Chapter 8 Periodic Properties of the Elements

Chapter 8Periodic

Properties of the Elements

2007, Prentice Hall

Chemistry: A Molecular Approach, 1st Ed.Nivaldo Tro

Roy KennedyMassachusetts Bay Community College

Wellesley Hills, MA

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

• 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

Tro, Chemistry: A Molecular Approach 7

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 propertiesTe & I


Mendeleev's Predictions

The Periodic Table

• Periodic Law - Both physical and chemical properties of the elements vary periodically with increasing atomic masso Exceptions Te/I, Ar/K, Co/Nio 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


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”


Electron Spin Experiment

Energy Shells and Subshellsn = 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 axisclockwise or counterclockwisespin up or spin down

• spins must cancel in an orbitalpaired

• ms can have values of ±½

Orbital diagram for H


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 sublevels sublevel has 1 orbital, therefore it can hold 2 electronsp sublevel has 3 orbitals, therefore it can hold 6 electronsd sublevel has 5 orbitals, therefore it can hold 10 electrons f sublevel has 7 orbitals, therefore it can hold 14 electrons


Allowed Quantum NumbersQuantum Number

Values Number of Values

Significance

Principal, n 1, 2, 3, ... - distance from nucleus

Azimuthal, l 0, 1, 2, ..., n-1 n shape of orbital

Magnetic, ml -l,...,0,...+l 2l + 1 orientation of orbital

Spin, ms -½, +½ 2 direction of electron 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

• since they are in the same orbital, they must have opposite spins

n l ml ms

first

electron1 0 0 +½

second

electron1 0 0 -½


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

http://wps.prenhall.com/wps/media/objects/166/170213/ElectronConfigurationsmov.html


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

orbital with1 electron

unoccupiedorbital

orbital with2 electrons


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 splitcaused by electron-electron repulsion

• the lower the value of the l quantum number, the less energy the sublevel hass (l = 0) < p (l = 1) < d (l = 2) < f (l = 3)


Penetration & Shielding

Consider bringing an e- close to Li+ 1s2


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

Ene

rgy

1s

7s

2s

2p

3s

3p3d

6s6p

6d

4s

4p4d

4f

5s

5p

5d5f

Notice the following:1. because of penetration, sublevels within

an energy level are not degenerate2. 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

3. the energy difference between levels becomes smaller for higher energy levels


Order of Subshell Fillingin Ground State Electron Configurations

1s

2s 2p

3s 3p 3d

4s 4p 4d 4f

5s 5p 5d 5f

6s 6p 6d

7s

start by drawing a diagramputting each energy shell ona row and listing the subshells, (s, p, d, f), for that shell in order of energy, (left-to-right)

next, draw arrows throughthe diagonals, looping back to the next diagonaleach time


Filling the Orbitals with Electrons

• energy shells fill from lowest energy to high• subshells fill from lowest energy to high

s → p → d → fAufbau Principle

• orbitals that are in the same subshell have the same energy

• no more than 2 electrons per orbitalPauli Exclusion Principle

• when filling orbitals that have the same energy, place one electron in each before completing pairsHund’s Rule


1. 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

Example 8.1 – Write the Ground State Electron Configuration and Orbital Diagram

and of Magnesium.


2. Draw 9 boxes to represent the first 3 energy levels s and p orbitals

a) since there are only 12 electrons, 9 should be plenty

1s 2s 2p 3s 3p

Example 8.1 – Write the Ground State Electron Configuration and Orbital

Diagram and of Magnesium.


3. 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






4. 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


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


Electron Configuration of Atoms in their Ground State

• Kr = 36 electrons1s22s22p63s23p64s23d104p6

there are 28 core electrons and 8 valence electrons

• Rb = 37 electrons1s22s22p63s23p64s23d104p65s1

[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 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


s1

s2

d1 d2 d3 d4 d5 d6 d7 d8 d9 d10

s2p1 p2 p3 p4 p5

p6

f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14 f14d1

1234567


Electron Configuration fromthe Periodic Table

P = [Ne]3s23p3

P has 5 valence electrons

3p3

P

Ne1234567

1A

2A 3A 4A 5A 6A 7A

8A

3s2

38

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

4s 3d

ZnZ = 30, Period 4, Group 2B[Ar]4s23d10

• for the f block metals, the principal energy level is two less than valence shell two less than the Period number they really belong tosometimes d electron in configurationEuZ = 63, Period 6[Xe]6s24f 7 6s 4f


Electron Configuration fromthe Periodic Table

As = [Ar]4s23d104p3

As has 5 valence electrons

As

1234567

1A

2A 3A 4A 5A 6A 7A

8A

4s2

Ar3d10

4p3


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)


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 5p4d

5s 4d


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


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


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 TableGroup 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 atoms gain enough electrons to have 8 valence electronsfilling the s and p sublevels of the valence shell

• the sulfur atom has 6 valence electronsS atom = 1s22s22p63s23p4

• in order to have 8 valence electrons, it must gain 2 moreS2- anion = 1s22s22p63s23p6


Electron Configuration of Cations in their Ground State

• cations are formed when an atom loses all its valence electronsresulting in a new lower energy level valence shellhowever the process is always endothermic

• the magnesium atom has 2 valence electronsMg atom = 1s22s22p63s2

• when it forms a cation, it loses its valence electronsMg2+ cation = 1s22s22p6


Electron Configuration of Atoms in their Ground State

• Electrons in some elements are “promoted” to form more stable configurations, e.g.

