CHAPTER 3 ATOMIC AND ELECTRONIC STRUCTURE Name: Prachayanee Chueamsuwanna date: October 6,2015

ATOMIC THEORY OF MATTER DEFINING AN ATOM  An atom is the smallest part of an element that retains its identity in a chemical reaction.

STRUCTURE OF THE ATOM: DISCOVERY OF SUBATOMIC PARTICLES  ELECTRONS : particles that have a negatively charged, and was discover by J.J. Thomson in 1890  PROTONS : particles that have a positively charged, and has a mass of x g. and was discover by Rutherford in  NEUTRON : has neutral charge which has a mass slightly bigger than the proton, but without charge x g.

MODEL OF ATOMIC STRUCTURE

1.DALTON’S MODEL  John Dalton proposed that all matter is composed of very small things which he called atoms. This was not completely new concept as the ancient Greeks had proposed that all matter is composed of small, indivisible objects. When Dalton proposed his model electrons and the nucleus were unknown.

2. THOMSON’S MODEL THE PLUM PUDDING MODEL  Initially it was supposed that electrons and positively charged particles were evenly distributed in a spherical atom. This prevailing theory was that of the “plum pudding” model, put forward by Thomson.

3. RUTHERFORD’S ATOMIC MODEL : RADIOACTIVITY  The discovery of radioactivity lead to discovery protons as well as the atomic model that the atoms consisted of the nucleus and which housed protons and that electrons surrounded the nucleus.  Radioactivity is the spontaneous emission of radiation by an atom.

4. BOHR’S MODEL  A new model was proposed, by Niels Bohr( ) in 1913, which says that an electron is found only in specific circular paths, or orbits around the nucleus.  To move from one energy level to the other, an electron must gain or lose just right amount of energy. A quantum of energy is the amount of energy needed to move an electron from one energy level to another energy level.

5. ERWIN SCHORDINGER : QUANTUM MECHANICAL MODEL  The quantum mechanical determined the allowed energies and electron can have and how likely it is to find the electron in various locations around the nucleus of an atom.  The Bohr was limited. Erwin Schodinger in 1926 developed a mathematical treatment into which both the wave and particle nature of matter could be incorporated. This is known as quantum mechanics.  Solving the wave equation gives a set of wave functions, or orbitals and their corresponding energies. Each orbitals describes a spatial distribution of electron density. An orbitals is described by a set of three quantum numbers.

ATOMIC ORBITALS  Solutions to Schordinger’s equations give the energies, or energy levels, an electron can have. For each energy level, the Schordinger’s equation also leads to a mathematical expression called an atomic orbital which describes the probability of finding an electron at various locations around the nucleus of. An atomic orbitals is represented pictorially as a region of space in which there is a high probably of finding an electron.  Each sublevel corresponds to one or more orbitals of different shapes. The orbitals describe where an electron is likely to be found.

S ORBITALS  Are sphere shaped and are concentric and bigger as the main shell number increases.  The value of l for S orbitals is 0.  They are spherical in shape.  The radius of the sphere increases with the value of n.

P ORBITALS  Two dumbbells on opposite sides of the nucleus, orientated along respective three axis of 3 space. Found in the principal shell 2 and up, not in 1.  The value of l for P orbitals is 1.  They have two lobes with a node between them.

D ORBITALS AND OTHER HIGHER ENERGY ORBITALS  Are 4 lobed(except one). These occur from the principal shell 3 and up. F orbitals are more complex( 8 lobed) and g…  The value of l for a d orbitals is 2.  Four of the five d orbitals have 4 lobes; the other resembles a p orbital with a doughnut around the center.

ELECTRONIC CONFIGURATIONS  There are three rules that tells the position of an electron namely, the Aufbau Principle, the Pauli Exclusion Principle and the Hand’s Rule.

THE AUFBAU PRINCIPLE (BUILDING UP PRINCIPLE)  According to principle, electrons occupy the orbitals of lowest energy first. It dictates that for every further proton in the nucleus, there is an electron in an orbital of that atom. This principle also dictates the chemical and physical properties of an element, and its position in the periodic table.

PAULI EXCLUSION PRINCIPLE  No two electrons in the same atom may have the same series of quantum numbers. An atomic orbital describes at most two electrons.  Electron Configurations - The electron configuration is a way of how electrons are arranged in various orbitals around the nuclei of an atom each component, e.g 4p 5, consists of: - A number denoting the energy level, (4) - A letter denoting the type of orbital, (p) - A superscript denoting the number of electron in those orbitals ( 5 )

HUND’S RULE  States that in subshells, as far as possible, electrons will half fill orbitals with parallel spins.  The electron configuration for nitrogen therefore is 1s 2 2s 2 2p 3 with the p orbitals with one electron each, spinning in the same direction.

ATOMIC NUMBER, MASS NUMBER, & ISOTOPE  Atomic number : which equals its number of electrons if it is uncharged. (number of protons or electrons)  Mass number : is the sum of the protons and neutrons in an atom.(number of protons plus neutrons)  Isotope : same number of protons but different number of neutrons or different mass number.

ATOMIC WEIGHT  Atomic and molecular can be measured with great accuracy using a mass spectrometer. Because in the real world we use large amounts of atoms and molecules, we use average masses in calculations. Average mass is calculated from the isotopes of an element weighted by their relative abundances.