Download Principles of Chemistry The Development of Periodic Table

Principles of Chemistry
A Molecular Approach,
1st Ed.
Nivaldo Tro
Chapter 8
Periodic Properties of
the Elements
Chem 1010
The Development of
Periodic Table
Dmitri Mendeleev(1834-1907)
Arranged elements in a table form by atomic mass
(modern PeriodicTable)
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 patterns to predict properties of undiscovered
elements
where atomic mass order did not fit other properties,
he reordered by different properties I.e. Te and I
Mendeleev’s Predictions
Quantum Theory and Periodic table
Mendeleev’s periodic law- predicts what the
properties of an element will be based on its position
on the table
Quantum theory explains why the periodic trends in
the properties exist.
Quantum theory describes the behaviour of
electrons in atoms
Helps us to understand the chemical behaviour
Chemical bonding involve the transfer or
sharing of electrons, and
Spin quantum number(m s): 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
with only two possibilities
spin up or spin down
Spin quantum number (m s) describes how the
electron spins on its axis.
ms = +½ ( )or (-½ ( ).
Opposite spins must cancel in an orbital(paired)
Orbital Diagram
Pauli Exclusion Principle
Pauli Exclusion Principle- No two electrons in an atom can have
the same four quantum numbers
Each orbital can have a maximum of only two electrons with
opposite spins only
Degenerate orbital- orbitals of equal energy e.g. three p orbitals,
five d orbitals or seven f orbitals
Subshell
Number of
Maximum
The maximum number of
Orbitals
Number of
Electrons
electrons in each subshell
s (l = 0)
1
p (l = 1)
3
2
6
d (l =2)
5
10
f (l =3)
7
14
Electron Configurations
Electron configuration- The distribution of electrons into the
various orbitals in an atom in its ground state
Ground State-of an atom is the lowest energy state
The notation for a configuration
Number designates the principal energy level.
The letter designates the sublevel and type of orbital.
The superscript designates the number of electrons in that
sublevel.
For example He (Z=2) has ground state electron
configuration
He = 1s 2
Orbital Diagram: We can represent an orbital as a
orbital with
box and the electrons in that orbital as arrows
2 electrons
Sublevel Splitting in Multielectron
Atoms
Single electron atoms (H and H like)-The sublevels in each
principal energy level of hydrogen all have the same energy
(degenerate)
Energy of hydrogen sublevels depends on principal quantum
number n
Multielectron atoms, the energies of the sublevels are split
Energy of the sublevels depends on n and l
caused by electron–electron repulsion
The lower the value of the quantum number l, the less energy the
sublevel has.
s (l = 0) < p (l = 1) < d (l = 2) < f (l = 3)
Penetrating and Shielding
The effective nuclear charge is the positive charge
an electron experiences from the nucleus minus any
“shielding effects” from intervening electrons
Orbital energies for multielectron
atoms
Important Points
1.
2.
3.
Because of penetration, sublevels within an energy
level are not degenerate.
Penetration of the fourth and higher energy levels is
so strong that 4s orbital lies lower in energy than the
3d orbitals and 5s orbital lies lower in energy than the
4d orbitals.
The energy difference between levels becomes
smaller for higher energy levels (and can cause
anomalous electron configurations for certain
elements).
Electron Configurations
Every atom has an infinite number of possible electron
configurations
Ground state electron configuration:The
configuration associated with the lowest energy level
of the atom
Other configurations correspond to “excited
states”
By filling orbital of lowest energy first you usually get
the lowest energy (ground state) of the atom
Electron Configurations for
Multielectron Atoms
Aufbau principle(build –up)- Electron occupy orbitals so as to
minimize the energy of the atom, therefore orbitals fill from
lowest energy to highest
Order of orbital filling
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f
Pauli Exclusion principle- No two electrons in one atom can
have the same four quantum numbers (max. 2 e- with opposite
spin)
Hund’s rule-When orbitals of same energy are available,
electron first occupy these orbitals singly with parallel spin
rather than in pairs
Once the orbitals of equal energy are half full, the electrons
start to pair.
Another way to remember the order
of orbital filling
Start by drawing a diagram
with each energy shell on one
row, and list the subshells
(s, p, d, f) for that shell in
order of energy (left to right).
Draw arrows down
through the diagonals, looping
back to the next diagonal
each time.
