For example, a number of compounds have a different colour in acid solutions than they have in basic solutions. These compounds are used in acid-base titration's as indicators, because they change colour during the course of the titration, to give a visual indication of the end-point of the titration.
Indicator Colour in Acid Colour in Base Litmus Red Blue Methyl Orange Red Yellow Methyl Red Red Yellow Phenolphthalene Colourless Red
These two series make up the f-block elements in the periodic table, and their chemical properties of the elements derive from the filling of the f atomic sub-orbitals. The electronic configuration of these elements are characterised as having full outer orbitals and full second outermost orbitals, while the second outermost orbitals are incompletely filled. Thus, in the case of the first inner transition metals series, the electronic configuration of the outermost and second outermost orbitals is 4s2 3d10, while the third outermost orbitals (i.e. the 4f level) are incompletely filled.
The present body of knowledge represented by chemistry was built on a development of the experimental methodologies first used by the alchemists. Although many of their theories were incorrect, and their interest in the transformation of base metals into gold, and the preparation of magic potions has no place in modern chemistry, the experimental approach which they adopted allowed later experimenters to deduce of the structure of matter
In 1828 Wohler was the first chemist to demonstrate the conversion of the inorganic compound (i.e. ammonium cyanate, NH4NCO) into an organic compound (i.e. Urea, NH2.CO.NH2). This experiment established that organic chemistry was just another branch of chemistry.
Organic compounds (e.g. wine, beer, dyes, perfumes, drugs, \ poisons, soaps, etc.) have been manufactured and used since early times. The first step in the development of organic chemistry was the preparation and description of individual organic compounds. Thus, distillation and other separation techniques were a major part of the early organic chemistry.
Inorganic Chemistry is the chemistry of all other substances. However, it is convenient to classify some common carbon compounds (e.g. carbon dioxide, carbon monoxide, sodium carbonates, etc.) as inorganic, so as to place carbon in its correct relationship to the other elements in the periodic system of classification. Inorganic chemistry is essentially the chemistry described by the periodic table.
By definition, carbon is always present in all organic compounds, and the vast majority of organic compounds also contain hydrogen. Compounds consisting of carbon and hydrogen only are called hydrocarbons. Many organic compounds contain oxygen, and some contain nitrogen, a halogens, sulphur, silicon and a few other elements. The classification of organic compounds will be illustrated for each functional group in this hypertext.
There are more than seven million chemical compounds known and more than three-quarters of them are carbon compounds. This is a remarkable proportion, particularly as very few other elements are included in these compounds. The main reason for the very large number of carbon compounds is the unique ability of the Carbon atom to combine with itself almost indefinitely, to form straight-chain or branched-chains or ring-systems, of all sizes and degrees of complexity.
Just as inorganic chemistry considers each element with regard to its position in the periodic table and then deals with the properties of the compounds of that element, organic chemistry is concerned with the properties of groups of organic compounds which have structural similarities (i.e. compounds which have certain functional groups in common).
An ionic bond results from the transfer of one or more electrons from the outer shell of one atom to the outer shell of another atom. This type of bond is usually formed between elements whose positions in the periodic table lie just before or just after the Nobel Gases.
For example, lithium fluoride, Li(+)F(-) is an ionic compound that is formed by the transfer of an electron from the lithium atom (thereby forming a positively charges lithium ion, Li(+)) to the fluorine atom (thereby forming a negatively charged fluoride ion, F(-)).
A lithium atom has two electrons in its inner shell, 1s(2), and one electron in its outer or valence shell, 2s(1). By the loss of one electron, the lithium atom can achieve the stable electron configuration of helium at the same time forming the lithium ion.
Li ==> Li(+) + e(-) 1s(2)2s(1) 1s(2) (2.1) (2)
Similarly, by gaining an electron the fluorine atom can achieve the neon configuration, thus forming the fluorine ion.
F + e(-) ==> F(-) 1s(2)2s(2)2p(5) 1s(2)2s(2)2p(6) (2.7) (2.8)
A metal atom converts its electronic configuration into that of an inert gas by the loss of one or more electrons. A non-metal converts its electronic configuration into that of an inert gas by gaining one or more electrons. The resulting particle in each instance is called an ion and carries a charge equal to the valency.
The lithium atom has 3 protons and 3 electrons, and therefore the atom has a neutral charge. However, the lithium ion has 3 protons and 2 electrons, and therefore the lithium ion has a positive charge.
