PF1009 2024 2 Atomic Structure PDF

Pharmaceutical Chemistry Atomic Structure Dr. J.J. Keating 1  Atomic Structure and Inorganic Compounds Atom – Smallest particle of an element that has the chemical properties of that element. Element – Substance that consists of atoms having th...

Pharmaceutical Chemistry Atomic Structure Dr. J.J. Keating 1  Atomic Structure and Inorganic Compounds Atom – Smallest particle of an element that has the chemical properties of that element. Element – Substance that consists of atoms having the same chemical properties. Elements with symbols taken from Latin and German names: Antimony Sb stibium Copper Cu cuprum Gold Au aurum Iron Fe ferrum Lead Pb plumbum Ag Mercury Hg hydrargyrum Potassium K kalium Silver Ag argentum Sodium Na natrium Tin Sn stannum Tungsten W wolfram Au 2 Periodic Table 3 Periodic Table 4  Periodic Table Groups – vertical columns. Periods – horizontal rows. Congeners – elements of the same group. Alkali metals – congeners of Group I (Na, Li, K, Rb, Cs). Alkaline earth metals –Group II (Be, Ca, Mg, Sr, Ba). Halogens – Group VII (F, Cl, Br, I). Noble gases – Group VIII (Kr, He, Ne, Xe, Rn). Transition metals – elements between Groups II and III. s-block – Group I and Group II columns p-block – Group III – Group VIII columns d-block – transition metals f-block – lanthanides and actinides 5 Minerals in Multivitamin Preparations and Baby Formula Ca P Mg Fe Zn Mn Cu I Se Na K Cl 6  Periodic Table Metal – substance that conducts electricity, has a metallic lustre, malleable and ductile. Non-metal – substance that does not conduct electricity and is neither malleable nor ductile. Metalloid – physical appearance and properties of a metal but behaves chemically like a non-metal (Si, Ge, As, Te). 7  Atom Electron Discovered by J.J. Thomson One negative charge Denoted by e– Charge = 1.6 x 10–19 C Mass = 9.1 x 10–28 g Proton Discovered by James Chadwick One positive charge Denoted by p 1836 times heavier than an electron Located in the nucleus 8  Atom Nucleus Proposed by Ernest Rutherford Positively charged Contains both protons and neutrons 9  Atomic Number Atomic Number Denoted by Z Number of protons in a nucleus Shown above chemical symbols Because an atom is electrically neutral, the number of protons in its nucleus must be the same as the number of electrons outside its nucleus. For Au, Z = 79: 79 protons and 79 electrons. Z is sometimes added as a subscript to the left of the chemical symbol. Mass number (A) = Z + N 10  Mass of an Atom / Isotope Mass of an atom Mass of an atom is measured using the technique of mass spectroscopy (MS). 1 H = 1.67 x 10–24 g 12 C = 1.99 x 10–23 g Isotope Atoms that have the same atomic number but different atomic masses. An atom with a fixed number of protons but different number of neutrons. Isotope = equal place (Greek) – although the atoms have different masses, they belong to an element that occupies one place in the periodic table. H D T 11  Atom Isotopic abundance Percentage (in terms of the number of atoms) of that isotope present in a sample of the element. Natural abundance Abundance in a sample of naturally occurring material. Natural abundance of Neon-20 = 91% Natural abundance of Uranium-235 = 0.7% Name Z No. A Exact Isotopic Abundance, % Symbol neutrons Mass, u Hydrogen 1 0 1 1.008 99.985 1 H Deuterium 1 1 2 2.014 0.015 2 H or D Tritium 1 2 3 3.016 Radioactive 3 H or T Carbon-12 6 6 12 12 98.90 12 C Carbon-13 6 7 13 13.003 1.10 13 C 12  Mass Number Mass number Denoted by A. Total number of protons and neutrons in the nucleus of an atom. If the mass number of an isotope and Z are known, then the number of neutrons can be determined. Isotope is named by writing its mass number after the name of the element: chlorine- 35, chlorine-37. A is written as a superscript to the left of the chemical symbol. 13  Average Atomic Mass (amu) Average atomic mass (atomic weight) Relative atomic mass of an element taking into account the natural abundances of the isotopes of the element. To calculate the average atomic mass of an element, you need the following information: the exact atomic mass for each naturally occurring stable isotope of the element and its percent abundance. To calculate the average atomic mass, each exact mass is multiplied by its percent abundance (expressed as a decimal). Then, add the results together and round off to an appropriate number of significant figures. Carbon Isotopic Abundance, % mass number Exact (A) Mass, u 12 12.000000 98.90 13 13.003355 1.10 For carbon, (12.000000)(0.9890) + (13.003355)(0.0110) = 12.011 amu 14  Relative Atomic Mass Relative atomic mass One atomic mass unit (1 u) = 1/12 the mass of an atom of carbon-12. Thus, the mass of one atom of carbon-12 = 12 u exactly. Mass of a single carbon-12 atom = 1.9926 x 10–23 g. As 1 u = 1/12 that mass, thus 1 u = (1.99 x 10–23 g) / 12 = 1.6605 x 10–24 g Name Z No. neutrons A Mass, u Hydrogen 1 0 1 1.008 Deuterium 1 1 2 2.014 The isotope against all other Tritium 1 2 3 3.016 atomic masses Carbon-12 6 6 12 12.0000 of isotopes are compared to. Carbon-13 6 7 13 13.003 15  Mole NA = 6.02214076(12) x 1023 mol–1 Avogadro’s Number 16  Mole 1 mole of atoms of any element = 6.022 x 1023 atoms = Avogadro’s number, NA. Number of atoms in 12 g of carbon-12 = 1 mole. SI unit – mol (not M = molar (concentration)) Molar mass Mass per mole of atoms of an element. Average mass per atom x number of atoms per mole The value for molar mass (in g) per mole is numerically identical to that for the average atomic mass in atomic mass units (amu). 1 mol Au = 196.97 g, 1 mol Cu = 63.54 g, 1 mol Hg = 200.59 g. mass of sample No. moles in a substance= mass of one mole of the substance 15 g Cl = (15 g)/(35.45 g/mol) = 0.423 mol Cl 17 Mole NA = 6.02214076(12) x 1023 mol–1 18