Chapter 8 Periodic Properties of the Elements

Nerve Transmission Movement of ions across cell membranes is the basis for the transmission of nerve signals Na+ and K+ ions are pumped across membranes in opposite directions through ion channels Na+ out and K+ in The ion channels can differentiate Na+ from K+ by their difference in size Ion size and other properties of atoms are periodic properties – properties whose values can be predicted based on the element’s position on the Periodic Table

Mendeleev (1834–1907) Order 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 Elements with similar properties  same column Used pattern to predict properties of undiscovered elements Where atomic mass order did not fit other properties, he re-ordered by other properties Te & I

a = acidic oxide, b = basic oxide, a/b = amphoteric oxide Periodic Pattern H NM H2O a/b 1.0 H2 Li Be B C N O F Na Mg Al Si P S Cl K Ca M M NM NM NM Li2O b BeO a/b B2O3 a CO2 a N2O5 a NM O2 NM 6.9 LiH 9.0 BeH2 10.8 BH3 12.0 CH4 14.0 NH3 16.0 H2O 19.0 HF M M M NM NM NM Na2O b MgO b Al2O3 a/b M/NM SiO2 a P4O10 a SO3 a Cl2O7 a 23.0 NaH 24.3 MgH2 27.0 AlH3 28.1 SiH4 31.0 PH3 32.1 H2S 35.5 HCl M M K2O b CaO b 39.1 KH 40.1 CaH2 M = metal, NM = nonmetal, M/NM = metalloid a = acidic oxide, b = basic oxide, a/b = amphoteric oxide

Mendeleev's Predictions

What vs. Why Mendeleev’s Periodic Law Predicts what the properties of an element will be based on its position on the table Does NOT explain why the pattern exists Quantum Mechanics theory that explains why the periodic trends in the properties exist and knowing WHY allows us to predict WHAT

Electron Configurations Quantum-mechanical theory describes the behavior of electrons in atoms The electrons in atoms exist in orbitals A description of the orbitals occupied by electrons is called an electron configuration number of electrons in the orbital 1s1 principal energy level of orbital occupied by the electron sublevel of orbital occupied by the electron

How Electrons Occupy Orbitals Calculations with Schrödinger’s equation show hydrogen’s one electron occupies the lowest energy orbital in the atom Schrödinger’s equation calculations for multi-electron atoms cannot be solved exactly due to additional terms added for electron-electron interactions Approximate solutions show the orbitals to be hydrogen-like Two additional concepts affect multielectron atoms: electron spin and energy splitting of sublevels

spinning charged particles generate a magnetic field Electron Spin Stern and Gerlach’s Experiment: a beam of silver atoms is split in two by a magnetic field Conclusion: the electrons spin on their axis As they spin, they generate a magnetic field spinning charged particles generate a magnetic field If there is an even number of electrons about half the atoms will have a net magnetic field pointing “north” the other half will have a net magnetic field pointing “south”

Electron Spin Experiment

The Property of Electron Spin Spin is a fundamental property of all electrons All electrons have the same amount of spin The orientation of the electron spin is quantized, it can only be in one direction or its opposite spin up or spin down Spin Quantum Number, ms Spin adds a fourth quantum number to the description of electrons in an atom not in the Schrödinger equation

Spin Quantum Number, ms, and Orbital Diagrams ms can have values of +½ or −½ Orbital Diagrams a square represents each orbital a half-arrow represents each electron A half-arrow pointing up represents an electron in an orbital with spin up Spins must cancel in an orbital Paired with reversed spins

Orbital Diagrams We often represent an orbital as a square and the electrons in that orbital as arrows the direction of the arrow represents the spin of the electron unoccupied orbital orbital with one electron orbital with two electrons

Pauli Exclusion Principle No two electrons in an atom may have the same set of four quantum numbers No orbital may have more than two electrons, and these must have with opposite spins Knowing the number orbitals in a sublevel allows us to determine the maximum number of electrons in the sublevel s sublevel has 1 orbital, therefore it can hold 2 electrons p sublevel has 3 orbitals, therefore it can hold 6 electrons d sublevel has 5 orbitals, therefore it can hold 10 electrons f sublevel has 7 orbitals, therefore it can hold 14 electrons