• Cu = 29 electronsExpected: 1s22s22p63s23p64s23d9

Actual: 1s22s22p63s23p64s13d10

• Cr = 24 electronsExpected: 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.


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


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 nucleusscreening effect

• effective nuclear charge is net positive charge that is attracting a particular electron

• Z is nuclear charge, S is electrons in lower energy levelselectrons in same energy level contribute to screening, but

very littleeffective nuclear charge on sublevels trend, s > p > d > f

Zeffective = Z - S

52

Screening & Effective Nuclear Charge


Trends in Atomic RadiusTransition Metals

• increase in size down the Group

• atomic radii of transition metals roughly the same size across the d blockmust less difference than across main group

elementsvalence shell ns2, not the d electronseffective nuclear charge on the ns2 electrons

approximately the same


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 largerGe is further down so largerAl is further down and left so largerOpposing trends


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


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


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 = Ar18 electrons

4s 3d

• Fe atom = [Ar]4s23d6

• unpaired electrons • paramagnetic

4s 3d

• Fe3+ ion = [Ar]4s03d5

• unpaired electrons • paramagnetic


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 anionsexcept 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

Periodic Pattern - Ionic Radius (Å)

Li

+1

Na

+1

K

+1

Rb

+1

Cs

+1

Be

+2

Mg

+2

Ca

+2

B

Tl

+3+1

In

+3+1

Ga

+3+1

Al

+3

Sn

+4+2

Ge

-4

Si-4

N

-3

Bi

Sb

As

-3

Te

-2

O

-2

P-3

S

-2

Se

-2

Br

-1

-1

F

I

-1

Cl

-1

H

+1-1

2A 3A 4A 5A 6A 7A

Sr

+2

Ba

+2

C-4

Pb

+4+2

0.68

0.97

1.33

1.47

0.31

0.66

0.99

1.13

1.40

1.84

1.98

2.21

1.33

1.81

1.96

2.20

Al

+3

+3

0.23

0.51

0.62

0.81

1A

1.69 1.35

1.71

2.12

2.22

0.71

0.840.95


Ionization Energy• minimum energy needed to remove an electron

from an atom gas stateendothermic processvalence electron easiest to removeM(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.

http://wps.prenhall.com/wps/media/objects/166/170446/IonizationEnergy.html


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 groupvalence electron farther from nucleus

• 1st IE generally increases across the periodeffective nuclear charge increases


Example 8.8 – Choose the Atom in Each Pair with the Higher First Ionization Energy

1) Al or S, Al is further left1) Al or S

2) As or Sb, Sb is further down

1) Al or S

2) As or Sb

3) N or Si, Si is further down & left

1) Al or S

2) As or Sb

3) N or Si

4) O or Cl? opposing trends


Irregularities in the Trend• Ionization Energy generally increases from left

to right across a Period

• except from 2A to 3A, 5A to 6A

Be 1s 2s 2p

B 1s 2s 2p

N 1s 2s 2p

O 1s 2s 2p

Which is easier to remove an electron from B or Be? Why?Which is easier to remove an electron from N or O? Why?


Irregularities in the First Ionization Energy Trends

Be 1s 2s 2p

B 1s 2s 2p

Be+ 1s 2s 2p

To ionize Be you must break up a full sublevel, cost extra energy

B+ 1s 2s 2p

When you ionize B you get a full sublevel, costs less energy


Irregularities in the First Ionization Energy Trends

To ionize N you must break up a half-full sublevel, cost extra energy

N+ 1s 2s 2p

O 1s 2s 2p

N 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 energyshrinkage in size due to having more

protons than electronsouter 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


Trends in Electron Affinity• energy released when an neutral atom gains an electron

gas stateM(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 rightnot absolute lowest EA in period = alkali earth metal or noble gashighest EA in period = halogen


Metallic Character• Metals

malleable & ductileshiny, lusterous, reflect lightconduct heat and electricitymost oxides basic and ionic form cations in solution lose electrons in reactions - oxidized

• Nonmetalsbrittle in solid statedullelectrical and thermal insulatorsmost oxides are acidic and molecular form anions and polyatomic anionsgain electrons in reactions - reduced

• metallic character increases left• metallic character increase down


Example 8.9 – Choose the More Metallic Element in Each Pair

1) Sn or Te

2) P or Sb

3) Ge or In, In is further down & left

1) Sn or Te

2) P or Sb

3) Ge or In

4) S or Br? opposing trends

1) Sn or Te, Sn is further left1) Sn or Te

2) P or Sb, Sb is further down


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 oxidizevery reactive, not found uncombined in nature react with nonmetals to form saltscompounds 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


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 reducevery reactive, not found uncombined in nature react with metals to form saltscompounds 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


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)



• 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)



• 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)


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


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

Chapter 8 Periodic Properties of the Elements

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Transcript of Chapter 8 Periodic Properties of the Elements