Chemistry I Periodic Properties and Periodicity Class Notes

Chemistry I Periodic Properties and Periodicity Class Notes History - Periodic Table Periodic Properties Review of electron configuration & quantum numbers I. Know what valence electrons are -those electrons involved in chemical reactions - most commonly are the outermost electrons, but may involve more (like transition metals) II. Degenerate - term for orbitals within a subshell that have the same amount of energy. ex. Px Py Pz all have the same energy -an electron may enter any one of them, or move from one to another with no energy gain or loss. III. Electron cloud - shape is based on the probability distribution curve for the location of the electron IV. In addition to number designations ( n = 1, 2, 3,...), shells can also be designated by the letters: K, L, M, N,... V. Diamagnetic vs. Paramagnetic - arrangement of electrons within an atom - whether there all electrons are paired, or if there are unpaired electrons. diamagnetic - no unpaired electrons (note: this does not simply mean that there are an even number of electrons!) paramagnetic - at least one unpaired electron VI. Isoelectronic - same number of electrons in species ex. the chloride ion, the sulfide ion, and the potassium cation all have the same number of electrons a neutral argon atom. VII. Cosmic Rays - higher in energy than even gamma rays - source - stellar explosions? ============================================================================= History of the Periodic Table Development of the Periodic Law 1770's - discovery of oxygen by Lavoisier start of modern chemistry - within the next 60 years over 50 elements are separated and identified as being elemental substances. (by 1830, 56 elements had been identified) Chemists started looking for grouping methods first separation was metals from nonmetals -still common today. 1829 - Dobereiner - grouped elements in triads -Law of Triads - elements can be grouped so that the "middle element" has the "average properties" of the extremes. ex. Si (MW =28) and Sn (MW = 119) average MW = 73.5 Molecular Weight of Ge is 72.5 -- only rough approximations

-2-1860s A. Newland - Law of Octaves - showed recurring similarities of physical properties for each eighth element - recurring based on increasing atomic weight - elements showed a similarity to those above and below it ** Noble gases had not been discovered yet, so the eighth element after Na was K ** B. de Chancourtois - produced a spiral arrangement - first attempt at putting all elements in one arrangement 1869 Meyer (German) and Mendeleev (Russian) - both worked independently and came up with similar idea at the same time - within months of each other -question of publication Mendeleev - father of periodic table 1. mainly because his table was based on both physical and chemical properties; Meyer's table was based on physical similarities 2. took bold steps to predict that "gaps" should be left in table to allow for elements that had not yet been discovered. 3. predicted some physical and chemical properties of elements as yet undiscovered. One notable example: Germanium - he called it eka-silicon (after silicon) predicted (1871) disc. by Winkler (1886) Property Mendeleev Prediction Actual appearance gray grayish-white atomic weight 72 72.6 density 5.5 5.4 specific heat 0.31 0.31 Formula of oxide EO2 GeO2 Mendeleev made predictions for about several elements - books disagree - some say as few as 4, while others say as many as 14. Remember that the noble gases had not been discovered yet. Sir William Ramsey (Cambridge) discovered all inert gases from 1890-1900. - No longer called inert gases because of the synthesis of compounds with the heavier elements. These were "stuck" on the end of the table - mainly as a place to put them. Note that the date is only around the time of the discovery of the electron - before nuclear composition was known. *** Periodic Table was based on increasing atomic weights. ** The Modern Periodic Table is based on increasing atomic number - progressive increase in the number of protons. This was achieved by the work of Moseley and X ray diffraction

-3- Periodic Properties: Work of G.N. Lewis - proposed the idea of the octet rule even before the electron configuration theory was proposed by Bohr. Stability of an element was related to its electronic stability Effective nuclear charge and shielding discuss (see additional reading on pg. 10) Nuclear charge is the charge based on the number of protons in the nucleus. Because the atom has shells of electrons, some of these electrons shield the outer electrons of the nuclear charge. The actual (or effective) charge on valence electrons is called the effective charge. Zeff = Z - S S = # of inner shell electrons Z = atomic number = # of protons Note that the effective charge on magnesium is greater than the effective charge on sodium Zeff for Na = 1 Zeff for Mg = 2 Sodium has 1 proton pulling on 1 electron While the ratio of electron: proton is 1:1 in both sodium and magnesium, in each magnesium atom each electron feels the positive pull of two electrons!... Divisions within the periodic table: Groups or families - vertical columns - similar in chemical properties Series or Periods - horizontal rows - show progressive variation 1. I-A alkali metal - very active metals oxid.# = +1 "alkali" from ashes soda ash = sodium carbonate potash = potassium carbonate 2. II-A alkaline earth metals - active metals - found in earthen or mineral form. 3. VI-A - chalcogens 4. VII-A - halogens or halides 5. VIII or 0 inert gases, noble gases 6. I-B - coinage metals - because of unreactivity 7. VIII-B - iron family Note: except for Fe, Co, and Ni, the bottom two rows of this group (metals #44, 45, 46, and #76, 77, 78) are also known as the noble metals. They are called this because of their lack of reactivity. It is interesting to note that gold is not included in this group, because it is probably one of the least reactive of all metals! 8. A elements - called representative elements 9. B elements - called transition metals those elements with "d" subshells being filled are transition those elements with "f" subshells being filled are "inner" transition metals Inner transition metals - lanthanide series (sometimes called rare-earth) -actinide series - those past U (# 93 and beyond) called transuranium elements Some similarity between I-A and I-B true for all A and B groups. Shielding effect minimizes periodic variations of properties for the transition and inner transition metals since electrons are being put in inner subshells and outer "valence" electrons are not changed. *** NOTE: Hydrogen could actually be grouped in either I-A or VII-A

