chapter 4 periodic table student
DESCRIPTION
periodic tableTRANSCRIPT
Lecturer:
Nor Fadilah Chayed
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CHAPTER 4CHAPTER 4PERIODIC TABLEPERIODIC TABLE
Upon completion of this course, students should be able to:1. Indicate period, group and block (s,p,d,f).2.Specify the position of metals, metalloids and non-metals.3.Deduce the position of elements from electronic configuration.4.Explain the variation in atomic and ionic radii.5.Explain the radius of isoelectronic species.6.Define first and second ionization energy and explain the
variations in the first ionization energy across period and down the group.
7.Define electron affinity and electronegativity.
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CLASSIFICATION OF ELEMENTS
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Dmitri Mendeleev (1869) and Lothar Meyer proposed the periodic law:
Elements were arranged based on the regular, periodic recurrence of properties
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Henry Moseley (1930) discovered the
‘atomic number’ which is later used
as the basis for classifying elements
in the Modern Periodic Table.
In the periodic table, the elements are
placed in increasing atomic
number, starting at the upper left and
arranged in a series of horizontal
rows.
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Classification of the Elements
Periodic table and electron configuration
Electron configuration play an important role in the construction
of the periodic table
The order of filling orbitals is as follows:
1s < 2s < 2p < 3s < 3p < 4s < 3d <4p < 5s < 4d < 5p < 6s < 4f <5d
< 6p < 7s < 5f < 6d < 7p
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Categories of Electrons
Inner (core) electrons are those an atom has in common with the previous noble gas and any completed transition series.
Outer electrons are those in the highest energy level (highest n value).
Valence electrons are those involved in forming compounds.- For main group elements, the valence electrons are the outer electrons.
- For transition elements, the valence electrons include the outer electrons and any (n -1)d electrons.
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Metals are found to the left of the metalloids
Nonmetals are found to the right of the metalloids.
1. Classification according to metallic property:
Metals located on the left part of the periodic table The alkali metals - Group 1AThe alkaline earth metal – Group 2A
Non metals are located on the right partThe halogens – Group 7A The Nobel gases – Group 8A
Metalloids possess both metallic and non metallic propertiesB, Si, Ge, As, Sb, Te, Po and At
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(a) Group
Vertical bars, ↓
Each group consists of elements that have the same number
of valence electrons in their valence (outermost) shells
There are a total of 18 groups from Group1 to Group 18
E.g. Group 1A (valence electron = 1)Li (3) : 1s2 2s1
Na (11) : ___________
Group 5A (valence electron = 2+3 = 5)N (7) : 1s2 2s2 2p3
P (15) : ___________
Parts of Periodic Table
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Elements with similarproperties are organizedin groups or families.
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Electron Configuration and Group
Similar outer electron configurations correlate with similar chemical behavior.
Elements in the same group of the periodic table have the same outer electron configuration.
Elements in the same group of the periodic table exhibit similar chemical behavior.
Potassium reacting with water. Chlorine reacting with potassium.12
(b) Period
Horizontal bars, →
Each period consists of elements that have the same
similar valence shells
E.g.
2nd Period (valence n=2) Li: 1s2 2s1
3rd Period (valence n=3) Al: 1s2 2s2 2p6 3s2 3p1
5th Period (valence n=5) Zr: [Kr] 5s2 4d2
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2. Classification according to valence shell electron configuration
(i) s-block elements
-Metal (group 1 and 2)
-Half filled s orbital (s1)or fully filled s orbital (s2)in the
valence shell
(ii) p-block elements
-Metals, non-metals and metalloids (group 13-18)
-Valence shell configuration varies from s2p1 to s2p6
(iii) d-block elements
-Metals with electronic configuration varies from s2d1 to s2d10
(iv) f-block elements
-Metals have incomplete f orbital
- Incomplete 4f orbital (The lanthanides)
-Incomplete 5f orbital (Actinides)14
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ns1
ns2
ns2
np1
ns2
np2
ns2
np3
ns2
np4
ns2
np5
ns2
np6
d1
d5 d10
4f
5f
Ground State Electron Configurations of the Elements
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Electron Configurations of Cations and Anions of Representative Element
Na : 1s22s22p5 3s1 or [Ne] 3s1 Na+: [Ne]
Ca : 1s22s22p6 3s23p6 4s2 or [Ar] 4s2Ca2+: [Ar]
Al :1s22s22p63s23p1 or [Ne]3s23p1 Al3+: [Ne]
Atoms lose electrons so that cation has a noble-gas outer electron configuration.
F : 1s22s22p5 F- : 1s22s22p6 or [Ne]
O : 1s22s22p4 O2-: 1s22s22p6 or [Ne]
N : 1s22s22p3 N3- : 1s22s22p6 or [Ne]
Atoms gain electrons so that anion has a noble-gas outer electron configuration.