Allowed Quantum Numbers Values Number of Values Significance Principal, n 1, 2, 3, … - Size and energy of the orbital Azimuthal , l s (l = 0) p (l = 1) d (l = 2) f (l = 3) 4 Shape of orbital Magnetic, ml -l, …, 0, …+l 2l+1 Orientation of Orbital Spin, ms -, + 2 Direction of e- spin

Quantum Numbers of Helium’s Electrons Helium has two electrons Both electrons are in the first energy level Both electrons are in the s orbital of the first energy level Because they are in the same orbital, they must have opposite spins

Sublevel Splitting in Multielectron Atoms The sublevels in each principal energy shell of Hydrogen all have the same energy or other single electron systems We call orbitals with the same energy degenerate For multi-electron atoms, the energies of the sublevels are split caused by charge interaction, shielding and penetration The lower the value of the l quantum number, the less energy the sublevel has s (l = 0) < p (l = 1) < d (l = 2) < f (l = 3)

Coulomb’s Law Describes attraction & repulsion between charged particles For like charges: the potential energy (E) is positive E decreases as the particles get farther apart (r increases) For opposite charges: E is negative E becomes more negative as the particles get closer together (r decreases) The strength of the interaction increases as the size of the charges increases electrons are more strongly attracted to a nucleus with a 2+ charge than a nucleus with a 1+ charge

Shielding & Effective Nuclear Charge Each electron in a multi-electron atom experiences both the attraction to the nucleus the repulsion by other electrons in the atom Repulsions cause electrons to have a net reduced attraction to the nucleus – it is shielded from the nucleus The total amount of attraction that an electron feels for the nucleus is called the effective nuclear charge of the electron

Penetration Movement of an electron toward the nucleus (r ) As r decreases: An electron is more attracted to the nucleus The better an outer electron is at penetrating through the electron cloud of inner electrons, the more attraction it will have for the nucleus The degree of penetration is related to the orbital’s radial distribution function in particular, the distance the maxima of the function are from the nucleus

Shielding & Penetration

Penetration and Shielding The radial distribution function shows 2s orbital penetrates more deeply into the 1s orbital than does the 2p The weaker penetration of the 2p sublevel Electrons in the 2p sublevel experience more repulsive force These electrons are more shielded from the attractive force of the nucleus The deeper penetration of the 2s electrons means electrons in the 2s sublevel experience a greater attractive force to the nucleus and are not shielded as effectively

Effect of Penetration and Shielding (Remember that orbitals with the same energy are said to be degenerate) Penetration causes the energies of sublevels in the same principal level to NOT be degenerate In the 4th and 5th principal levels, the effects of penetration become so important, causing the s orbital to lie lower in energy than the d orbitals of the previous principal level The energy separations between one set of orbitals and the next become smaller beyond the 4s the ordering can vary among elements creates variations in the e- configurations of the transition metals and their ions

6s 6p 6d 7s 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d Energy Notice the following: because of penetration, sublevels within an energy level are not degenerate penetration of the 4th and higher energy levels is so strong that their s sublevel is lower in energy than the d sublevel of the previous energy level the energy difference between levels decreases for higher energy levels causing anomalous electron configurations for certain elements 2s 2p 1s

Filling the Orbitals with Electrons Energy levels and sublevels fill from lowest energy to high s → p → d → f Aufbau Principle Orbitals that are in the same sublevel have the same energy No more than two electrons per orbital Pauli Exclusion Principle When filling orbitals that have the same energy, place one electron in each before completing pairs Hund’s Rule

Electron Configuration of Atoms in their Ground State The electron configuration is a listing of the sublevels in order of filling with the number of electrons in that sublevel written as a superscript. Kr = 36 electrons = 1s22s22p63s23p64s23d104p6 A short-hand way of writing an electron configuration is to use the symbol of the previous noble gas in [Ng] to represent all the inner electrons, then just write the last set. Rb = 37 electrons = 1s22s22p63s23p64s23d104p65s1 = [Kr]5s1

Order of Sublevel Filling in Ground State Electron Configurations Start by drawing a diagram putting each energy shell on a row and listing the sublevels, (s, p, d, f), for that shell in order of energy (left-to-right) 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s Next, draw arrows through the diagonals, looping back to the next diagonal each time

Electron Configurations

Practice — write the full ground state orbital diagram and electron configuration of potassium.

