George Mason University General Chemistry 211 Chapter 8

George Mason University General Chemistry 211 Chapter 8
Electron Configuration and Chemical Periodicity Acknowledgements Course Text: Chemistry: the Molecular Nature of Matter and Change, 7th edition, 2011, McGraw-Hill Martin S. Silberberg & Patricia Amateis The Chemistry 211/212 General Chemistry courses taught at George Mason are intended for those students enrolled in a science /engineering oriented curricula, with particular emphasis on chemistry, biochemistry, and biology The material on these slides is taken primarily from the course text but the instructor has modified, condensed, or otherwise reorganized selected material. Additional material from other sources may also be included. Interpretation of course material to clarify concepts and solutions to problems is the sole responsibility of this instructor.

Electron Configuration & Chemical Periodicity
Development of the Periodic Table Characteristics of Many-Electron Atoms The Electron Spin Quantum Number The Pauli Exclusion Principle Electrostatic Effects and Energy-Level Splitting Development of the Quantum Mechanical Model of the Periodic Table Building up of Periods 1 & 2 Building up of Period 3 Electron Configuration Within Groups Building up Period 4 General Principles of Electron Configuration Unusual Configurations: Transition and Inner Transition Elements

Electron Configuration & Chemical Periodicity
Trends in Three Key Atomic Properties Trends in Atomic Size Trends in Ionization Energy (IA) Trends in Electron Affinity (EA) Atomic Structure and Chemical Reactivity

Electron Configuration
The work of Balmer (Rydberg equation) and Bohr (Bohr Postulates) invoked the idea that spectra lines of compounds represented different wavelengths (Balmer) and energy levels (Bohr) of the radiation. Each energy level was designated by the whole number integer, n. “n” could have any value from 1 -  (infinity) The higher the value of “n,” the smaller the wavelength, which equates to higher frequency and higher energy

The work of Balmer, Bohr, and Einstein, de Broglie, Heisenberg and many others lead to the development of Quantum Mechanics Quantum mechanics, also known as quantum physics or quantum theory, is a branch of physics providing a mathematical description of the wave-particle duality of matter and energy In quantum chemistry each atom is distinguished by its unique number of electrons, which is matched by an equal number of protons in the nucleus (atomic number, Z) Each new element has one more electron than its predecessor Hydrogen (H) 1 e-; Helium (He) 2 e-

Electronic Configuration
The electrons are configured (distributed) into “orbitals,” which represent energy levels Electrons are able to move from one orbital (energy level) to another by emission or absorption of a quantum of energy, in the form of a photon Knowledge of the electron configuration of different atoms is useful in understanding the structure of the periodic table of elements The concept is also useful for describing the chemical bonds that hold atoms together In bulk materials this same idea helps explain the peculiar properties of lasers and semiconductors

According to quantum mechanics each electron is described by 4 Quantum numbers Principal Quantum Number (n) Angular Momentum Quantum Number (l) Magnetic Quantum Number (ml) Spin Quantum Number (ms) The first 3 quantum numbers define the wave function of the electron’s atomic orbital, i.e., it size and general energy level The fourth quantum number refers to the Spin Orientation of the 2 electrons that occupy an atomic orbital

Quantum Numbers and Atomic Orbitals The Principal Quantum Number (n) represents the “Shell Number” in which an electron “resides” It represents the relative size of the orbital Equivalent to periodic chart Period Number Defines the principal energy of the electron The smaller “n” is, the smaller the orbital size The smaller “n” is, the lower the electron energy n can have any positive value from 1, 2, 3, 4 …  (Currently, n = 7 is the maximum known)

Quantum Numbers and Atomic Orbitals (Con’t) The Angular Momentum Quantum Number (l) distinguishes “sub shells” within a given shell Each main “shell,” designated by quantum number “n,” is subdivided into: l = n - 1 “sub shells” (l) can have any integer value from 0 to n - 1 The different “l” values correspond to the s, p, d, f designations used in the electronic configuration of the elements Letter s p d f (g) l value (4)

Quantum Numbers and Atomic Orbitals (Con’t) The Magnetic Quantum Number (ml) defines the atomic orbitals within a given sub-shell Each value of the angular momentum number (l) determines the number of atomic orbitals For a given value of “l,” ml can have any integer value from –l to +l ml = –l to +l Each orbital has a different shape and orientation (x, y, z) in space Each orbital within a given angular momentum number sub shell (l) has the same energy

Quantum Numbers and Atomic Orbitals (Con’t) The Spin Quantum Number (ms) refers to the two possible spin orientations of the electrons residing within a given atomic orbital Each atomic orbital can hold only: two (2) electrons Each electron has a “spin” orientation value The spin values must oppose one another The possible values of ms spin values are: +1/2 and –1/2 2

Stern-Gerlach Experiment
A beam of H atoms can be separated into 2 beams of opposite electron spin in a magnetic field ms ( –1/2 ) electrons have a slightly greater energy than ms ( +1/2 ) electrons