Writing Electron Configuration of
Atoms in Their Ground State
Filling order
s p d f
The # of e- in an atom = Z (atomic number
Full configuration: C (Z= 6) 6 electrons
C: 1s22s22p2
Inner electron configuration -A shorthand way of writing an
electron configuration
use the symbol of the previous noble gas(noble gas core) in
brackets to represent all the inner electrons, then just write
the last set
e.g. Na= 11 electrons = 1s22s22p63s1
Na
[ Ne]3s1
Electron Configurations (Z= 3-10)
Electron configuration, Valence
electrons, and the Periodic Table
Valence electrons-The electrons in the highest principal
energy level(outermost shell)
e-s outside the “noble gas core”—are called valence
electrons.
Important in chemical bonding
Core electrons-Electrons in lower energy shells (inner shells)
Both chemical and physical properties of an atom, are
determined by the number of valence electrons
Sample Problem
Write electron configuration for the following and
identify the valence electron and core electron
Al Z= 13
1s22s22p63s23p1
Kr Z= 36
Rb Z= 37
Electron Configuration and the
Periodic Table
The periodic table is divided into four blocks corresponding to
the filling of the four quantum sublevels (s,p, d, and f)
The number of columns in each “block” is the maximum number
of electrons that sublevel can hold
For the main group elements (recall chapt.2)
group # = Total #of valence electrons
Period # of element = Principal quantum # of the outer shell
Periodic table and electron
configuration
Periodic Table can be divided into four regions or
blocks according to the orbitals being fil led
1A 2A
3A 4A 5A 6A 7A 8A
Ne
P
Periodic table and electron
configuration
s- block elements - Group 1A and 2A (filling of s
orbital)
p-block elements - Group 3A through 8A (filling of p
orbital; ns orbitals are already filled)
d-block elements- Transition metals (filling of (n-1)d
orbitals)
f-block elements- Inner transition metals also called
Lanthanide and actinide (filling of (n-2)f orbital
Sample Problem
Which of the following orbital filling diagrams
represent :
1) ground state or 2) excited state or 3) forbidden
for Boron (Z= 5)
1s
2s
2p
( )
( )
( ) ( ) ( )
(
)
(
)
(
)
(
)
(
)
(
)
(
) ( ) ( )
( ) ( ) ( )
(
) ( ) ( )
Chem 1010
Transition elements
Overlap of principal energy levels- After Ar the
next electron will enter into the lowest sublevel of 4th
principal level 4s instead of the highest sub level of
third principal level (3d) e.g.
Ca (Z = 20) Ca [Ar], 4s2
Elements with Z = 21 –30 (Sc to Zn)
After 4s is filled , 3d will start
Energy of 4s orbital < 3d orbital and so on
1s, 2s, 2p ,3s, 3p, 4s, 3d
At Z= 36 Kr, 4p is completely filled.
4p6
Anomalous Electron Configurations
Half filled and filled subshells have an unusual stability
Leads to anomalies in electron configuration
Sometimes s electron is “promoted” to d subshell (anamolous
configuration)
Results in lowering the total energy of the atom.
Due to decrease in electron- electron repulsions
Anomalies occur where the energy differences between subshells are small i.e. 4s and 3d or 5s and 4d
Z > 40
Example: Give an abbreviated electronic configuration of Ag
Solution: Ag is one of the elements that will transfer an electron from one
sub-shell to another in order to lower the total energy of the atom.
Therefore e- configuration of Ag
[Kr] 5s1 4d10
Exceptions to the Aufbau principle
Some of the transition metals have anomalous electron
configurations in which the ns only partially fills before the
(n-1)d or doesn’t fill at all
Therefore, their electron configurations must be found
experimentally (XPS spectroscopy)
Expected (Aufbau principle)
Experimental
Cr = [Ar]4s23d4
Cr = [Ar]4s13d5
Cu = [Ar]4s23d9
Cu = [Ar]4s13d10
Mo = [Kr]5s24d4
Mo = [Kr]5s14d5
Ru = [Kr]5s24d6
Ru = [Kr]5s14d7
Pd = [Kr]5s24d8
Pd = [Kr]5s04d10
Quantum mechanical Model and
chemical properties
Quantum mechanical model accounts for the
chemical properties of the elements
The chemical properties of elements are largely
determined by their # of valence e-s
Elements in the same column have similar
chemical and physical properties?
same # of valence e-s
Noble Gas
Electron Configuration
Quantum-mechanical calculations show that
an atom with eight valence electrons should
be unreactive.
