Periodic Trends. Slide 1 / 102. Slide 2 / 102. Slide 3 / 102. Slide 4 / 102. Slide 6 / 102. Slide 5 / 102. AP Chemistry.

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1 Slide 1 / 10 Slide / 10 New Jersey enter for Teaching and Learning Progressive Science Initiative P hemistry This material is made freely available at 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. NJTL 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. Periodic Trends lick to go to website: Slide 3 / 10 Slide 4 / 10 The Periodic Law 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 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. charge of common ion formed by that element tomic/ionic Radii Distance from the nucleus to outermost electron Density Ratio of Mass/Volume Ionization Energy Energy required to remove valence electron Metallic haracter Disposition to have metallic characteristics - ie. conduct electricity Electronegativity Measure of attraction for electrons when the atom is sharing electrons in a molecule. Slide 5 / 10 Slide 6 / 10 The Periodic Law The Periodic Law 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. Let's look at the first eleven elements to illustrate this. tomic Number The pattern can be easily visualized on a graph, particularly as we move past the first 11 elements! +4 H He Li e N O F Ne Na N +1 Ionic +1,-1 N harges DEFINITION Ionic harge ion +3 charge Notice that neither He or Ne form ions. lso, 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! atomic number

2 Slide 7 / 10 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. Principal Quantum Element Number (N) of valence electrons Electron onfiguration Lose/ Gain electrons -1 lose 1 +1 N N [He]s1 lose 1 +1 [He]s lose + [He]sp1 lose 3 +3 [He]sp lose 4 +4 N [He]sp3 gain 3-3 O [He]sp4 gain - F [He]sp5 gain 1-1 Ne [He]sp6 N N Na 3 [Ne]3s1 lose s1 He 1 1s Li e 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? Ionic harge gain 1 H Slide 8 / 10 oth have a full principal energy level 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! Slide 9 / 10 The Periodic Law and the Quantum Model Question 3: Explain why P would be expected to have the same ionic charge as N? oth 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]sp3 gain 3 e- --> Ne P = [Ne]3s3p3 gain 3 e- --> r Question 4: fter 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 move for answer energy level with 1 electron, just as sodium was beginning the 3rd with 1 electron. move He for answer = 1s Ne = [He]sp6 Question : Why do both Li and Na have the same charge? oth require only a small amount of energy to lose 1 electron to become a noble gasfor with a full principal energy level. move answer Slide 10 / 10 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 transition elements Slide 11 / 10 The Periodic Law and the Quantum Model d orbital Slide 1 / 10 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 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. fter the Mn group, the charges decrease, one of the reasons being that the stability of the "d" orbital increases as it becomes full. Question 1: Elements within the Fe group can form ions of both + and +3 charges. Explain why the +3 charge is more common: Fe = [r]4s3d6 The 4s electrons areforreadily lost yielding the + ion. move answer half-full "d" orbital is quite stable so Fe will lose 1 d orbital electron as well to yield the +3 ion.

3 Slide 13 / 10 Let's use quantum theory to explain the trends we see among the charges of the transition elements. Question : Why do the elements in the zinc group tend to only form ions with a + charge? Zn = [r]4s3d10 move for answer The "d" orbital is full so only the outer "s" electrons are lost. 1 The trends in chemical and physical properties tend to recur as atoms Fill a new principal energy level Gain more neutrons Decrease in mass The Periodic Law and the Quantum Model Slide 14 / 10 D Increase in atomic number E oth and D Slide 14 () / 10 1 The trends in chemical and physical properties tend to recur as atoms Fill a new principal energy level False E D Increase in atomic number E oth and D Slide 15 () / 10 n atom with a + charge must be in the same group as barium. Slide 16 / 10 3 Which of the following EST explains why O and S both form ions with a - charge? They both have the same atomic number False They are both in the same period True FLSE They both have the same electron configuration D They both have the same number of valence electrons E They both have the same mass n atom with a + charge must be in the same group as barium. True Gain more neutrons Decrease in mass Slide 15 / 10