-4- Periodic Properties: - atomic radius/ volume - ionic size compared to atomic size - ionization energy - electron affinity - electronegativity - density - colors - melting points/ boiling points - metallic character - properties of metals, nonmetals, metalloids I. Atomic/Ionic Radius/Volume - atomic radius - measured from the nucleus to the outermost electron - not necessarily the highest energy electron. - volume is directly related to the radius - atom is considered to be spherical Periodic trends: - as we move down the table within a group, the atomic radius increases because we are filling shells that are successively farther away from the nucleus. This increase is most noticeable for the first four shells, but drops off as we move even further out. - as we move across the periodic table within a period, the radius tends to decrease. This decrease is due to the increased effective nuclear charge. Within a period, the distance to the outer shell is not increased, but the number of protons in the nucleus is increasing. - 'lanthanide contraction" -note that the radii of the 5d elements are similar to those of the 4d elements - they don't show the expected increase in size - this is due to the increased Zeffective of the element Ionic Radius For cations, the ions are smaller than the atoms. ex. Na + is smaller than the neutral Na atom For anions, the ions are larger than the atoms. ex. Cl - is larger than the neutral Cl atom Reasons: Cations are smaller than the neutral atoms because (1) the cation has lost its outermost electrons, making the new outermost electrons those of the previously filled shell; (2) the cation now has an excess of protons, so all electrons are pulled closer to the nucleus. Anions are larger than the neutral atoms because (1) the anion has gained electrons in order to achieve an octet configuration, making it more stable. This filled stability increases the electron repulsion, so electrons move further away; (2) the anion now has an excess of electrons, so all electrons are able to move away from the nucleus. Be able to compare ions to their isoelectronic noble gas in terms of size.

-5- II. IONIZATION ENERGY (IONIZATION POTENTIAL) - amount of energy required to completely remove an electron from an atom or ion. Important points: 1. Energy must always be added to remove electrons. 2. May have as many Ionization Potentials as there are electrons in a species. ie. Na would have 11 I.E's. + - M(g) -------------> M(g) + 1 e H is a positive value (endothermic) Note: 1. I.E. increases as you go across a period 2. I.E. decreases as you go down a family or group Lowest Ionization Energy - bottom left corner Highest Ionization Energy - top right corner Reminder: if an atom has a filled or half-filled subshell (or shell), there is a notable irregularity due to greater shielding and/or increased electron repulsion. Periodic properties will reflect these changes (jumps). There are some noticeable irregularities: 1. from II-A to III-A 2. from V-A to VI-A 3. at VIII (inert gases) 4 from II-B to III-A These are in units of kcal/mole: 1) Na (119); ** Mg (176); ** Al (138) K (100); ** Ca (141); ** Ga (138) 2) for C (260); ** for N (335) **; for O (314) Note that the magnitude of the half-filled irregularity decreases as n increases ex. for 4th shell Ge (192); As (226); Se (225) 3) F (402) Ne high at 497 (not actually an irregularity - predictably high value because of stable octet configuration) 4) Ni (176); Cu (178); ** Zn (217); ** Ga (138) NOTE also the similarities of I.E. for transition metals For 4th series, elements show an average jump of about 50 kcal - for transition metals in row #4 Sc (151); Ti (157); V (155); Cr (156); Mn (171)? reason for small jump? THE IONIZATION ENERGIES FOR THE LANTHANIDE SERIES ARE EVEN CLOSER

