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Packet 8 Periodic Properties of the Elements last updated: 5/24/2017 CHE 170 Packet 8 - 1 Concept Area I: Terminology atomic orbital energy diagram Aufbau principle degenerate diamagnetic effective nuclear charge, Z* electron affinity electronic configuration ferromagnetic filled subshell ground state half-filled subshell Hund’s rule inner-shell/core electrons ionization energy Noble gas notation octet rule orbital paramagnetic shell shields, shielding subshell valence electrons CHE 170 Packet 8 - 2 Concept Area II: Magnetism and Atomic Structure a. You should know how to determine if a substance is paramagnetic or diamagnetic. b. You should understand how magnets can be formed from ferromagnetic elements. CHE 170 Packet 8 - 3 Paramagnetic and Diamagnetic paramagnetic: Liquid oxygen is attracted to strong magnets because oxygen is paramagnetic – it has unpaired electrons. similar to Tro page 379 CHE 170 Packet 8 - 4 Paramagnetic and Diamagnetic diamagnetic: Liquid nitrogen is not attracted to strong magnets because nitrogen is diamagnetic – it has no unpaired electrons. CHE 170 Packet 8 - 5 Paramagnetic and Diamagnetic How can we show that a sample is paramagnetic? iron(III) oxide Nifty mnemonic devise for para and diamagnetic? CHE 170 Packet 8 - 6 How does paramagnetism allow the formation of magnets for ferromagnetic elements? Well, if all of those unpaired electrons can be aligned…… all of their little magnetic fields would add up to one larger field – a magnet! We can align all the little electrons by exposing them to a very strong magnetic field. CHE 170 Packet 8 - 7 Concept Area III: Atomic Orbital Energy Level Diagrams a. You need to remember what the four quantum numbers tell you (names of orbitals, how many degenerate orbitals). b. You need to know how many orbitals of each type are present in a shell or subshell of specific energy. c. You need to know how many electrons can go into each orbital, subshell and shell. d. You should know that the different energies of atomic orbitals control the order in which they fill up with electrons. e. You should be able to use the Aufbau Principle and know when and why to break it for certain elements. f. You must be able to draw an atomic orbital energy level diagram or write an electron configuration for any atom or ion CHE 170 Packet 8 - 8 Lots of Things to Remember! shell: atomic orbitals of similar energy are grouped together by their principal quantum number, n. The energy level is shown by counting numbers: . subshell: atomic orbitals of the same type/shape are grouped together, this is given by the angular momentum quantum number, ℓ. What are the names of the different orbitals? How many degenerate orbitals for each? . . atomic orbital: this is where the electrons live, this is given by the magnetic quantum number, mℓ. How many electrons can live in each kind of orbital? . . valence electrons: those electrons that are in the outermost shell of the atom CHE 170 Packet 8 - 9 Let’s review everything! CHE 170 Packet 8 - 10 Generic Atomic Orbital Energy Diagram Handout Generic Atomic Orbital Energy Diagram 6d 5f 7 6 7s 6p 5d 5 4f 6s 5p 4d 4 5s 4p 3d Energy 3 4s 3p 3s 2 1 2p 2s 1s Note that all the degenerate orbitals are written on the same line to indicate that they have the same energy and are in the same subshell. Note that the subshells have a number by them to indicate what shell they are in. Note that the energies of the different shells overlap. Note: These energies are not plotted to scale. In fact, the further up the diagram we go, the orbitals get closer in energy. CHE 170 Packet 8 - 11 How do we draw one of those things?! First, we have to remember the order that we fill the orbitals: Second, easy as counting! One subshell in the first shell: Two subshells in the second shell: Three subshells in the third shell: Four subshells in the fourth shell: Finally, we have to remember how many degenerate orbitals for each kind: Now, we use this mnemonic device… CHE 170 Packet 8 - 12 How to Easily Create an Atomic Orbital Energy Diagram Handout How do we know the energy order to list the orbitals? This handy mnemonic device! 7s 6s 5s 4s 3s 2s 1s 7p 6p 5p 4p 3p 2p 7d 6d 5d 4d 3d 7f 7g 7h 7i 6f 6g 6h 5f 5g 4f CHE 170 Packet 8 - 13 Looking at the periodic table, when does a new subshell start filling? CHE 170 Packet 8 - 14 The Aufbau Principle 1. Electrons are placed in the lowest energetically available (i.e. unfilled) orbital. 