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Periodic Table
Larry SchefflerLincoln High School
The Periodic Table-Key QuestionsWhat is the periodic table ?What information is obtained from the table ?How can elemental properties be predicted based on the Periodic Table?
Periodic Table• The development of the periodic table brought a
system of order to what was otherwise an collection of thousands of pieces of information
• The periodic table is a milestone in the development of modern chemistry. It not only brought order to the elements but it also enabled scientists to predict the existence of yet undiscovered
elements.
Early Attempts to Classify Elements
Dobreiner’s Triads (1827)• Classified elements in sets of three
having similar properties.• Found that the properties of the middle
element were approximately an average of the other two elements in the triad.
Dobreiner’s TriadsElement Atomic
MassAverage Density Average
ClBrI
35.579.9126.9
81.21.563.124.95
3.25
CaSrBa
40.187.6137.3
88.71.552.63.5
2.53
Note: In each case, the numerical values for the atomic mass and density of the middle element are close to the averages of the other two elements
Newland’s Octaves -1863John Newland attempted to classify the then 62 known elements of his day.He observed that when classified according to atomic mass, similar properties appeared to repeat for about every eighth elementHis Attempt to correlate the properties of elements with musical scales subjected him to ridicule.In the end his work was acknowledged and he was vindicated with the award of the Davy Medal in 1887 for his work.
Dmitri Mendeleev Dmitri Mendeleev is
credited with creating the modern periodic table of the elements.
He gets the credit because he not only arranged the atoms, but he made predictions based on his arrangement which were shown to be quite accurate.
Mendeleev’s Periodic Table
• Mendeleev organized all of the elements into one comprehensive table.
• Elements were arranged in order of increasing mass.
• Elements with similar properties were placed in the same row.
Mendeleev’s Periodic Table
Mendeleev’s Periodic Table
Mendeleev left some blank spaces in his periodic table. At the time the elements gallium and germanium were not known. He predicted their discovery and estimated their properties
Periodic Table
The Periodic Table has undergone several modifications before it evolved in its present form. The current form is usually attributed to Glenn Seaborg in 1945
Periodic Table Expanded ViewThe Periodic Table can be
arranged by energy sub levels The s-block is Group IA and & IIA, the p-block is Group IIIA - VIIIA. The d-block is the transition metals, and the f-block are the Lanthanides and Actinide metals
The way the periodic table usually shown is a compressed view. The Lanthanides and actinides (F block)are cut out and placed at the bottom of the table.
Periodic Table: Metallic Arrangement
Layout of the Periodic Table: Metals vs. nonmetals
1IA
18VIIIA
12IIA
13IIIA
14IVA
15VA
16VIA
17VIIA
2
33
IIIB4IVB
5VB
6VIB
7VIIB
8 9VIIIB
10 11IB
12IIB
4
5
6
7
MetalsNonmetals
The Three Broad Classes Are Main, Transition, Rare Earth
Main (Representative), Transition metals, lanthanides and actinides (rare earth)
Reading the Periodic Table: Classification
Nonmetals, Metals, Metalloids, Noble gases
Periodic Table: The electron configurations are inherent in the
periodic table
B2p1
1IA
18VIIIA
12IIA
13IIIA
14IVA
15VA
16VIA
17VIIA
2
33
IIIB4IVB
5VB
6VIB
7VIIB
8 9VIIIB
10 11IB
12IIB
4
5
6
7
H1s1
Li2s1
Na3s1
K4s1
Rb5s1
Cs6s1
Fr7s1
Be2s2
Mg3s2
Ca4s2
Sr5s2
Ba6s2
Ra7s2
Sc3d1
Ti3d2
V3d3
Cr4s13d5
Mn3d5
Fe3d6
Co3d7
Ni3d8
Zn3d10
Cu4s13d10
B2p1
C2p2
N2p3
O2p4
F2p5
Ne2p6
He1s2
Al3p1
Ga4p1
In5p1
Tl6p1
Si3p2
Ge4p2
Sn5p2
Pb6p2
P3p3
As4p3
Sb5p3
Bi6p3
S3p4
Se4p4
Te5p4
Po6p4
Cl3p5
Be4p5
I5p5
At6p5
Ar3p6
Kr4p6
Xe5p6
Rn6p6
Y4d1
La5d1
Ac6d1
Cd4d10
Hg5d10
Ag5s14d10
Au6s15d10
Zr4d2
Hf5d2
Rf6d2
Nb4d3
Ta5d3
Db6d3
Mo5s14d5
W6s15d5
Sg7s16d5
Tc4d5
Re5d5
Bh6d5
Ru4d6
Os5d6
Hs6d6
Rh4d7
Ir5d7
Mt6d7
Ni4d8
Ni5d8
Periodic Table Organization------ Groups or Families
Vertical columns in the periodic table are known as groups or families The elements in a group have similar electron configurations
Periodic Table Organization ---- Periods
Horizontal Rows in the periodic table are known as Periods The Elements in a period undergo a gradual change in properties as one proceeds from left to right
Periodic PropertiesElements show gradual changes in certain physical properties as one moves across a period or down a group in the periodic table. These properties repeat after certain intervals. In other words they are PERIODIC
Periodic properties include:
-- Ionization Energy-- Electronegativity-- Electron Affinity-- Atomic Radius-- Ionic Radius
Ionization energy increases across a period because the positive charge increases.Metals lose electrons more easily than nonmetals.Nonmetals lose electrons with difficulty (they like to GAIN electrons).
