Slide 1 / 102 Slide 2 / 102 New Jersey Center for Teaching and Learning Progressive Science Initiative This material is made freely available at www.njctl.org and is intended for the non-commercial use of students and teachers. These materials may not be used for any commercial purpose without the written permission of the owners. NJCTL maintains its website for the convenience of teachers who wish to make their work available to other teachers, participate in a virtual professional learning community, and/or provide access to course materials to parents, students and others. AP Chemistry Periodic Trends Click to go to website: www.njctl.org Slide 3 / 102 The Periodic Law www.njctl.org Slide 4 / 102 The Periodic Law Over the course of this unit, we will use our knowledge of the atom to explain the periodic trends we see regarding the following properties: PROPERTY Ionic Charge Atomic/Ionic Radii Density DEFINITION charge of common ion formed by that element Distance from the nucleus to outermost electron Ratio of Mass/Volume Recall that the periodic law states that the physical and chemical properties of the elements tend to recur in a systematic way when arranged by increasing atomic number. Slide 5 / 102 Ionization Energy Metallic Character Electronegativity Energy required to remove valence electron Disposition to have metallic characteristics - ie. conduct electricity Measure of attraction for electrons when the atom is sharing electrons in a molecule. Slide 6 / 102 Recall that the periodic law states that the physical and chemical properties of the elements tend to recur in a systematic way when arranged by increasing atomic number. Atomic Number Ionic Charges The Periodic Law Let's look at the first eleven elements to illustrate this. H He Li Be B C N O F Ne Na 1 2 3 4 5 6 7 8 9 10 11 +1,-1 NA +1 +2 +3 +4-3 -2-1 NA +1 Notice that neither He or Ne form ions. Also, notice that in both cases the atom that precedes them can form a -1 ion and the atom that succeeds them forms a +1 ion. There is definitely a systemic pattern here! The pattern can be easily visualized on a graph, particularly as we move past the first 11 elements! +4 ion +3 charge +2 +1-1 -2-3 The Periodic Law 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 atomic number
Slide 7 / 102 The Periodic Law and the Quantum Model This trend in ionic charge can be easily explained if we apply the quantum model of the atom. Element Principal Quantum Number (N) of valence electrons H 1 1s 1 Electron Configuration Lose/ Gain electrons gain 1 lose 1 Ionic Charge -1 +1 He 1 1s 2 NA NA Li 2 [He]2s 1 lose 1 +1 Be 2 [He]2s 2 lose 2 +2 B 2 [He]2s 2 2p 1 lose 3 +3 C 2 [He]2s 2 2p 2 lose 4 +4 N 2 [He]2s 2 2p 3 gain 3-3 O 2 [He]2s 2 2p 4 gain 2-2 F 2 [He]2s 2 2p 5 gain 1-1 Ne 2 [He]2s 2 2p 6 NA NA Na 3 [Ne]3s 1 lose 1 +1 Slide 9 / 102 The pattern recurs with every increase in the principal quantum number. This means every time a new shell of electrons is filled, the pattern repeats! The Periodic Law and the Quantum Model Question 3: Explain why P would be expected to have the same ionic charge as N? Both have the same number of valence electrons (5) so both need to gain three electrons to fill their outer principal energy level. move for answer N = [He]2s 2 2p 3 gain 3 e- --> Ne P = [Ne]3s 2 3p 3 gain 3 e- --> Ar Question 4: After sodium, which element would most likely form an ion with +1 charge and why? Potassium (K), because it is beginning to fill the 4th principal energy level with 1 electron, move for just answer as sodium was beginning the 3rd with 1 electron. Slide 8 / 102 The Periodic Law and the Quantum Model Let's use to quantum model to answer some questions about these ionic charges. Question 1: Why do both He and Ne not form ions? Both have a full principal energy level move He for answer = 1s 2 Ne = [He]2s 2 2p 6 Question 2: Why do both Li and Na have the same charge? Both