Ch_02_summary - CHAPTER 2 ATOMIC STRUCTURE (IB TOPICS 2 AND...

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Unformatted text preview: CHAPTER 2 ATOMIC STRUCTURE (IB TOPICS 2 AND 12) SUMMARY Introduction • Relative masses: p = 1, n = 1, e = 1/1840; charges: p = +1, n = 0, e– = –1. • Protons and neutrons are present in the nucleus of an atom, electrons are in orbits or shells around the nucleus. • Atomic number, Z = number of protons; the fundamental characteristic of an element. • Mass number, A = number of (protons + neutrons). • Isotopes: same atomic number, different mass number OR same number of protons, different number of neutrons OR atoms of the same element with different masses. • Isotopes differ in physical properties that depend on mass such as density, rate of diffusion etc. Chemical properties are the same because of the same electronic configuration or arrangement. • Atomic mass of an element is the average of the atomic masses of its isotopes; depends on isotopes relative abundance; leads to non-integer atomic masses. Mass Spectrometer • Stages of Operation: Vaporization of sample, ionization to produce M+ ions, acceleration of ions by electric field, deflection of ions by magnetic field, vacuum, detection of ions. • Degree of deflection: • Lower the mass, higher the deflection. • Higher the charge, higher the deflection. • Deflection reflects mass/charge ratio; for charge of +1, deflection depends on mass. • For an element, the mass spectrum gives two important pieces of information: the number of isotopes, and the abundance of each isotope; thus the relative average atomic mass, Ar can be calculated. • For a molecule, the highest peak represents the molecular (parent) ion and its mass gives the relative molecular mass, Mr of the compound (and the fragmentation pattern can help determine its structure). o A continuous spectrum contains light of all wavelengths in the visible range. o A line spectrum consists of a few lines of different wavelengths. o When electrons are excited, they jump to higher energy levels. o Electrons fall back to lower energy levels, and the energy equivalent to the difference in energy level is emitted in the form of photons. o Energy levels come together in terms of energy the farther away they are from the nucleus; this explains the convergence of lines in a line spectrum. o The maximum number of electrons in a main energy level n is 2n2: 1st energy level, n = 1; maximum 2 e–; n = 2, maximum 8 e–; n = 3, maximum 18 e–. The electron arrangement (or configuration) indicates the number of electrons and their energy distribution. This determines an element’s physical and chemical properties. © IBID Press 2007 1 CHAPTER 2 ATOMIC STRUCTURE (IB TOPICS 2 AND 12) SUMMARY • Main (or principal) energy levels, sub-levels and orbitals: The main energy levels, n are assigned whole number integers, n = 1, 2, 3, 4… . n = 1 represents the lowest energy level. Each main energy level contains n sub-levels and a total of n2 orbitals. o s, p, d, f etc. is the common notation for sub-levels and orbitals within sub-levels. An orbital is an area of space around the nucleus in which an electron moves. o Orbitals have characteristic shapes. There is one s orbital which is spherical in shape, three p orbitals which are dumbbell shaped, called px, py pz, and arranged in the x, y, and z directions respectively, five d orbitals and seven f orbitals (both with complex shapes). The relative energies of s, p, d, and f orbitals with in a sub-level are: s < p < d < f. o Each orbital can have a maximum of 2 electrons. n = 1 has one sub-level which is called an s sub-level and which contains one s orbital. n = 2 has two sub-levels: 2s and 2p; n = 3 has 3 sub-levels: 3s, 3p and 3d; n = 4 has 4 sub-levels:4s, 4p, 4d and 4f, etc. • The Aufbau (‘building-up’) Principle: Electrons are placed in orbitals in order of increasing energy, starting with the lowest energy level, and in general, filling each sublevel completely before beginning the next. This is due to the fact that systems in nature prefer minimum energy in order to achieve maximum stability. • Hund’s Rule: Occupation of sub-levels takes place singly as far as possible before pairing starts. • Pauli exclusion principle: No two electrons in an atom can be in exactly the same state; no two electrons in a given atom can have the same four quantum numbers (that is, these can not be in the same place at the same time) • nlx notation is used to describe the electron configuration of an element: n is the main energy level, l the sub-level, and x is the number of electrons in the sub-level. • The ionisation energy of an atom is the minimum amount of energy required to remove a mole of electrons from a mole of gaseous atoms to form a mole of gaseous ions. (N.B. Shading indicates Topic 12 (AHL) material.) © IBID Press 2007 2 ...
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