11/13/2008

1

Anomalous Electron Configurations

• A few exceptions to the Aufbauprinciples exist. Stable configuration:–half-filled d shell:

• Cr has [Ar]4s13d5;

• Mo has [Kr] 5s14d5

– filled d subshell:

• Cu has [Ar]4s13d10

• Ag has [Kr]5s14d10.

• Au has [Xe]6s14f145d10

• Exceptions occur with larger elements where orbital energies are similar.

Anomalous electron configuration of some elements!!!

Element Atomic number

Expected configuration

Actual Configuration

Cr 24 3d4 4s2 3d5 4s1

Cu 29 3d9 4s2 3d10 4s1

Mo 42 4d4 5s2 4d5 5s1

*Pd 46 4d8 5s2 4d10 5s0

Ag 47 4d9 4s2 4d10 5s1

*Pt 78 5d8 6s2 5d9 6s1

Au 79 5d9 6s2 5d10 6s1

The explanation for this deviation lies in the superior stability The explanation for this deviation lies in the superior stability of completely filled or all halfof completely filled or all half--filled filled orbitalsorbitals than nearly filled than nearly filled of nearly halfof nearly half--filled filled orbitalsorbitals..****They are exception to this rule also. So, remember. ****They are exception to this rule also. So, remember. Some other of this kinds are Some other of this kinds are NbNb(41), (41), RuRu(44), W(74), (44), W(74), SgSg(106)**** (106)****

An application of Electronic An application of Electronic Configuration of AtomConfiguration of Atom

Locating elements in The Periodic TableLocating elements in The Periodic Table

• In The Periodic Table elements are organized on the basis of their electronic configuration

• The highest ‘n’ value in electronic configuration determines the period of that element.

• No of outer shell e determines the Group

§ If at s & p orbital of the outer shell have 1 to 7 e (s1 to s2

p5) then they are elements of Group I to VII

Ex: Na (11) – 1s22s22p63s1 : Group I

Cl (17) – 1s22s22p63s23p5 : Group VII

§ If at s & p orbital of outer shell have s2 p6 then the element is in group O. Ex: Ne (10)

§ Again at d & s orbital of outer shell determines the Group

Ex: Sc (21) - 1s22s22p63s23p63d14s2 : Group III

Cont’dCont’d

§But at d & s orbital of outer shell, if there are 8/9/10 e then it is in the Group VIII

Ex: Fe (26) - 1s22s22p63s23p63d64s2 : Group VIII

§If you have more then 10 e at d & s of outer shell then only ‘s’ orbital’s e will give the Group

Ex: Cu (29) - 1s22s22p63s23p63d104s1 : Group I

• ‘A’ Sub Group: if at the outer shell you do not have ‘d’ orbital or ‘d’ is filled only then it in ‘A’ Sub

Ex: Cl (17) – 1s22s22p63s23p5 : Group VIIA

Ga (31) - 1s22s22p63s23p63d104s24p1 : Group IIIA

• ‘B’ Sub Group: if you have e in ‘d’ orbital at outer shell (d1–d10) then it is in ‘B’ Sub Group

Ex: Sc (21) - 1s22s22p63s23p63d14s2 : Group IIIB

Zn (30) - 1s22s22p63s23p63d104s2 : Group IIB

11/13/2008

2

Valence Electrons

• The electrons in all the sub-shells with the highest principal energy shell are called the valence electrons

• e in lower energy shells are called core electrons

• Starting with one valence electron for the first element in a period, the number of electrons increases as you move from left to right across a period.

• 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 Anions Electron Configuration of Anions

• 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 electronsS atom = 1s22s22p63s23p4

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

When an atom or molecule gain or loses an eWhen an atom or molecule gain or loses an e-- it becomes an it becomes an ionion

•• A A cationcation has lost ehas lost e-- and therefore has positive chargeand therefore has positive charge

•• An An anionanion has gained an ehas gained an e-- and therefore has a negative and therefore has a negative charge.charge.

Electron Configuration of Cations• 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 electronsMg atom = 1s22s22p63s2

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

• 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

DevelopmentDevelopment ofof thethe PeriodicPeriodic TableTable• There were 114 elements known by 1999.• The majority of the elements were discovered

between 1735 and 1843.• How do we organize 114 different elements in a

meaningful way that will allow us to make predictions about undiscovered elements?

• Arrange elements to reflect the trends in chemical and physical properties.

• First attempt (Mendeleev and Meyer) arranged the elements in order of increasing atomic weight.

The Periodic Table of ElementsThe Periodic Table of Elements

Week -7

Lecture 13 &14

11/13/2008

3

The elements were first arranged in this way by Dmitri Mendeleev, a professor at St. Petersburg University, in 1869. His arrangement was based on atomic mass.

When Mendeleev was setting out the table, only 63 elements had been discovered. His big idea was to leave gaps for yet to be discovered elements. He was able to predict the properties of some of these elements, including silicon and boron. When his predictions were shown to be accurate his table became accepted, and it is the basis of the one we use today.