require only a small amount of energy to lose 1 electron to become a noble move gas for with answer a full principal energy level. Slide 10 / 102 The Periodic Law and the Quantum Model We have seen that the quantum model explains the periodic trend with regard to ionic charges for the main group elements in the first three periods. Quantum theory can also explain the periodic trends amongst the transition elements that are in the midst of filling their "d" orbitals. d orbital +3 +4 +5 +6 +7 +3 +3 +2 +1 +2 +3 +4 +5 +6 +7 +3 +3 +2 +1 +2 Slide 11 / 102 transition elements Slide 12 / 102 The Periodic Law and the Quantum Model d orbital +3 +4 +5 +6 +7 +3 +3 +2 +1 +2 +3 +4 +5 +6 +7 +3 +3 +2 +1 +2 transition elements The charges increase from left to right as the atoms lose both their two valence "s" electrons and however many "d" electrons they have also. The Periodic Law and the Quantum Model Let's use quantum theory to explain the trends we see amongst the charges of the transition elements. Question 1: Elements within the Fe group can form ions of both +2 and +3 charges. Explain why the +3 charge is more common: Fe = [Ar]4s 2 3d 6 The 4s electrons move are for readily answer lost yielding the +2 ion. A half-full "d" orbital is quite stable so Fe will lose 1 d orbital electron as well to yield the +3 ion. After the Mn group, the charges decrease, one of the reasons being that the stability of the "d" orbital increases as it becomes full.
Slide 13 / 102 The Periodic Law and the Quantum Model Let's use quantum theory to explain the trends we see among the charges of the transition elements. Question 2: Why do the elements in the zinc group tend to only form ions with a +2 charge? Zn = [Ar]4s 2 3d 10 The "d" orbital is full move so only for answer the outer "s" electrons are lost. Slide 14 / 102 1 The trends in chemical and physical properties tend to recur as atoms A Fill a new principal energy level B Gain more neutrons C Decrease in mass D Increase in atomic number E Both A and D Slide 15 / 102 2 An atom with a +2 charge must be in the same group as barium. Slide 16 / 102 3 Which of the following BEST explains why O and S both form ions with a -2 charge? True False A They both have the same atomic number B They are both in the same period C They both have the same electron configuration D They both have the same number of valence electrons E They both have the same mass Slide 17 / 102 4 An atom with the electron configuration of [Kr]5s 2 4d 2 would be in the same group as and have a likely charge of. A Sc, +1 B Hf, +4 C Ti, +3 D Zn, +2 E Y, +1 Slide 18 / 102 5 Atoms on the right side of the chart tend to form negative ions because... A Their principal energy level is almost empty B Their principal energy level is almost full C Their atomic number is less than other elements in that period D Both B and C E A, B, and C
Slide 19 / 102 The Periodic Law and Atomic/Ionic Radii The atomic/ionic radii of an atom can be measured and or calculated a number of different ways. We will be using values calculated via the Clementi method (E. Clementi, D.L.Raimondi, and W.P. Reinhardt, J. Chem. Phys. 1963, 38, 2686.) The atomic radius of an atom or ion can be thought of as the distance between the nucleus and the region of space where the outermost valence electrons would be most likely found. radius **Note: Remember an electron is not in orbit round the nucleus like a planet. The radius therefore is determined out to the point where the electron charge density starts to diminish Slide 21 / 102 The Periodic Law and Atomic Radii The distance the electrons can be from the nucleus is governed by Coulomb's law of attraction. The greater the charge, the greater the attraction between the charges, and the shorter the distance. As atomic number increases across a period, so does the nuclear charge (Z) resulting in a greater attraction and a smaller distance between the