4 Slide 16 () / 10 They both have the same atomic number They are both in the same period They both have the same electron configuration D They both have the same numberd of valence electrons 4 n atom with the electron configuration of [Kr]5s4d would be in the same group as and have a likely charge of. Sc, +1 Hf, +4 3 Which of the following EST explains why O and S both form ions with a - charge? Slide 17 / 10 Ti, +3 D Zn, + E Y, +1 E They both have the same mass Slide 17 () / 10 4 n atom with the electron configuration of [Kr]5s4d would be in the same group as and have a likely charge of. Slide 18 / 10 5 toms on the right side of the chart tend to form negative ions because... Hf, +4 Ti, +3 Their principal energy level is almost full D Zn, + Their atomic number is less than other elements in that period D oth and E Y, +1 Slide 18 () / 10 5 toms on the right side of the chart tend to form negative ions because... Their principal energy level is almost empty Their principal energy level is almost full Their atomic number is less than other elements in that period D oth and E,, and Slide 19 / 10 The Periodic Law and tomic/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 lementi method (E. lem enti, D.L.Raim ondi, and W.P. Reinhardt, J. hem. Phys. 1963,3 8, 686.) 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 E,, and **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 Their principal energy level is almost empty Sc, +1

5 Slide 0 / 10 Slide 1 / 10 The Periodic Law and tomic Radii Let's examine the trend in atomic radii for the first 18 elements. Na 00 radius (pm) Li H r Ne He The Periodic Law and tomic Radii The distance the electrons can be from the nucleus is governed by oulomb's law of attraction. The greater the charge, the greater the attraction between the charges, and the shorter the distance. s 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 atomic number We clearly see two trends! 1. s atomic number increases down a group, the radii increase. H < Li < Na. s atomic number increases across a period, the radii decrease. Lithium (Z=3) arbon (Z=6) Neon (Z=10) radii = 167 pm radii = 67 pm 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. Li > e > > > N > O > F > Ne Slide / 10 Slide 3 / 10 The Periodic Law and tomic Radii The Periodic Law and Ionic 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) 1s1 radii = 53 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. Lithium (Z=3) 1ss1 radii = 167 pm Sodium (Z=11) 1ssp63s1 radii = 190 pm The core electrons shield the valence electrons from the nucleus thus diminishing the coulombic attraction and increasing the atomic radii. Slide 4 / 10 The Periodic Law and Ionic Radii l Mg If an atom loses electrons, the radii will decrease. a --> a pm e- If an atom gains electrons, the radii will increase. F + e- --> F4 pm 99 pm When electrons are lost, the remaining electrons feel a stronger coulombic attraction from the nucleus. 136 pm When electrons are gained, the nuclear charge is spread over a larger number of electrons, resulting in a weaker coulombic attraction. Slide 5 / 10 Let's rank a series of atoms and ions in order of increasing radii. l3+ When electrons are gained or lost, the effect on the radii can be dramatic or slight but there are some certainties. The Periodic Law and Ionic Radii Recall that in an isoelectronic series, the atoms/ions have the same number of electrons. Mg+ Whenever comparing radii, use the following procedure: 1. Determine the energy level of the atom/ion.. For atoms in the same energy level, use the nuclear charge (Z) to determine the radii. l3+ l Mg Mg+ Energy Level 3 3 "Z" l3+ < Mg+ < l < Mg radius (pm) 50 < 65 < 118 < 145 In this case, Na+, Mg+, l3+, O-, and F- are all isoelectronic with Ne. s a result, they all experience the same core shielding. The ionic radii then decreases with an increasing nuclear charge. l3+ < Mg+ < Na+ < F- < OZ=