Practice — write the full ground state orbital diagram and electron configuration of potassium, answer K Z = 19, therefore 19 e− s sublevel holds 2 e− p sublevel holds 6 e− d sublevel holds 10 e− f sublevel holds 14 e− 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 1s 2s 2p 3s 3p 4s                 2 e− +2 = 4e− +6 +2 = 12e− Therefore the electron configuration is 1s22s22p63s23p64s1 +6 +2 = 20e− Based on the order of sublevel filling, we will need the first six sublevels

Core Electrons = The electrons in lower energy shells OBSERVATION: Valence Electrons Valence Electrons = The electrons in all the sublevels with the highest principal energy shell Core Electrons = The electrons in lower energy shells OBSERVATION: one of the most important factors in the way an atom behaves, both chemically and physically, is the number of valence electrons

Electron Configuration of Atoms in their Ground State Kr = 36 electrons 1s22s22p63s23p64s23d104p6 there are 28 core electrons and 8 valence electrons Rb = 37 electrons 1s22s22p63s23p64s23d104p65s1 [Kr]5s1 For the 5s1 electron in Rb the set of quantum numbers is n = 5, l = 0, ml = 0, ms = +½ For an electron in the 2p sublevel, the set of quantum numbers is n = 2, l = 1, ml = −1 or (0,+1), and ms = −½ or (+½)

Electron Configuration & the Periodic Table The Group number corresponds to the number of valence electrons The length of each “block” is the maximum number of electrons the sublevel can hold The Period number corresponds to the principal energy level of the valence electrons

s1 s2 1 2 3 4 5 6 7 p1 p2 p3 p4 p5 p6 s2 d1 d2 d3 d4 d5 d6 d7 d8 d9 d10 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14 f14d1

Electron Configuration from the Periodic Table 1 2 3 4 5 6 7 2A 3A 4A 5A 6A 7A Ne 3s2 P 3p3 P = [Ne]3s23p3 P has five valence electrons

Transition Elements For the d block metals, the principal energy level is one less than valence shell one less than the Period number sometimes s electron “promoted” to d sublevel Zn Z = 30, Period 4, Group 2B [Ar]4s23d10 4s 3d For the f block metals, the principal energy level is two less than valence shell two less than the Period number they really belong to sometimes d electron in configuration Eu Z = 63, Period 6 [Xe]6s24f 7 6s 4f

Electron Configuration from the Periodic Table 1 2 3 4 5 6 7 2A 3A 4A 5A 6A 7A 3d10 Ar As 4s2 4p3 As = [Ar]4s23d104p3 As has five valence electrons

Na (#11) Te (#52) Tc (#43) [Ne]3s1 [Kr]5s24d105p4 [Kr]5s24d5 Practice – Use the Periodic Table to write the short electron configuration and short orbital diagram for each of the following Na (#11) Te (#52) Tc (#43) 3s [Ne]3s1 5s 5p 4d [Kr]5s24d105p4 5s 4d [Kr]5s24d5

Irregular Electron Configurations Because of sublevel splitting: the 4s sublevel is lower in energy than the 3d the 4s fills before the 3d But the difference in energy is not large Some transition metals have irregular e- configurations the ns only partially fills before the (n−1)d or doesn’t fill at all Therefore, their electron configuration must be found experimentally

Irregular Electron Configurations Expected Cr = [Ar]4s23d4 Cu = [Ar]4s23d9 Mo = [Kr]5s24d4  Ru = [Kr]5s24d6 Pd = [Kr]5s24d8 Found Experimentally Cr = [Ar]4s13d5 Cu = [Ar]4s13d10 Mo = [Kr]5s14d5 Ru = [Kr]5s14d7 Pd = [Kr]5s04d10

