Chapter 8 Periodic Properties of the Elements

Chapter 8 Periodic Properties of the Elements
Chemistry: A Molecular Approach, 1st Ed. Nivaldo Tro Chapter 8 Periodic Properties of the Elements Roy Kennedy Massachusetts Bay Community College Wellesley Hills, MA 2007, Prentice Hall

Discovery Elements are organized according to their physical and chemical properties in the Periodic Table Historically, several people contributed to this effort in the late 19th century: Dobereiner- “triads”: Ca,Sr,Ba; Li, Na, K (1817) Newlands- similarity between every eighth element (1864) Mendeleev (1870) - Arranged elements according to atomic mass. Similar elements were arranged together in a “group”

Question Dobereiner’s “triads” (1817)
Determine the Atomic mass of Sr by averaging the masses of Ca and Ba 87.62 amu

Newlands Law of Octaves
In 1863, he suggested that elements be arranged in “octaves” because he noticed (after arranging the elements in order of increasing atomic mass) that certain properties repeated every 8th element Law of Octaves

Newlands “Law of Octaves”
Newlands (1864) arranged the elements as follows: Did not contain the Noble gases (yet to be discovered) Not valid for elements with Z > 20 His table had 7 groups instead of 8 Li Be B C N O F Na Mg Al Si P S Cl K Ca

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 pattern to predict properties of undiscovered elements where atomic mass order did not fit other properties, he re-ordered by other properties Te & I Tro, Chemistry: A Molecular Approach

Mendeleev's Predictions
Tro, Chemistry: A Molecular Approach

The Periodic Table Periodic Law - Both physical and chemical properties of the elements vary periodically with increasing atomic mass Exceptions Te/I, Ar/K, Co/Ni Placed Te (M = 127.6) ahead of I (M=126.9) because Te was similar to Se and S, and I was similar to Cl and Br Moseley showed listing elements based on atomic no. resolved these issues

What vs. Why 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 Laws summarize behavior while theories explain them! Quantum Mechanics is a theory that explains why the periodic trends in the properties exist Tro, Chemistry: A Molecular Approach

Electron Spin experiments by Stern and Gerlach showed a beam of silver atoms is split in two by a magnetic field the experiment reveals that 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” and the other half will have a net magnetic field pointing “South” Tro, Chemistry: A Molecular Approach

Electron Spin Experiment

Energy Shells and Subshells
n = principal quantum no. (overall size) l = angular momentum quantum no. (shape of orbital) ml = magnetic quantum no. (orientation of orbital)

Spin Quantum Number, ms spin quantum number describes how the electron spins on its axis clockwise or counterclockwise spin up or spin down spins must cancel in an orbital paired ms can have values of ±½ Orbital diagram for H Tro, Chemistry: A Molecular Approach

Pauli Exclusion Principle
no two electrons in an atom may have the same set of 4 quantum numbers therefore no orbital may have more than 2 electrons, and they 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 Tro, Chemistry: A Molecular Approach

Allowed Quantum Numbers

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 since they are in the same orbital, they must have opposite spins n l ml ms first electron 1 +½ second -½ Tro, Chemistry: A Molecular Approach

Electron Configurations
the ground state of the electron is the lowest energy orbital it can occupy the distribution of electrons into the various orbitals in an atom in its ground state is called its electron configuration the number designates the principal energy level the letter designates the sublevel and type of orbital the superscript designates the number of electrons in that sublevel He = 1s2 Tro, Chemistry: A Molecular Approach

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 1 electron orbital with 2 electrons Tro, Chemistry: A Molecular Approach

Sublevel Splitting in Multielectron Atoms
the sublevels in each principal energy level (n = 2…) of Hydrogen all have the same energy (empty in lowest state) – called degenerate for multielectron atoms, the energies of the sublevels are split caused by electron-electron repulsion 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) Tro, Chemistry: A Molecular Approach

Penetration & Shielding
Consider bringing an e- close to Li+ 1s2 Tro, Chemistry: A Molecular Approach

