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Chapter 8

Atomic Electron Configurations and

Chapter 8

Chemical Periodicity

Chapter 8Evolving model of the atom

1803 (Dalton): All matter is composed of tiny, indivisible, indestructible particles called atom.

1903 (Thompson): Subatomic particles: electrons and positive charges. Plum-pudding model.

1911(Rutherford): Protons (positively charge) and neutrons (neutral) are located in the centre of the atomneutrons (neutral) are located in the centre of the atom. Electrons are somewhere outside the nucleus.

1913 (Bohr): Electrons are moving in a circular orbitaround the nucleus. Only certain orbits with fixed energy are permissible.

1932 (Schrodinger): The region of space (ORBITAL)outside the nucleus where the probability (likelihood) of finding an electron with a given energy is maximum.

Chapter 8

ORBIT: The circular path in which electrons move around the nucleus

ORBITAL: The region in space where an electron is most likely to be found

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Chapter 8

Orbitals- Home of ElectronsFirst three quantum numbers (n, l, and ml) describe orbitals

shell 1 1sshell 2 2s 2pshell 3 3s 3p 3dshell 4 4s 4p 4d 4fp

shell: Each shell with a designated n has many subshellssubshell: Each subshell with a designated l

has many orbitalsorbital: Each orbital with a designated by ml has a specific orientation and has room for TWO electrons

What determines the relative energies of these orbitals? Which are lower in energy, which are higher in energy?

Chapter 8

Orbital Energies

What general principle explains orbital energies?

Which orbital has higher energy, 1s, 2s or 3s? Why?

Which orbital has higher energy, 2s or 2p?Why?

Which orbital has higher energy, 2px, 2py or 2pz?Why?

Chapter 8

Rad

ial p

roba

bilit

y

Orbital Energies

E1s< E2s < E3s

Distance from nucleus

Rad

ial p

roba

bilit

y

Distance from nucleus

E2s< E2p

3

Chapter 8

Orbital EnergiesChapter 8

Effective Nuclear ChargeZeff: the positive charge actually felt by a valence electron

Zeff = Z – s Z = atomic numbers = shielding parameterZeff increases across the period of periodic tableeff p p

LithiumZeff = 3 – 1.72 = 1.28

NitrogenZeff = 7 – 3.15 = 3.85

Which electron will be easy to remove, the one from Lithium or Nitrogen?

Chapter 8 Effective Nuclear Charge

Orbital stability

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Chapter 8

Effective Nuclear ChargeOrbital stability

Chapter 8 Effective Nuclear Charge

Zeff: the positive charge actually felt by a valence electron

Zeff = Z – s

Orbital stability

A quantity that comes due to electron-electron repulsion

Chapter 8

Magnetic Properties: Electron

A physical phenomenon: spinning, charged particles produce magnetic fields

Spinning electronsSpinning electrons produce tiny magnetic fields

Electrons can spin in one of two directions

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Chapter 8

Diamagnetic: substances repelled by a strong magnetic fieldPaired electrons

Magnetic Properties of ElectronPaired electrons are more stable

Paramagnetic: substances attracted to a strong magnetic field

Unpaired electrons

Chapter 8

The 4th Quantum Number

Electron spin, ms: ms = ½ or -½

Pauli exclusion principle:

Aligned or opposed to the magnetic field

No two electrons in an atom can have the same set of four quantum numbers n, l, ml, and ms.

In order to put more than one electron in an orbital, electrons must have different values of ms. i.e. they must have different spins.

Maximum of 2 electrons per orbital

Chapter 8

Quantum Mechanical Model andPeriodic Table

Li ground state

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Chapter 8

Energy of OrbitalsFor the same type of orbital (same ______), energy increases as n increases

(1s < 2s < 3s < 4s…)

For the same n, energy increases s < p < d < f (3s < 3p < 3d)(3s 3p 3d)

All orbitals of the same subshell have the same energy (degenerate)

(3px = 3py = 3pz)

1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p

Chapter 8

Energy of Orbitals: n+l rule

Draw this diagram and by hand and start filling out electrons.

This diagram will beThis diagram will be counted as oneproblem i.e. 1/4th extra credit

Chapter 8

Orbital Diagrams

n=33s 3p 3d

3s 3p 3d

orbital

subshell

shell

n=3p

n=33s 3p 3d

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Chapter 8

Electron Configuration RulesElectrons fill the lowest energy orbital first (Aufbau principle)

1s

2s 2pDiagonal Diagram:

This diagram and any 10 elements’ electron-filled orbital diagram will be counted as oneproblem i.e. 1/4th extra credit

3s 3p 3d

4s 4p 4d 4f

5s 5p 5d 5f

6s 6p 6d 6f

7s 7p 7d 7f

Diagonal Diagram:

a guide used to determine the relative energies of subshells in multi-electron atoms

Chapter 8

Electron Configuration Rules

Two electrons (max) per orbital

Pauli exclusion principleNo two electrons in an atom can have the same set

of four quantum numbers n, l, ml, and ms.