nucleus and the outermost electrons. Lithium (Z=3) radii = 167 pm Carbon (Z=6) radii = 67 pm Neon (Z=10) radii = 38 pm **Note: The size of an atom is NOT determined by the size of the nucleus. It is the electron cloud that contains most of the volume of an atom and therefore determines the radii. Slide 23 / 102 The Periodic Law and Ionic Radii When electrons are gained or lost, the effect on the radii can be dramatic or slight but there are some certainties. If an atom loses electrons, the radii will decrease. Ca --> Ca 2+ + 2e- 194 pm 99 pm When electrons are lost, the remaining electrons feel a stronger coulombic attraction from the nucleus. If an atom gains electrons, the radii will increase. F + e - --> F - 42 pm 136 pm When electrons are gained, the nuclear charge is spread over a larger number of electrons, resulting in a weaker coulombic attraction. Slide 20 / 102 The Periodic Law and Atomic Radii Let's examine the trend in atomic radii for the first 18 elements. radius (pm) 200 Li 100 H Ar Ne He 0 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 atomic number Na We clearly see two trends! 1. As atomic number increases down a group, the radii increase. H < Li < Na 2. As atomic number increases across a period, the radii decrease. Li > Be > B > C > N > O > F > Ne Slide 22 / 102 The Periodic Law and Atomic Radii Why don't the radii continue to get smaller as the atomic number and nuclear charge increase. The quantum model explains why. Hydrogen (Z=1) 1s 1 radii = 53 pm Lithium (Z=3) 1s 2 2s 1 radii = 167 pm Sodium (Z=11) 1s 2 2s 2 2p 6 3s 1 radii = 190 pm Only a certain number of electrons are permitted within a given energy level, so additional ones must be added to higher energy levels farther from the nucleus. The core electrons shield the valence electrons from the nucleus thus diminishing the coulombic attraction and increasing the atomic radii. Slide 24 / 102 The Periodic Law and Ionic Radii Let's rank a series of atoms and ions in order of increasing radii. Al 3+ Al Mg Mg 2+ Whenever comparing radii, use the following procedure: 1. Determine the energy level of the atom/ion. 2. For atoms in the same energy level, use the nuclear charge (Z) to determine the radii. Al 3+ Al Mg Mg 2+ Energy Level 2 3 3 2 "Z" 13 13 12 12 Al 3+ < Mg 2+ < Al < Mg radius (pm) 50 < 65 < 118 < 145
Slide 25 / 102 The Periodic Law and Ionic Radii Recall that in an isoelectronic series, the atoms/ions have the same number of electrons. Slide 26 / 102 The Periodic Law and Ionic Radii Let's try a few more together. 1. Explain why Si has an atomic radii of 111 pm while C has an atomic radii of 67 pm despite Si having a higher nuclear charge (Z)? In this case, Na +, Mg 2+, Al 3+, O 2-, and F - are all isoelectronic with Ne. As a result, they all experience the same core shielding. The ionic radii then decreases with an increasing nuclear charge. Al 3+ < Mg 2+ < Na + < F - < O 2- Z = 13 12 11 9 8 Si has an additional energy level, so the valence electrons are farther move away for and answer more shielded than those of C resulting in a smaller coulombic attraction. Slide 27 / 102 The Periodic Law and Ionic Radii Let's try a few more together. 2. Explain why iron (Fe) has a smaller atomic radii - 154 pm - than does scandium (Sc) - 184 pm. Although both have the same amount of shielding, Fe has a larger Z creating move a stronger for answer coulombic attraction and a smaller radii. Slide 28 / 102 6 Which of the following influences the atomic/ionic radii? A the number of neutrons B the amount of core electrons between the nucleus and the valence electrons C the number of protons D A and B E B and C Slide 29 / 102 7 The atomic radius of main-group elements generally increases down a group because. Slide 30 / 102 8 Of the following, which gives the correct order for atomic radius for Be, Li, N, C and Ne? A effective nuclear charge increases down a group A Be > Li > N > C > Ne B effective nuclear charge decreases down a group B Ne > C > N > Li > Be C effective nuclear charge zigzags down a group C C > N > Ne > Li > Be D E the principal quantum number of the valence orbitals increases both effective nuclear charge increases down a group and the principal quantum number of the valence orbitals increases D E Li > Be > C > N > Ne Ne > N > C > Be > Li