Properties & Electron Configuration The properties of the elements follow a periodic pattern elements in the same column have similar properties the elements in a period show a pattern that repeats The Explanation from the quantum-mechanical model: the number of valence electrons and the types of orbitals they occupy are also periodic

The Noble Gas Electron Configuration The noble gases have eight valence electrons. except for He, which has only two electrons We know the noble gases are especially non-reactive He and Ne are practically inert The reason the noble gases are so non- reactive is that the electron configuration of the noble gases is especially stable

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 one electron, resulting in the same electron configuration as a noble gas forming a cation with a 1+ charge

Everyone Wants to Be Like a Noble Gas! The Halogens The electron configurations of 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 Forming an anion with charge 1− 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

Eight Valence Electrons Quantum mechanical calculations show that eight valence electrons should result in a very unreactive atom an atom that is very stable the noble gases have eight valence electrons and are all very stable and unreactive He has two valence electrons, but that fills its valence shell Conversely, elements that have either one more or one less electron should be very reactive the halogen atoms have seven valence electrons and are the most reactive nonmetals the alkali metals have one more electron than a noble gas atom and are the most reactive metals as a group

Electron Configuration & 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

Electron Configuration of Anions in Their Ground State Anions are formed when nonmetal atoms gain enough electrons to have eight valence electrons filling the s and p sublevels of the valence shell The sulfur atom has six valence electrons S atom = 1s22s22p63s23p4 To have eight valence electrons, sulfur must gain two more S2− anion = 1s22s22p63s23p6

Electron Configuration of Cations in Their Ground State Cations are formed when a metal atom loses all its valence electrons resulting in a new lower energy level valence shell however the process is always endothermic The magnesium atom has two valence electrons Mg atom = 1s22s22p63s2 When magnesium forms a cation, it loses its valence electrons Mg2+ cation = 1s22s22p6

Trend in Atomic Radius – Main Group There are several methods for measuring the radius of an atom, and they give slightly different numbers van der Waals radius = nonbonding covalent radius = bonding radius atomic radius is an average radius of an atom based on measuring large numbers of elements and compounds Atomic Radius INCREASES down group valence shell farther from nucleus effective nuclear charge fairly close Atomic Radius DECREASES across period (left to right) adding electrons to same valence shell effective nuclear charge increases valence shell held closer

Outer electrons are shielded from nucleus by the core electrons Shielding In a multi-electron system, electrons are simultaneously attracted to the nucleus and repelled by each other Outer electrons are shielded from nucleus by the core electrons screening or shielding effect outer electrons do not effectively screen for each other The shielding causes the outer electrons to not experience the full strength of the nuclear charge

Effective Nuclear Charge The effective nuclear charge is net positive charge that is attracting a particular electron Z is the nuclear charge, S is the number of electrons in lower energy levels electrons in same energy level contribute to screening, but very little so are not part of the calculation trend is s > p > d > f Zeffective = Z − S

Screening & Effective Nuclear Charge 55

Quantum-Mechanical Explanation for the Group Trend in Atomic Radius The size of an atom is related to the distance the valence electrons are from the nucleus The larger the orbital an electron is in, the farther its most probable distance will be from the nucleus and the less attraction it will have for the nucleus Traversing down a group adds a principal energy level The larger the principal energy level an orbital is in, the larger its volume  quantum-mechanics predicts the atoms should get larger down a column

Quantum-Mechanical Explanation for the Period Trend in Atomic Radius The larger the effective nuclear charge an electron experiences, the stronger the attraction it will have for the nucleus The stronger the attraction the valence electrons have for the nucleus, the closer their average distance will be to the nucleus Traversing across a period increases the effective nuclear charge on the valence electrons  quantum-mechanics predicts the atoms should get smaller across a period

Example 8.5: Choose the Larger Atom in Each Pair N or F, N is farther left N or F C or Ge, Ge is farther down N or F C or Ge N or Al, Al is farther down & left N or F C or Ge N or Al Al or Ge? opposing trends