Maximize parallel spins when filling a subshellIf more than one orbital in a subshell is available, electrons will fill empty orbitals in the subshell first.(Hund’s Rule)

Alternately….Electrons preferred to be unpaired as long as an empty orbital with the same energy is available

Chapter 8

Energy of Orbitals: Summary

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Chapter 8

Electron Configurations

Three notations for the arrangement of electrons in atoms

Orbital box diagrams

spdf notation

noble gas notation

Chapter 8

Electron Configurations

Hydrogen

Orbital Box Notation

1s1

number of electrons

orbital type (l)

electron shell (n)Lithium # of es =3

spdf NotationLithium # of es =3Α. 1s22s1

B. 1s12s12p1

C. 2p3

D. 1s3

Chapter 8

Electron Configurations

Hydrogen

Orbital Box Notation

1s1

number of electrons

orbital type (l)

electron shell (n)Oxygen: # of es =8

spdf NotationOxygen: # of es =8Α. 2s22p6

B. 1s12s12p6

C. 1s22s22p4

D. 1s22s32p3

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Chapter 8

Electron Configurations

Hydrogen

Orbital Box Notation

1s1

number of electrons

orbital type (l)

electron shell (n)

spdf Notation

Chlorine: # of es =17Α. 1s22s22p63s23p33d3

B. 1s22s22p63s23p5

C. 1s22s22p53s23p6

D. 1s22s32p63s13p6

Chapter 8

More Examples

Provide the electron configurations (in orbital box, spdf and noble gas notation)(a) P

(b) V

(c) I

Chapter 8

H

Li Be

Na Mg

CaK Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr

Ar

Ne

He

F

ClSPSiAl

B C N O

1A

2A

3B 4B 5B 6B 7B 8B 8B 8B 1B 2B

3A 4A 5A 6A 7A

8ATransition Metals

Rb

Cs

Fr Ra

Ba

Sr XeITeSbSnInCdAgPdRhRuTcMoNbZrY

La

Ac Rf

Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn

Db Sg Bh Hs Mt

LuYbTmErHoDyTbGdEuSmPmNdPrCe

Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr

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Chapter 8

Some Anomalies?

Chromium and copper

Transition metal ions

Half-filled and fully filled d-subshells have extra stability (lower energy).

Chapter 8

More Examples: Ions

(a) S2–

So does S2– = Ar?

(b) Br –

Isoelectronic species

( )

(c) Al3+

Chapter 8

Periodic Table Organization

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Chapter 8

Periodic Table Organizations-block atoms where an s subshell is being filledp-block atoms where a p subshell is being filledd-block atoms where a d subshell is being filled

Valence electrons

Core electrons: electrons included in the noble gas notation

Li (3): 1s2 2s1 Na(11): 1s22s22p63s1

[He] 2s1 [Ne] 3s1

Same group = same number and type of valence electrons

Chapter 8

Effective Nuclear Charge

Take the case of Li1s22s1

l pro

babi

lity

Distance from nucleus

Rad

ia

Chapter 8

Electron ConfigurationsValence electrons: electrons in the outermost shells responsible for all macroscopic properties

Core electrons: electrons included in the noble t tigas notation

Li (3): 1s2 2s1 Na(11): 1s22s22p63s1

[He] 2s1 [Ne] 3s1

Same group = same number and type of valence electrons Similarity of properties

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Chapter 8

Electron Configurations: Atoms and Ions

Noble gas elementsHe (2) : 1s2

Ne (10) : [He] 2s2 2p6

Ar (18) : [Ne] 3s2 3p6

Kr (36): [Ar] 4s2 4p6

K+ (19-1= 18) ≡ [Ar] or [Ne] 3s2 3p6

Br- (35 +1= 36) ≡ [Kr] or [Ar] 4s2 4p6

Chapter 8

Periodic PropertiesYou will need to know the following:

1. Definitions and chemical equations where appropriate

2. Periodic trends moving up and down and left to right across the periodic table

3. Explanations of the trends

4. How the atomic properties affect chemical properties

Chapter 8

Effective Nuclear ChargeValence electrons don’t “feel” the full charge of the nucleus

Valence electrons are shielded

But … valence electrons “feel” a charge that is greater than Z – core electrons

Valence electrons are not completely shielded

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Chapter 8

The distance from the nucleus to the edge of the outermost electron

Periodic trend:

Atomic Size

Explanation:

Chapter 8

Atomic Size

Decrease

Effective nuclear charge increases across the group

Chapter 8 Atomic Size

Decrease across a period

Decrease across a period

Decrease across a period

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Chapter 8 Atomic Size

The best way to explain the increase of atomic size as one goes downward through groups

Α. The electrons in a shell repel more, therefore the atom expands

B. The nucleus becomes bigger in size as it has more protons and neutrons

C Down the group new shells (i e n is increased by 1)C. Down the group, new shells (i.e. n is increased by 1) are added; each new shell is further and further away from the nucleus

D. The nucleus expands and the shells (filled with electrons) expands

Chapter 8 Atomic SizeThe best way to explain the decrease of atomic size as

one goes across periods Α. The electrons repel less, therefore the atom shrinksB. The electrons are put on a same shell . The nuclear

effective charge increases and the effective pull of the nucleus on its outermost shell electrons increases many fold y