Slide 31 / 102 9 Which of the following atoms would have a smaller atomic radii than Ar and why? Slide 32 / 102 10 Which ion below has the largest radius? A Fe - It has more core electrons B Si - It has fewer core electrons C O - It has fewer core electrons D Ne - it has a higher nuclear charge (Z) A O 2- B Li + C I - D N 3- E K + E Ca - it has a higher nuclear charge (Z) Slide 33 / 102 11 Which of the following pairs correctly shows the proper relationship between the two atoms/ions in terms of atomic/ionic radii? A Ca < Ca 2+ B F < F - C V < Mn D Ca < Be E He > Li Slide 34 / 102 12 Which of the following correctly states why the atomic radii do not consistently decrease as the atomic number rises throughout the periodic table? A The nuclear charge (Z) does not always increase with atomic number B The number of neutrons start to influence the atomic radii C Filled energy levels shield the nucleus and diminish coulombic forces D Electrons become less negative the more there are E A higher atomic number increases the size of the radii, not decreases it. Slide 35 / 102 13 Which of the following would correctly rank the following in order of decreasing atomic/ionic radii? A V 4+ > V 5+ > F > F - B V 4+ > V 5+ > F- > F Slide 36 / 102 14 Isotopes of an element, like C-12 and C-13, are likely to have different atomic radii? Yes No C V 5+ > V 4+ > F - > F D V 5+ > V 4+ > F > F - E F > F - > V 4+ > V 5+
Slide 37 / 102 The Periodic Law and Ionization Energy Ionization energy is the amount of energy required to remove an electron from an atom. This creates an ion, hence the name! The stronger the Coulombic attraction between the valence electron and the nucleus, the greater the energy required to remove an electron. Element Li + IE --> Li + + e- Na + IE --> Na + + e- Ionization Energy 520 kj/mol 496 kj/mol Less energy is required to remove sodium's electron than lithium's because sodium has a full energy level more of core electrons shielding the nuclear charge. Slide 39 / 102 The Periodic Law and Ionization Energy The chart below clearly shows the impact of being isoelectronic with a noble gas on the ionization energy. Slide 38 / 102 The Periodic Law and Ionization Energy Unless you're hydrogen, you've got multiple electrons that can be lost. As a result we have to distinguish between 1st, 2nd, 3rd, etc. ionization energies. Each successive ionization energy is always higher than the previous. This is due to the higher nuclear charge felt by the remaining electrons. Ionization Ionization Energy 1st: Na + IE --> Na + + e- 496 kj/mol 2nd: Na + + IE --> Na 2+ + e- 4560 kj/mol 3rd: Na 2+ + IE --> Na 3+ + e- 6,900 kj/mol 4th: Na 3+ + IE --> Na 4+ + e- 9540 kj/mol Note the huge jump in ionization energy from the 1st to the 2nd. After sodium loses it's first electron, it is isoelectronic with [Ne], with an extremely stable full s and p orbital and minimal shielding. Slide 40 / 102 The Periodic Law and Ionization Energy The trend in first ionization energies mostly matches what we would expect. Ionization Energy (kj/mol) Na + Mg 2+ Al 3+ Si 4+ P 5+ S 6+ The ionization energy increases across a period with increasing atomic number. ( Li < Ne) The ionization energy decreases down a group with increasing atomic number due to additional core electrons from each filled energy level shielding the nucleus. ( He > Ne) Slide 41 / 102 The Periodic Law and Ionization Energy Slide 42 / 102 The Periodic Law and Ionization Energy There are however a few hiccups that need to be explained. Let's look carefully at the ionization energies of Be and B as well as N and O indicated in the circles. Be: [He]2s 2 N: [He]2s 2 2p 3 Shouldn't the ionization energy increase with increasing atomic number across a period? Quantum theory will explain. B: [He]2s 2 2p 1 O: [He]2s 2 2p 4 More energy is required to remove an electron from Be's full "s" orbital More energy is required to remove an electron from N's 1/2 full "p" orbital
