chapter 8 periodic properties of the elements

Copyright 2011 Pearson Education, Inc.

Chapter 8Periodic

Properties of the Elements

Roy KennedyMassachusetts Bay Community College

Wellesley Hills, MA

Chemistry: A Molecular Approach, 2nd Ed.Nivaldo Tro


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 channelsNa+ 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

2Tro: Chemistry: A Molecular Approach, 2/e


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

• 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



Periodic Pattern

a = acidic oxide, b = basic oxide, a/b = amphoteric oxide

M = metal, NM = nonmetal, M/NM = metalloid

Li Be B C N O F

Na Mg Al Si P S Cl

K Ca

HNM H2O

a/b

H21.0

M Li2Ob

LiH6.9

M Na2Ob

NaH23.0

M K2Ob

KH39.1

M BeOa/b

BeH29.0

M MgOb

MgH224.3

M CaOb

CaH240.1

NM B2O3

a

BH310.8

M Al2O3

a/b

AlH327.0

NM CO2

a

CH412.0

M/NM SiO2

a

SiH428.1

NM N2O5

a

NH314.0

NM P4O10

a

PH331.0

NM

H2S32.1

SO3

a

NM

H2O16.0

O2

NM Cl2O7

a

HCl35.5

NM

HF19.0



Mendeleev's Predictions



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• Quantum Mechanics is a theory that explains

why the periodic trends in the properties existand 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

7

1s1principal energy level of orbital

occupied by the electron

sublevel of orbital occupied by the

electron

number of electrons in the orbital

Tro: Chemistry: A Molecular Approach, 2/e


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 multielectron atoms cannot be exactly solveddue 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



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 fieldspinning 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



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 oppositespin up or spin down

• The electron’s spin adds a fourth quantum number to the description of electrons in an atom, called the Spin Quantum Number, msnot in the Schrödinger equation



Spin Quantum Number, ms, and Orbital Diagrams

• ms can have values of +½ or −½• Orbital Diagrams use a square to

represent each orbital and a half-arrow to represent each electron in the orbital

• By convention, a half-arrow pointing up is used to represent an electron in an orbital with spin up

• Spins must cancel in an orbitalpaired



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 withone electron

unoccupiedorbital

orbital withtwo electrons



Pauli Exclusion Principle


• No two electrons in an atom may have the same set of four quantum numbers

• Therefore no orbital may have more than two 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, ... - size and energy of the

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

1 n shape of

orbital Magnetic,

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

orbital Spin, m s -_ , +_ 2 direction of

electron sp in



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 multielectron atoms, the energies of the

sublevels are splitcaused by charge interaction, shielding and penetration

• 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)



Coulomb’s Law• Coulomb’s Law describes the attractions and

repulsions between charged particles• For like charges, the potential energy (E) is positive

and decreases as the particles get farther apart as r increases

• For opposite charges, the potential energy is negative and becomes more negative as the particles get closer together

• 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 multielectron atom

experiences both the attraction to the nucleus and repulsion by other electrons in the atom

• These repulsions cause the electron 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

• The closer an electron is to the nucleus, the more attraction it experiences

• 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 functionin particular, the distance the maxima of the

function are from the nucleus



Shielding & Penetration



Penetration 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



Effect of Penetration and Shielding• Penetration causes the energies of sublevels in the

same principal level to not be degenerate• In the fourth and fifth principal levels, the effects of

penetration become so important that the s orbital lies 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 therefore vary among elementscausing variations in the electron configurations of

the transition metals and their ions23Tro: Chemistry: A Molecular Approach, 2/e


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 (and can cause anomalous electron configurations for certain elements)



Filling the Orbitals with Electrons


• Energy levels and sublevels fill from lowest energy to highs → p → d → fAufbau Principle

• Orbitals that are in the same sublevel have the same energy

• No more than two electrons per orbitalPauli Exclusion Principle

• When filling orbitals that have the same energy, place one electron in each before completing pairsHund’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 [] to represent all the inner electrons, then just write the last set.

