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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
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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