Representation of electron spin
A spinning charged particle aligns in a magnetic field depending on spin state

Summary of Quantum Numbers
Name Symbol Permitted Values Property principal n positive integers (1, 2, 3, …) orbital energy (size) angular momentum l integers from 0  n -1 orbital shape The l values , 1, 2, and 3 correspond to s, p, d, f orbitals, respectively magnetic ml integers from –l  0  +l orbital orientation spin ms +1/2  or – 1/2  direction of e- spin

Summary of Quantum Numbers
Name, Symbol (Property) Allowed Values Quantum Numbers Principal, n (size, energy) Positive integer (1, 2, 3, ...) n=3 n=1 n=2 -1 +1 +2 -2 l=0 l=1 l=2 (1s) (2s) (2p) (3s) (3p) (3d) Angular momentum, l (shape) l = 0  n -1 0(s), 1(p), 2(d), 3(f) Magnetic, ml (orientation) -l,…,0,…,+l

An electron configuration of an atom is a particular distribution of electrons among available sub shells The configuration notation lists the subshell symbols (s, p, d, f…) sequentially with a superscript indicating the number of electrons occupying that subshell Ex: lithium (Period (n) = 2, Atomic No 3) has 2 electrons in the “1s” sub shell 1 electron in the “2s” sub shell 1s2 2s1 Fluorine (Period (n) 2, Atomic No 9) has 2 electrons in the “1s” sub shell 2 electrons in the “2s” sub shell 5 electrons in the “2p” subshell 1s2 2s2 2p5 2

A unique set of the first 3 quantum numbers (n, l, m l) defines an “Orbital” An orbital can contain a maximum of 2 electrons, each with a different “spin” (+1/2 or -1/2) An orbital diagram is notation used to show how the orbitals of a sub shell are occupied by electrons Each orbital is represented by a circle Each orbital can have a maximum of 2 electrons Each group of orbitals is labeled by its Sub Shell Notation (s, p, d, f) Electrons are represented by arrows: up () for ms = +1/2 and down () for ms = -1/2 1s 2s 2p 2

The Pauli Exclusion Principle
No two electrons in an atom can have the same four quantum numbers An orbital (unique combination of n, l, m l ) can hold, at most, two electrons Two electrons in the same Orbital have opposite spins +1/2  -1/2  2

The maximum number of electrons and their orbital diagrams are: Sub Shell No Values Orbitals (-l to +l) Max No. Electrons s (l = 0) (0) 2 p (l = 1) (-1,0,+1) 6 d (l =2) (-2,-1,0,+1,+2) 10 f (l =3) 7 (-3,-2,-1,0,+1,+2,+3) 14 2

n l (n-1) ml 1 0 (1s) (2s) 0 1 (2p) (3s) 0 1 (3p) (3d) (4s) 0 1 (4p) (4d) (4f) (5s) 0 1 (5p) (5d) (5f) (5g)

Noble Gas Electronic Configurations
In the following slides electronic configurations of the elements in the Periodic Table will be discussed Electronic configurations can become quite complex (lengthy) as the Atomic Number increases A condensed form of the Electronic Configuration of a given element or ion is often used A symbol, [X], representing the electron configuration of the Noble gas in the period just above the element of interest is substituted for the detail configuration The following slide illustrates the Noble Gas configurations and the “Condensed Form” symbol used with other elements 2

Configurations and the Periodic Table
Condensed Electronic Configurations Full Electronic Configuration Condensed Electronic Configuration Helium 1s2 [He] 2 e- Neon 1s22s22p6 [Ne] 10 e- Argon 1s22s22p63s23p6 [Ar] 18 e- Krypton 1s22s22p63s23p63d104s24p6 [Kr] 36 e- Xenon 1s22s22p63s23p63d104s24p64d105s25p6 [Ze] 54 e- Beryllium 1s22s2 [He] 2s2 4 e- Magnesium 1s22s22p63s2 [Ne] 3s e- Calcium 1s22s22p63s23p64s2 [Ar] 4s e- Sodium Ion (Na) 1s22s22p63s1  (Na+) 1s22s22p6 + 1e- [Ne] + 1e- 2

Quantum Number n = 1 (Period 1) l values = 0 to (n-1) = 0 to (1 -1) = 0  l = 0 (s orbital) ml values = -l,…0,…+l = 0 (1 s orbital) ms values = -1/2 & +1/2 = (2 e- per orbital) 1s orbital Z = 1 Hydrogen 1s1 Z = 2 Helium 1s2 Thus, for n = 1 there is one orbital (s) which can accommodate 2 elements – Hydrogen & Helium 2

