Chapter 8 Periodic Table Mendeleev
Transcription
Chapter 8 Periodic Table Mendeleev
Periodic Table Mendeleev • ordered elements by atomic mass • 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. • put elements with similar properties in the same column • used patterns to predict properties of undiscovered elements • where atomic mass order did not fit other properties, he reordered by different properties Chapter 8 Periodic Properties of the Elements The beginning of a periodic Trend NM H 1.0 H2 M Li 6.9 Li2O b M LiH 9.0 Be Na2O M b M Na 23.0 NaH 24.3 M K2O b K 39.1 What vs. Why H2O a/b KH BeO NM a/b B BeH2 10.8 MgO b Al MgH2 27.0 M Ca 40.1 CaO b CaH2 C BH3 12.0 M Mg B2O3 NM a Al2O3 a/b M/NM Si AlH3 28.1 ? CO2 NM a N N2O5 NM a CH4 14.0 NH3 16.0 NM P4O10 NM a SiO2 a P NM F H2O 19.0 S SO3 a HF NM Cl Cl2O7 a PH3 32.1 H2S 35.5 HCl M/NM As2O5 NM a/b SeO3 NM a Br2O7 a SiH4 31.0 ? O2 O As 74.9 AsH3 79.0 As H2Se 79.9 Br HBr M = metal, NM = nonmetal, M/NM = metalloid • Mendeleev’s periodic law allows us to predict what the properties of an element will be based on its position on the table. • It doesn’t explain why the pattern exists. • Quantum mechanics and its theory of electron configurations is what explains why the periodic trends in the properties exist. a = acidic oxide, b = basic oxide, a/b = amphoteric oxide 2 4 First: the Electron Spin • Electron pairs residing in the same orbital are required to have opposing spins. – This causes electrons to behave like tiny bar magnets. Electron Configurations • An electron configuration of an atom is a particular distribution of electrons among available subshells. – The notation for a configuration lists the subshell symbols sequentially with a superscript indicating the number of electrons occupying that subshell. – For example, lithium (at. # 3) has two electrons in the “1s” subshell and one electron in the “2s” sub shell is: Lithium: 1s2 2s1 Pauli Exclusion Principle • No two electrons can have the same four quantum numbers. – In other words, an orbital can hold at most only two electrons, and then only if the electrons have opposite spins. 6 Summary of Orbital Types • The maximum number of electrons and their orbital shells are: Maximum Number of Number of Sub shell Orbitals Electrons s 1 2 p 3 6 d 5 10 f 7 14 8 Order for Filling Atomic Subshells The Aufbau Principle • Every atom has an infinite number of possible electron configurations. – The configuration associated with the lowest energy level of the atom is called the “ground state.” – Other configurations correspond to “excited states.” – The Aufbau Principle is a scheme used to reproduce the ground state electron configurations of atoms by following the “building up” order. – You need to remember the order of filling the orbitals (next slide) 9 Another Way to Remember the Filling Order 1s 2s 3s 4s 5s 6s or 2p 3p 4p 5p 6p 3d 4d 4f etc 5d 5f etc 6d 6f etc 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s etc Example of Subshell Filling Order • With sodium (Z = 11), the 3s subshell begins to fill. Z=11 Sodium 1s22s22p63s1 or [Ne]3s1 Z=12 Magnesium 1s22s22p23s2 or [Ne]3s2 • Then the 3p subshell begins to fill. • Progressing from left to right across each consecutive row will give the correct filling order… 1s,2s,2p,3s,3p… 11 Z=13 : : Z=18 Aluminum 1s22s22p63s23p1 or [Ne]3s23p1 Argon 1s22s22p63s23p6 or [Ne]3s23p6 12 Abbreviated form Using the Aufbau Principle • The “building up” order corresponds, for the most part, to increasing energy of the subshells. – By filling orbitals of the lowest energy first, you usually get the lowest total energy (“ground state”) of the atom. – Remember, the number of electrons in the neutral atom equals the atomic number, Z. “Abbreviated form” of electron configurations uses the closet noble gas to indicate that its inner electrons have filled that particular noble gas configuration. 