• Moseley later discovered that the periodic nature of the elements was associated with atomic number, not atomic mass

• The Periodic Law

– When elements are arranged in order of increasing atomic number, there is a periodic pattern in their physical and chemical properties.

The Father of the Periodic Table —Dimitri Mendeleev

–The periodic table is made up of rows of elements and columns.

–An element is identified by its chemical symbol.

–The number above the symbol is the atomic number

–The number below the symbol is the atomic weight of the element.

–A row is called a period

–A column is called a family or group

–Elements are arranged left to right and top to bottom in order of increasing atomic number

–This order usually coincides with increasing atomic mass

Introduction of The Periodic Table

Key Concepts:

Atoms & Elements

Elements

consist of

Atoms

that have

Subatomic particles

Mass Number

has a

Nucleus

ElectronsNeutronsProtons

determine

Atomic Number

Make up the

are

Metals Non metalsor

that have

Chemical symbols

arranged in the

Periodic Table

by

GroupsPeriods

Energy levels

Outer shell electrons

Periodic law Group number

are in

with

determine

• Periodic Patterns– The chemical behavior of elements is determined by

its electron configuration

– The first three periods contain just A families. Each period begins with a single electron in a new outer electron shell.

– Each period ends with a completely filled outer shell that has the maximum number of electrons for that shell.

– The outer shell electrons are responsible for chemical reactions. Elements in the same family have the same number of outer shell electrons; so they will have similar chemical properties.

– Group A elements are called representative elements

– Group B elements are called transition elements.

11/13/2008

4

• Chemical Families– IA are called alkali metals because the react with

water to from an alkaline solution, very soft metals. (except H2)

– Group IIA are called the alkali earth metals because mostly we found them in soils as salts/minerals & they are also reactive, but not as reactive as Group IA.

• They are also soft metals, though not as soft as alkali metals

– Group VIIA are the halogens

• These need only one electron to fill their outer shell

• They are very reactive.( disinfectants, bleach)

– Group VIIIA are the noble gases as they have completely filled outer shells

• They are almost non reactive.

Metals, Nonmetals, and MetalloidsMetals, Nonmetals, and Metalloids

MetalsMetals•Metallic character refers to the properties of metals

– Shiny or lustrous– Malleable (can be hammered into shape)– Ductile (can be drawn out into wires)– All except mercury are solids at room temperature– They are sonorous (make a ringing sound when hit)– In solution lose electrons in reactions - oxidized– Most oxides are basic and ionic

Ex: Metal oxide + water → metal hydroxideNa2O(s) + H2O(l) → 2NaOH(aq)

– Tends to form cation in aqueous solution• Metallic character increases down a group.• Metallic character decreases across a period.

Only a few metals are magnetic.Magnetism is not a property of most metals!

Metals, Nonmetals, and MetalloidsMetals, Nonmetals, and Metalloids

MetalsMetals• When metals are oxidized they tend to form

characteristics cations.• All group 1A metals form M+ ions.• All group 2A metals form M2+ ions.• Most transition metals have variable charges.

Properties of Non-Metals

§ They are poor conductors of electrical energy§ They are poor conductors of thermal

energy§Many of them are gases§ They are brittle if they are solid§ Form anions§Most oxides are acidicEx: nonmetal oxide + water → acid

P4O10(s) + H2O(l) → 4H3PO4(aq)§Gain electrons in reactions –

reduced§When nonmetals react with metals,

nonmetals tend to gain electrons:metal + nonmetal → salt

2Al(s) + 3Br2(l) → 2AlBr3(s)

Both

a d

iam

ond a

nd a

pen

cil

‘lea

d’ a

re

mad

e of

the

sam

e el

emen

t –

carb

on.

11/13/2008

5

Metals,Metals, Nonmetals,Nonmetals, andand MetalloidsMetalloids

MetalloidsMetalloidsMetalloids have properties that are intermediate between metals and nonmetals.Example: Si (shown here) has a metallic luster but it is brittle.

Metalloids have found fame in the semiconductor industry.

The Groups of the Periodic Table

• Group 1: The Alkali Metals

– Most reactive metals on the PT

– Rarely found free in nature

– Charge of +1 = 1 valence electron

• Group 2: The Alkaline Earth Metals– Still quite reactive

– Charge of +2 = 2 valence electrons

• Groups 3-12: Transition Metals– Found freely and in compounds in nature

– Charge is usually +2 but can vary = usually 2 valence electrons

• Group 13: Boron Family

– Charge is +3 = 3 valence electrons

The Groups of the Periodic Table

• Group 14: The Carbon Family

– Contains elements that can form unusual bonds (carbon and silicon)

– Charge is +4 or -4 = contains 4 valence electrons

• Group 15: The Nitrogen Family

– Charge is -3 = contains 5 valence electrons• Group 16: The Oxygen Family

– Also known as the chalcogens– Charge is -2 = 6 valence electrons

• Group 17: The Halogens– Most reactive nonmetals– charge is -1 = 7 valence electrons

• Group 18: The Noble Gases (The Inert Gases)– Nonreactive – Charge is 0 = 2 or 8 valence electrons

Periodic Properties

• Periodic law = elements arranged by atomic number gives physical and chemical properties varying periodically.