Quantum Number n = 2 (Period 2) l values = 0 to (n-1) = 0 to (2-1) = 0 to 1  l = 0(s), 1(p) For l = 0 (s) ml = -l 0 + l = 0 (one 2s orbital, 2 electrons) ms values = -1/2 & +1/2 For l = 1 (p) ml = (three 2p orbitals, 6 electrons) ms values = -1/2 & +1/2 in each orbital 2s orbitals Z=3 Lithium 1s22s1 or [He]2s1 Z=4 Beryllium 1s22s2 [He]2s2 Z=5 Boron 1s22s22p1 [He]2s22p1 Z=6 Carbon 1s22s22p2 [He]2s22p2 Z=7 Nitrogen 1s22s22p3 [He]2s22p3 Z=8 Oxygen 1s22s22p4 [He]2s22p4 Z=9 Fluorine 1s22s22p5 [He]2s22p5 Z=10 Neon 1s22s22p6 [He]2s22p6 2p orbitals 2

With Sodium (Z = 11), the 3s sub shell begins to fill Z=11 Sodium 1s22s22p63s1 or [Ne]3s1 Z=12 Magnesium 1s22s22p63s2 or [Ne]3s2 Starting with Z = 13, the 3p sub shell begins to fill Z=13 Aluminum 1s22s22p63s23p1 or [Ne]3s23p1 Z=18 Argon 1s22s22p63s23p6 or [Ne]3s23p6 2

Electrostatic Effects and Energy-Level Splitting The principal quantum number (n) defines the energy level of an atom The higher the “n” value, the higher the energy level The unique values of the principal quantum numbers of multi-electron atoms (n, l, ml) define a unique energy level for the orbital of a given electron The energy of a given orbital depends mostly on the value of the principal quantum number (n), i.e. its size, and to a lesser degree on the shape of the orbital represented by the various values of the magnetic quantum number (l) 2

The energy states of multi-electron atoms arise from 2 counteracting forces: Nucleus – Positive protons attract Negative electrons Electron – Negative electrons repulse each other Nuclear protons create a pull (attraction) on electrons Higher nuclear charge (Z) lowers orbital energy (stabilizes system) by increasing proton-electron attractions The energy required to remove the 1s electron from Hydrogen (H), Z =1, is much less than the energy to remove the 1s electron from the Li2+ ion, Z = 3

Effect of Nuclear Charge (Z) on Orbital Energy Greater Nuclear Charge lowers orbital energy making it more difficult to remove the electron from orbit The absolute value of the 1s orbital energy is related directly to Z2 Energy required to remove 1s electron from H 1311 kJ/mol (Z= +1, Least stable) Energy required to remove 1s electron from He+ 5250 kJ/mol (Z = +2) Energy required to remove 1s electron from Li+ 11815 kJ/mol (Z = +3, Most stable)

Shielding – Effect of Electron Repulsions on Orbital Energy Electrons feel repulsion from other electrons somewhat shielding (counteracting) the attraction of the nuclear protons Shielding (screening) lowers the full nuclear charge to an “Effective Nuclear Charge (Zeff) The lower the Effective Nuclear Charge, the easier it is to remove an electron It takes less than half as much energy to remove an electron from Helium (He) (2373 kJ/mol) than from He+ (5250 kJ/mol) because the second electron in He repels the first electron and effectively shields the first electron from the full nuclear charge (lower Zeff)

Penetration: Effects of orbital shape The shape of an atomic orbital affects how close an electron moves closer to nucleus, i.e., the level of penetration Penetration and the resulting effects of shielding on a atomic orbital causes the energy level (n) to be split into sublevels of differing energy representing the various values of the magnetic quantum number (l) The lower the value of the magnetic quantum number (l), the more its electrons penetrate Order of Sublevel Energies s (l=0) < p(l=1) < d(l=2) < f(l=3) Each of the orbitals for a given value of l (ml = -l 0 +l) has the same energy

Aufbau Principle Aufbau Principle – scheme used to reproduce the ground state electron configurations of atoms by following the “building up” order based on relative energy levels of quantum subshells 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

You need not memorize this order
Aufbau Principle Listed below is the order in which all the possible subshells fill with electrons Note the order does NOT follow the strict numerical subshell order shown on slide 20 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f You need not memorize this order The next slide provides a pictorial providing an easier way of the viewing the ‘build-up” order 2

Order for Filling Atomic Subshells
Principal Quantum No. (n) Angular Momentum (l) Setup rows for each Principal Quantum No. (n) Set columns for each Angular Momentum (l) Draw a series of diagonals Order of filling is the order in which diagonals strike subshells Note the 4s subshell is filled before the 3d subshell because the 4s electrons are at lower energy levels than the 3d electrons 1s 2s 3s 4s 5s 6s 7s n = 1 n = 2 n = 3 n = 4 n = 5 n = 6 n = 7 2p 3p 4p 5p 6p 7p 3d 4d 5d 6d 7d 4f 5f 6f 7f