The fluorine atom has 10 protons and 10 electrons, and therefore atom has a neutral charge. However, the fluorine ion has 10 protons and 11 electrons, and therefore the fluoride ion has a negative charge.
Since the lithium and fluorine ions have opposite charges, they attract one another and this constitutes the ionic bond between the ions.
The second, and third ionisation potentials, etc. relate to the removal of the second, and third electrons respectively, and are higher than the first ionisation potential, as each successive electron is removed against increasing positive charge. The ionisation potential is usually expressed in electron volts and is determined from spectra.
Since it is possible to remove more than one electron from some atoms, an atom can have more than one ionization potential (IP).
M ==> M(+) + e(-) (I1 = First Ionisation Potential ) M+ ==> M(++) + e(-) (I2=Second Ionisation Potential ) M++ ==> M(+++) + e(-) (I3=Third Ionisation Potential)
The magnitude of the ionisation potential values depend on various factors, including
As the atomic number increases, there is an increase in nuclear charge. However, this effect is counterbalanced by an increase in the distance of the outer electron from the nucleus and by the screening effect of the outer electron by the inner electron shells present in the atom. The values for the first ionisation potential show that it is relatively easy to remove electrons from these atoms to produce ions which have the electronic configuration of the atoms of helium, neon, argon, krypton and xenon.
The Nobel Gases have the highest ionisation potentials so that a large amount of energy is required to remove an electron from a Noble Gas, while the alkali metals have the lowest ionisation potentials, so that a smaller amount of energy is required to remove an electron from an alkali metal.
First Ionisation Potentials (eV) H 13.6 He 24.4 Li 5.4 Be, 9.3 B, 8.3 C, 11.3 N, 14.6 O, 13.6 F, 17.4 Ne, 21.6 Na, 5.1 Mg, 7.6 Al, 6.0 Si, 8.1 P, 11.0 S, 10.4 Cl, 13.0 Ar, 15.8 K, 4.3 Ca, 6.1 Rb, 4.1 Sr, 5.7 Cs, 3.9 Ba, 5.2
A further amount of energy, the second ionisation potential, is required to remove a second electron from the Group I elements. A comparison of the values for the first and second ionisation potentials along with the atomic number, the electron configuration of the group 1 elements is given below.
Electronic Atomic First Second Configuration Number IP(eV) IP(eV) Lithium 2.1 3 5.39 75.62 Sodium 2.8.1 11 5.14 47.29 Potassium 22.214.171.124 19 4.34 31.81 Rubidium 126.96.36.199.1 37 4.18 27.36 Caesium 188.8.131.52.8.1 55 3.89 23.40
It will be observed that there is a large increase between the first and the second ionisation potentials for each element. This is because the first ionisation potential involves removal of an electron from a complete octet.
Ionisation potentials are always positive, since energy must be provided to remove an electron against the attraction of the nucleus.
Isomers are substances which have the same molecular formula but which have different structural formulae. Isomers are distinct chemical compounds, which have different physical and chemical properties. Structural isomerism are very common in organic chemistry.
The isotopes of an element are the atom of different atomic mass for that element. All atoms of an element must have the same atomic number (i.e. the same number of protons in its nucleus) However, they may have different atomic mass (i.e. due to numbers of neutrons in the nucleus). The isotopes of an element have different physical properties. However, because they all have the same number of electrons, their chemical properties are identical.
Isotopes are separated by the principal described in Graham's Law which uses diffusion as the method of separation.
By using a simple mass spectrometer, F.W.Aston discovered that all atoms of a particular element do not have the same mass. In the Mass Spectrometer, atoms are given a charge and accelerated in an electric field so that they move at high speeds. When this beam of charged atoms passes through an electric field and a magnetic field, all particles of the same mass are focused on a line. The particles strike a photographic plate or detector so that their positions are recorded. These atoms of different mass number but having the same atomic number are called Isotopes.
The second isotope (called Deuterium) has a mass number of 2, and has one proton and one neutron in the nucleus.
The third isotope (called Tritium) has a mass number of 3, and has one proton and two neutrons in its nucleus.
The second isotope has a mass number of 37 which has 20 neutrons in the nucleus.
The observed atomic weight of chlorine is 35.46, which indicates that chlorine is a mixture of these two isotopes, and that the isotope with mass number 35 is the more abundant atom present. Indeed, the isotope of mass number 35 is approximately three times more abundant than the isotope of mass number 37.