13 Electronic config. can be written in 1) order of filling or 2) in order of increasing principal quantum #. The table above follows the latter. Ground State Electron Configuration Examples Sample Problem Valid and Invalid Electron Configurations • Which ground-state electron configurations are INCORRECT? WHY? Z=4 Beryllium 1s22s2 or [He]2s2 Z=3 Lithium Z=5 Z=6 Z=7 Z=8 Z=9 Z=10 Boron Carbon Nitrogen Oxygen Fluorine Neon 1s22s1 or [He]2s1 1s22s22p1 1s22s22p2 1s22s22p3 1s22s22p4 1s22s22p5 1s22s22p6 or or or or or or [He]2s22p1 [He]2s22p2 [He]2s22p3 [He]2s22p4 [He]2s22p5 [He]2s22p6 1. 2. 3. 4. 5. 15 Abbreviated forms 15 Cr: [Ar] 3d6 Ca: [Ar] 4s2 Na: 1s2 2s2 2p6 3s1 Zn: [Ar] 3d10 4s2 Kr: [Ar] 3d10 4s2 4p6 14 Electron Configurations of elements and their orbital diagram Orbital Filling Diagrams • An orbital filling diagram is used to show how the orbitals of a subshell are occupied by electrons. – Each orbital is represented by a box or circle. – Each group of orbitals is labelled by its subshell notation. Electrons are represented by arrows (up or down) – : up for spin= +1/2 and down for spin = -1/2 For He 1s 17 Sample Problem Hund’s Rule • Give the expected ground-state electron configurations (full as well as abbrev.) and the orbital-filling diagrams for: 1. P (Z=15) 2. Zn (Z=30) 3. Ca (Z=20) • Consider carbon (Z = 6) with the ground state configuration 1s2 2s2 2p2. • Hund’s rule states that the lowest energy arrangement (the “ground state”) of electrons in a subshell is obtained by putting electrons into separate orbitals of the subshell with the same spin before pairing electrons. Carbon's orbital diagram is 19 ↑↓ ↑↓ ↑ ↑ 1s 2s 2p 20 20 Valence Electrons 21 • The electrons in all the subshells with the highest principal energy shell are called the valence electrons. • Electrons in lower energy shells are called core electrons. • Chemists have observed that one of the most important factors in the way an atom behaves, both chemically and physically, is the number of valence electrons. The # of valence electrons = group # 22 Electron Configuration and the Periodic Table • The group number corresponds to the number of valence electrons. • The number of columns in each “block” is the maximum number of electrons that sublevel can hold. • The period number corresponds to the principal energy level of the valence electrons. 24 Writing electron configurations from their position in the Periodic Table • Progressing from left to right across each consecutive row will give the correct filling order… 1s,2s,2p,3s,3p… Anomalous Electron Configurations • We know that because of sublevel splitting, the 4s sublevel is lower in energy than the 3d, and therefore the 4s fills before the 3d. • But the difference in energy is not large. • Some of the transition metals have anomalous electron configurations in which the (n)s only partially fills before the (n−1)d or doesn’t fill at all. • Therefore, their electron configurations must be found experimentally. 27 Anomalous Electron Configurations Expected • Cr = [Ar]4s23d4 • Cu = [Ar]4s23d9 • Mo = [Kr]5s24d4 • Pd = [Kr]5s24d8 Found experimentally • Cr = [Ar]4s13d5 • Cu = [Ar]4s13d10 • Mo = [Kr]5s14d5 • Pd = [Kr]5s04d10 That is, if the “d” orbital can be half filled or completely filled , it will take from the “s” orbital 28 Trends in Periodic Properties • The periodic law states that when the elements are arranged by atomic number, their physical and chemical properties vary periodically. • We will look at three periodic properties: A) Atomic radius B) Ionization energy C) Electron affinity Effective Nuclear Charge Why the Trends? Shielding and Effective Nuclear Charge • In a multielectron system, electrons are simultaneously attracted to the nucleus and repelled by each other. • Outer electrons are shielded from the nucleus by the core electrons. – Thus a screening effect – But outer electrons do not effectively screen for each other. • Because of this shielding, the outer electrons do not experience the full strength of the nuclear charge. 30 Screening and Effective Nuclear Charge • Effective nuclear charge (Zeffective) is the net positive charge that is attracting a particular electron. • Z is the nuclear charge; S is the charge due to electrons in lower energy levels. – Electrons in same energy level do contribute to screening, but very little, so are not part of the calculation. – Trend is s > p > d > f Zeffective = Z − S 32 32 A) Atomic Radius Trend Figure Representation of atomic radii (covalent radii) of the maingroup elements • Within each period (horizontal row), the atomic radius tends to decrease with increasing atomic number (or increasing effective nuclear charge). • Within each group (vertical column), the atomic radius tends to increase with the period number. 