Aufbau Principle Every atom has an infinite number of possible electron configurations (electrons can be raised to any number of energy (n) levels) The configuration associated with the lowest energy level of the atom is called the “ground state” Other configurations correspond to “excited states” Tables on the next 3 slides list the ground- state configurations of atoms up to krypton 2

Partial Orbital Diagrams

Chromium (Cr) relative to Vanadium (V) The Cr 4s1 subshell is filled before the 3d subshell is completed An [Ar]3d44s2 orbital configuration would be expected for ground state Cr, but the [Ar]3d54s1 orbital is lower in energy Partial Orbital Diagrams and Electron Configurations* for the Elements in Period 4 Cr

Copper (Cu) relative to Nickel (Ni) Copper would be expected to have a ground state configuration of [Ar]3d94s2 The [Ar]3d104s1 configuration is actually lower in energy Partial Orbital Diagrams and Electron Configurations* for the Elements in Period 4 Ni Cu

Orbital Energy Levels in Multi-Electron Systems
2p 3s 3p 3d 4s 3d orbitals would be expected to be filled before 4s orbitals Actual order of filling depends on total ground state energy of the atom 3d and 4s orbitals are very close in energy Selected 4s, 5s, 6s, 7s levels are filled before 3d, 4d, 4f, 5f, respectively (see slides 31 & 32) Energy 3

Electrons that reside in the outermost shell of an atom - or in other words, those electrons outside the “noble gas core” - are called valence electrons These electrons are primarily involved in chemical reactions Elements within a given group have the same valence shell configuration This accounts for the similarity of the chemical properties among groups of elements n = 2 Li – 2s Be – 2s2 n = 3 Na – 3s Mg – 3s2 n = 4 K – 4s Ca – 4s2 n = 5 Rb – 5s Sr – 5s2 2

Noble gas core: an inner shell configuration resembling one of the noble gases (He, Ne, Ar, Kr, Xn) Pseudo-noble gas core: noble gas core + (n-1)d10 electrons: Ex Sn  Sn+4 Sn ([Kr] 5s2 4d10 5p2)  Sn+4 ([Kr] 4d e-

Configurations of Main Group Ions Noble gases have filled outer energy levels (ns2np6), have very high Ionization Energies (IEs), and positive (endothermic) Electron Affinities (EAs); thus do not readily form ions Elements in Groups 1A, 2A, 6A, 7A that readily form ions by gaining electrons (1A & 2A) or losing electrons (6A & 7A) attain a filled outer level conforming to a Noble Gas configuration Such ions are said to be “Isoelectronic” with the nearest Noble gas configuration Na (1s22s22p63s1)  Na+ (1s22s22p6) + 1e- Isoelectronic with [Ne] + 1e-

The energy needed to remove the electrons from metals in groups 1A, 2A, 6A, 7A, is supplied during exothermic reactions with nonmetals Attempts to remove more than 1 electron from group 1A or 2 electrons from group 2A metals would mean removing core (not valence) electrons requiring significantly more energy than is available from a reaction with a non-metal

The larger metals from Groups 3A, 4A, and 5A form cations through a different process It would be energetically impossible for them to lose enough electrons to attain a noble gas configuration Ex: Tin (Sn), Z = 50 would have to lose 14 electrons (two 5p, ten 4d, two 5s) to be isoelectronic with Krypton: Z =36 Instead, tin loses fewer electrons and still attains one or more stable pseudo-noble gas configurations Sn ([Kr] 5s24d105p2)  Sn4+ ([Kr] 4d10) + 4e- Stability comes from empty 5s & 5p sublevels and a filled inner 4d sublevel (n-1)d10 configuration Pseudo-Noble Gas Configuration

Practice Problem Which of the following electron configurations represents an excited state? a. He: 1s2 b. Ne: 1s2 2s2 2p6 c. Na: 1s2 2s2 2p6 3s1 d. P: 1s2 2s2 2p6 3s2 3p2 4s1 e. N: 1s2 2s2 2p3 Ans: d Ground state for Phosphorus is: 1s2 2s2 2p6 3s2 3p3 The 3p subshell would continue to fill before the 4s subshell would start to fill

Practice Problem What is the electron configuration for the valence electrons of Technetium (Tc, Z = 43)? a. 4d55s2 b. 5s25d4 c. 4s24d4 d. 4d65s2 e. 3d44s2 Ans: a 4d55s = 7 valence electrons Technetium (atomic no. = 43 = 43 total electrons) Select “Noble Gas” Configuration prior to Technetium (Kr) 1s22s22p63s23p63d104s24p6  [Kr] 36 e- [Kr] + 4d55s2 = = 43 = Technetium Note: 4d orbitals filled before 5p orbitals (Aufbau)