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+1
+2
+3 -1-2-3
Cations and Anions of Representative Elements
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Electron Configurations of Cations of Transition Metals
When a cation is formed from an atom of a transition metal, electrons are always removed first from the ns orbital and then from the (n – 1)d orbitals.
Fe: [Ar]4s23d6
Fe2+: [Ar]4s03d6 or [Ar]3d6
Fe3+: [Ar]4s03d5 or [Ar]3d5
Mn: [Ar]4s23d5
Mn2+: [Ar]4s03d5 or [Ar]3d5
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Problem 1
Determine the period, block and group for each element with the following configuration:
A: 1s2 2s2 2p3B: 1s2 2s2 2p6 3s2 3p6C: 1s2 2s2 2p6 3s2D: 1s2 2s2 2p6 3s2 3p6 4s2 3d3
Answer:
Metallic Behavior Metallic character decreases across a period (increase non
metallic character) and increase down a group.
On descending a group (downwards), the metallic character
increase:
(i) ionization energy decrease
(ii) reactivity increase.
Metallic bonding is stronger and the melting point is higher if:
1) The ionic size is smaller
2) The number of valence electron is greater
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Trends in metallic behavior
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Periodic Physical Properties
The periodic law: When elements are arranged in the periodic table in order of increasing atomic number, a regular change in the outer electronic configuration and a periodic variation of properties is observed.
The following physical properties show a periodic trend:1. Atomic size2. Ionization energy3. Electron affinity4. Electronegativity
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1. Atomic size The atomic radius
= ½ of the distance between 2 nuclei of two adjacent
atoms
= ½ d
Atomic radius of an element is determined by two factor:
Screening effect
Nuclear charge24
Screening effectThe atomic radius increase downwards in a group.Electrons fill up a new shell downwards. Outer electrons are
shielded from the nucleus by electrons in inner shells (mutual repulsions between electrons in different shell) and are less tightly held.
Nucleus chargeThe atomic radius decrease towards the right across a
period.Electrons fill up the same shell, cause the effective nuclear
charge (Zeff) of the atom increases which pulls all electrons closer to nucleus, thus the electrons are held more tightly.
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Effective nuclear charge (Zeff) is the “positive charge” felt by an electron.
Na
Mg
Al
Si
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12
13
14
10
10
10
10
1
2
3
4
186
160
143
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ZeffCoreZ Radius
Zeff = Z - 0 < < Z ( = shielding constant)
Zeff Z – number of inner or core electrons
8.3
Zeff = Z - σ
As the distance from nucleus increases, σ increases andZeff decreases
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Effective Nuclear Charge (Zeff)
increasing Zeff
incr
easi
ng Z
eff
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8.328
Atomic radius across d-block elements d-block elements consist of three series Period 4, 5 and 6
The atomic radii of first row of d-block elements tend to be
approximately constant across the period
Reason:
additional electrons go into inner electron subshell (3d).
At the same time, the number of electrons in the
outermost subshell (4s) remain constant.
3d electrons shield the outer 4s electrons from nuclear
charge more effectively than the outer shell can shield
one another.
Effect of increased nuclear charge cancelled by the
screening effect electron in 3d orbital29
Size of cations
Sizes of cations are smaller than their corresponding parent atom .
The atoms lose the valence electrons, leaving electrons of the inner shells which
tend to attract the nucleus more, thus decreasing size.
E.g. Li+ ion < Li atom
Size of anions
Sizes of anions are larger compared to their corresponding parent atom.
The electrons are added into the same shell and tend to repel each other and so
increase size.
E.g. F- ion > F atom
Ionic Radius
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Cation is always smaller than atom from which it is formed.Anion is always larger than atom from which it is formed.
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Comparison of Atomic Radii with Ionic Radii
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Trends in Atomic Size
Atomic size increases as the principal quantum number n increases.- As n increases, the probability that the outer electrons will be farther from the nucleus increases.
Atomic size decreases as the effective nuclear charge Zeff increases.- As Zeff increases, the outer electrons are pulled closer to the nucleus.
For main group elements:- atomic size increases down a group in the periodic table and decreases across a period.
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Isoelectronic: a series of ions or atoms which have the same number of electrons and same ground-state electron configuration.
Sizes of the ion/atom decrease as the number of protons increases.
Reason: higher nuclear charge (protons) are pulling in the same number of electrons
The larger number of electrons, the greater repulsion between electrons, the larger ionic or atomic radius.