Rb = 37 electrons = 1s22s22p63s23p64s23d104p65s1

= [Kr]5s1



Order of Sublevel 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 sublevels, (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



Electron Configurations



Example: Write the full ground state orbital diagram and electron configuration of manganese

Mn Z = 25, therefore 25 e−

Based on the order of sublevel filling, we will need the first seven sublevels

1s

2s 2p

3s 3p 3d

4s 4p 4d 4f

s sublevel holds 2 e−

p sublevel holds 6 e−

d sublevel holds 10 e−

f sublevel holds 14 e−

2 e−

+2 = 4e−

+6 +2 = 12e−

+6 +2 = 20e−

1s 2s 2p 3s 3p 4s

Therefore the electron configuration is 1s22s22p63s23p64s23d5

+10 = 30e−3 d



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−

Based on the order of sublevel filling, we will need the first six sublevels

1s

2s 2p

3s 3p 3d

4s 4p 4d 4f

s sublevel holds 2 e−

p sublevel holds 6 e−

d sublevel holds 10 e−

f sublevel holds 14 e−

2 e−

+2 = 4e−

+6 +2 = 12e−

+6 +2 = 20e−

1s 2s 2p 3s 3p 4s

Therefore the electron configuration is 1s22s22p63s23p64s1



Valence Electrons


• The electrons in all the sublevels 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 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

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 five valence electrons

3p3

PNe

1234567

1A2A 3A 4A 5A 6A 7A

8A

3s2



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

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 to sometimes d electron in configuration

EuZ = 63, Period 6[Xe]6s24f 7 6s 4f



Electron Configuration fromthe Periodic Table

As = [Ar]4s23d104p3

As has five valence electrons

4s2

Ar3d10

4p3

As

1234567

1A2A 3A 4A 5A 6A 7A

8A



Practice – Use the Periodic Table to write the short electron configuration and short orbital diagram for

each of the following

• Na (at. no. 11)

• Te (at. no. 52)

• Tc (at. no. 43)

3s[Ne]3s1

5s 5p4d

[Kr]5s24d105p4

5s 4d

[Kr]5s24d5



Irregular Electron Configurations • We know that because of sublevel splitting, the

4s sublevel is lower in energy than the 3d; and therefore the 4s fills before the 3d

• But the difference in energy is not large• Some of the transition metals have irregular

electron configurations in which 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



• The properties of the elements follow a periodic patternelements in the same column

have similar propertiesthe elements in a period show a

pattern that repeats• The quantum-mechanical

model explains this because the number of valence electrons and the types of orbitals they occupy are also periodic

43

Properties & Electron Configuration



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-reactiveHe 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 gasforming 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 gasForming 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 unreactiveHe 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 metalsas 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 nonmetal atoms gain enough electrons to have eight valence electronsfilling the s and p sublevels of the valence shell

• The sulfur atom has six valence electronsS atom = 1s22s22p63s23p4

• To have eight valence electrons, sulfur must gain two moreS2− anion = 1s22s22p63s23p6



Electron Configuration of Cations in Their Ground State

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

• The magnesium atom has two valence electronsMg atom = 1s22s22p63s2

• When magnesium forms a cation, it loses its valence electronsMg2+ 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


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 electronsscreening or shielding effectouter 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 levelselectrons in same energy level contribute to

screening, but very little so are not part of the calculation

trend is s > p > d > fZeffective = Z − S



Screening & Effective Nuclear Charge



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



1. N or F2. C or Ge3. N or Al4. Al or Ge? opposing trends

1. N or F2. C or Ge3. N or Al, Al is farther down & left

1. N or F2. C or Ge, Ge is farther down

Example 8.5: Choose the Larger Atom in Each Pair

1. N or F, N is farther left

Tro: Chemistry: A Molecular Approach, 2/e 60


Practice – Choose the Larger Atom in Each Pair

• 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

• C or O• Li or K• C or Al• Se or I



Trends in Atomic RadiusTransition Metals

• Atoms in the same group increase in size down the column

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

elementsvalence shell ns2, not the (n−1)d electronseffective 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 electronsFe atom = 1s22s22p63s23p64s23d6

• When iron forms a cation, it first loses its valence electronsFe2+ cation = 1s22s22p63s23p63d6