6 Slide 6 / 10 Slide 7 / 10 The Periodic Law and Ionic Radii The Periodic Law and Ionic Radii Let's try a few more together. Let's try a few more together.. Explain why iron (Fe) has a smaller atomic radii pm than does scandium (Sc) pm. 1. Explain why Si has an atomic radii of 111 pm while has an atomic radii of 67 pm despite Si having a higher nuclear charge (Z)? lthough both have the same amount of shielding, Fe has a for answer larger Z creatingmove a stronger coulombic attraction and a smaller radii. Si has an additional energy level, so the valence move forand answer electrons are farther away more shielded than those of resulting in a smaller coulombic attraction. Slide 8 / 10 Slide 8 () / 10 6 Which of the following influences the atomic/ionic radii? 6 Which of the following influences the atomic/ionic radii? the amount of core electrons between the nucleus and the valence electrons the amount of core electrons between the nucleus and the valence electrons the number of protons D and the number of neutrons the number of neutrons the number of protons D and E and E E and Slide 9 / 10 Slide 9 () / 10 effective nuclear charge increases down a group effective nuclear charge increases down a group effective nuclear charge decreases down a group effective nuclear charge decreases down a group effective nuclear charge zigzags down a group D the principal quantum number of the valence orbitals increases 7 The atomic radius of main-group elements generally increases down a group because. 7 The atomic radius of main-group elements generally increases down a group because. effective nuclear charge zigzags down a group D the principal quantum number of the valence orbitals increases both effective nuclear charge increases down a group E and the principal quantum number of the valence orbitals increases both effective nuclear charge increases down a group E and the principal quantum number of the valence orbitals increases

7 Slide 30 / 10 Slide 30 () / 10 8 Of the following, which gives the correct order for atomic radius for e, Li, N, and Ne? 8 Of the following, which gives the correct order for atomic radius for e, Li, N, and Ne? e > Li > N > > Ne Ne > > N > Li > e Ne > > N > Li > e > N > Ne > Li > e e > Li > N > > Ne D > N > Ne > Li > e D Li > e > > N > Ne D Li > e > > N > Ne E Ne > N > > e > Li E Ne > N > > e > Li Slide 31 / 10 Slide 31 () / 10 9 Which of the following atoms would have a smaller atomic radii than r and why? 9 Which of the following atoms would have a smaller atomic radii than r and why? Si - It has fewer core electrons Si - It has fewer core electrons O - It has fewer core electrons Fe - It has more core electrons Fe - It has more core electrons O - It has fewer core electrons D Ne - it has a higher nuclear charge (Z) D Ne - it has a higher nuclear charge (Z) E a - it has a higher nuclear charge (Z) E a - it has a higher nuclear charge (Z) Slide 3 / 10 Slide 3 () / Which ion below has the largest radius? O- O- Li+ Li+ ID N3E K+ ID N3-10 Which ion below has the largest radius? E K+

8 Slide 33 / 10 Slide 33 () / Which of the following pairs correctly shows the proper relationship between the two atoms/ions in terms of atomic/ionic radii? 11 Which of the following pairs correctly shows the proper relationship between the two atoms/ions in terms of atomic/ionic radii? F < F- F < F- V < Mn V < Mn D a < e D a < e E He > Li E He > Li Slide 34 / 10 a < a+ a < a+ Slide 34 () / 10 1 Which of the following correctly states why the atomic radii do not consistently decrease as the atomic number rises throughout the periodic table? 1 Which of the following correctly states why the atomic radii do not consistently decrease as the atomic number rises throughout the periodic table? The number of neutrons start to influence the atomic radii The number of neutrons start to influence the atomic radii The nuclear charge (Z) does not always increase with atomic number The nuclear charge (Z) does not always increase with atomic number Filled energy levels shield the nucleus and diminish coulombic forces Filled energy levels shield the nucleus and diminish coulombic forces D Electrons become less negative the more there are D Electrons become less negative the more there are E higher atomic number increases the size of the radii, not decreases it. E higher atomic number increases the size of the radii, not decreases it. Slide 35 / 10 Slide 35 () / Which of the following would correctly rank the following in order of decreasing atomic/ionic radii? 13 Which of the following would correctly rank the following in order of decreasing atomic/ionic radii? V4+ > V5+ > F- > F V4+ > V5+ > F- > F V5+ > V4+ > F- > F D V5+ > V4+ > F > FE F > F- > V4+ > V5+ V4+ > V5+ > F > F- V4+ > V5+ > F > F- V5+ > V4+ > F- > F D V5+ > V4+ > F > FE F > F- > V4+ > V5+