E.g. 10 electron series: 10Ne>11Na+ >12Mg2+ >13Al3+
Isoelectronic ions
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Ne : 1s22s22p6 Mg2+ 1s22s22p6 or [Ne]
Al3+: 1s22s22p6 or [Ne]
Na+, Mg2+, Al3+are all isoelectronic with Ne
Problem 2 : Ranking Elements by Atomic Size
Using only the periodic table (not Figure 8.15), rank each set of main-group elements in order of decreasing atomic size:
(a) Ca, Mg, Sr (b) K, Ga, Ca
(c) Br, Rb, Kr (d) Sr, Ca, Rb
Answer:
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Answer:
Problem 3: Ranking Ions by Size
Rank each set of ions in order of decreasing size, and explain your ranking:
(a) Ca2+, Sr2+, Mg2+ (b) K+, S2−, Cl− (c) Au+, Au3+
PLAN: Find the position of each element on the periodic table and apply the trends for ionic size.
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2. Ionization energy Ionization energy is the minimum energy (kJ/mol) required to remove an electron from a gaseous atom in its ground state.
I1 + X (g) X+
(g) + e-
I2 + X (g) X2+(g) + e-
I3 + X (g) X3+(g) + e-
I1 first ionization energy
I2 second ionization energy
I3 third ionization energy
I1 < I2 < I3
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First ionization energy = the energy needed to remove 1 mole of the outermost electrons from 1 mole of neutral atoms in the gas phase
M(g) → M+(g) + e- ΔH1 = +X kJ mol-1
Second ionization energy = energy required to remove 1 mole of electrons from 1 mole of unipositive ions in gaseous state
M+(g) → M2+
(g) + e- ΔH2 = +Y kJ mol-1
Factor that affect ionization energy:Atomic radiusNuclear chargeScreening effect
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Atomic radius Atomic radius increase (distance outer electron from nucleus increase), the ionization decrease (easy to lose electron).
Nuclear charge Nuclear charge becomes more positive, attraction on the outershell electrons increase, cause ionization energy increase.
Screening effect (repulsion effect)Valence electron are shielded from the attraction of the nucleus by the screening effects of the electrons in inner shells.Lower quantum number (n) have stronger shielding effect.Radius of positive ion is smaller than its atom, attraction between nucleus and electrons left become stronger (screening effect decrease).IE always increase in the order: 1st IE < 2nd IE < 3rd IE…..
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Factor that affect ionization energy:
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Filled n=1 shell
Filled n=2 shell
Filled n=3 shell
Filled n=4 shell
Filled n=5 shell
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General Trend in First Ionization Energies
Increasing First Ionization Energy
Incr
ea
sing
Firs
t Io
niz
atio
n E
ner
gy
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Answer:
Problem 4: Ranking Elements by First Ionization Energy
Using the periodic table only, rank the elements in each of the following sets in order of decreasing IE1:
(a) Kr, He, Ar
(b) Sb, Te, Sn
(c) K, Ca, Rb (d) I, Xe, Cs
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3. Electron affinity The electron affinity is the energy involved when 1 mole of electrons
is gained (accepted) by 1 mole of neutral atoms in the gas phase.
The process would be represented by the following equation:
X (g) + e- → X- (g) ΔHEA = -A kJ mol-1
The addition of one electron to a neutral atom is exothermic for
nearly all atoms
When an electron is more easily accepted into an atom, more
energy is given off, thus the higher the electron affinity and more –ve
is the value of ΔHEA 44
F (g) + e- F-(g) H = -328 kJ/mol EA = +328 kJ/mol
Electron affinities of the main-group elements.
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General trend:
EA become less negative (decreases) downwards in the
periodic table
Reason: Electrons are added less easily into the atom
because of the increase in size and there is greater repulsion
from electrons already present
EA become more negative (increases) towards the right
across the period
Reason: Electrons are added more easily into the atom
because of the smaller size
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Trends in three atomic properties.
` Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
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4. Electronegativity Electronegativity is the ability of the atom in a covalent bond to attracts
pairs of shared electrons to itself.
The higher the electronegativity of an atom, the greater its attraction for
bonding electrons.
Elements with low ionization energies have low electronegativity
(electropositive elements).
Elements with high ionization energies have high electronegativity
(electronegative elements).
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General trend:
In a group (downwards) the electronegativity decrease.
Reason: Increased distance between the valence electrons and the nucleus weakens the pull of the nucleus on the electrons.
The a period (towards the right), the electronegativity increase.
Reason: Decreased distance between the valence electrons and the nucleus, thus stronger pull of the nucleus on the electrons.
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LEFT TO RIGHT(Across period)
TOP TO BOTTOM (Down the group)
Atomic Radius Decreases Increases
Ionization energy Increases Decreases
Electron affinity Increases (more –ve)
Decreases (less –ve)
Electronegativity Increases Decreases
Summary of physical properties
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