Slide 43 / 102 The Periodic Law and Ionization Energy Let's look at another hiccup in the trend. Slide 44 / 102 The Periodic Law and Ionization Energy Let's practice ranking atoms/ions in terms of ionization energy: 1. Rank the following in terms of increasing ionization energy: C Al Na+ Ne Na Notice that a lot less energy is required to remove an electron from Ga (Z=31) than from Zn (Z=30). How can this be? Zinc has a full "s" and "d" orbital conferring extra stability while in gallium, the electron is being taken from a "p" orbital which is heavily shielded from the nucleus by the "d" orbital itself. As with atomic radii, determine their outermost principal energy level and nuclear charge. C Al Na+ Ne Na Valence "N" 2 3 2 2 3 move for answer "Z" 6 13 11 10 11 Na < Al < C < Ne < Na + IE(kJ/mol) 496 578 1086 2081 4560 Slide 45 / 102 15 What is the ionization energy? A Energy change associated with the gain of an electron B Amount of energy that is required to move an electron from an s to a p orbital C Measure of the attraction of an atom for electrons when in a compound D Pull of the neutrons on the electrons E Amount of energy required to remove an electron from an atom or ion Slide 46 / 102 16 Which of the following would NOT influence the ionization energy? A The shielding from core electrons B The extent to which an orbital is full C The nuclear charge D The number of principal energy levels between the valence electrons and the nucleus E All of these influence the ionization energy Slide 47 / 102 17 Which of the following elements would be expected to have a higher ionization energy than magnesium (Mg)? A Al B Ca C Na D K E B Slide 48 / 102 18 Which of the following correctly ranks the elements below in order of decreasing ionization energy? A Ne > O > N B Ne > N > O C H > He > Ne D Li > Mg > Ga E Zn > Ga > Br
Slide 49 / 102 19 Which of the following elements best fits the data provided below? Slide 50 / 102 20 Which of the following pairs are correct in terms of relative first ionization energy and why? A Li B C C Be D Ne E O Ionization Ionization Energy 1st: X + IE --> X + + e- 900 kj/mol 2nd: X + + IE --> X 2+ + e- 1757 kj/mol 3rd: X 2+ + IE --> X 3+ + e- 14,850 kj/mol Answer A O 2- < Ne, due to smaller nuclear charge on oxide ion B Li > Na, due to increased shielding in the Na atom C Zn > Cu, due to a higher nuclear charge in zinc D Cl > S, due to the smaller nuclear charge in sulfur E All of these Slide 51 / 102 Slide 52 / 102 21 The second ionization energy will always be higher than the first. 22 have the lowest first ionization energies of the groups listed. True False A B C D E Alkali metals Transition elements Halogens Alkaline eath metals Noble gases Slide 53 / 102 23 Of the choices below, which gives the order for decreasing first ionization energies? A B C D E Cl > S > Al > Ar > Si Ar > Cl > S > Si > Al Al > Si > S > Cl > Ar Cl > S > Al > Si > Ar S > Si > Cl > Al > Ar Slide 54 / 102 Ionization Energy and PES Ionization energy data can be determined from PES (photoelectron spectroscopy). Recall that PES looks at the energy of light required to remove electrons from an atom. Each orbital appears as a peak on the spectrum. Intensity Li (2s) Be (2s) binding energy Li (1s) Be (1s) The PES spectrum clearly shows that the core electrons require the most energy to remove. It also shows that Be has a higher 1st IE for the removal of the valence electrons than does Li. This is expected as Be has a higher "Z".