• Various Elemental Properties change fairly smoothly going across a period or down a group.

• We will study the following periodic trends:– Atomic radii

– Ionization energy

– Electron affinity

– Melting Points and Boiling Points

– Density

11/13/2008

6

ElectronElectron ShellsShells andand thethe SizesSizes ofofAtomsAtoms

Atomic SizesAtomic Sizes• As a consequence of the ordering in the periodic table, properties of elements vary periodically.

• Atomic size varies consistently through the periodic table.

• As we move down a group, the atoms become larger.

• As we move across a period, atoms become smaller.

There are two factors at work:•principal quantum number, n, and•the effective nuclear charge, Zeff.

ElectronElectron ShellsShells andand thethe SizesSizes ofof AtomsAtoms

• As the principle quantum number increases (i.e., we move down a group), the distance of the outermost electron from the nucleus becomes larger. Hence, the atomic radius increases.

• As we move across the periodic table, the number of core electrons remains constant. However, the nuclear charge increases. Therefore, there is an increased attraction between the nucleus and the outermost electrons. This attraction causes the atomic radius to decrease.

• increase in size down the Group

• atomic radii of transition metals roughly the same size across the d block– valence shell ns2, not the d electrons

– effective nuclear charge on the ns2 electrons approximately the same

Trends in Atomic Radius - Transition Metals

Atomic Radius

Fig

. 8.1

5 A

tom

ic R

adii

for

Mai

n G

rou

p E

lem

ents

See Figure 8.16

11/13/2008

7

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

36

37

11/13/2008

8

39

Ionization Energy

• minimum energy needed to remove an electron from an atom – gas state

– endothermic process

– valence electron easiest to remove

– 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.

• IE increases (irregularly) as you move from left to right across a period.

• IE decreases (irregularly) as you move down a group. 40

IonizationIonization EnergyEnergy

• The first ionization energy, I1, is the amount of energy required to remove an electron from a gaseous atom:

Na(g) → Na+(g) + e-.•The second ionization energy, I2, is the energy required to remove an electron from a gaseous ion:

Na+(g) → Na2+(g) + e-.

The larger ionization energy, the more difficult it is to remove the electron.There is a sharp increase in ionization energy when a core electron is removed.

41

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

42

11/13/2008

9

45

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

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 electronfrom 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

B, Al, Ga, etc.: their ionization energies are slightly less than the ionization energy of the element preceding them in their period.

• Before ionization ns2np1. • After ionization is ns2. Higher energy ⇒ smaller radius.

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 energyGroup 6A elements.

• Before ionization ns2np4. • After ionization ns2np3 where each p electron in different orbital (Hund’s

rule).

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

11/13/2008

10

HIGHER IONIZATION ENERGIES

51

ElectronElectron AffinitiesAffinities

• Electron affinity is the opposite of ionization energy.• Electron affinity is the energy change when a gaseous atom gains an electron to form a gaseous ion:

Cl(g) + e- → Cl-(g)• Electron affinity can either be exothermic (as the above example) or endothermic:• more energy released (more -); the larger the EAgenerally increases across period

becomes more negative from left to rightnot absolutelowest EA in period = alkali earth metal or noble gashighest EA in period = halogen

The added electron in Cl is placed in the 3p orbital to form the stable 3p6

electron configuration.54

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 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 a magnetic field

• 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

11/13/2008

11

55

Example 8.6 – 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

• Fe atom = [Ar]4s23d6

• unpaired electrons • paramagnetic

4s 3d

• Fe3+ ion = [Ar]4s03d5

• unpaired electrons • paramagnetic

Melting Points and Boiling Points•Trends in melting Points and boiling points can be used as a measure of the attractive forces between atoms or molecules.

•Within the halogens (group 17 or VIIA) melting points and boiling points increase so that at room temperature fluorine and chlorine are gases, bromine is a liquid, and iodine is a solid as you go down this periodic group.

•This indicates that the intermolecular forces become stronger going down a group.

•In the second period, melting points increase, going from left to right across the period for the first four elements.

• Melting points then decrease drastically for nitrogen, oxygen, and fluorine, which are all diatomic molecules.

• The lowest melting point is for neon, which is monatomic.

Melting Point Trends in Period

Melting Points of Elements Trends in Density

•Densities of elements increase in a group as atomic number increases.

•In periods, going from left to right, densities increase, then decrease.

•Elements with the greatest densities are at the center of period 6.

11/13/2008

12

Densities of Elements