Practice – Choose the Larger Atom in Each Pair C or O Li or K C or Al Se or I C or O C is farther left in the period Li or K K is farther down the column C or Al Al is farther left and down Se or I ? opposing trends

Trends in Atomic Radius Transition Metals Atoms in the same group increase in size down the column Atomic radii of transition metals roughly the same size across the d block much less difference than across main group elements valence shell ns2, not the (n−1)d electrons effective nuclear charge on the ns2 electrons approximately the same

Electron Configurations of Main Group Cations in Their Ground State Cations form when the atom loses electrons from the valence shell Al atom = 1s22s22p63s23p1 Al3+ ion = 1s22s22p6

Electron Configurations of Transition Metal Cations in Their Ground State When transition metals form cations, the first electrons removed are the valence electrons, even though other electrons were added after Electrons may also be removed from the sublevel closest to the valence shell after the valence electrons The iron atom has two valence electrons Fe atom = 1s22s22p63s23p64s23d6 When iron forms a cation, it first loses its valence electrons Fe2+ cation = 1s22s22p63s23p63d6 It can then lose 3d electrons Fe3+ cation = 1s22s22p63s23p63d5

Magnetic Properties of Transition Metal Atoms & Ions Paramagnetism = Electron configurations that result in unpaired electrons mean that the atom or ion will have a net magnetic field will be attracted to a magnetic field Diamagnetism = Electron configurations that result in all paired electrons mean that the atom or ion will have no magnetic field slightly repelled by a magnetic field

Transition Metal Atoms and Ions: Electron Configuration & Magnetic Properties Both Zn atoms and Zn2+ ions are diamagnetic showing that the two 4s electrons are lost before the 3d Zn atoms [Ar]4s23d10 Zn2+ ions [Ar]4s03d10 Ag forms both Ag+ ions and, rarely, Ag2+ Ag atoms [Kr]5s14d10 are paramagnetic Ag+ ions [Kr]4d10 are diamagnetic Ag2+ ions [Kr]4d9 are paramagnetic

Fe3+ ion = [Ar]4s03d5 Unpaired electrons Paramagnetic Example 8.6c: write the electron configuration and determine whether the Fe atom and Fe3+ ion are paramagnetic or diamagnetic Fe Z = 26 Previous noble gas = Ar 18 electrons 4s 3d Fe3+ ion = [Ar]4s03d5 Unpaired electrons Paramagnetic 4s 3d Fe atom = [Ar]4s23d6 Unpaired electrons Paramagnetic

Mn = [Ar]4s23d5 paramagnetic Practice – determine whether the following are paramagnetic or diamagnetic Mn Sc3+ Mn = [Ar]4s23d5 paramagnetic 4s 3d Sc = [Ar]4s23d1 Sc3+ = [Ar] diamagnetic

Trends in Ionic Radius Ions in same group have same charge Ion size increases down the column higher valence shell, larger Cations smaller than neutral atoms; anions larger than neutral atoms Cations smaller than anions except Rb+ & Cs+ bigger or same size as F− and O2− Larger positive charge = smaller cation for isoelectronic species isoelectronic = same electron configuration Larger negative charge = larger anion

Periodic Pattern – Ionic Radius (Å)

Quantum-Mechanical Explanation for the Trends in Cation Radius When atoms form cations, the valence electrons are removed The farthest electrons from the nucleus then are the p or d electrons in the (n − 1) energy level This results in the cation being smaller than the atom These “new valence electrons” also experience a larger effective nuclear charge than the “old valence electrons,” shrinking the ion even more Traversing down a group increases the (n − 1) level, causing the cations to get larger Traversing to the right across a period increases the effective nuclear charge for isoelectronic cations, causing the cations to get smaller

Quantum-Mechanical Explanation for the Trends in Anion Radius When atoms form anions, electrons are added to the valence shell The “new valence electrons” experience a smaller effective nuclear charge than the “old valence electrons,” The anion ends up being larger than the atom Traversing down a group increases the n level, causing the anions to get larger Traversing to the right across a period decreases the effective nuclear charge for isoelectronic anions, causing the anions to get larger