Slide 55 / 102 Ionization Energy and PES Let's interpret another PES spectra, this one of nitrogen and oxygen. O (2p) O (2s) O (1s) Intensity Slide 56 / 102 Ionization Energy and PES Click to go to an interactive PES spectra database and answer the questions. N (2p) binding energy N (2s) N (1s) Why is the N (2p) peak greater than the O (2p) peak? N has a half-full "p" orbital increasing the move ionization for answer energy Why is the N(2s) peak less than the O (2s) peak? O has the higher nuclear charge move for answer Slide 57 / 102 24 The following PES spectrum shows the valence "p" orbital peaks for Si and for C. Which of the following would be TRUE? Intensity binding energy A The Si peak is of lower energy due to it's higher nuclear charge B The Si peak is of higher energy due to the increased shielding from core electrons C The Si peak is of lower energy due to the increased shielding from core electrons D The Si peak is of higher energy due to its higher nuclear charge Slide 59 / 102 26 Below is an actual PES spectrum of palladium (Pd). Which of the following would be TRUE? (Note: the outer 5s and 4d peaks are not shown) 3s 3p A Compared to Pd, the 3d peak in Cd would be found to the left of the 3d Pd peak B Compared to Pd, the 3d peak in Rb would be of a higher binding energy due to lower nuclear charge C Compared to Pd, the 3p peak in Kr should be found to the left of the 3p peak in Pd 3d 4p 4s Why is the binding energy of the electrons greater in He than H? Similar shielding move for but answer greater "Z" Which peak in the Li spectra represents the valence electrons? Peak with move lower for binding answer energy Why is the valence peak binding energy less in Li than in H? Increased shielding due to move core 1s for electrons, answerlessens coulombic force Why is the core peak (1s) binding energy greater in Li than in H? Lithium has a higher nuclear move charge for "Z" answer so higher coulombic attractions Slide 58 / 102 25 The 3s peak for magnesium should have a higher binding energy than that of the 4s peak in calcium due to calcium's higher amount of shielding by core electrons? True False Slide 60 / 102 27 Based on the PES data below, what would be TRUE regarding atoms 1 and 2? Intensity 10 1.09 0 1.72 28.6 1 2 Binding Energy A I only B II and III only C 1 and III only D II and IV only 100 10 Intensity 10 0 1.40 2.45 39.6 Binding Energy I. Atom 1 has a smaller atomic radii II. Atom 2 has a larger first ionization energy III. Both atoms are in the same period IV. Both atoms are in the same group E I, II, III, and IV 100 10
Slide 61 / 102 Ionization Energy and Metallic Character Metals are generally described as being able to lose electrons readily which promotes conductivity. Slide 62 / 102 Ionization Energy and Metallic Character We can predict, based on ionization energies, where the metals and non-metals are on the periodic table. Since metals lose electrons easily, they must have low ionization energies compared to non-metals. Element Metal or Non-metal 1st Ionization Energy (kj/mol) Na metal 496 O non-metal 1314 semi-metals or metalloids Notice that an element becomes more metallic as the shielding increases and as the nuclear charge - for a given level of shielding - decreases. Slide 63 / 102 Ionization Energy and Metallic Character Let's answer a few questions regarding metallic character. 1. Why is lead considered a metal and carbon a non-metal despite being in the same group? Pb has much more shielding due to more move levels for of answer core electrons thereby causing it's electrons to be lost far more easily than that of C. C Si Ge Sn Pb Slide 64 / 102 Ionization Energy and Metallic Character Let's answer a few questions regarding metallic character. 2. Which metal would we expect to be a better conductor of electricity? Ag or Cu Ag due to the higher amount of shielding, causing it to ionize more easily, thereby creating move mobile for answer electrons. Cu Ag Slide 65 / 102 Ionization Energy and Metallic Character Application: Elements of Life The most common elements in living things are C,H,N,O,P, and S. Interestingly, these are all non-metals. Slide 66 / 102 28 Which of the following is the LEAST metallic of those below? A F B At Serotonin - brain hormone Interestingly, all metal atoms found in living things are in their ionic form (Mg 2+, Ca 2+, Zn 2+, etc.) In order to form large stable, yet complex, molecules, the elements must not be able to lose electrons easily. 2+ 2+ C Ne D Xe E Ba
