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Chapter 4. The Periodic table. Mendeleev’s Table. Mendeleev’s table grouped elements with similar properties into vertical columns called “groups” or “families” example the halogen group. Mendeleev found only 63 elements known at that time. There are 2 interesting things to note.
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Chapter 4 The Periodic table
Mendeleev’s Table • Mendeleev’s table grouped elements with similar properties into vertical columns called “groups” or “families” example the halogen group.
Mendeleev found only 63 elements known at that time. There are 2 interesting things to note. • The elements do not always fit neatly in the order of increasing atomic mass. • There were gaps in his table where elements with a particular atomic mass should occur. He was able to predict the properties of these undiscovered elements which he called ekaaluminium , ekaboron and ekasilicon. • Eka meaning “one beyond” in Sanskrit.
Moseley’s theory • Mosley an England scientist worked with x-ray spectra to study atomic structure which helped him determine atomic numbers for chemical elements. He discovered isotopes, explaining how atomic mass did not order the elements appropriately. Moseley was the catalyst for the periodic law, a rule stating that the properties of elements are a periodic function of their atomic numbers. Using the atomic numbers, He was able to more accurately position the elements in the periodic table. Moseley's change to the periodic table allowed the few problems with Mendeleev's periodic table to disappear. The modern periodic table is now based on atomic numbers.
. • We now follow the periodic table recommended by IUPAC. • The main group elements: • Groups 1,2, and 13 through 18 are referred to as the main group elements. • Group 1 the alkali metals. These metals react with water to form alkaline solution. They are excellent conductors of electricity.
Group 2- The alkaline earth metals. These are harder than alkali metals. They are less reactive than alkali metals. Mg from this group when reacts with oxygen from the air forms magnesium oxide which slows corrosion of the magnesium metals underneath. Forms good alloys.
Lanthanides and actinides: the lanthanides have atomic numbers from 58 to 71.The name originated from the element 57. they are shiny, reactive metals that have irregular electronic configuration. • Actinides are unique in that their nuclear structures are of more importance than their electron configurations. They have an unstable arrangement of protons and neutrons in the nucleus so they are called radioactive. Example Uranium. Nuclear disintegration of the uranium nucleus releases sufficient energy to run power parts,submarines also aircrafts.
Group 17-Halogens Derived from the Greek word meaning “salt former”. Most reactive group of nonmetals. • Group 18- Noble gases. Formerly called inert gases because they were thought to be completely unreactive. Except for helium these atoms are characterized by an octet of electrons, ns2 and np6 in the outermost energy level.
Homework • Page 123 # 7,9,11,13
Metals are good conductors of heat and electricity. Metals in the form of crystals . • A crystal is a substance in which atoms or molecules are arranged in an orderly geometric fashion.
Crystal and conduction band • The outer electrons in the atoms of each element form the bonds that bind the atoms together in a crystal. In a crystal, each atom is bonded to all neighboring atoms. These electrons exist in orbitals just as electrons do in atoms. There are vacant orbitals in the bond just as in atoms. Electrons are free to move in three dimensions through the crystal, forming the conduction band.
Conduction band is a band within which electrons must move to allow electrical conduction. In metals the orbitals of the conduction band actually overlap and are continous with the bond orbitals so even at room temp. there are many electrons in the conduction band.
The metal atoms are vibrating back and forth with thermal motion, interfering with the motion of electrons. The vibration increases as the temperature increases so the conductivity decreases.
In semiconductors the energy of the higher orbitals is not very high compared with the energy of the filled bonding orbitals.less energy is required to excite electrons to these orbitals, so even at room temperature they carry a certain amount of current. At higher temperature there are more electrons excited in the conducting band, so conductivity of semiconductors increases as the temperature increases.
In typical non metals these orbitals have very high energy so it is practically impossible to excite electrons to the orbitals of the conduction band example diamond is a nonconductor.
Trends in periodic table • Trend is a periodic change in a particular direction. • Periodic trends in atomic radii: The size of an atom is determined by the volume occupied by the electrons surrounding the nucleus. Radii are usually determined for atoms that are chemically bonded or close together in the solid state, as in Bromine and Iodine.
Bond radii: Half the distance between the nuclei of the atoms in each molecule is the bond radius. • The distance between the nuclei in adjacent non bonded molecules, which is twice the distance, called the van der Waals radius.
Van der waal’s radii are not very precise because of the fuzziness of the atoms. Bond radius is usually considered to dfind the size of the atom.
Atomic radius increases as you move down a group. This is caused by the addition of another principal energy level from one period to the next.
Electron shielding: The electrons in the inner energy levels are between the nucleus and the outermost electrons and therefore shield the outer electrons from the full charge of the nucleus
Atomic radius decreases as you move across a period. • As you move from left to right across a period, each atom has one more proton and one more electron than the atom before it has. All additional electrons go into the same principal energy level—no electrons are being added to the inner levels. • Electron shielding does not play a role as you move across a period. • As the nuclear charge increases across a period, theeffective nuclear charge acting on the outer electronsalso increases.
Ionization Energy: The energy used to remove the electron is the ionization energy of the atom. • Each element has more occupied energy levels than the one above it has. • The outermost electrons are farthest from the nucleus in elements near the bottom of a group. • As you move down a group, each successive element contains more electrons in the energy levels between the nucleus and the outermost electrons.
Ionization energy tends increases as you move from left to right across a period. • From one element to the next in a period, the number of protons and the number of electrons increase by one each. • The additional proton increases the nuclear charge. • A higher nuclear charge more strongly attracts the outer electrons in the same energy level, but the electron-shielding effect from inner-level electrons remains the same.
Electron affinity: The ability of an atom to attract and hold an electron is called electron affinity. Although the atom is neutral the electrons in the orbitals generally do not shield the nuclear charge a 100%. The approaching electron may experience a net pull because the effective nuclear charge is greater than zero. The electron enters a vacant orbital.
For example, when a neutral chlorine atom in the gaseous form picks up an electron to form a Cl- ion, it releases an energy of 349 kJ/mol or 3.6 eV/atom. It is said to have an electron affinity of -349 kJ/mol and indicates that it forms a stable negative ion
Electron affinity becomes more negative across a period and tends to decrease from top to bottom. • Electron affinity trends within groups of elements are not as regular as trends for ionization energy. Alkaline earth metals are an exception. An electron added in the alkaline earth metals must go to the p orbital and is shielded to a greater extent by the s electrons.
Periodic trends in melting and boiling points • As the number of electrons increases across the period the m.p and b.p. increases. This indicates stronger bonding.
However when the d orbitals become half filled the m.p and b.p decreases. This indicates that added electrons become less involved in bonding. As the p electrons are added we see a similar rise and fall with the maximum near the stage where p orbitals are half filled. A minimum is reached when the p orbitals are filled and the atoms no longer bond. The noble gases are monatomic and have no chemical bonding forces between atoms. Hence their m.p. and b.p. are unusually low.
Homework • Page 151 Term review all • Page 155 test prep all.
