Slide 1 / 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. Click to go to website: www.njctl.org Slide 2 / 102 AP Chemistry Periodic Trends www.njctl.org Slide 3 / 102 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.
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 DEFINITION Ionic Charge charge of common ion formed by that element Atomic/Ionic Radii Distance from the nucleus to outermost electron Density Ratio of Mass/Volume Ionization Energy Energy required to remove valence electron Disposition to have metallic characteristics - ie. conduct Metallic Character electricity Measure of attraction for electrons when the atom is Electronegativity sharing electrons in a molecule. Slide 5 / 102 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. H He Li Be B C N O F Ne Na Atomic 1 2 3 4 5 6 7 8 9 10 11 Number Ionic +1,-1 NA +1 +2 +3 +4 -3 -2 -1 NA +1 Charges 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! Slide 6 / 102 The Periodic Law The pattern can be easily visualized on a graph, particularly as we move past the first 11 elements! +4 +3 ion charge +2 +1 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 -1 atomic number -2 -3
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. Principal Quantum Lose/ Electron Ionic Number (N) Element Gain Configuration Charge of valence electrons electrons gain 1 -1 H 1 1s 1 The pattern recurs with lose 1 +1 He 1 1s 2 NA NA every increase in the Li 2 [He]2s 1 lose 1 +1 principal quantum Be 2 [He]2s 2 lose 2 +2 number. This means B 2 [He]2s 2 2p 1 lose 3 +3 every time a new shell C 2 [He]2s 2 2p 2 lose 4 +4 of electrons is filled, the pattern repeats! 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 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 for answer He = 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 gas with a full principal energy level. move for answer Slide 9 / 102 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 move for answer energy level with 1 electron, just as sodium was beginning the 3rd with 1 electron.
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 transition elements Slide 11 / 102 The Periodic Law and the Quantum Model d orbital +3 +4 +5 +6 +7 +3 +3 +2 +1 +2 +4 +5 +3 +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. 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 12 / 102 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 are readily lost yielding the +2 ion. move for answer A half-full "d" orbital is quite stable so Fe will lose 1 d orbital electron as well to yield the +3 ion.
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 move for answer The "d" orbital is full so only 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 Answer 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. True False Answer
Slide 16 / 102 3 Which of the following BEST explains why O and S both form ions with a -2 charge? A They both have the same atomic number B They are both in the same period C They both have the same electron configuration Answer 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 Answer 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 Answer 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 radius 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. **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 20 / 102 The Periodic Law and Atomic Radii Let's examine the trend in atomic radii for the first 18 elements. Na 200 Li radius (pm) 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 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 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) Carbon (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.
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