Slide 67 / 102 29 Which of the following would be TRUE? A The higher the ionization energy, the less metallic an element will be B The lower the ionization energy, the less metallic an element will be C For a given amount of core electron shielding, the higher the nuclear charge, the more metallic an element will be D Both A and C Slide 68 / 102 30 Which of the following has the elements correctly ordered by increasing metallic character? A Li < Be < B B Ca < K < Ga C Ga < Ca < K D Rb < Cs < As E Ga < As < Ba E Both B and C Slide 69 / 102 Ionization Energy and Light As we have seen, EM radiation can provide the necessary energy to ionize an electron from an atom. photon The higher the ionization energy, the higher the frequency of light needed to ionize the electron. e- Slide 70 / 102 Ionization Energy and Light Which of the following elements would require the shortest wavelength to lose an electron? Si C N Short wavelength means high energy so this would be the element with the largest ionization energy. Si C N "N" 3 2 2 move for answer "Z" 14 6 8 N has similar shielding as carbon but a higher nuclear charge so it would require the shortest wavelength to ionize an electron. Slide 71 / 102 Ionization Energy and Light What would be the necessary wavelength required to remove one of Neon's outermost p electrons? 1. Look up 1st IE of Neon = 2081 kj/mol 2. Convert to kj/atom = 2081 kj x 1 mol = 3.46 x 10-21 kj mol 3. Convert to J = 3.46 x 10-18 J 6.022 x 10 23 atoms 4. Convert to v via E=hv --> v = E/h = 3.46 x 10-18 J = 5.2 x 10 15 1/s 6.3 x 10-34 J*s 5. Convert to wavelength via v = c --> = c/v 3 x 10 8 m*s = 5.77 x 10-8 m = 57.7 nm 5.2 x 10 15 s move for answer Slide 72 / 102 31 Which of the following orbitals of calcium would require the highest frequency of light to ionize? A 2s B 2p C 3s D 3p E 4s
Slide 73 / 102 32 Based on the table of 1st ionization energies below, which element is likely to ionized by light with wavelength of 214 nm? Slide 74 / 102 33 What frequency of light would be required to ionize the first electron of cesium (1st IE = 376 kj/mol)? A I B Ga C In D He E Rb Element Ionization Energy (kj/mol) I 1009 Ga 579 In 558 He 2372 Rb 403 Answer Slide 75 / 102 Periodic Law and Electronegativity As we know, atoms do not often exist in isolation. They form bonds with other atoms to make molecules and compounds. Slide 76 / 102 Periodic Law and Electronegativity Let's compare the electronegativities of H and O within the water molecule. H O H water H O H Recall that electronegativity is defined as a measure of an atom's attraction for electrons in a bond. The greater the nuclear charge and the smaller the shielding, the greater the electronegativity. O has more shielding but a much higher nuclear charge so it will have the higher electronegativity. Therefore the electrons get pulled unevenly toward the oxygen atom. H O H Slide 77 / 102 Periodic Law and Electronegativity Trends in electronegativity for periods 2-4. Slide 78 / 102 Periodic Law and Electronegativity Trends in electronegativity for periods 2-4. O S Se Li Na K What is the trend in electronegativity down a group? decreases, due to additional shielding from each new energy move for answer level What is the trend in electronegativity across a period from left to right? increases, due to increasing nuclear charge with steady amount of move shielding for answer Why do the noble gases not have published electronegativity values? They have a full move outer for "s" answer and "p" system and do not form compounds.
