterisk): Cl– →hn Cl* + e– + The electrons are transferred to Ag ions, forming silver metal: Ag+ + e– → Ag Silver metal is reflective, producing the photo-gray color. How might substitution of AgBr for AgCl affect the light sensitivity of photo-gray lenses? In answering this question, consider whether more or less energy is needed to remove an electron from a Br– ion than from a Cl– ion.

More energy is required to remove an electron in Br- ion than from a Cl- ion hence the color of the visible light will change (shorter wavelength visible light)...