2. An orbital can hold at most two electrons. 3. If two or more degenerate orbitals are available (for example: p, d etc.) then electrons should be spread out before they are paired up (Hund's rule). CHE 170 Packet 8 - 15 Hund’s Rule Just like most folks prefer to sit alone on a city bus, electrons prefer to be alone (fewer repulsions) in an orbital. So, if they have the choice, they stay alone! CHE 170 Packet 8 - 16 Time to use Hund’s Rule! 1. So, put 3 electrons in a p subshell: 2. Now, 4 electrons in a p subshell: 3. Now, try 7 electrons in a d subshell: CHE 170 Packet 8 - 17 Putting it all together…. Writing atomic orbital energy diagrams and electronic configurations! 1. Determine how many electrons the atom has (same as atomic number on periodic table) 2. Use an energy level diagram and put in electrons one by one using the Aufbau principle spdf notation for electron configurations 3. If requested, con- for H, atomic number = 1 vert diagram to an 1 # of s 1 electron configuraelectrons tion or write it value of ℓ value of n directly. CHE 170 Packet 8 - 18 Let’s write the atomic orbital energy diagram and then write the electron configuration! Let’s use Chlorine atom as our example: How many electrons? Diagram: 3d 4s 3 3p 3s 2p 2 Energy 1 2s 1s So, it would be: Or with noble gas notation: CHE 170 Packet 8 - 19 Now try a few with a neighbor… Write atomic orbital energy diagram (use the handout) and the electronic configuration for the following atoms: 1. Carbon C 2. Fluorine F 3. Sodium 4. Iron Na 5. Zirconium CHE 170 Packet 8 - 20 What is the correct orbital order for electronic configurations? Let’s use krypton to answer this. Its complete electronic configuration listed on the previous slide was… 1s22s22p63s23p63d104s24p64d25s2 The orbitals were listed from innermost to outermost shell. Instead, we could list them in the order that they energetically fill… 1s22s22p63s23p64s23d104p65s24d2 Thus, it really doesn’t matter which of these two ways that the orbitals are listed! However, wait a few slides and we’ll see when we do cations that the shell order is better. CHE 170 Packet 8 - 21 Exceptions to the Rule! Remember the Aufbau principle? Guess what, some elements break that rule! There’s another competing rule: full subshells and half-full subshells are more stable than other configurations. So, for d or higher subshells, we can fill one electron out of order to get to a “full” or “half-full” subshell. Note: Cr, Cu, Mo, Ag, etc. Some of the other transition metals do further rule-breaking stuff (like Pd and Pt). Don’t worry about why but do realize there are some more details we aren’t learning! CHE 170 Packet 8 - 22 Valence Electrons Where are the valence electrons? Why are valence electrons important? Because they are located on the outside of the atom. These are the electrons that are removed to make cations – need to know how many valence electrons so you know the maximum amount that you can remove. To make anions, electrons are added to the valence shell – need to know how many valence electrons so you know how much room you’ve got! Is there an easy way to determine number of valence electrons? Yes! Let’s go to the next slide to see how… CHE 170 Packet 8 - 23 How many valence electrons? The periodic table is very useful to tell how many valence electrons we have in an atom! For main group elements, simply count from left to right: 1,2,3,4,5,6,7,8. (Note, use the group numbers with “A” after them.) Transition metals are a bit trickier because of the d orbitals. They should all be two, but because of the rule breaking they do, they have to be determined in context of a compound. Notes page has the labeled periodic table. CHE 170 Packet 8 - 24 Electron Configurations of Ions How do we write an electron configuration for an anion? Simple, really! Anions are just like atoms. Write the atom’s configuration, and then, add in the additional electron(s) to the next available spot(s). Let’s try one! CHE 170 Packet 8 - 25 What is the electron configuration for a chloride ion? How many electrons in the atom?17 Diagram for the atom: 3d 4s 3 3p 3s 2p 2 2s Energy 1 1s So, e.c. for atom would be: We actually have a chloride ion. What is the symbol & charge? So, we need to add how many electrons? So, e.c. for ion is: Or in noble gas notation: Or with noble gas notation: CHE 170 Packet 8 - 26 Electron Configurations of Ions How do we write an electron configuration for an cation? This is a bit more complicated! In a cation we remove electrons