The noble gases have eight valence electrons
(ns2np6) very stable configuration
Except for He, which has only two electrons
Properties- noble gases are especially
nonreactive
He and Ne are practically inert
Everyone Wants to Be Like a Noble
Gas!
The alkali and alkaline earth metals (Group IA & IIA) have one
and two more electron than the previous noble gas
Very reactive metals
Valence shell configuration ns1 and ns2
In their reactions,
They tend to lose their extra electrons, resulting in the same
electron configuration as a noble gas e.g. Na+ and Ca2+
The halogens (Group7A)- valence shell e-configuration (ns2np5)
All are one electron short of the next noble gas configuration
Most reactive nonmetals- form anion with –1 charge
with nonmetals, they tend to share electrons
Trend in Atomic Radius—Main Group
van der Waals radius (nonbonding atomic radius)One–half the distance between adjacent nuclei in the
atomic solid
Covalent radius (bonding atomic radius) – One-half
the distance between the nuclei of two identical
covalently bonded atoms
Approximate bond length = sum of the atomic radii of the two
covalently bonded atoms
Periodic Trends in size of atomsMain Group
The general trends in the atomic radii of main group
elements in the periodic table
Atomic radius increases down a column (or group)
The principle quantum # n increases resulting in larger
orbitals
effective nuclear charge fairly close
Atomic radius decreases across period (from left to
right)
adding electrons to same valence shell
effective nuclear charge increases
valence shell held closer
Trends in Atomic Radius
Chem 1010
Factors determining the size of an
atom
Two factors determine the size of an atom
1. The principal quantum number, n-As the value of
n increases the size of the outer orbital increases
(resulting in larger atoms)
2.The effective nuclear charge (Zeff)- The positive
charge an electron experiences from the nucleus
minus any “shielding effects” from intervening
electrons
Atomic Radii and the transition
elements
In general the atomic radii of the transition elements
do not follow the same trend as main group elements
increase in size down the group
Atomic radii of transition metals stay roughly
constant across each row. Why?
The # of electrons in the outermost principal
energy level (highest n value) is nearly constant
valence shell ns2, not the d electrons
effective nuclear charge on the ns2 electrons
approximately the same
Electron Configuration and
Ion Charge
Ions are formed by loss (cation + ve) or gain (anion -ve) of
electrons by a neutral atom
the charge on an ion is predictable based on its position on
the periodic table.
Electron configuration of the ions is the same as the nearest
noble gas.
Isoelectronic - species with same number of electrons. Ne and
Al3+
Ions of the main-group elements
The common monatomic ions found in compounds of the main
group elements fall into three catagories
Cations of group IA-IIIA: Positively charged ions
The ion charges = group #
2. Cations of group IIIA-VA (e- config. ns2):
The ion charges = Group #
e.g. Tl+ ’ Sn2+, Pb2+ and Bi3+
1.
3. Anions of Groups VA –VIIA:The Ion charges = group# - 8
Cations form when the atom loses electrons from the valence
shell.
Al (Z=13) = 1s22s22p63s23p1
Al3+ ion =
1s22s22p6
Electron Configurations of Ions
Transition metals- The ‘s’ electrons are “first in”
with the atoms and first out” with the cations.
Lose electrons from the valence shell first, which is
not the last sublevel to fill according to the Aufbau
sequence
For example, zinc generally loses two electrons
from its 4s sublevel to adopt a “pseudo”-noble
gas configuration
[Ar]4s23d10
[Ar]3d10
Magnetic Properties of
Transition Metal Atoms and Ions
Only atoms with unpaired electrons exhibit magnetic
susceptibility
A paramagnetic substance is one that is weakly
attracted by a magnetic field, usually the result of
unpaired electrons
A diamagnetic substance is not attracted by a
magnetic field generally because it has only paired
electrons
slightly repelled by a magnetic field
Zn atom and Zn2+ ions are diamagnetic.
Zn
[ Ar]4s23d10
Zn2+
[ Ar]3d1
Example 8.6
Write the electron configuration and determine
whether the Fe atom and Fe3+ ion are paramagnetic
or diamagnetic.
previous noble gas = Ar 18 e
Fe atom = [Ar]4s23d6
unpaired electrons
4s
3d
• paramagnetic
•
Fe3+ ion = [Ar]4s03d5
• unpaired electrons
• paramagnetic
Chem 1010
4s
3d
Periodic Properties
The periodic law states that when the elements are
arranged by atomic number, their physical and
chemical properties vary periodically.