Practice Problem What is the electron configuration for the valence electrons of Polonium (Po, Z=84)? a. 6s26p b. 6s25d106p4 c. 6s25d106p d. 6s26p e. 7s26p4 Ans: b 6s25d106p = 16 valence electrons Polonium (atomic no. 84 = 84 total electrons) Select “Noble Gas” Configuration prior to Polonium Ze(54) Xenon 1s22s22p63s23p63d104s24p64d105s25p6 [Ze] 54 e- 84 – 54 = 30 electrons which must include 14 electrons that fill in the 4f orbitals that start with Lanthanum = 16 – 10 (filled 5d10) = 6 = 6s25d106p4

The following slide illustrates how the periodic table provides a sound way to remember the Aufbau sequence In many cases you need only the configuration of the outer electrons You can determine this from their position on the periodic table The total number of valence electrons for an atom equals its group (vertical column) number 2

Periodic Table (Subshells)

Main block = s + p blocks s block p block d block Transition Elements f block Inner Transition Elements   2

Orbital Diagrams Diagram 1: Diagram 2: Diagram 3:
Consider carbon (Z = 6) with the ground state configuration 1s22s22p2 Three possible arrangements are given in the following orbital diagrams. 1s s p Diagram 1: Diagram 2: Diagram 3: Each state has a different energy and different magnetic characteristics 2

Orbital Diagrams Hund’s rule states that the lowest energy arrangement (the “ground state”) of electrons in a sub-shell is obtained by putting electrons into separate orbitals of the sub shell with the same spin before pairing electrons Looking at carbon again, we see that the ground state configuration corresponds to diagram 1 when following Hund’s rule Note: The 2 e- in the 2p orbitals are shown as “up” arrows representing the +1/2 spin state, which has lower energy the -1/2 spin state 1s s p 2

Orbital Diagrams To apply Hund’s rule to Oxygen, whose ground state configuration is 1s22s22p4, place the first seven electrons as follows 1s s p The last electron is paired with one of the 2p electrons to give a doubly occupied orbital, i.e., a +½ spin state and a – ½ spin state 1s s p 2

Recall: +1/2 spin has lower energy then -1/2 spin
Summary Pauli Exclusion principle: no 2 e-s in an atom can have the same four quantum numbers Aufbau Principle: obtain electron configurations of the ground state of atoms by successively filling subshells with electrons in a specific order Hunds Rule: the lowest energy arrangement of electrons in a subshell is obtained by putting electrons into separate orbitals of the subshell with the same spin before paring them Recall: +1/2 spin has lower energy then -1/2 spin

Periodic Properties Two factors determine the size of an atom
One factor is the principal quantum number, n. The larger “n” is , the larger the size of the orbital The other factor is the effective nuclear charge (slide 28), which is the positive charge an electron experiences from the nucleus minus any “shielding effects” from intervening electrons The Periodic Law states that: When the elements are arranged by atomic number, their physical and chemical properties vary periodically – across the periodic chart row 2

Periodic Properties – Atomic Size
Atomic Size, Ionization Energy, Electron Affinity Atomic radius Within each Period (across horizontal row), the atomic radius tends to decrease with increasing atomic number (nuclear charge more dominant than electron repulsion) Within each Group (down a vertical column), the atomic radius tends to increase with increasing period number (electron repulsion dominates nuclear charge increase) 2

Representation of atomic radii (covalent radii) of the main- group elements (neutral atoms)

Elements vs Ions Ionic Size increase down a group Number of energy levels increases Ionic Size becomes more complicated across a period Decreases among cations Increase dramatically with first anion Decreases within anions

Ionic Size and Atomic Size Cations are smaller than their parent atoms Electrons are removed from the outer level Resulting decrease in electron repulsions allows nuclear charge to pull remaining electrons closer Anions are larger than their parent atoms Electrons added to outer level Resulting in increased electron repulsion allowing them to occupy more space

Periodic Properties – Ionization Energy
The first ionization energy of an atom is the minimal energy needed to remove the highest energy (outermost) electron from the neutral atom For a Lithium atom, the first ionization energy is illustrated by: Li(1s22s1) → Li+(1s2) + e- IE = 520 kJ/mol Endothermic (requires energy input) 2

Ionization energy (IE) 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 descending down a column of elements 2

Ionization Energy vs Atomic Number
Noble gases have highest IE’s Alkali metals have lowest IE’s

Successive Ionization Energies of the First Ten Elements (kJ / mol* Ionization Energies to the “Right” of the a vertical line correspond to removal of electrons from the “Core” of the atom

Ionization energy (IE) The electrons of an atom can be removed successively The energies required at each step are known as the first ionization energy, the second ionization energy, and so forth Successive Ionization Energies increase because each electron is pulled away from an ion with a progressively higher positive charge, i.e., a more effective nuclear charge 2

Exceptions to Ionization Energy Trends
A IIIA element , such as Boron (2s22p1), has a smaller ionization energy (IE) than the preceding IIA element Beryllium (2s2) because one np electron is more easily removed than the second ns electron A VIA element, such as oxygen (2s22p4), has smaller ionization Energy than the preceding VA element nitrogen (2s22p3). As a result of repulsion it is easier to remove an electron from the doubly occupied 2p orbital of the VI element that from a singly occupied p orbital of the preceding VA element Nitrogen 2s22p3 Oxygen 2s22p4