Penetrating and Shielding
the radial distribution function shows that the 2s orbital penetrates more deeply into the 1s orbital than does the 2p the weaker penetration of the 2p sublevel means that electrons in the 2p sublevel experience more repulsive force, they 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 by 1s e- the result is that the electrons in the 2s sublevel are lower in energy than the electrons in the 2p 2p is shielded Tro, Chemistry: A Molecular Approach

6d 7s 5f 6p 5d 6s 4f 5p 4d 5s 4p 3d Energy 4s 3p 3s 2p 2s 1s
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 becomes smaller for higher energy levels 2s 2p 1s

Order of Subshell Filling in Ground State Electron Configurations
start by drawing a diagram putting each energy shell on a row and listing the subshells, (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 Tro, Chemistry: A Molecular Approach

Tro, Chemistry: A Molecular Approach

Filling the Orbitals with Electrons
energy shells fill from lowest energy to high subshells fill from lowest energy to high s → p → d → f Aufbau Principle orbitals that are in the same subshell have the same energy no more than 2 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 Tro, Chemistry: A Molecular Approach

Mg Z = 12, so Mg has 12 protons and 12 electrons
Example 8.1 – Write the Ground State Electron Configuration and Orbital Diagram and of Magnesium. Determine the atomic number of the element from the Periodic Table This gives the number of protons and electrons in the atom Mg Z = 12, so Mg has 12 protons and 12 electrons Tro, Chemistry: A Molecular Approach

Example 8.1 – Write the Ground State Electron Configuration and Orbital Diagram and of Magnesium.
Draw 9 boxes to represent the first 3 energy levels s and p orbitals since there are only 12 electrons, 9 should be plenty 1s 2s 2p 3s 3p Tro, Chemistry: A Molecular Approach

Add one electron to each box in a set, then pair the electrons before going to the next set until you use all the electrons When pair, put in opposite arrows          1s 2s 2p 3s 3p Tro, Chemistry: A Molecular Approach

Use the diagram to write the electron configuration Write the number of electrons in each set as a superscript next to the name of the orbital set 1s22s22p63s2 = [Ne]3s2 1s 2s 2p 3s 3p  Tro, Chemistry: A Molecular Approach

Valence Electrons 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 Tro, Chemistry: A Molecular Approach

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) Tro, Chemistry: A Molecular Approach

Electron Configurations

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 Tro, Chemistry: A Molecular Approach

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

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 5 valence electrons Tro, Chemistry: A Molecular Approach

Transition Elements for the d block metals, the principal energy level of the d orbital being filled 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 5 valence electrons Tro, Chemistry: A Molecular Approach

Na (at. no. 11) Te (at. no. 52) Tc (at. no. 43)
Practice – Use the Periodic Table to write the short electron configuration and orbital diagram for each of the following Na (at. no. 11) Te (at. no. 52) Tc (at. no. 43) Tro, Chemistry: A Molecular Approach

Na (at. no. 11) [Ne]3s1 Te (at. no. 52) [Kr]5s24d105p4
Practice – Use the Periodic Table to write the short electron configuration and orbital diagram for each of the following Na (at. no. 11) [Ne]3s1 Te (at. no. 52) [Kr]5s24d105p4 Tc (at. no. 43) [Kr]5s24d5 3s 5s 4d 5p 5s 4d Tro, Chemistry: A Molecular Approach

Properties & Electron Configuration
elements in the same column have similar chemical and physical properties because they have the same number of valence electrons in the same kinds of orbitals Tro, Chemistry: A Molecular Approach

Electron Configuration & Element Properties
the number of valence electrons largely determines the behavior of an element chemical and some physical since the number of valence electrons follows a Periodic pattern, the properties of the elements should also be periodic quantum mechanical calculations show that 8 valence electrons should result in a very unreactive atom, an atom that is very stable – and the noble gases, that have 8 valence electrons are all very stable and unreactive conversely, elements that have either one more or one less electron should be very reactive – and the halogens are the most reactive nonmetals and alkali metals the most reactive metals as a group Tro, Chemistry: A Molecular Approach

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 Tro, Chemistry: A Molecular Approach

Electron Configuration of Anions in their Ground State
anions are formed when atoms gain enough electrons to have 8 valence electrons filling the s and p sublevels of the valence shell the sulfur atom has 6 valence electrons S atom = 1s22s22p63s23p4 in order to have 8 valence electrons, it must gain 2 more S2- anion = 1s22s22p63s23p6 Tro, Chemistry: A Molecular Approach