from an atom. We have to remove the outermost electrons on the atom. Now, the outermost electrons sometimes are at lower energy. Therefore, we’re not always removing the last electrons we added! Let’s try one! CHE 170 Packet 8 - 27 What is the electron configuration for a iron(II) ion? How many electrons in the atom? Diagram for the atom: 3d 4s 3 3p 3s 2p 2 2s Energy 1 1s So, e.c. for atom would be: Or with noble gas notation: We actually have a iron(II) ion. What is the symbol & charge? So, we need to remove how many electrons? Now, what is the outermost subshell? So, e.c. for ion is: Or noble gas notation: CHE 170 Packet 8 - 28 Why don’t you try a few… Write the electronic configurations for the following. 1. fluoride ion 2. oxide ion 3. calcium ion 4. germanium(II) ion 5. vanadium(II) ion Now that we’ve done those. Are any of these ions paramagnetic? CHE 170 Packet 8 - 29 Concept Area IV: Applications a. You should remember that the periodic table is arranged so that elements with similar properties are by each other. b. You need to be able to calculate an approximation of the effective nuclear charge of an atom or ion and use it along with atomic theory to explain some of trends on the periodic table. c. You need to recognize the role that ionization energy and electron affinity play in the chemistry of the elements and be able to explain the observed trends. d. You should be able to predict whether an atom becomes larger or smaller when it becomes an ion. e. You should be able to explain why we can use the periodic table to predict the most commonly formed ion for main group elements. CHE 170 Packet 8 - 30 Mendeleev (from Russia, 1834-1907) and the Periodic Table He ordered elements by atomic mass. When he did so, he saw a repeating pattern of properties. Periodic law: When the elements are arranged in order of increasing atomic mass, certain sets of properties recur periodically. So, he arranged elements with similar properties in the same column. Then he used the patterns he saw to predict properties of undiscovered elements (Ga and Ge). When atomic mass order did not fit observed properties, he reordered (Te and I) CHE 170 Packet 8 - 31 Periodic Trend for the Reactivity of Alkali Metals with Water Lithium Sodium Potassium And, for more alkali metal fun: http://youtu.be/uixxJtJPVXk http://youtu.be/uixxJtJPVXk And, for more sodium fun: http://youtu.be/HY7mTCMvpEM http://youtu.be/HY7mTCMvpEM CHE 170 Packet 8 - 32 What versus Why Mendeleev’s periodic law allows us to predict what the properties of an element will be based on its position on the table. But, it doesn’t explain why the pattern exists. Side note – true of all laws. Tells what we observe without explaining why it is that way. Quantum mechanics is a theory that explains why the periodic trends in the properties exist. electron configurations/atomic theory effective nuclear charge, Z* CHE 170 Packet 8 - 33 Effective Nuclear Charge Effective nuclear charge, Z*, of an electron is how much of a “pull” from the nucleus it feels. The greater the pull from the nucleus, the less repulsive force it feels from the other electrons. Electrons with less repulsive force will be at lower energy. The text goes into great detail about how electrons in s orbitals shield better than those in p orbitals shield better than d orbitals and so on. Tro page 274 CHE 170 Packet 8 - 34 Calculating Effective Nuclear Charge We’re going to ignore that fact that different electrons shield differently for our calculations and discussions. However, whenever an atom seems to buck the trend, if we account for different electrons shielding differently, those atoms aren’t so strange after all. Now, the effective nuclear charge, Z*, of a valence electron in an atom can be approximated by subtracting the number of inner-shell or core electrons from the number of protons: Z*=Z – S (where Z is actual nuclear charge and S is the charge screened by other electrons) CHE 170 Packet 8 - 35 Calculating Effective Nuclear Charge What’s an inner-shell or core electron? Remember, the atomic number tells us the total number of protons and the total number of electrons in an atom. Remember to determine number of valence electrons, we count the main group elements left to right. CHE 170 Packet 8 - 36 Effective Nuclear Charge explains why atoms get smaller across a period. Z* on Li’s atomic mass = 6.941amu. C’s atomic mass = 12.011 amu. Ne’s atomic mass = 20.179 amu. CHE 170 Packet 8 - 37 Why do atoms get smaller from