C. Across a period, the total positive charge at the nucleus remains constant

D. The nucleus shrinks as it accommodates more neutrons

Chapter 8

#1:Identify the one which is correctly arranged in order of increasing (smallest to largest) atomic size:

a. Be, C, Ob. Be, O, Cc. O, C, Bed. C,O, Be

#2:Identify the one which is correctly arranged in order of increasing (smallest to largest) atomic size:

a. Cl, K, Sb. Cl, S, Kc. K, S, Cld. K, Cl, S

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Chapter 8

Ionization Energy (IE)

The energy required to remove an electron from a gaseous atom

A(g) + energy A+(g) + e-

Energy inputinput required

Chapter 8

Sign Conventions

Energy absorbed (in) = a positive value + 165 kJEnergy required (input, raw material)

Energy released (out) = a negative value - 165 kJEnergy produced (output, product)

The sign tells us which way energy is goingThe magnitude tells us how much energy is required

Chapter 8

Ionization Energies

Dec

reas

e

Effective nuclear charge increasesacross the group

IE (Be) > IE (B)Be(4): 1s2 2s2

B(5) : 1s2 2s2 2p1

IE (N) > IE (O)

N (7): 1s2 2s2 2p3

O (8) : 1s2 2s2 2p4

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Chapter 8 First Ionization Energy Chapter 8

Successive Ionizations

IE1 IE2 IE3 IE4 IE5 IE6 IE7

Na 495 4560

Mg 735 1445 7730

Al 580 1815 2740 11600

Si 780 1575 3220 4350 16100Si 780 1575 3220 4350 16100

P 1060 1890 2905 4950 6270 21200

S 1005 2260 3375 4565 6950 8490 27000

Example: Na(g) + IE1 Na+(g) + e-

Na+(g) + IE2 Na2+(g) + e-

Chapter 8

Successive Ionizations

For Mg, 2nd IE > 1st IEFor Al, 3rd IE > 2nd IE > 1st IEWhy?

For Mg, 3rd IE >>> 2nd IEFor Al, 4th IE >>> 3rd IEWhy?

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Chapter 8

Ionization Energies: Summary

First ionization energies generally increase across a period and decrease down a group

Effective nuclear charge increases thacross the group

Chapter 8

#3:Arrange each set of atoms in increasing IE1:

a. Sr, Ca, Bab. Ba, Sr, Cac. Ca, Sr, Bad. Ba, Ca, Sr

#4:#4:Arrange each set of atoms in increasing IE1:

a. Br, Rb, Seb. Br, Se, Rbc. Rb, Br, Sed. Rb, Se, Br

Chapter 8

Electron Affinity

The energy released when an electron is added to a gaseous atom

A(g) + e- A-(g) + energy

A free electron is not a stable. It would always be associated with an atom.

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Chapter 8

A(g) + e- A-(g) + energy

Across a period: Should it get easier or harder to add an electron?

Electron Affinity Predictions

Down a group: Should it get easier or harder to add an electron?

If it’s easy to add an electron, is the EA a large negative number or a small negative number?

Deviations from the general trends

Chapter 8

Electron Affinity Trends

Exception

Chapter 8

Electron Affinity Summary

An element with a high ionization energies generally has a high affinity for an electron.

Effective nuclear charge increases across the group and decreases down a group

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Chapter 8

Trends in Metallic Behavior

Relative tendencies to lose and gain electrons

Elements at the left form cations easilyElements at the right form anions easily

Chapter 8 Acid-base Behaviors of Elemental OXides

Metals donate electrons to oxygen

Nonmetals share electrons to oxygen

Covalent

Metal oxides react with water to produce hydroxides (OH-) that are basic

Ionic

Nonmetal oxides react with water to produce acids that releases proton in solution H+

Covalent

Chapter 8

Ionization: Change in Size

Why does the size decrease?

3 p+ and 3 e-

3 p+ and 2 e-

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Chapter 8

Ionization: Change in Size

Why does the size increase?

9 p+ and 9 e-9 p+ and 10 e-

Chapter 8Review

Zeff: the positive charge actually felt by a valence electron Atomic size: The distance from the nucleus to the edge of the outermost electronIE: The energy required to remove an electron from a gaseous atom.

Successive ionization

EA: The energy released when an electron is added to a gaseous atomIon sizes

Chapter 8

The Reaction of Na and ClIE EA

Na 495 EA > 04560

Cl 1251 -348

How can we use these numbers to explain the product of the reaction?

Is NaCl2 a reasonable product?

Is Na2Cl a reasonable product?

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Chapter 8

Periodic trends and Chemical Properties

• Reactivity of metals

• Reactivity of nonmetals

Chapter 8

Chemical Reactivity Summary

Noble gases high IE, low EA do not react

Metals low IE, low EA lose electrons

Non-metals high IE, high EA add electrons

Metal + non-metal metal loses e-’s and non-metal gains e-’s

non-metal + non-metal shared e-’s