We will look at the following periodic properties:
Atomic or ionic radius
Ionization energy
Electron affinity
Metallic character
Chem 1010
Trends in Ionic Radius
Ion size increases down the group and decreases
across a row Why?
higher valence shell = larger size
Cations are smaller than the neutral atom
Na
[ Ne] 3s1 and Na+ [ Ne]
Anions are bigger than the neutral atom Why?
cations smaller than anions
for isoelectronic species-species with same
number of electrons.
larger positive charge = smaller cation
larger negative charge = larger anion
Chem 1010
Chem 1010
Order the following sets by size (smallest to largest).
Zr4+, Ti4+, Hf4+
same column and charge;therefore,
Ti4+ < Zr4+ < Hf4+
Na+, Mg2+, F- , Ne
isoelectronic; therefore,
Mg2+ < Na+ < Ne < F-
I- , Br- , Ga3+, In+
Ga3+ < In+ < Br- < I-
Periodic Trends: Ionization Energy
Ionization Energy-Minimum energy needed to
remove an electron from an atom or ion in the
gaseous state Units (eV) 1 eV = 96.5 kJ mol-1
Endothermic process
First ionization energy (IE1)–energy required to
remove electron from neutral atom
M(g)
M+ (g) + 1 eIE1 =
kJ mol-1
Second IE2 – energy required to remove an
electron from 1+ ion
M+ (g)
M2+(g) + 1 eIE2 =
kJ mol-1
Chem 1010
General Trends in
First Ionization Energy
For main group elements:
Ionization energy decreases down the group
valence electrons farther from nucleus
experience less Zeff
Ionization energy generally increases across the
period
effective nuclear charge increases and
electrons are held tightly by the nucleus
Chem 1010
First ionization Energy versus At. #
Chem 1010
Irregularities in the Trend
Which is easier to remove, an electron from B or Be?
Why?
Be(Z = 4) and B(Z= 5)
Which is easier to remove an electron from, N or O?
Why?
N(Z =7) and O(Z=8)
Chem 1010
Trends in Successive
Ionization Energies
Removal of each successive
electron costs more energy.
Regular increase in energy
for each successive v. e-
Large increase in energy when start removing core
electrons
Chem 1010
Trends in Successive Ionization
Energies
Chem 1010
Electron Affinity
The electron affinity is the energy change for the
process of adding an electron to a neutral atom in
the gaseous state to form a negative ion
For a chlorine atom, the first electron affinity is
illustrated by:
Cl(g) + 1 eCl-(g)
EA= -349 kJ mol-1
Defined as exothermic (-), but may actually be
endothermic (+)
Some alkali earth metals and all noble gases are
endothermic. Why?
Chem 1010
Trends in Electron Affinity
The more energy that is released, the larger the
electron affinity of the atom.
The more negative the number, the larger the EA,
and the more stable the negative ion that is
formed
Most groups in P.T do not exhibit any definite trend
in EA
Generally EA become more negative (adding an ebecome more exothermic) as you move to the right
across a period
Highest EA in period = halogen
Chem 1010
Chem 1010
Properties of Metals and Nonmetals
Metals
malleable and ductile
shiny, lustrous, reflect light
conduct heat and electricity
most oxides are basic and ionic
form cations in solution
lose electrons in reactions—oxidized
Nonmetals
brittle in solid state
dull, nonreflective solid surface
electrical and thermal insulators
most oxides are acidic and molecular
form anions and polyatomic anions
gain electrons in reactions—reduced
Chem 1010
General Trends in Metallic character
Metallic character – The tendency to lose electrons
decreases from left to right across a period.
Metals are found at the left side and middle of
the period Table
increases down the column.
Nonmetals are found at the top right of the P.T
Chem 1010
Operational Skills
Applying the Pauli exclusion principle and applying
Hund’s rule to write electron configuration
Writing electron configurations of atoms and ions
using the Aufbau principle
Determining the electron configuration using the
period and group numbers
Writing orbital diagrams
Identifying valence electrons and core electrons
Applying periodic trends to predict ion size, relative
ionization energies and metallic character
Chem 1010
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