-6- What we have been discussing are the first ionization energies; there is an ionization energy for each electron in an atom. The electron with the highest ionization energy is the one closest to the nucleus 2 reasons 1. those closest to the nucleus have the farthest distance to go to be separated from an atom. 2. as successive electrons are removed, with the number of protons staying the same, the electrons must overcome progressively larger nuclear charges. Comparison of different ionization energies: all measurements are in kcal/mol 1st 2nd 3rd 4th Na 119 1090 1652 2281 Mg 176 347 1848 2519 Al 138 434 656 2767 ==========================================================================. III. Electron Affinity Electron affinity: Pay close attention because my definition and values are different from those in the book. Electron affinity: A measure of the "affinity" that an atom has for an electron. The number value is an indication of the atom's preference for the electron. In contrast to ionization energy, the values of electron affinity may be positive, negative, or zero. Even the magnitude of the positive values indicates how strongly the atom wants to accept an electron PERIODIC TRENDS: 1. Metals have a tendency to lose electrons; nonmetals tend to gain electrons. Nonmetals would be expected to have a high electron affinity. 2. As we move across the periodic table, we move from metals to nonmetals. The electron affinity should increase as we move from left to right. 3. Because of the large size of the atoms at the bottom of a group compared to the atomic size at the top of the group, it would be easier (take less energy) to attach an electron to a metal with a large atomic radius than one with a small radius. It also means there is less affinity for an electron. One periodic trend, then, is that electron affinity decreases as we move down a group.

-7- Electron Affinity (cont.) For nonmetals, there is a natural tendency to want to gain electrons. Even for many metals, their electronic structure allows them to accept electrons even if it isn't their natural tendency. 1. If the E.A. is a large positive value, there is a strong attraction (affinity) for an electron. 2. If the E.A. is a small positive value, there is a weak attraction (affinity) for an electron, but it will accept an electron. 3. If the E.A. is a negative value, energy must be added to force an electron to be accepted by an atom. 4. A value of zero simply indicates that there is no tendency to gain the electron. EMPHASIZE: DIFFERENCE FROM BOOK - LOOKS AT THE D ENERGY INVOLVED IN THE PROCESS: If A + e - -----------> A - + energy (EXO) If A + e - + energy -----------> A - (ENDO) IMPORTANT POINTS: 1. Even though the natural tendency for metals is to lose electrons, we must consider if there are orbitals free to accept electrons. Consider the sodium atom: configuration is [Ne] 3s' This means there is one electron in the single s orbital, but it also means there is room for one more electron. Na would be predicted to have a small positive value. 2. In contrast, the configuration of Mg is [Ne] 3s 2 This means there are two electrons in the single s orbital, and it also means there no room for any more electrons. To add an electron to magnesium means that there must be enough energy supplied to force the electron into one of the as-yet-unoccupied 3p orbitals. Note: AREAS OF NEGATIVE OR ZERO VALUES 1. at II-A 2. at II-B 3. at VIII All of these are due to the completion of a subshell. Note the low values at V-A and VII-B - half-filled subshells.

-8- IV. ELECTRONEGATIVITY - the tendency for an atom to attract electrons to itself when it is chemically combined with another element. - developed by Linus Pauling - based on the values of electron affinity and ionization energies - not experimental data like EA and IE. - high EN - nonmetals -low EN - metals -Note: there are no EN values for noble gases because these gases have no tendency to react! Trends: 1. lowest value - bottom left corner (Fr = 0.70) highest value - top right corner (F = 4.0) Actually the values were set with fluorine at 4.0 and all other values were calculated based on fluorine. 2. EN values tend to increase as you move across a period. 3. EN values tend to decrease as you move down a group or family. 4. Note the lack of change as you progress through the transition metals. V. DENSITY - remember that density is the ratio of mass to volume - mass increases regularly as you progress across a period - volume increases as you move down a group, but decreases you move across a period, but the amount of volume change is greater for the first four periods, but is not as great for the larger periods. - highest density then would be predicted to be in the bottom right corner, but because of the lanthanide contraction and the increased mass due to addition of the "f" subshell elements, the elements with the greatest density fit in the middle of the bottom, ie. Ir and Os have the greatest density. - review the significance of the lanthanide contraction Trends: V-formation - most dense elements are at the bottom of the V - lightest elements are hydrogen and helium VI. Melting / Boiling point Related to (1) size of atom; (2) electron arrangement and number of protons; (3) amount of attraction between atoms of the same element. Trends are not very regular: 1. Note that for halogens - as you move down the group, the elements move from the gaseous state to solids. 2. As you move across a period, you move from metals to nonmetals - MP tends to decrease 3. For the most part, MP and BP are related - high MP generally means high BP. 4. Elements of note: carbon - highest melting point of all elements tungsten - highest melting point of all metals Note: while tungsten also has a high BP, carbon's BP is not much higher than its MP.