Ionization energy is the energy required toRemove an electron from an atom
Trends in Ionization Energy
The ionization energy increases UP a group
Because size increases due to an effect known as the Shielding Effect
Trends in Ionization Energy
Ionization Energies
Ionization Energies are Periodic
Electronegativity Electronegativity
is a measure of the ability of an atom in a molecule to attract electrons to itself.
This concept was first proposed by Linus Pauling (1901-1994). He later won the Nobel Prize for his efforts
This concept was first proposed by Linus Pauling (1901-1994). He later won the Nobel Prize for his efforts
Periodic Trends: Electronegativity
In a group: Atoms with fewer energy levels can attract electrons better (less shielding). So, electronegativity increases UP a group of elements.
In a period: More protons, while the energy levels are the same, means atoms can better attract electrons. So, electronegativity increases RIGHT in a period of elements.
Trends in Electronegativity
Electronegativity
Electronegativity
Electron Affinities
Electron Affinities Are Periodic
Electron Affinity v Atomic Number
The Electron Shielding EffectElectrons between the nucleus and the valence electrons repel each other making the atom larger.
The radius increases on going down a group.
Because electrons are added further from the nucleus, there is less attraction. This is due to additional energy levels and the shielding effect. Each additional energy level “shields” the electrons from being pulled in toward the nucleus.
The radius decreases on going across a period.
The radius increases on going down a group.
Because electrons are added further from the nucleus, there is less attraction. This is due to additional energy levels and the shielding effect. Each additional energy level “shields” the electrons from being pulled in toward the nucleus.
The radius decreases on going across a period.
Atomic Radius
Atomic RadiusAtomic Radius
The radius decreases across a period owing to increase in the positive charge from the protons. Each added electron feels a greater and greater + charge because the protons are pulling in the same direction, where the electrons are scattered.
Large Small
Atomic Radius
Atomic Radius
Trends in Ion SizesTrends in Ion SizesRadius in pm
Cations
Cations (positive ions) are smaller than their corresponding atoms
Does the size go up or down when gaining an electron to form an anion?
Does the size go up or down when gaining an electron to form an anion?
F,64 pm9e and 9p
F- , 136 pm10 e and 9 p
-
Ion Sizes
CATIONS are SMALLER than the atoms from which they come.The electron/proton attraction has gone UP and so the radius DECREASES.
Li,152 pm3e and 3p
Li+, 78 pm2e and 3 p
+
Ionic RadiusForming a cation.
Ionic Radius for CationsPositve ions or cations are smaller than the corresponding atoms.
Cations like atoms increase as one moves from top to bottom in a group.
Anions
Anions (negative ions) are larger than their corresponding atoms
Ionic Radius-AnionsIonic Radius-Anions
ANIONS are LARGER than the atoms from which they come.The electron/proton attraction has gone DOWN and so size INCREASES.Trends in ion sizes are the same as atom sizes.
Forming an anion.Forming an anion.
F 64 pm9e- and 9p+
F-, 133 pm10 e- and 9 p+
-
Ionic Radii for Anions
Negative ions or anions are larger than the corresponding atoms.
Anions like atoms increase as one moves from top to bottom in a group.
Ionic Radius for an Isoelectronic Group
Isoelectronic ions have the same number of electrons.
The more negative an ion is the larger it is and vice versa.
Summary of Periodic Trends
Properties of the Third Period Oxides
Properties of the Third Period Chlorides
The D Block Elements
The d block elements fall between the s block and the p block.