34 33 B) Ionization Energy Trend i.e, Na (g) → Na +(g) + 1 e- Ionization Energy Trend IE1= 496 kJ/mol • Ionization energy trends – There is a general trend that ionization energies increase with atomic number within a given period. – This follows the trend in size, as it is more difficult to remove an electron that is closer to the nucleus. – For the same reason, we find that ionization energies, again following the trend in size, decrease as we descend a column of elements. 35 C) Electron Affinity Electron Affinity Trend • The electron affinity is the energy change for the process of adding an electron to a neutral atom in the gaseous state to form a negative ion. • The more negative the electron affinity, the more stable the negative ion that is formed. • The general trend goes from lower left to upper right as electron affinities become more negative. • The next Table gives the electron affinities of the main-group elements. – For a chlorine atom, the first electron affinity is illustrated by: Cl (g) + e- → Cl - (g) EA= -349 kJ/mol Also by showing the species” electron configuration Cl([Ne]3s 2 3p 5 ) + e − → Cl − ([Ne]3s 2 3p 6 ) Electron Affinity = -349 kJ/mol 37 38 The Main-Group Elements • The physical and chemical properties of the main-group elements clearly display “periodic behaviour’. e- higher affinities – Variations of metallic-non-metallic character. – Basic-acidic behaviour of the oxides - atomic radii, electron affinity, etc – Reaction with water – Readily oxidized or reduced to form ions 39 39 40 Properties of Groups IA and IIA • Group IA (Alkali Metals) – Largest atomic radii – React violently with water to form H2 – Readily oxidized to 1+ – Oxides dissolved in water form basic solutions • Group IIA (Alkaline Earth Metals) – Readily ionized to 2+ – React with water to form H2 – Oxides dissolved in water form basic solutions Everyone Wants to Be Like a Noble Gas! The Alkali Metals • The alkali metals have one more electron than the previous noble gas. • In their reactions, the alkali metals tend to lose their extra electron, resulting in the same electron configuration as a noble gas. – forming a cation with a 1+ charge Also the case for The Alkali Earth Metals which tend to loss two electrons, resulting in a noble gas electron configuration i.e., Ca 2+ Properties of Groups VIA, VIIA, VIIIA Properties of Groups IIIA, IVA and VA • Group VI A • Group III A – Metals (except for boron) – Several oxidation states (commonly 3+) • Group IV A (i.e., C) – Form the most covalent compounds – Oxidation numbers vary between 4+ and 4- • Group V A (i.e., N, P) – Form anions generally ( 3-), though positive oxidation states are possible 42 (i.e., O, S) – anions generally, though positive oxidation states are possible – React vigorously with alkali and alkali earth metals and nonmetals • Group VII A (halogens) – – – – Form mono anions High electronegativity (electron affinity) Diatomic gases Most reactive of the nonmetals (F2 in particular) • Group VIII A (noble gases) – Minimal reactivity – Monatomic gases; Closed (fill) shell 44 Everyone Wants to Be Like a Noble Gas!— The Halogens • The halogens all have one fewer electron than the next noble gas. • In their reactions with metals, the halogens tend to gain an electron and attain the electron configuration of the next noble gas. Noble Gas Electron Configuration • The noble gases have eight valence electrons. – Except for He, which has only two electrons. • We know that the noble gases are especially nonreactive. – forming an anion with charge 1−. – He and Ne are practically inert. • In their reactions with nonmetals, they tend to share electrons with the other nonmetal so that each attains the electron configuration of a noble gas. • The noble gases are so nonreactive because the electron configuration of the noble gases is especially stable. 