Download Atomic Electron Configurations and Chapter 8 Chemical Periodicity

Chapter 8
Chapter 8
Evolving model of the atom
Chapter 8
¾ 1803 (Dalton): All matter is composed of tiny, indivisible,
indestructible particles called atom.
Chapter 8
Atomic Electron Configurations and
Chemical Periodicity
¾ 1903 (Thompson): Subatomic particles: electrons and
positive charges. Plum-pudding model.
¾ 1911(Rutherford): Protons (positively charge) and
neutrons (neutral) are located in the centre of the atom
atom.
Electrons are somewhere outside the nucleus.
¾ 1913 (Bohr): Electrons are moving in a circular orbit
around the nucleus. Only certain orbits with fixed energy
are permissible.
¾ORBIT: The circular path in which
electrons move around the nucleus
¾ORBITAL: The region in space where an
electron is most likely to be found
¾ 1932 (Schrodinger): The region of space (ORBITAL)
outside the nucleus where the probability (likelihood) of
finding an electron with a given energy is maximum.
1
Chapter 8
¾ First three quantum numbers (n, l, and ml) describe
orbitals
shell 1
shell 2
shell 3
shell 4
1s
2s
3s
4s
2p
3p
4p
p
3d
4d
4f
¾ shell: Each shell with a designated n has many
subshells
¾ subshell: Each subshell with a designated l
has many orbitals
¾ orbital: Each orbital with a designated by ml has a
specific orientation and has room for TWO electrons
Chapter 8
Orbital Energies
¾ What general principle explains orbital energies?
¾ Which orbital has higher energy, 1s, 2s or 3s?
Why?
Orbital Energies
Radial probability
Orbitals- Home of Electrons
E1s< E2s < E3s
Distance from nucleus
¾ Which orbital has higher energy, 2s or 2p?
Why?
¾ Which orbital has higher energy, 2px, 2py or 2pz?
Why?
Radial probability
Chapter 8
E2s< E2p
What determines the relative energies of these orbitals? Which are
lower in energy, which are higher in energy?
Distance from nucleus
2
Chapter 8
Chapter 8
Orbital Energies
Chapter 8
Effective Nuclear Charge
¾ Zeff: the positive charge actually felt by a valence
electron
Zeff = Z – s
¾ Z = atomic number
¾ s = shielding parameter
¾ Zeff increases across the p
period of p
periodic table
Effective Nuclear Charge
Orbital stability
¾ Lithium
¾ Zeff = 3 – 1.72 = 1.28
¾ Nitrogen
¾ Zeff = 7 – 3.15 = 3.85
Which electron will be easy to remove, the one from Lithium or Nitrogen?
3
Chapter 8
Chapter 8
Effective Nuclear Charge
Orbital stability
Effective Nuclear Charge
Chapter 8
Magnetic Properties: Electron
Orbital stability
¾Zeff: the positive charge actually felt by a
valence electron
Zeff = Z – s
A quantity that comes due
to electron-electron repulsion
¾ A physical phenomenon:
spinning, charged particles
produce magnetic fields
¾ Spinning electrons
produce tiny magnetic
fields
¾ Electrons can spin in one
of two directions
4
Chapter 8
Chapter 8
Magnetic Properties of Electron
Chapter 8
The 4th Quantum Number
Paired electrons are more stable
¾ Diamagnetic: substances repelled by a strong magnetic field
¾ Paired electrons
¾ Electron spin, ms: ms = ½ or -½
Quantum Mechanical Model and
Periodic Table
Li ground state
Aligned or opposed to the magnetic field
¾ Pauli exclusion principle:
¾ Paramagnetic: substances attracted to a strong magnetic
field
¾ Unpaired electrons
No two electrons in an atom can have the same set of
four quantum numbers n, l, ml, and ms.
¾In order to put more than one electron in an orbital,
electrons must have different values of ms. i.e. they must
have different spins.
¾Maximum of 2 electrons per orbital
5
Chapter 8
Chapter 8
Energy of Orbitals
¾ For the same type of orbital (same ______), energy
increases as n increases
(1s < 2s < 3s < 4s…)
¾ For the same n, energy increases s < p < d < f
(3s < 3p < 3d)
¾ 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
Orbital Diagrams
orbital
Draw this diagram and
by hand and start
filling out electrons.
This diagram will be
counted as one
problem i.e. 1/4th
extra credit
n=3
3s
3p
3s
3p
p
3d
subshell
n=3
3d
shell
n=3
3s
3p
3d
6
Chapter 8
Chapter 8
Electron Configuration Rules
¾ Electrons fill the lowest energy orbital first (Aufbau
principle)
This diagram and any 10
elements’ electron-filled orbital
diagram will be counted as one
problem i.e. 1/4th extra credit
1s
2s
2p
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
Energy of Orbitals: Summary
¾Pauli exclusion principle
No two electrons in an atom can have the same set
of four quantum numbers n, l, ml, and ms.
¾ Two electrons (max) per orbital
¾ Maximize parallel spins when filling a subshell
¾ If more than one orbital in a subshell is available,
electrons will fill empty orbitals in the subshell first.