Periodic Properties – Electron Affinity
Electron Affinity (EA): the energy change for the process of adding an electron to a neutral atom in the gaseous state to form a negative ion, i.e., an Anion 1st Electron Affinity – Formation of 1 mole of monovalent (1-) gaseous ions Atoms(g) + e-  ion-(g) E = EA1 For the formation of the Chloride ion (Cl-) from the Chlorine atom, the first electron affinity is illustrated by: Electron Affinity = EA1 = 349 kJ/mol Exothermic (releases energy) 2

Electron Affinity (EA) The more negative the electron affinity, the more stable the negative ion that is formed Broadly speaking, the general trend goes from lower left to upper right as electron affinities become more negative Highest electron affinities occur for halogens, F and Cl Negative values indicate that energy is released when the Anion forms Note: Electron Affinity is not the same as Electronegativity – relative ability of a bonded atom to attract shared electrons 2

Atomic Size (neutral atoms & ions) increases down a main group Atomic Size (neutral atoms & ions) decreases across a Period Atomic Size remains relatively constant across a transition series Ionization Energy First Ionization Energy (remove outermost e-) is inversely related to atomic size 1st Ionization Energy decreases down a group 1st Ionization Energy increases across period Successive IEs show very large increases after 1st electron is removed Electron Affinity Similar patterns (with many exceptions) to ionization Energy (lower left to upper right) Highest electron affinities occur for halogens, F and Cl

Periodic Properties - Summary

Atomic Structure / Chemical Reactivity
Metals Metals are located in the left and lower three- quarters of the Periodic Table Typical Properties Shiny Solids High Melting Points Good Thermal & Electrical Conductors Malleable – Drawn into wires and rolled into sheets Lose electrons to non-metals

Non-Metals Non-metals are located in the upper right quarter of the Periodic Table Not Shiny Low Melting Points Poor Thermal & Electrical Conductors Crumbly Solids or gases Gain Electrons from Metals

Metalloids (semi-metals) Located between Metals & Non-Metals in the Periodic Chart boron, silicon, germanium, arsenic, antimony, tellurium, and polonium An element that exhibits the external characteristics of a metal, but behaves chemically more as a nonmetal Arsenic, for example, is a metalloid that has the visual appearance of a metal, but is a poor conductor of electricity The intermediate conductivity of metalloids means they tend to make good semiconductors

Metalloids (semi-metals) The electronegativities and ionization energies of the metalloids are between those of the metals and nonmetals, so the metalloids exhibit characteristics of both classes The reactivity of the metalloids depends on the element with which they are reacting Ex. Boron Acts as a nonmetal when reacting with Sodium Acts as a metal when reacting with Fluorine

Metalloids (semi-metals) The boiling points, melting points, and densities of the metalloids vary widely As a rule, metalloids do not form multiple bonds Compounds containing these elements will often show an incomplete octet around the central atom

Metallic Behavior decrease from left to right and increases from top to bottom in Periodic Tables Non-Metals Metals Transition Elements f-block Inner Transition Elements

Metallic Behavior Metals tend to “Lose” electrons Metals tend to lose electrons during chemical reactions because they have “Low” ionization energies compared to non-metals Elements generally tend to increase their metallic character going down a Periodic Table group The greatest contrast in changing metallic character is in groups 3A – 6A Elements at the top tend to form “Anions”, i.e., more non-metallic character, while those at the bottom tend to form metallic “Cations”

Metallic Behavior (Con’t) Nitrogen (N) & Phosphorus (P), both non-metals tend to form 3- anions Arsenic (As) (period 4) & Antimony (Sb) (period 5) are metalloids and generally do not form ions Bismuth (Bi) (period 6) is a typical metal forming mostly ionic compounds as a 3+ cation

Metal Behavior (Con’t) Metallic behavior decreases going from left to right across the Period table Increasing group number (left to right) Ability to lose electrons (form cations) becomes more difficult with as Ionization Energy (IE) increases Ability to gain electrons (form anions) increases as Electron Affinity (EA) decreases (becomes more negative) Elements on the left (more metallic) tend to form positively charged “Cations” Elements on the right (more non-metallic) tend to form negatively charged “Anions”

Metallic Behavior (Con’t) Sodium (Na) group 1 – Very Metallic Readily loses electron (Na+ ion) which reacts immediately with oxygen to form an oxide Aluminum (Al) group 3 – Metalloid Form some Al3+ ionic compounds, but is covalently bonded in others Silicon (Si) group 4 – Metalliod Does not occur as a monoatomic ion Phosphorus (P) group 5 – non-metal Forms a few 3- ions Sulfur (S) group 6 – non-metal Forms 2- anions, such as Sulfide