Electron Configuration of Cations in their Ground State
cations are formed when an 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 2 valence electrons Mg atom = 1s22s22p63s2 when it forms a cation, it loses its valence electrons Mg2+ cation = 1s22s22p6 Tro, Chemistry: A Molecular Approach

Electron Configuration of Atoms in their Ground State
Electrons in some elements are “promoted” to form more stable configurations, e.g. Cu = 29 electrons Expected: 1s22s22p63s23p64s23d9 Actual: 1s22s22p63s23p64s13d10 Cr = 24 electrons Expected: 1s22s22p63s23p64s23d4 Actual: 1s22s22p63s23p64s13d5 The usual reason given for this configuration is that the half-filled or completely filled d-orbitals are more stable as compared to those which are not. Tro, Chemistry: A Molecular Approach

Trend in Atomic Radius – Main Group
Different methods for measuring the radius of an atom, and they give slightly different trends 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 Tro, Chemistry: A Molecular Approach

Effective Nuclear Charge
in a multi-electron system, electrons are simultaneously attracted to the nucleus and repelled by each other outer electrons are shielded from full strength of nucleus screening effect effective nuclear charge is net positive charge that is attracting a particular electron Z is nuclear charge, S is electrons in lower energy levels electrons in same energy level contribute to screening, but very little effective nuclear charge on sublevels trend, s > p > d > f Zeffective = Z - S Tro, Chemistry: A Molecular Approach

Screening & Effective Nuclear Charge

Trends in Atomic Radius Transition Metals
increase in size down the Group atomic radii of transition metals roughly the same size across the d block must less difference than across main group elements valence shell ns2, not the d electrons effective nuclear charge on the ns2 electrons approximately the same Tro, Chemistry: A Molecular Approach

Example 8.5 – Choose the Larger Atom in Each Pair
N or F? C or Ge? N or Al? Al or Ge? N is further left so larger Ge is further down so larger Al is further down and left so larger Opposing trends Tro, Chemistry: A Molecular Approach

Electron Configuration of Cations in their Ground State
cations form when the atom loses electrons from the valence shell for transition metals electrons, may be removed from the sublevel closest to the valence shell Al atom = 1s22s22p63s23p1 Al+3 ion = 1s22s22p6 Fe atom = 1s22s22p63s23p64s23d6 Fe+2 ion = 1s22s22p63s23p63d6 Fe+3 ion = 1s22s22p63s23p63d5 Cu atom = 1s22s22p63s23p64s13d10 Cu+1 ion = 1s22s22p63s23p63d10 Tro, Chemistry: A Molecular Approach

Magnetic Properties of Transition Metal Atoms & Ions
Ferromagnetism – certain mateirals form permanent magnets electron configurations that result in unpaired electrons mean that in the presence of an externally applied magnetic field the atom or ion will have a net magnetic field – this is called paramagnetism 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 slightly repelled by an externally applied magnetic field both Zn atoms and Zn2+ ions not magnetic (diamagnetic), showing that the two 4s electrons are lost before the 3d Zn atoms [Ar]4s23d10 Zn2+ ions [Ar]4s03d10 Tro, Chemistry: A Molecular Approach

Fe Z = 26 (elemental iron is magnetic!) previous noble gas = Ar
Example 8.6 – Write the Electron Configuration and Determine whether the Fe atom and Fe3+ ion are Paramagnetic or Diamagnetic Fe Z = 26 (elemental iron is magnetic!) previous noble gas = Ar 18 electrons 4s 3d Fe3+ ion = [Ar]4s03d5 unpaired electrons paramagnetic 4s 3d Fe atom = [Ar]4s23d6 unpaired electrons paramagnetic Tro, Chemistry: A Molecular Approach

Trends in Ionic Radius Ions in same group have same charge
Ion size increases down the group higher valence shell, larger Cations smaller than neutral atom; Anions bigger than neutral atom Cations smaller than anions except Rb+1 & Cs+1 bigger or same size as F-1 and O-2 Larger positive charge = smaller cation for isoelectronic species isoelectronic = same electron configuration Larger negative charge = larger anion for isoelectronic series Tro, Chemistry: A Molecular Approach