(Hund’s Rule)
Alternately….
¾ Electrons preferred to be unpaired as long as an
empty orbital with the same energy is available
7
Chapter 8
Chapter 8
Electron Configurations
¾ Three notations for the arrangement of
electrons in atoms
¾ Hydrogen
Orbital Box Notation
¾
noble gas notation
number of electrons
Electron Configurations
¾ Hydrogen
1s1
Orbital box diagrams
spdf notation
Chapter 8
Electron Configurations
Lithium # of es =3
Α. 1s22s1
B. 1s12s12p1
C. 2p3
D. 1s3
number of electrons
1s1
orbital type (l)
electron shell (n)
spdf Notation
Orbital Box Notation
¾
Oxygen: # of es =8
Α. 2s22p6
B. 1s12s12p6
C. 1s22s22p4
D. 1s22s32p3
orbital type (l)
electron shell (n)
spdf Notation
8
Chapter 8
Chapter 8
Chapter
1A 8
Electron Configurations
¾ Hydrogen
number of electrons
1s1
Orbital Box Notation
¾
orbital type (l)
More Examples
¾ Provide the electron configurations (in orbital box,
spdf and noble gas notation)
(a) P
8A
Transition
Metals
2A
3A
4A 5A 6A 7A He
Li Be
B
C
N
O
F
Ne
Na Mg
Al Si
P
S
Cl
Ar
K
3B 4B 5B 6B 7B 8B
Ca Sc Ti
V
Cr Mn Fe
8B 8B 1B 2B
Co Ni
Cu Zn Ga Ge As
Se Br Kr
electron shell (n)
Rb
spdf Notation
(b) V
Chlorine: # of es =17
Α. 1s22s22p63s23p33d3
B. 1s22s22p63s23p5
C. 1s22s22p53s23p6
D. 1s22s32p63s13p6
H
(c) I
Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb
Te I
Cs Ba La Hf
Sr Y
Zr
Ta W
Po At Rn
Fr Ra Ac Rf
Db Sg Bh Hs
Re Os
Ir
Pt
Au Hg Tl Pb Bi
Xe
Mt
Ce Pr
Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
Th Pa
U
Np Pu Am Cm Bk Cf
Es Fm Md No Lr
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Chapter 8
Chapter 8
Chapter 8
More Examples: Ions
Some Anomalies?
Periodic Table Organization
(a) S2–
Chromium and
copper
So does S2– = Ar?
Isoelectronic species
((b)) Br –
Transition metal
ions
(c) Al3+
¾Half-filled and fully filled d-subshells have extra
stability (lower energy).
10
Chapter 8
Chapter 8
Periodic Table Organization
atoms where an s subshell is being filled
atoms where a p subshell is being filled
atoms where a d subshell is being filled
¾ Valence electrons
¾ Core electrons: electrons included in the noble gas
notation
Li (3): 1s2 2s1
Na(11): 1s22s22p63s1
[He] 2s1
[Ne] 3s1
¾ Same group = same number and type of valence
electrons
¾ Take the case of Li
1s22s1
Radial probability
s-block
p-block
d-block
Chapter 8
Effective Nuclear Charge
Electron Configurations
¾ Valence electrons: electrons in the outermost
shells Î responsible for all macroscopic
properties
¾ Core electrons: electrons included in the noble
gas notation
t ti
Li (3): 1s2 2s1
Na(11): 1s22s22p63s1
Distance from nucleus
[He] 2s1
[Ne] 3s1
¾ Same group = same number and type of
valence electrons Î Similarity of properties
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Chapter 8
Chapter 8
Electron Configurations: Atoms and
Ions
Chapter 8
Periodic Properties
¾ You will need to know the following:
Noble gas elements
He (2) : 1s2
1. Definitions and chemical equations where
appropriate
Effective Nuclear Charge
¾ Valence electrons don’t “feel” the full charge of the
nucleus
¾ Valence electrons are shielded
Ne (10) : [He] 2s2 2p6
Ar (18) : [Ne] 3s2 3p6
Kr (36): [Ar] 4s2 4p6
K+ (19-1= 18) ≡ [Ar] or [Ne] 3s2 3p6
2. Periodic trends moving up and down and left to right
across the periodic table
¾ But … valence electrons “feel” a charge that is
greater than Z – core electrons
¾ Valence electrons are not completely shielded
3. Explanations of the trends
4. How the atomic properties affect chemical properties
Br- (35 +1= 36) ≡ [Kr] or [Ar] 4s2 4p6
12
Chapter 8
Chapter 8
Atomic Size
Chapter 8
Atomic Size
¾ The distance from the nucleus to the edge of the
outermost electron
¾ Periodic trend:
Atomic Size
Decrease
across a Decrease
period
Decrease
across a
across a
period
period
Decrease
¾ Explanation:
Effective nuclear charge
increases across the group
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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
C.
group, new shells (i
(i.e.