9 Slide 36 / 10 Slide 36 () / Isotopes of an element, like -1 and -13, are likely to have different atomic radii? 14 Isotopes of an element, like -1 and -13, are likely to have different atomic radii? No No Yes Yes NO Slide 37 / 10 Slide 38 / 10 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 Periodic Law and Ionization Energy Unless you're hydrogen, you've got multiple electrons that can be lost. s a result we have to distinguish between 1st, nd, 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. The stronger the oulombic attraction between the valence electron and the nucleus, the greater the energy required to remove an electron. Ionization Ionization Energy Element Ionization Energy 1st: Na + IE --> Na+ + e- 496 kj/mol Li + IE --> Li+ + e- 50 kj/mol nd: Na+ + IE --> Na+ + e kj/mol Na + IE --> Na+ + e- 496 kj/mol 3rd: Na+ + IE --> Na3+ + e- 6,900 kj/mol 4th: Na3+ + IE --> Na4+ + e 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 / 10 Note the huge jump in ionization energy from the 1st to the nd. fter 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 / 10 The Periodic Law and Ionization Energy The chart below clearly shows the impact of being isoelectronic with a noble gas on the ionization energy. The Periodic Law and Ionization Energy The trend in first ionization energies mostly matches what we would expect. Ionization Energy (kj/mol) Na+ Mg+ l3+ Si4+ P5+ S6+ 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)

10 Slide 41 / 10 Slide 4 / 10 The Periodic Law and Ionization Energy 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 e and as well as N and O indicated in the circles. e: [He]s N: [He]sp3 : [He]s p More energy is required to remove an electron from e's full "s" orbital Shouldn't the ionization energy increase with increasing atomic number across a period? Quantum theory will explain. O: [He]sp4 1 Slide 43 / 10 More energy is required to remove an electron from N's 1/ full "p" orbital Slide 44 / 10 The Periodic Law and Ionization Energy The Periodic Law and Ionization Energy Let's look at another hiccup in the trend. Let's practice ranking atoms/ions in terms of ionization energy: 1. Rank the following in terms of increasing ionization energy: l Na+ Ne Na s with atomic radii, determine their outermost principal energy level and nuclear charge. Valence "N" 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? "Z" 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. Na IE(kJ/mol) Slide 45 / Na+ Ne Na move for answer < 496 l 578 < < Ne < Na What is the ionization energy? Energy change associated with the gain of an electron mount of energy that is required to move an electron from an s to a p orbital l Slide 45 () / What is the ionization energy? Energy change associated with the gain of an electron E mount of energy that is required to move an electron from an s to a p orbital Measure of the attraction of an atom for electrons when in a compound Measure of the attraction of an atom for electrons when in a compound D Pull of the neutrons on the electrons D Pull of the neutrons on the electrons E mount of energy required to remove an electron from an atom or ion E mount of energy required to remove an electron from an atom or ion

11 Slide 46 / 10 Slide 46 () / Which of the following would NOT influence the ionization energy? 16 Which of the following would NOT influence the ionization energy? The extent to which an orbital is full The extent to which an orbital is full The nuclear charge The shielding from core electrons The shielding from core electrons E The nuclear charge D The number of principal energy levels between the valence electrons and the nucleus D The number of principal energy levels between the valence electrons and the nucleus E ll of these influence the ionization energy E ll of these influence the ionization energy Slide 47 / 10 Slide 47 () / Which of the following elements would be expected to have a higher ionization energy than magnesium (Mg)? l a a l Na D K D K E E Slide 48 / 10 Slide 48 () / Which of the following correctly ranks the elements below in order of decreasing ionization energy? 18 Which of the following correctly ranks the elements below in order of decreasing ionization energy? Ne > O > N Ne > N > O Ne > N > O Ne > O > N H > He > Ne E H > He > Ne D Li > Mg > Ga D Li > Mg > Ga E Zn > Ga > r E Zn > Ga > r Na 17 Which of the following elements would be expected to have a higher ionization energy than magnesium (Mg)?