They share common characteristics since the orbitals of d sublevel of the atom are being filled.
The D Block ElementsThe D block elements include the transition metals. The transition metals are those d block elements with a partially filled d sublevel in one of its oxidation states.
Since the s and d sublevels are very close in energy, the d block elements show certain special characteristics including:
1. Multiple oxidation states
2. The ability to form complex ions
3. Colored compounds
4. Catalytic behavior
5. Magnetic properties
The D Block ElementsThe d electrons are close in energy to the s electrons. D block elements may lose 1 or more d electrons as well as s electrons. Hence they often have multiple oxidation states
Some common D block oxidation states
Multiple Oxidation StatesThere is no sudden sharp increase in ionization energy as one proceed through the d electrons as there would be with the s block. D block elements can lose or share d electrons as well as s electrons, allowing for multiple oxidation states.Most d Block elements have a +2 oxidation State which corresponds to the loss of the two s electrons. This is especially true on the right side of the d block, but less true on the left.
---- For example Sc+2 does not exist, and Ti+2 is unstable, oxidizing in the presence of any water to the +4 state.
Complex IonsThe ions of the d block and the lower p block have unfilled d or p orbitals.
These orbitals can accept electrons either an ion or polar molecule, to form a dative bond. This attraction results in the formation of a complex ion.
A complex ion is made up of two or more ions or polar molecules joined together.
The molecules or ions that surround the metal ion donating the electrons to form the complex ion are called ligands.
Complex IonsCompounds that are formed with complex ions are called coordination compoundsCommon ligands
Complex ions usually have either 4 or 6 ligands.
K3Fe(CN)6 Cu(NH3)42+
Complex IonsThe formation of complex ions stabilizes the oxidations states of the metal ion and they also affect the solubility of the complex ion.
The formation of a
complex ion often has
a major effect on the
color of the solution of
a metal ion.
The D Block Colored Compounds
In an isolated atom all of the d sublevel electrons have the same energy.
When an atom is surrounded by charged ions or polar molecules, the electric field from these ions or molecules has a unequal effect on the energies of the various d orbitals and d electrons.
The colors of the ions and complex ions of d block elements depends on a variety of factors including:
– The particular element– The oxidation state– The kind of ligands bound to the element
Various oxidation states of Nickel (II)
Colors in the D BlockThe presence of a partially filled d sublevels in a transition element results in colored compounds. Elements with completely full or completely empty subshells are colorless, – For example Zinc which has a full d subshell. Its
compounds are whiteA transition metal ion is colored, if it absorbs light in the visible range (400-700
nanometers).If the compound absorbs a
particular wavelength of light its color will be the composite of those wavelengths that it does not absorb.
In other words it shows its complimentary color.
Colors and d Electron TransitionsWhen ligands are attached to transition metal ions, the d orbitals may split into two groups. Some of the orbitals are at a lower energy than the othersThe difference in energy of these orbitals varies slightly with the nature of the ligand or ion surrounding the metal ion
The energy of the transition: ∆E =hn may occur in the visible region. When white light passes through a compound of a transition metal, light of a particular frequency is absorbed as an electron is promoted from a lower energy d orbital to a higher one. The result is a colored compound
Magnetic PropertiesParamagnetism --- Molecules with one or more unpaired electrons are attracted to a magnetic field. The more unpaired electrons in the molecule the stronger the attraction. This type of behavior is calledDiamagnetism --- Substances with no unpaired electrons are weakly repelled by a magnetic field. Transition metal complexes with unpaired electrons exhibit simple paramagnetism. The degree of paramagnetism depends on the number of unpaired electrons
Catalytic BehaviorMany D block elements are catalysts for various chemical reactions
Catalysts speed up the rate of a reaction with out being consumed.
The transition metals form complex ions with ligands that can donate lone pairs of electrons.
This results in close contact between the metal ion and the ligand.
Transition metals also have a wide variety of oxidation states so they gain and lose electrons in oxidation- reduction reactions
Some Common D Block Catalysts
Examples of D block elements that are used as catalysts:
1. Platnium or rhodium in a catalytic converter
2. MnO2 decomposition of hydrogen peroxide
3. V2O5 in the contact process
4. Fe in Haber process 5. Ni in conversion of
alkenes to alkanes
The Periodic Table--Summary
The periodic table is a classification system. Although we are most familiar with the periodic table that Seaborg proposed more than 60 years ago, several alternate designs have been proposed.
Alternate Periodic Tables
Alternate Periodic Tables II
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