Metallic Behavior (Con’t) Acid-Base Behavior of Element Oxides Metals Most main group metals transfer electrons to oxygen forming ionic oxides Ionic oxides act as bases producing OH- (hydroxide) ions from O2- Non-metals Share electrons with oxygen to form covalent oxides Covalent oxides act as acids producing H+ ions (protons)

Metallic Behavior (Con’t) Amphoteric Behavior Some metals and many metalloids form oxides that can act as either an acid or a base

Acid-Base behavior of common oxides As elements become more metallic going down a group, the oxides become more basic Nitrogen Pentoxide (N2O5) Period 2 non-metallic forms nitric acid, a strong acid N2O5(s) H2O(l) 2HNO3(aq) Tetraphosphorus decaoxide (P4O10) Slightly more metallic Period 3 non-metal forms a weaker acid P4O10(s) H2O  4H3PO4(aq)

Acid-Base behavior of common oxides Arsenic Pentoxide (As2O5) Group 4 metalloid (more metallic) is weakly basic Bismuth Pentoxide (Bi2O5) Group 5 metalloid (most metallic in group) Basic oxide, insoluble in water, forms salt & water with an acid Bi2O3(s) + 6HNO3(aq)  2Bi(NO3)3(aq) + 3H2O(l)

Acid-Base behavior of common oxides Across a group Elements become less metallic across a group Oxides becomes more acidic Metallic Sodium (Na) (group 1) & Magnesium (Mg) (group 2) form strongly basic oxides Metallic Aluminum (group 3) forms amphoteric aluminum oxide (Al2O3), which can act as a base to react with an acid or as an acid to react with a base Al2O3(s) + 6HCl(aq)  2AlCl3(aq) + 3H2O(l) Al2O3(s) NaOH(aq) + 3H2O(l)  2NaAl(OH)4(aq)

Acid-Base behavior of common oxides Across a group (Con’t) Silicon Dioxide (SiO4) group 4 Weakly acidic forming salt & water with a base SiO2(s) NaOH(aq)  Na2SiO3(aq) + H2O(l) Common oxides of Phosphous (group 5) Sulfur (group 6) and Chlorine (group 7) are increasingly acidic forming increasingly stronger acids Acidity H3PO4 < H2SO4 < HClO4

Acid-Base behavior of common oxides Trends in acid-base behavior of Group 5 and Period 3 oxides Red – Acidic (non-metal oxides) Blue – Basic (metal oxides) Other – Metalloid oxides (note gradations ) Atomic Size Oxide Basicity More Metallic Ionization Energy

Properties of Monoatomic Ions Electron Configuration of Main-Group ions Recall: Elements in Groups 1 & 2 readily lose electrons to form cations and elements in groups 6 & 7 readily gain electrons to form anions The formation of the anions or cations result in a filled outer shell, i.e., the nearest noble gas configuration Na(1s22s22p63s1)  Na+ (1s22s22p6)  Ne + e- Br ([Ar] 4s23d104p5) + e-  Br- ([Ar] 4s23d104p6) ([Ar] 4s23d104p6)  [Kr]

Properties of Monoatomic Ions (Con’t) Electron Configuration of Main-Group ions Energy to remove the outer valence shell electrons (Ionization Energy) is supplied during the exothermic reaction of a metal with a non-metal Removing more than one electron from Na to form Na2+ or two electrons from Mg to form Mg3+ means removing core (non-valence) electrons, which requires much more energy than is available from the chemical reaction Similarly, adding 2 electrons to Fluorine to form F2- means adding electrons to the next energy level, which would require a large amount of energy to overcome the shielding of the nuclear charge by the 18 inner core electrons Thus, compounds such as Na2F & Mg3O2 do not exist

Properties of Monoatomic Ions Electron Configuration of Main-Group ions Larger Metals of Groups 3, 4, 5 Energetically impossible for them to lose enough electrons to attain noble gas configuration Tin (Sn) [Kr] 5s25p24d10 would have to lose 14 electrons (two 5p, ten 4d, and two 5s) to be isoelectronic with Krypton (Kr) – [Ar] 4s24p6 Cations formed through a different process Sn+4 – loss of two 5s & two 5p electrons, attaining stability from the filled in 4d sublevel Sn+2 – loss of two 5p electrons, attaining stability from the filled 5s & 4d sublevels

Properties of Monoatomic Ions Electron Configuration of Main-Group ions Larger Metals of Groups 3, 4, 5 Carbon Would have to either lose 4 electrons to attain the C4+ Helium configuration or gain 4 electrons to attain the C4- Neon configuration In either case, the energy requirements are extremely high, i.e. sun-like temperatures of 106k

Properties of Monoatomic Ions Electron Configuration of Main-Group ions Most elements that form Monatomic ions that are Isoelectronic with a noble gas lie in the four groups that flank group 8