12 Slide 49 / 10 Slide 50 / 10 0 Which of the following pairs are correct in terms of relative first ionization energy and why? 19 Which of the following elements best fits the data provided below? Ionization e O- < Ne, due to smaller nuclear charge on oxide ion Ionization Energy 1st: X + IE --> X+ + e- 900 kj/mol nd: X+ + IE --> X+ + e kj/mol 3rd: X+ + IE --> X3+ + e- 14,850 kj/mol D Ne Li > Na, due to increased shielding in the Na atom E O Zn > u, due to a higher nuclear charge in zinc Li D l > S, due to the smaller nuclear charge in sulfur E ll of these Slide 50 () / 10 O- < Ne, due to smaller nuclear charge on oxide ion Li > Na, due to increased shielding in the Na atom 1 The second ionization energy will always be higher than the first. True False 0 Which of the following pairs are correct in terms of relative first ionization energy and why? Slide 51 / 10 Zn > u, due to a higher nuclear charge in zinc E D l > S, due to the smaller nuclear charge in sulfur E ll of these Slide 51 () / 10 True False TRUE have the lowest first ionization energies of the groups listed. lkali metals Transition elements Halogens D lkaline eath metals E Noble gases 1 The second ionization energy will always be higher than the first. Slide 5 / 10

13 Slide 5 () / 10 3 Of the choices below, which gives the order for decreasing first ionization energies? lkali metals l > S > l > r > Si Transition elements r > l > S > Si > l l > Si > S > l > r D l > S > l > Si > r E S > Si > l > l > r Halogens D lkaline eath metals E Noble gases have the lowest first ionization energies of the groups listed. Slide 53 / 10 Slide 53 () / 10 3 Of the choices below, which gives the order for decreasing first ionization energies? l > S > l > r > Si r > l > S > Si > l l > Si > S > l > r D l > S > l > Si > r E S > Si > l > l > r Slide 54 / 10 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 e (1s) e (s) Li (s) Li (1s) binding energy The PES spectrum clearly shows that the core electrons require the most energy to remove. It also shows that e has a higher 1st IE for the removal of the valence electrons than does Li. This is expected as e has a higher "Z". Slide 55 / 10 Ionization Energy and PES Let's interpret another PES spectra, this one of nitrogen and oxygen. Slide 56 / 10 Ionization Energy and PES lick to go to an interactive PES spectra database and answer the questions. O (p) O (s) Intensity N (p) O (1s) N (s) N (1s) binding energy Why is the binding energy of the electrons greater in He than H? Why is the N (p) peak greater than the O (p) peak? Similar shielding but answer greater "Z" move for N has a half-full "p" orbital increasing Which peak in the Li spectra represents the valence electrons? for answer themove ionization energy Why is the N(s) peak less than the O (s) peak? O has the higher nuclear charge move for answer move answer Peak with lowerfor binding energy Why is the valence peak binding energy less in Li than in H? answerlessens coulombic force Increased shielding due tomove core 1sfor electrons, Why is the core peak (1s) binding energy greater in Li than in H? Lithium has a higher nuclear charge so higher coulombic attractions move for"z" answer