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 Size
The best way to explain the decrease of atomic size as
one goes across periods
Α. The electrons repel less, therefore the atom shrinks
B. 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
manyy fold
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:
¾ Identify the one which is correctly arranged in order of
increasing (smallest to largest) atomic size:
a. Be, C, O
b. Be, O, C
c. O, C, Be
d. C,O, Be
#2:
¾ Identify the one which is correctly arranged in order of
increasing (smallest to largest) atomic size:
a. Cl, K, S
b. Cl, S, K
c. K, S, Cl
d. K, Cl, S
14
Chapter 8
Ionization Energy (IE)
¾ The energy required to remove an electron from a
gaseous atom
A(g) + energy Æ A+(g) + e-
Chapter 8
Sign Conventions
Ionization Energies
¾ Energy absorbed (in) = a positive value + 165 kJ
¾ Energy required (input, raw material)
¾ Energy released (out) = a negative value - 165 kJ
¾ Energy produced (output, product)
¾ Energy input required
Decrease
Chapter 8
Effective nuclear charge increases
across the group
The sign tells us which way energy is going
The magnitude tells us how much energy is required
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
Chapter 8
Successive Ionizations
IE1
IE2
IE3
IE4
IE5
IE6
Na
495
4560
Mg
735
1445
7730
Al
580
1815
2740
Si
780
1575
3220
4350
16100
P
1060
1890
2905
4950
6270
21200
S
1005
2260
3375
4565
6950
8490
Successive Ionizations
IE7
¾ For Mg, 2nd IE > 1st IE
¾ For Al, 3rd IE > 2nd IE > 1st IE
¾ Why?
11600
27000
¾ For Mg, 3rd IE >>> 2nd IE
¾ For Al, 4th IE >>> 3rd IE
Example:
Na(g) + IE1 Æ Na+(g) + eNa+(g)
+ IE2 Æ
Na2+(g)
+
¾ Why?
e-
16
Chapter 8
Chapter 8
Ionization Energies: Summary
¾First ionization energies generally increase across a period
and decrease down a group
¾Effective nuclear charge increases
across the
th group
#3:
¾ Arrange each set of atoms in increasing IE1:
a. Sr, Ca, Ba
b. Ba, Sr, Ca
c. Ca, Sr, Ba
d. Ba, Ca, Sr
#4:
¾ Arrange each set of atoms in increasing IE1:
a. Br, Rb, Se
b. Br, Se, Rb
c. Rb, Br, Se
d. Rb, Se, Br
Chapter 8
Electron Affinity
¾ The energy released when an electron is added to a
gaseous atom
A(g) + e- Æ A-(g) + energy
¾ A free electron is not a stable. It would always be
associated with an atom.
17
Chapter 8
Chapter 8
Electron Affinity Predictions
Chapter 8
Electron Affinity Trends
Electron Affinity Summary
A(g) + e- Æ A-(g) + energy
Exception
¾ Across a period: Should it get easier or harder to add
an electron?
¾ Down a group: Should it get easier or harder to add an
electron?
¾ If it’s easy to add an electron, is the EA a large
negative number or a small negative number?
¾ Deviations from the general trends
¾An element with a high ionization energies
generally has a high affinity for an electron.
¾Effective nuclear charge increases across the
group and decreases down a group
18
Chapter 8
Chapter 8
Trends in Metallic Behavior
Relative tendencies to lose and gain electrons
Acid-base Behaviors of
Elemental OXides
Chapter 8
Ionization: Change in Size
¾ Why does the size decrease?
Metals donate
electrons to oxygen
Nonmetals share
electrons to oxygen
3 p+ and 3 e3 p+ and 2 e-
Ionic
Metal oxides react with
water to produce hydroxides
(OH-) that are basic
Covalent
Nonmetal oxides react with
water to produce acids
that releases proton in solution H+
Elements at the left form cations easily
Elements at the right form anions easily
19
Chapter 8
Chapter 8
Ionization: Change in Size
¾ Why does the size increase?
9 p+ and 10 e9 p+ and 9 e-
Review
¾ Zeff: the positive charge actually felt by a valence electron
¾ Atomic size: The distance from the nucleus to the edge of
the outermost electron
¾ IE: The energy required to remove an electron from a
gaseous atom.
¾ Successive ionization
¾ EA: The energy released when an electron is added to a
gaseous atom
¾ Ion sizes
Chapter 8
The Reaction of Na and Cl
Na
IE
495
4560
EA
EA > 0
Cl
1251
-348
¾ How can we use these numbers to explain the
product of the reaction?
¾ Is NaCl2 a reasonable product?
¾ Is Na2Cl a reasonable product?
20
Chapter 8
Chapter 8
Periodic trends and Chemical
Properties
• Reactivity of metals
Chemical Reactivity Summary
¾ Noble gases
high IE, low EA
do not react
¾ Metals
low IE, low EA
lose electrons
¾ Non-metals
high IE, high EA
add electrons
• Reactivity of nonmetals
¾ Metal + non-metal Æ metal loses e-’s and non-metal
gains e-’s
¾ non-metal + non-metal Æ shared e-’s
21