-9- VII. Colors - transition metal salts are colored due to the movement of "d" electrons; all other salts are white - only two colored metals - Cu and Au - all others are silvery Note: all metals when ground into a fine powder appear black; only aluminum remains silvery. VIII. metallic / nonmetallic character - tendency of a metal is to lose electrons / nonmetals tend to gain electrons - would look at properties such as ionization energy and electron affinity francium should be the most metallic because of its low ionization energy, while fluorine would be the best nonmetal. Note: because of the EN values, all elements have a relative tendency for metallic and nonmetallic properties. In terms of electrical conductivity, silver is the best with copper coming in a close second. Gold is the third best electrical conductor. Notice that these three make up the coinage metals. Aluminum is 4th, by the way. IX. acid/base properties - metals tend to be basic while nonmetals tend to be acidic - especially true of oxides of the elements. - metalloids would then be amphoteric - able to react with either acids or bases. Remember that metallic oxides are classified as basic anhydrides; while nonmetallic oxides are acidic anhydrides. - for metals that have more than one charge, the higher oxidation state produces acidic properties Ex. chromium has 3 charges - Cr +6 is acidic; Cr +3 is amphoteric; Cr +2 is basic X. Oxidation Numbers (ionic charge) & Valence Electrons - Oxidation # s: the number of electrons that are typically gained/lost in a chemical bond in order to achieve an octet - Valence Electrons: # of electrons in outermost energy level - According to the Octet Rule, atoms are most stable (energetically) when 8 electrons occupy the valence shell - Ox. No and valence electrons for transition metals will not need to be memorized. Group # Ox. # # Valence e - 1 +1 1 2 +2 2 13 +3 3 14 +/-4 4 15-3 5 16-2 6 17-1 7 18 0 8

-10- An atom of Hydrogen consists of a single proton surrounded by an electron that resides in a spherical 1s orbital. Recall that orbitals represent probability distributions meaning that there is a high probability of finding this electron somewhere within this spherical region. This electron, being a negatively charged particle, is attracted to the positively charged proton. Now consider an atom of helium containing two protons (there are neutrons in the nucleus, too, but they are not pertinent to this topic) surrounded by two electrons both occupying the 1s orbital. In this case, and for all other many electron atoms, we need to consider not just the proton - electron attractions, but also the electron - electron repulsions. Because of this repulsion, each electron experiences a nuclear charge that is somewhat less than the actual nuclear charge. Essentially, one electron shields, or screens the other electron from the nucleus. The positive charge that an electron actually experiences is called the effective nuclear charge, Z eff, and Z eff is always somewhat less than the actual nuclear charge. Z eff = Z - S, where S is the shielding constant The two figures represent different instantaneous positions for the two electrons in an atom of helium. The electrons both occupy a 1s orbital meaning that the time-averaged positions of the electrons can be represented by a sphere. At any moment, however, we could envision the two electrons as being on opposite sides of the nucleus in which case they poorly shield each other from the positive charge. At a different moment, one electron may be between the nucleus and the other electron, in which case the electron farther from the nucleus is rather effectively shielded from the positive charge. Considering helium, you might initially think that Z eff would be one for each electron (i.e., each electron is attracted by two protons and shielded by one electron, 2-1 = 1), but Z eff is actually 1.69. To determine Z eff, we do not subtract the number of shielding electrons from the actual nuclear charge (that would always result in a value of Z eff of one for any electron in any atom!), but rather we subtract the average amount of electron density that is between the electron we are concerned with and the nucleus. There are rules we can use to estimate the value of S and thus determine Z eff, but we do not need to go into such details for this course. Penetration An electron in an s orbital has a finite, albeit very small, probability of being located quite close to the nucleus. An electron in a p or d orbital on the other hand has a node (i.e., a region where there cannot be any electron density) at the nucleus. Comparing orbitals within the same shell, we say that the s orbital is more penetrating than the p or d orbitals, meaning that an electron in an s orbital has a greater chance of being located close to the nucleus than an electron in a p or d orbital. For this reason, electrons in an s orbital have a greater shielding power than electrons in a p or d orbital of that same shell. Also, because they are highly penetrating, electrons in s orbitals are less effectively shielded by electrons in other orbitals. For example, consider an atom of carbon whose electron configuration is 1s 2 2s 2 2p 2. The two electrons in the 1s orbital of C will do a better job of shielding the two electrons in the 2p orbitals than they will of shielding the two electrons in the 2s orbital. This means that for electrons in a particular shell, Z eff will be greater for s electrons than for p electrons. Similarly, Z eff is greater for p electrons than for d electrons. As a result, within a given shell of an atom, the s subshell is lower in energy than the p subshell which is in turn lower in energy than the d subshell.