14 Slide 57 / 10 Slide 57 () / 10 4 The following PES spectrum shows the valence "p" orbital peaks for Si and for. Which of the following would be TRUE? 4 The following PES spectrum shows the valence "p" orbital peaks for Si and for. Which of the following would be TRUE? binding energy The Si peak is of lower energy due to it's higher nuclear charge Intensity Intensity binding energy The Si peak is of lower energy due to it's higher nuclear charge The Si peak is of higher energy due to the increased shielding from core electrons The Si peak is of higher energy due to the object is a pull tab] increased shielding from core [This electrons The Si peak is of lower energy due to the increased shielding from core electrons 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 D The Si peak is of higher energy due to its higher nuclear charge Slide 58 / 10 Slide 58 () / 10 5 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? 5 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? False False True True TRUE Slide 59 / 10 Slide 59 () / 10 6 elow 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) 6 elow 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) 4p 4s 3s 3p 3d 3d 3s 3p 4p 4s ompared to Pd, the 3d peak in d would be found to the left of the 3d Pd peak ompared to Pd, the 3d peak in d would be found to the left of the 3d Pd peak ompared to Pd, the 3d peak in Rb would be of a higher binding energy due to lower nuclear charge ompared to Pd, the 3d peak in Rb would be of a higher binding energy due to lower nuclear charge ompared to Pd, the 3p peak in Kr should be found to the left of the 3p peak in Pd ompared to Pd, the 3p peak in Kr should be found to the left of the 3p peak in Pd

15 Slide 60 / 10 Slide 60 () / inding Energy inding Energy inding Energy 100 Intensity Intensity Intensity Intensity ased on the PES data below, what would be TRUE regarding atoms 1 and? 7 ased on the PES data below, what would be TRUE regarding atoms 1 and? inding Energy 10 I. tom 1 has a smaller atomic radii I. tom 1 has a smaller atomic radii II. tom has a larger first ionization energy II. tom has a larger first ionization energy [This object is a pull tab] III. oth atoms are in the same period III. oth atoms are in the same period IV. oth atoms are in the same group IV. oth atoms are in the same group II and III only I only D II and IV only 1 and III only E I, II, III, and IV Slide 61 / 10 I only II and III only D II and IV only 1 and III only E I, II, III, and IV Slide 6 / 10 Ionization Energy and Metallic haracter Metals are generally described as being able to lose electrons readily which promotes conductivity. Ionization Energy and Metallic haracter 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 Slide 63 / 10 Let's answer a few questions regarding metallic character. Pb has much more shielding due to moremove levelsfor of answer core electrons thereby causing it's electrons to be lost far more easily than that of. Notice that an element becomes more metallic as the shielding increases and as the nuclear charge - for a given level of shielding - decreases. Slide 64 / 10 Ionization Energy and Metallic haracter 1. Why is lead considered a metal and carbon a non-metal despite being in the same group? semi-metals or metalloids Si Ge Sn Pb Ionization Energy and Metallic haracter Let's answer a few questions regarding metallic character.. Which metal would we expect to be a better conductor of electricity? g or u u g g due to the higher amount of shielding, causing it to ionize more easily, thereby movemobile for answer creating electrons.

16 Slide 65 / 10 Slide 66 / 10 Ionization Energy and Metallic haracter pplication: Elements of Life F In order to form large stable, yet complex, molecules, the elements must not be able to lose electrons easily. + Ne D Xe E a + Slide 66 () / 10 8 Which of the following is the LEST metallic of those below? F Slide 67 / 10 9 Which of the following would be TRUE? The higher the ionization energy, the less metallic an element will be t Ne D Xe Interestingly, all metal atoms found in living things are in their ionic form (Mg+, a+, Zn+, etc.) t The most common elements in living things are,h,n,o,p, and S. Interestingly, these are all non-metals. Serotonin - brain hormone 8 Which of the following is the LEST metallic of those below? The lower the ionization energy, the less metallic an element will be For a given amount of core electron shielding, the higher the nuclear charge, the more metallic an element will be E a D oth and E oth and Slide 67 () / 10 The higher the ionization energy, the less metallic an element will be The lower the ionization energy, the less metallic an element will be For a given amount of core electron shielding, the higher the nuclear charge, the metallic an [Thismore object is a pull tab] element will be D oth and E oth and 30 Which of the following has the elements correctly ordered by increasing metallic character? Li < e < a < K < Ga Ga < a < K D Rb < s < s E Ga < s < a 9 Which of the following would be TRUE? Slide 68 / 10