Periodic Pattern - Ionic Radius (Å)
H +1 -1 2A A A A A A Li +1 Be +2 B +3 C -4 N -3 O -2 -1 F 0.68 0.31 0.23 1.71 1.40 1.33 Na +1 Mg +2 Al +3 Al +3 Si -4 P -3 S -2 Cl -1 0.66 0.51 2.12 0.97 1.84 1.81 K +1 Ca +2 Ga +3 +1 0.62 Ge -4 As -3 Se -2 Br -1 1.33 0.99 2.22 1.98 1.96 Rb +1 Sr +2 0.81 In +3 +1 0.71 Sn +4 +2 Sb Te -2 I -1 1.13 1.47 2.21 2.20 Cs +1 Ba +2 0.95 Tl +3 +1 Pb +4 +2 0.84 Bi 1.69 1.35

Ionization Energy minimum energy needed to remove an electron from an atom gas state endothermic process valence electron easiest to remove M(g) + IE1  M+(g) + 1 e- M+(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. Tro, Chemistry: A Molecular Approach

General Trends in 1st Ionization Energy
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 Tro, Chemistry: A Molecular Approach

Example 8.8 – Choose the Atom in Each Pair with the Higher First Ionization Energy
Al or S As or Sb N or Si O or Cl? opposing trends Al or S As or Sb N or Si, Si is further down & left Al or S, Al is further left Al or S As or Sb, Sb is further down Tro, Chemistry: A Molecular Approach

Irregularities in the Trend
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 B or Be? Why? Which is easier to remove an electron from N or O? Why? Tro, Chemistry: A Molecular Approach

Irregularities in the First Ionization Energy Trends
   Be Be+ 1s 2s 2p 1s 2s 2p To ionize Be you must break up a full sublevel, cost extra energy B  1s 2s 2p  B+  1s 2s 2p When you ionize B you get a full sublevel, costs less energy Tro, Chemistry: A Molecular Approach

Irregularities in the First Ionization Energy Trends
 1s 2s 2p  N+  1s 2s 2p  To ionize N you must break up a half-full sublevel, cost extra energy  O 1s 2s 2p   O+  1s 2s 2p When you ionize O you get a half-full sublevel, costs less energy Tro, Chemistry: A Molecular Approach

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 Tro, Chemistry: A Molecular Approach

Trends in Electron Affinity
energy released when an neutral atom gains an electron gas state M(g) + 1e-  M-(g) + EA defined as exothermic (-), but may actually be endothermic (+) alkali earth metals & noble gases endothermic, WHY? more energy released (more -); the larger the EA generally increases across period becomes more negative from left to right not absolute lowest EA in period = alkali earth metal or noble gas highest EA in period = halogen Tro, Chemistry: A Molecular Approach

Metallic Character Metals Nonmetals metallic character increases left
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 electrical and thermal insulators most oxides are acidic and molecular form anions and polyatomic anions gain electrons in reactions - reduced metallic character increases left metallic character increase down Tro, Chemistry: A Molecular Approach

Example 8.9 – Choose the More Metallic Element in Each Pair
Sn or Te P or Sb, Sb is further down Sn or Te, Sn is further left Sn or Te P or Sb Ge or In, In is further down & left Sn or Te P or Sb Ge or In S or Br? opposing trends Tro, Chemistry: A Molecular Approach

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 Tro, Chemistry: A Molecular Approach

2 Na(s) + 2 H2O(l)  2 NaOH(aq) + H2(g)

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 increases down the column density increases down the column in general, the increase in mass is greater than the increase in volume Tro, Chemistry: A Molecular Approach

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) Tro, Chemistry: A Molecular Approach

reaction between rubidium metal and liquid water Rb(s) + 2H2O(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) Tro, Chemistry: A Molecular Approach

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 found in all states at room temperature) Tro, Chemistry: A Molecular Approach

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 undesirable Tro, Chemistry: A Molecular Approach

Trends in the Noble Gases
melting point and boiling point increases 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 Tro, Chemistry: A Molecular Approach

Chapter 8 Periodic Properties of the Elements

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