left to right across the periodic table? So, size decreases across a period owing to increase in Z*. Notice that neon had an effective nuclear charge of +8 versus lithium at +1. Therefore, the valence electrons on Ne “feel” more of a pull towards the nucleus, and the valence electrons on Li are more shielded from the nucleus and “feel” less of a pull. Since the valence electron for Li is more shielded, Li is larger than Ne, even though Ne has more electrons and more mass than Li! Tro page 285 CHE 170 Packet 8 - 38 Atomic Radii Trends The atomic radii are given here in picometers for main group elements. So, atomic radii . across a period. This is because the effective nuclear charge, Z*, . across a period. CHE 170 Packet 8 - 39 Atomic Radii Trends - plotted Notice the periodic trend in atomic radius, starting at a peak with each alkali metal and falling to a minimum with each noble gas. Tro page 285 CHE 170 Packet 8 - 40 Smaller atoms hang on to their electrons more, why? If two atoms have the same effective nuclear charge, why is it more difficult to remove an electron from the smaller atom? Well, opposites . Plus, the closer they are together, the that is! Thus, smaller atoms hang on to their electrons than larger atoms of the same effective nuclear charge because the valence electrons are to the nucleus in the smaller atoms. This is also manifested by smaller atoms having ionization energies. Note: although ionization energies are determined in the gas phase, we can use those values to decide whether an atom in the solid/liquid phase would be more likely to gain or lose an electron. We just don’t have a specific value for solid/liquid. Tro page 295 CHE 170 Packet 8 - 41 Smaller atoms hang on to their electrons more. CHE 170 Packet 8 - 42 So, why do we have dips in our ionization energies? Remember how completely full and half-full sublevels are more stable? So, each of those dips occurs where if we remove an electron, we’ll have a more stable configuration! Similar image on Tro page 294 that includes transition metals CHE 170 Packet 8 - 43 Why do atoms on the top-right side of the table (ignoring noble gases) have a higher electron affinity? Well, the more positive the Z*, the more likely the atom can capture a wayward electron if it is not energetically bad for it to do so. Again, even though these values are measured in the gas state, we can generalize that information to other states. Tro page 297 CHE 170 Packet 8 - 44 CHE 170 Packet 8 - 45 Ionic Size How about if we add valence electrons; would the ionic radii for an anion be larger or smaller than the atomic radii? Why? CHE 170 Packet 8 - 46 Ionic Size How about if we take away valence electrons; would the ionic radii for a cation be larger or smaller than the atomic radii? Why? CHE 170 Packet 8 - 47 Ion size compared to atom size. Similar images on Tro pages 291 and 292 CHE 170 Packet 8 - 48 What charge doth an ion have? Using the periodic table to predict the charge on an ion formed from an atom. Be2+ B3+ C4± Si4+ Ga3+ Ge4+ Here are some charges on some common monatomic cations and anions. Metals usually form cations and nonmetals usually form anions. Learn how to get the periodic table to tell you the most commonly formed ion. The transition metals are trickier because of the d shells; their charges must be determined in context. CHE 170 Packet 8 - 49 Noble Gases and the Octet Rule The noble gases have eight valence electrons. Except for He, which has only two electrons. We know that the noble gases are especially nonreactive. He and Ne are practically inert. The noble gases are so nonreactive because the electron configuration of the noble gases is especially stable which means low energy. Other atoms would like to be at lower energy too, so… CHE 170 Packet 8 - 50 They try and look like one: Noble Gases and the Octet Rule Once we have the atomic orbital diagram (or electronic configuration), simply decide if it would be easier to add or take away electrons to look like a noble gas and reach stability/lower energy. We call this the octet rule because most noble gas configurations have eight valence electrons. Eight, two or zero electrons in the valence shell is usually the magic number, why? Elements in the 3rd period can sometimes act like they have more than eight valence electrons, why? CHE 170 Packet 8 - 51 The End of Packet 8 Any Questions? Homework assignment! The Lewis Dot Handout is due when we get to Lewis Dot structures in the next chapter! CHE 170 Packet 8 - 52