17 Slide 68 () / 10 Slide 69 / Which of the following has the elements correctly ordered by increasing metallic character? Ionization Energy and Light s we have seen, EM radiation can provide the necessary energy to ionize an electron from an atom. Ga < a < K e- photon D Rb < s < s The higher the ionization energy, the higher the frequency of light needed to ionize the electron. E Ga < s < a Slide 70 / 10 Slide 71 / 10 Ionization Energy and Light Ionization Energy and Light What would be the necessary wavelength required to remove one of Neon's outermost p electrons? Which of the following elements would require the shortest wavelength to lose an electron? Si 1. Look up 1st IE of Neon = 081 kj/mol N. onvert to kj/atom = 081 kj mol Short wavelength means high energy so this would be the element with the largest ionization energy. "N" Si = 3.46 x 10-1 kj 4. onvert to v via E=hv --> v = E/h = 3.46 x J = 5. x /s 6.3 x J*s 8 5. onvert to wavelength via N has similar shielding as carbon but a higher nuclear charge so it would require the shortest wavelength to ionize an electron. v = c --> = c/v 3 x 10 m*s = 5.77 x 10 m = 57.7 nm x 1015 s Slide 7 / 10 Slide 7 () / Which of the following orbitals of calcium would require the highest frequency of light to ionize? 31 Which of the following orbitals of calcium would require the highest frequency of light to ionize? s p p s 3s 1 mol move for answer N move for answer "Z" x 6.0 x 103 atoms 3. onvert to J = 3.46 x J 3s D 3p D 3p E 4s E 4s a < K < Ga Li < e <

18 Slide 73 / 10 Slide 74 / 10 I Element Ionization Energy (kj/mol) Ga In In 558 He 37 Rb 403 D He I Ga 33 What frequency of light would be required to ionize the first electron of cesium (1st IE = 376 kj/mol)? 3 ased on the table of 1st ionization energies below, which element is likely to ionized by light with wavelength of 14 nm? E Rb Slide 74 () / What frequency of light would be required to ionize the first electron of cesium (1st IE = 376 kj/mol)? Slide 75 / 10 Periodic Law and Electronegativity s we know, atoms do not often exist in isolation. They form bonds with other atoms to make molecules and compounds. H 9.43 X /s O H water 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. Slide 76 / 10 Periodic Law and Electronegativity Let's compare the electronegativities of H and O within the water molecule. H O O Periodic Law and Electronegativity Trends in electronegativity for periods -4. H 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 Slide 77 / 10 H O Li S Na Se 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 for answer shielding

19 Slide 78 / 10 Slide 79 / 10 Periodic Law and Electronegativity Periodic Law and Electronegativity Trends in electronegativity for periods -4. The following electronegativity values will need to be memorized as this will aid in understanding bonding later on. H. N O F S 4.0 l.6 3. r Why do the noble gases not have published electronegativity values? They have a fullmove outerfor "s"answer and "p" system and do not form compounds. 3.0 Slide 80 / 10 Slide 80 () / Of the atoms below, is the most electronegative. Si l Si l Rb Rb D a D a E S E S 34 Of the atoms below, is the most electronegative. Slide 81 / 10 Slide 81 () / Which of the following EST explains why fluorine has a higher electronegativity than oxygen? 35 Which of the following EST explains why fluorine has a higher electronegativity than oxygen? F has a higher nuclear charge and similar shielding of O F has a higher nuclear charge and similar shielding of O F has a higher nuclear charge and less shielding than O F has a higher nuclear charge and less shielding than O F has the equivalent nuclear charge and less shielding than O F has the equivalent nuclear charge and less shielding than O D F has the equivalent nuclear charge and more shielding than O D F has the equivalent nuclear charge and more shielding than O E None of these E None of these