• It can then lose 3d electronsFe3+ cation = 1s22s22p63s23p63d5



Magnetic Properties of Transition Metal Atoms & Ions

• Electron configurations that result in unpaired electrons mean that the atom or ion will have a net magnetic field – this is called paramagnetismwill 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 diamagnetismslightly 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 paramagneticAg+ ions [Kr]4d10 are diamagneticAg2+ ions [Kr]4d9 are paramagnetic



4s 3d

• Fe atom = [Ar]4s23d6

• Unpaired electrons • Paramagnetic

4s 3d

• 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



Practice – Determine whether the following are paramagnetic or diamagnetic

4s 3d

• Mn

• Sc3+

Mn = [Ar]4s23d5 paramagnetic

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 for isoelectronic species


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

• These “new valence electrons” experience a smaller effective nuclear charge than the “old valence electrons,” increasing the size

• The result is that the anion is 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 iongas stateendothermic processvalence electron easiest to remove, lowest IEM(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 groupvalence electron farther from nucleus

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



80



Quantum-Mechanical Explanation for the Trends in First Ionization Energy

• The strength of attraction is related to the most probable distance the valence electrons are from the nucleus and the effective nuclear charge the valence electrons experience

• 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

• 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



1. Al or S2. As or Sb, As is farther up1. Al or S2. As or Sb3. N or Si, N is farther up & right

1. Al or S2. As or Sb3. N or Si4. O or Cl? opposing trends

Example 8.8: Choose the atom in each pair with the larger first ionization energy

1. Al or S, S is farther right



• Mg or P

• Ag or Cu

• Ca or Rb

• P or Se

Practice – Choose the atom with the largest first ionization energy in each pair

?



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

Exceptions in the 1st IE Trends

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?


• First Ionization Energy generally increases from left to right across a Period

• Except from 2A to 3A, 5A to 6A


Exceptions in the First Ionization Energy Trends,

Be and B

Be 1s 2s 2p

B 1s 2s 2p

Be+ 1s 2s 2p

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

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

To ionize N you must break up a half-full sublevel, costs 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



Electron Affinity


• Energy released when an neutral atom gains an electrongas stateM(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 affinitythe more negative the number, the larger the EA


Trends in Electron Affinity• Alkali metals decrease electron affinity down the

columnbut not all groups dogenerally irregular increase in EA from 2nd period to 3rd

period• “Generally” increases across period

becomes more negative from left to rightnot absoluteGroup 5A generally lower EA than expected because

extra electron must pairGroup 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


• Metalsmalleable & ductileshiny, lusterous, reflect lightconduct heat and electricitymost oxides basic and ionicform cations in solutionlose electrons in reactions – oxidized

• Nonmetalsbrittle in solid statedull, non-reflective solid surfaceelectrical and thermal insulatorsmost oxides are acidic and molecularform anions and polyatomic anionsgain electrons in reactions – reduced


Metallic Character• Metallic character is how closely an element’s

properties match the ideal properties of a metalmore malleable and ductile, better conductors, and

easier to ionize• Metallic character decreases left-to-right

across a periodmetals 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 affinitiesexcept 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



1. Sn or Te2. P or Sb, Sb is farther down1. Sn or Te2. P or Sb3. Ge or In, In is farther down & left

Example 8.9: Choose the more metallic element in each pair

1. Sn or Te2. P or Sb3. Ge or In4. S or Br? opposing trends

1. Sn or Te, Sn is farther left



Practice – Choose the more metallic element in each pair

• Mg or Al

• Si or Sn

• Br or Te

• Se or I?


Copyright 2011 Pearson Education, Inc.99

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 volumeTro: Chemistry: A Molecular Approach, 2/e


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 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 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)




• 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)




• Reaction between chlorine gas and solid iodineCl2(g) + I2(s)

Cl2(g) + I2(s) IClwrite the halogen lower in the column firstassume 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 natureused as “inert” atmosphere when reactions with other

gases would be undersirable• Melting point and boiling point increase down the

columnall gases at room temperaturevery low boiling points

• Density increases down the column in general, the increase in mass is greater than the

increase in volumeTro: Chemistry: A Molecular Approach, 2/e


Noble Gases


chapter 8 periodic properties of the elements

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