Practice Problem Using condensed electron configurations, write reactions for the formation of the common ions of the following: Iodine: I ([Kr] 5s24d105p5) + e-  I- ([Kr] 5s24d105p6)  I- [Xe] Potassium: K ([Ar] 4s1)  K+ ([Ar]) + e- Indium: Group 3A – loses 3 electrons or loses 1 electron In ([Kr] 5s24d105p1)  In3+ ([Kr] 4d10) + 3e- In ([Kr] 5s24d105p1)  In+ ([Kr] 5s24d10) + 1e-

Electron Configurations of Transition Metal Ions Transition metal ions rarely attain noble gas configurations Energy required to attain noble gas configuration is very high Exceptions Scandium – forms Sc3+; Titanium – forms Ti4+ In Periods 4 & 5, a transition metal can form more than one cation by losing all of its ns and some of the (n-1)d electrons

Electron Configurations of Transition Metal Ions Aufbau electron build-up At the beginning of Period 4, the 4s orbital is nearer the nucleus than the 3d orbital making it more stable than the empty 3d orbital The first & second electrons fill the 4s orbital before filling the empty 3d orbitals At the beginning of the transition elements (group 3B), however, the previously filled 4s orbitals do not do a very good job of shielding the 3d electrons The 3d orbitals now become more stable than the 4s orbitals and begin to fill under the influence of increased nuclear charge - a cross-over in orbital energy

Electron Configurations of Transition Metal Ions Aufbau electron build-up (Con’t) The 4s electrons, which were added before the 3d electrons, are now lost preferentially before the 3d electrons to form the transition metal electrons Simple Rules for forming the ion of any “Main Group” or “Transition” Group element Electrons with the highest “n” value are removed first For main-group, s block metals, remove all electrons with the highest “n” value For main-group, p-block metals, remove “np” electrons before “ns” electrons For transition (d-block) metals, remove “ns” electrons before “(n-1)d” electrons For non-metals, add electrons to the “p” orbitals of the highest “n” value

Magnetic Properties A spinning electron behaves like a tiny magnet generating a magnetic field A single electron (unpaired) in an orbital can be affected by an externally applied magnetic field A Paramagnetic element (or ion) has 1 or more orbitals with unpaired electrons and is weakly attracted by a magnetic field Titanium [Ar]4s23d2 A Diamagnetic element (or ion) has only paired electrons and is not attracted by a magnetic field Copper ion Cu+ [Ar]4s23d10 4s 3d 4p 4s 3d 4p 2

Magnetic Properties of Transition Metal Ions Ag (Z=47) [Kr] 5s14d10 Unpaired – Paramagnetic – split by applied magnetic field Cd (Z=48) [Kr] 5s24d10 Paired – Diamagnetic – not split by applied magnetic field 5s 4d 5p 5s 4d 5p

Using Paramagnetism to verify electron configuration Titanium (Ti) [Ar] 4s23d2 Titanium (II) Ion (Ti2+) ([Ar] 3d2) e- Paramagnetic Unpaired e- Ti 4s 3d 4p Paramagnetic Unpaired e- Ti2+ 4s 3d 4p If Titanium had lost its two “3d” electrons, the Titanium Ion would have been “diamagnetic (all electrons shared) The Titanium Ion actually shows properties of “Paramagnetism” (The mass of the titanium ion is affected when placed in a magnetic field)

Increasing “Paramagnetism” Iron (Fe)  Iron III (Fe3+) Fe ([Ar] 4s23d6)  Fe3+ ([Ar] 3d5) e- Fe 4s 3d 4p Fe3+ 4s 3d 4p The loss of the 2 4s electrons and one of its paired 3d electrons results in “increased” paramagnetism

Practice Problem Use condensed electron configuration to write the reaction for the formation of Mn2+ ion, and predict whether the ion is paramagnetic Manganese Mn (Z = 25) Mn ([Ar] 4s23d5) → Mn2+ ([Ar] 3d5) + 2e- Mn 4s 3d 4p Mn2+ 4s 3d 4p Rule: Remove “ns” electrons first The Mn2+ ion is “paramagnetic”

Practice Problem Use condensed electron configuration to write the reaction for the formation of Cr3+ ion, and predict whether the ion is paramagnetic Chromium - Cr (Z = 24) Cr ([Ar] 4s13d5) → Cr3+ ([Ar] 3d3) + 3e- Cr 4s 3d 4p Cr3+ 4s 3d 4p Note irregularity for Cr: 4s subshell fills before 3d subshell is complete Rule: Remove “ns” electrons first The Cr3+ ion is “paramagnetic”

Summary Equations N = Principal quantum number (size, energy) values = 1, 2, 3 …. l = Angular Momentum (orbital shape) values = 0  n (s), 1(p), 2(d), 3(f) ml = magnetic (orbital orientation) values = -l l Ms = Spin (direction) values = -1/2 & +1/2

George Mason University General Chemistry 211 Chapter 8

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George Mason University General Chemistry 211 Chapter 8

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