20 Slide 8 / 10 Slide 8 () / Which of the following groups of elements are ranked properly from lowest to highest electronegativity? 36 Which of the following groups of elements are ranked properly from lowest to highest electronegativity? H < < Li H < < Li < Si < Ge < Si < Ge D I < r < l D I < r < l E F < S < s E F < S < s Slide 83 / 10 H < Li < Na H < Li < Na D Slide 83 () / n element with a small electronegativity value is likely to have n element with a small electronegativity value is likely to have... high nuclear charge and a low amount of shielding high nuclear charge and a low amount of shielding low nuclear charge and a high amount of shielding Valence shell PES peaks with high binding energies Valence shell PES peaks with high binding energies low nuclear charge and a high amount of shielding D oth and D oth and E oth and E oth and Slide 84 / 10 Specific Groups of Periodic Table We will now examine six groups of the periodic table in more detail. Slide 85 / 10 lkali Metals They are highly reactive due to their extremely low ionization energies. s a result, they are found only in compounds in nature, not in their elemental forms. Group 1: lkali Metals Group : lkaline Earth Metals Group 3-1: Transition Metals Group 13/14/15: Metalloids Group 17: Halogens Group 18: Noble Gases They have low densities and melting points. In fact Li, Na, and K have densities so low, they'll float on water!

21 Slide 86 / 10 lkaline Earth Metals lkaline earth metals have higher densities and melting points than alkali metals. Their ionization energies are low, but not as low asthose of alkali metals so they are slightly less reactive. Slide 87 / 10 lkaline Earth Metals eryllium 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. an you explain why that would be? Slide 88 / 10 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 (u, Pt, g) are in this group. Slide 89 / 10 Transition Metals Some complex ions formed from transition metals and their colors. 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 90 / 10 Slide 91 / 10 Metalloids Metalloids There are six elements that are classified as metalloids: oron () 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. Silicon (Si) rsenic (s) Tellurium (Te) Germanium (Ge) ntimony (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. ircuits that form the basis for modern electronics are composed of doped metalloids like Si and Ge.

22 Slide 9 / 10 Halogens The halogens are prototypical nonmetals. They only require one more electron to have a full "s" and "p" and are therefore highly reactive. Slide 93 / 10 Halogens (at standard temp and pressure) Flourine is a colorless gas The name comes from the Greek words halos and gennao: salt formers. hlorine is a greenish gas romine is a brownish liquid Iodine is a purplish solid Slide 94 / 10 Noble Gases Slide 95 / n atom with a very high ionization energy and is a liquid at room temperature is most likely a: The noble gases have very high ionization energies. s a result, unlike the diatomic halogens, they are found as monatomic gases lkali metal lkaline earth metal transition metal Therefore, they are relatively unreactive. D Halogen E Noble gas Slide 95 () / Which of the following ranks the metals in order of increasing reactivity? Li < Na < Mg < K lkaline earth metal Mg < Li < Na < K lkali metal transition metal D D Halogen E Noble gas K < Li < Na < K D Li < Fe < Zn < u E None of these 38 n atom with a very high ionization energy and is a liquid at room temperature is most likely a: Slide 96 / 10

23 Slide 96 () / Which of the following elements would form colored complex ions? Li < Na < Mg < K F Mg < Li < Na < K o K < Li < Na < K D Li < Fe < Zn < u 39 Which of the following ranks the metals in order of increasing reactivity? Slide 97 / 10 a D l E None of these E Na Slide 97 () / Which of the following elements would form colored complex ions? D l F D Pb E Na E Y Slide 98 () / Which of the following elements would serve as a semiconductor? Slide 99 / 10 4 What would be the alkaline earth metal with the highest ionization energy? Li l Ge F D Pb E Y a e D E Ra o 41 Which of the following elements would serve as a semiconductor? Ge F Slide 98 / 10

24 Slide 99 () / 10 4 What would be the alkaline earth metal with the highest ionization energy? F e 43 Which would be the halogen with the smallest atomic radii? Ne Li l Slide 100 / 10 t D Pb D E Fr E Ra Slide 100 () / 10 Slide 101 / Which would be the halogen with the smallest atomic radii? Ne t D Pb F E Fr Slide 10 / 10 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.