45 46 Electron Configuration and Ion Charge • We have seen that many metals and nonmetals form one ion, and that the charge on that ion is predictable based on its position on the periodic table. – group 1A = 1+, group 2A = 2+, – group 7A = 1−, group 6A = 2−, etc. • These atoms form ions that will result in an electron configuration that is the same as the nearest noble gas. 48 Electron Configuration of Anions in Their Ground State Electron Configuration of Cations in Their Ground State • Anions are formed when atoms gain enough electrons to have eight valence electrons. • Cations are formed when an atom loses all its valence electrons. – filling the s and p sublevels of the valence shell – resulting in a new lower energy level valence shell – However, the process is always endothermic. • The sulfur atom has six valence electrons. S atom = 1s22s22p63s23p4 • The magnesium atom has 2 valence electrons. • In order to have eight valence electrons, it must gain two more. Mg atom = 1s22s22p63s2 • When it forms a cation, it loses its valence electrons. S2− anion = 1s22s22p63s23p6 Mg2+ cation = 1s22s22p6 50 Magnetic Properties of Transition Metal Atoms and Ions Electron Configuration of Cations in Their Ground State • Electron configurations that result in unpaired electrons mean that the atom or ion will have a net magnetic field. This is called paramagnetism. • Cations form when the atom loses electrons from the valence shell. • For transition metals, electrons may also be removed from the sublevel closest to the valence shell. – will be attracted to a magnetic field • Electron configurations that result in all paired electrons mean that the atom or ion will have no magnetic field. This is called diamagnetism. 1s22s22p63s23p1 Al atom = Al3+ ion = 1s22s22p6 Fe atom =1s22s22p63s23p64s23d6 Fe2+ ion = 1s22s22p63s23p63d6 Fe3+ ion = 1s22s22p63s23p63d5 =1s22s22p63s23p64s13d10 Cu atom Cu+ ion = 1s22s22p63s23p63d10 – slightly repelled by a magnetic field 51 52 Practice—Determine whether the following are paramagnetic or diamagnetic. another Term to know: isoelectronic = ions with same electron configuration i.e., Ca2+ is isoelectronic with K+ D) Trends in Ionic Radius • Mn Mn = [Ar]4s23d5 4s paramagnetic ! 3d • Sc3+ Sc = [Ar]4s23d1 paramagnetic! Sc3+ = [Ar] diamagnetic ! • 1) ion size increases down the group – higher valence shell = larger • • • • 2) cations are smaller than the neutral atom 3) anions are bigger than the neutral atom 4) cations smaller than anions 5) larger positive charge = smaller cation – for isoelectronic species i.e., Na+ > Mg2+ > Al3+ • 6) larger negative charge = larger anion – for isoelectronic species N3 - > O2 - > F - Main Group 1A, 2A &3A 54 Main group 6A & 7A 55 56 56 Ionic vs. neutral atom radii Choose the larger of each pair • S or S2− – S2− is larger because there are more electrons (18 e−) for the 16 protons to hold. – The anion is larger than the neutral atom. • Ca or Ca2+ – Ca is larger because its valence shell electrons have been lost to form Ca2+. – The cation is always smaller than the neutral atom. • Br− or Kr – Br− is larger because it has fewer protons (35 p+) to hold the 36 electrons than does Kr (36 p+). – For isoelectronic species, the more negative the charge, the larger the atom or ion. 58 57 Ionic Radii Ionic Radii • Within an isoelectronic group of ions, the one with the greatest nuclear charge will be the smallest. • In this group, calcium has the greatest nuclear charge and is, therefore, the smallest. – For example, look at the ions listed below: 20 Ca 2+ 19 K + 20 Ca 18 Ar 17 Cl - 16S All have 18 electrons – Note that they all have the same number of electrons, but different numbers of protons. 2- 2+ < 19K + < 18 Ar < 17Cl - < 16S 2All have 18 electrons – Sulfur has only 16 protons to attract its 18 electrons and, therefore, has the largest radius. 60 Practice—Order the following sets by size (smallest to largest). Zr4+, Ti4+, Hf4+ same column and charge; therefore, Ti4+ < Zr4+ < Hf4+ isoelectronic; Na+, Mg2+, F−, Ne therefore, Mg2+ < Na+ < Ne < F− I−, Br−, Ga3+, In+ Ga3+ < In+ < Br− < I−