Example 8.7: 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 has been lost from Ca2+ the cation is always smaller than the neutral atom Br− or Kr the 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

Practice – Order the following sets by size (smallest to largest) Zr4+, Ti4+, Hf4+ Na+, Mg2+, F−, Ne I−, Br−, Ga3+, In+ same column & charge, therefore Ti4+ < Zr4+ < Hf4+ isoelectronic, therefore Mg2+ < Na+ < Ne < F− Ga3+ < In+ < Br− < I−

Ionization Energy Minimum energy needed to remove an electron from an atom or ion gas state endothermic process valence electron easiest to remove, lowest IE M(g) + IE1  M1+(g) + 1 e- M+1(g) + IE2  M2+(g) + 1 e- first ionization energy = energy to remove electron from neutral atom; 2nd IE = energy to remove from 1+ ion; etc.

General Trends in 1st Ionization Energy The larger the effective nuclear charge on the electron, the more energy it takes to remove it The farther the most probable distance the electron is from the nucleus, the less energy it takes to remove it 1st IE decreases down the group valence electron farther from nucleus 1st IE generally increases across the period effective nuclear charge increases

79

Quantum-Mechanical Explanation for the Trends in First Ionization Energy The strength of attraction is related to the most probable distance the valence e- are from the nucleus the effective nuclear charge the valence e- experience The larger the orbital an electron is in, the farther its most probable distance will be from the nucleus the less attraction it will have for the nucleus  quantum-mechanics predicts the atom’s first ionization energy should get lower down a column Traversing across a period increases the effective nuclear charge on the valence electrons  quantum-mechanics predicts the atom’s first ionization energy should get larger across a period

Example 8.8: Choose the atom in each pair with the larger first ionization energy Al or S, S is farther right Al or S As or Sb N or Si O or Cl? opposing trends Al or S As or Sb N or Si, N is farther up & right Al or S As or Sb, As is farther up

Practice – Choose the atom with the largest first ionization energy in each pair Mg or P Ag or Cu Ca or Rb P or Se ?

Exceptions in the 1st IE Trends First Ionization Energy generally increases from left to right across a Period Except from 2A to 3A, 5A to 6A  N 1s 2s 2p  O Be  1s 2s 2p B  Which is easier to remove an electron from, N or O? Why? Which is easier to remove an electron from B, or Be? Why?

Exceptions in the First Ionization Energy Trends, Be and B    Be Be+ 1s 2s 2p 1s 2s 2p To ionize Be, you must break up a full sublevel, costs extra energy B  1s 2s 2p  B+  1s 2s 2p When you ionize B you get a full sublevel, costs less energy

Exceptions in the First Ionization Energy Trends, N and O  1s 2s 2p  N+  1s 2s 2p  To ionize N you must break up a half-full sublevel, costs extra energy  O 1s 2s 2p   O+  1s 2s 2p When you ionize O you get a half-full sublevel, costs less energy

Trends in Successive Ionization Energies Removal of each successive electron costs more energy shrinkage in size due to having more protons than electrons outer electrons closer to the nucleus, therefore harder to remove Regular increase in energy for each successive valence electron Large increase in energy when start removing core electrons

Electron Affinity Energy released when an neutral atom gains an electron gas state M(g) + 1e−  M1−(g) + EA Defined as exothermic (−), but may actually be endothermic (+) some alkali earth metals & all noble gases are endothermic, WHY? The more energy that is released, the larger the electron affinity the more negative the number, the larger the EA

Trends in Electron Affinity Alkali metals decrease electron affinity down the column but not all groups do generally irregular increase in EA from 2nd period to 3rd period “Generally” increases across period becomes more negative from left to right not absolute Group 5A generally lower EA than expected because extra electron must pair Group 2A and 8A generally very low EA because added electron goes into higher energy level or sublevel Highest EA in any period = halogen