Slide 79 / 102 Periodic Law and Electronegativity The following electronegativity values will need to be memorized as this will aid in understanding bonding later on. H 2.2 C 2.5 N 3.0 O 3.5 S 2.6 F 4.0 Cl 3.2 Br 3.0 Slide 80 / 102 34 Of the atoms below, is the most electronegative. A B C D E Si Cl Rb Ca S Slide 81 / 102 35 Which of the following BEST explains why fluorine has a higher electronegativity than oxygen? A F has a higher nuclear charge and less shielding than O B F has a higher nuclear charge and similar shielding of O C F has the equivalent nuclear charge and less shielding than O D F has the equivalent nuclear charge and more shielding than O Slide 82 / 102 36 Which of the following groups of elements are ranked properly from lowest to highest electronegativity? A H < Li < Na B H < C < Li C C < Si < Ge D I < Br < Cl E F < S < As E None of these Slide 83 / 102 37 An element with a small electronegativity value is likely to have... A Valence shell PES peaks with high binding energies B A high nuclear charge and a low amount of shielding C A low nuclear charge and a high amount of shielding D Both A and B Slide 84 / 102 Specific Groups of Periodic Table We will now examine six groups of the periodic table in more detail. Group 1: Alkali Metals Group 2: Alkaline Earth Metals Group 3-12: Transition Metals Group 13/14/15: Metalloids Group 17: Halogens Group 18: Noble Gases E Both A and C
Slide 85 / 102 Alkali Metals They are highly reactive due to their extremely low ionization energies. As a result, they are found only in compounds in nature, not in their elemental forms. Slide 86 / 102 Alkaline Earth Metals Alkaline earth metals have higher densities and melting points than alkali metals. Their ionization energies are low, but not as low as those of alkali metals so they are slightly less reactive. They have low densities and melting points. In fact Li, Na, and K have densities so low, they'll float on water! Slide 87 / 102 Alkaline Earth Metals Beryllium does not react with water and magnesium reacts only with steam, but the others react readily with water. Reactivity tends to increase as you go down the group. Can you explain why that would be? Slide 88 / 102 Transition Metals The transition metals vary somewhat in properties but we can simplify to say that they are less reactive than either of the first two groups. In fact, the least reactive metals (Au, Pt, Ag) are in this group. Transition metals tend also to have higher densities and melting points than the first two groups. Due to their "d" orbitals, they can form ions with much higher charges than the first two groups which will allow them to form colored complex ions with water and other species. Slide 89 / 102 Transition Metals Some complex ions formed from transition metals and their colors. Slide 90 / 102 Metalloids There are six elements that are classified as metalloids: Boron (B) Silicon (Si) Arsenic (As) Tellurium (Te) Germanium (Ge) Antimony (Sb) These have some characteristics of metals and some of nonmetals. For instance, silicon looks shiny like a metal, but is brittle and a fairly poor conductor.
Slide 91 / 102 Metalloids Metalloids like Si, although they are not particularly conductive due to higher ionization energies than metals can be made to be by "doping" them with certain elements to increase their conductivity. Slide 92 / 102 Halogens The halogens are prototypical nonmetals. They only require one more electron to have a full "s" and "p" and are therefore highly reactive. The name comes from the Greek words halos and gennao: salt formers. Circuits that form the basis for modern electronics are composed of doped metalloids like Si and Ge. Slide 93 / 102 Halogens (at standard temp and pressure) Flourine is a colorless gas Chlorine is a greenish gas Bromine is a brownish liquid Iodine is a purplish solid Slide 94 / 102 Noble Gases The noble gases have very high ionization energies. Therefore, they are relatively unreactive. As a result, unlike the diatomic halogens, they are found as monatomic gases Slide 95 / 102 38 An atom with a very high ionization energy and is a liquid at room temperature is most likely a: Slide 96 / 102 39 Which of the following ranks the metals in order of increasing reactivity? A Alkali metal B Alkaline earth metal C transition metal D Halogen E Noble gas A Li < Na < Mg < K B Mg < Li < Na < K C K < Li < Na < K D Li < Fe < Zn < Au E None of these
Slide 97 / 102 40 Which of the following elements would form colored complex ions? Slide 98 / 102 41 Which of the following elements would serve as a semiconductor? A F B Co C Ca D Al E Na A Ge B C C F D Pb E Y Slide 99 / 102 42 What would be the alkaline earth metal with the highest ionization energy? A Li B Al C Be D B E Ra Slide 100 / 102 43 Which would be the halogen with the smallest atomic radii? A Ne B F C At D Pb E Fr Slide 101 / 102 Slide 102 / 102 Now that we have a good understanding of some of the properties of various elements, we will now examine how they react and what they produce when they do in the next chapter.