Properties of Metals & Nonmetals malleable & ductile shiny, lusterous, reflect light conduct heat and electricity most oxides basic and ionic form cations in solution lose electrons in reactions – oxidized Nonmetals brittle in solid state dull, non-reflective solid surface electrical and thermal insulators most oxides are acidic and molecular form anions and polyatomic anions gain electrons in reactions – reduced

Metallic Character Metallic character is how closely an element’s properties match the ideal properties of a metal more malleable and ductile, better conductors, and easier to ionize Metallic character decreases left-to-right across a period metals are found at the left of the period and nonmetals are to the right Metallic character increases down the column nonmetals are found at the top of the middle Main Group elements and metals are found at the bottom

Quantum-Mechanical Explanation for the Trends in Metallic Character Metals generally have smaller first ionization energies and nonmetals generally have larger electron affinities except for the noble gases  quantum-mechanics predicts the atom’s metallic character should increase down a column because the valence electrons are not held as strongly  quantum-mechanics predicts the atom’s metallic character should decrease across a period because the valence electrons are held more strongly and the electron affinity increases

Example 8.9: Choose the more metallic element in each pair Sn or Te, Sn is farther left Sn or Te P or Sb Ge or In S or Br? opposing trends Sn or Te P or Sb, Sb is farther down Sn or Te P or Sb Ge or In, In is farther down & left

Practice – Choose the more metallic element in each pair Mg or Al Si or Sn Br or Te Se or I ?

Trends in the Alkali Metals Atomic radius increases down the column Ionization energy decreases down the column Very low ionization energies good reducing agents, easy to oxidize very reactive, not found uncombined in nature react with nonmetals to form salts compounds generally soluble in water  found in seawater Electron affinity decreases down the column Melting point decreases down the column all very low MP for metals Density increases down the column except K in general, the increase in mass is greater than the increase in volume

Alkali Metals

Trends in the Halogens Atomic radius increases down the column Ionization energy decreases down the column Very high electron affinities good oxidizing agents, easy to reduce very reactive, not found uncombined in nature react with metals to form salts compounds generally soluble in water  found in seawater Reactivity increases down the column React with hydrogen to form HX, acids Melting point and boiling point increase down the column Density increases down the column in general, the increase in mass is greater than the increase in volume

Halogens

Reactions of Alkali Metals with Halogens Alkali metals are oxidized to the 1+ ion Halogens are reduced to the 1− ion The ions then attach together by ionic bonds The reaction is exothermic

Example 8.10: Write a balanced chemical reaction for the following Reaction between potassium metal and bromine gas K(s) + Br2(g)  K(s) + Br2(g)  K+ Br 2 K(s) + Br2(g)  2 KBr(s) (ionic compounds are all solids at room temperature)

Reactions of Alkali Metals with Water Alkali metals are oxidized to the 1+ ion H2O is split into H2(g) and OH− ion The Li, Na, and K are less dense than the water – so float on top The ions then attach together by ionic bonds The reaction is exothermic, and often the heat released ignites the H2(g)

Example 8.10: Write a balanced chemical reaction for the following Reaction between rubidium metal and liquid water Rb(s) + H2O(l)  Rb(s) + H2O(l)  Rb+(aq) + OH(aq) + H2(g) 2 Rb(s) + 2 H2O(l)  2 Rb+(aq) + 2 OH(aq) + H2(g) (alkali metal ionic compounds are soluble in water)

Example 8.10: Write a balanced chemical reaction for the following Reaction between chlorine gas and solid iodine Cl2(g) + I2(s)  Cl2(g) + I2(s)  ICl write the halogen lower in the column first assume 1:1 ratio, though others also exist 2 Cl2(g) + I2(s)  2 ICl(g) (molecular compounds are found in all states at room temperature, so predicting the state is not always possible)

Trends in the Noble Gases Atomic radius increases down the column Ionization energy decreases down the column very high IE Very unreactive only found uncombined in nature used as “inert” atmosphere when reactions with other gases would be undersirable Melting point and boiling point increase down the column all gases at room temperature very low boiling points Density increases down the column in general, the increase in mass is greater than the increase in volume

Noble Gases