Models Of The Atom PDF

MODELS OF THE ATOM  SECOND QUARTER 00 ENABLING COMPETENCY Describe how Bohr’s Model of the atom improved Rutherford’s Atomic Model. 06 MOST ESSENTIAL LEARNING COMPETENCY Explain how the Quantum Mechanical Model of the atom describes the energies and positions o...

MODELS OF THE ATOM  SECOND QUARTER 00 ENABLING COMPETENCY Describe how Bohr’s Model of the atom improved Rutherford’s Atomic Model. 06 MOST ESSENTIAL LEARNING COMPETENCY Explain how the Quantum Mechanical Model of the atom describes the energies and positions of the electrons  MODELS OF THE ATOM What is an atom? Atoms are the building blocks of matter. “All matter, be it solid, liquid or gas or any other form is made up of atoms of different elements.”  Element pertains to specific types of atoms. Examples of elements are: gold, oxygen, copper, sulfur, carbon and hydrogen Sub-atomic particles are particles that are much smaller than the atom. 1. Proton – has positive charge; massive. 2. Electron – has negative charge; light. 3. Neutron – has no charge; same mass as proton. Atomic number – the number of protons in an atom. “A neutral atom has an equal number of protons and electrons so that the positive and negative charges exactly balance.” How does an atom look like? QUANTUM MECHANICAL MODEL Erwin Schrodinger 1 PLANETARY MODEL PLUM PUDDING MODEL Neils Bohr 4 2 Joseph John Thomson 3 NUCLEAR MODEL Ernest Rutherford  NUCLEAR MODEL by Ernest Rutherford It describes the atom as having: 1. A massive positive core called nucleus at the center of the atom. QUANTUM MECHANICAL MODEL Erwin Schrodinger 1 PLANETARY MODEL PLUM PUDDING MODEL Neils Bohr 4 2 Joseph John Thomson 3 NUCLEAR MODEL Ernest Rutherford  NUCLEAR MODEL by Ernest Rutherford It describes the atom as having: 1. A massive positive core called nucleus at the center of the atom. 2. Electrons that occupy most of the volume of the atom and move around the nucleus in an arbitrary distance. QUANTUM MECHANICAL MODEL Erwin Schrodinger 1 PLANETARY MODEL PLUM PUDDING MODEL Neils Bohr 4 2 Joseph John Thomson 3 NUCLEAR MODEL Ernest Rutherford PLANETARY MODEL by Neils Bohr He asserted that: 1. Electrons move around the nucleus in fixed circular orbits or energy levels. 2. These energy levels are found at definite distances from the nucleus. 3. Energy of electrons in each orbit is quantized /fixed. 4. No absorption or emission of energy if electrons stay in its orbit. PLANETARY MODEL by Neils Bohr When an atom absorbs extra energy (heat/electric current), its electrons moves from lower energy level to a higher energy level. At this point the atom is at its excited state. PLANETARY MODEL by Neils Bohr Once excited, the atom is unstable. The same electron will return to the lower energy levels releasing the same amount of energy it absorb in the form of infrared radiation and visible light. QUANTUM MECHANICAL MODEL Erwin Schrodinger 1 PLANETARY MODEL PLUM PUDDING MODEL Neils Bohr 4 2 Joseph John Thomson 3 NUCLEAR MODEL Ernest Rutherford QUANTUM MECHANICAL MODEL by Erwin Schrodinger This model… 1. views an electron as a cloud of negative charge having certain geometrical shapes. 2. shows how likely an electron could be found in various locations around the nucleus. QUANTUM MECHANICAL MODEL by Erwin Schrodinger The volume or region of space around the nucleus where the electron is most likely to be found is called an atomic orbital.  QUANTUM MECHANICAL MODEL by Erwin Schrodinger Basis for this model: 1. Louis De Broglie – electron is both a wave and a particle. 2. Werner Heisenberg – electrons cannot be located exactly 3. Erwin Schrodinger – developed a mathematical equation to find the probable location of an electron. “… no single image has been entirely satisfactory at visualizing the atom’s various characteristics.” (britannica.com)  ELECTRONIC STRUCTURE OF THE ATOM  SECOND QUARTER 00 ENABLING COMPETENCY Describe how Bohr’s Model of the atom improved Rutherford’s Atomic Model. QUANTUM MECHANICAL MODEL – views electron as a cloud of negative charge. Orbital – region in space where electrons are most likely to be found.  Specifically, the Quantum Mechanical Model states the following: 1. Electrons that surround the nucleus are confined to regions called principal energy levels or shells. Principal Energy Level or Shell – region of space around a nucleus containing electrons having approximately the same energy. Specifically, the Quantum Mechanical Model states the following: 1.1 Shells are numbered, n = 1,23,4,5,6,7 or named as K, L, M, N, O, P, Q Specifically, the Quantum Mechanical Model states the following: 1.2 Electrons closer to the nucleus are held tightly and are lower in energy. 1.3 Electrons farther from the nucleus are held less tightly and are higher in energy. Specifically, the Quantum Mechanical Model states the following: 1.4 It takes an energy to move an electron away from the nucleus to an outer circle. Specifically, the Quantum Mechanical Model states the following: 1.5 The farther the shell from the nucleus, the more electron is can hold. Specifically, the Quantum Mechanical Model states the following: AUFBAU PRINCIPLE - electrons fill atomic orbitals of the lowest energy levels before occupying higher levels.  Specifically, the Quantum Mechanical Model states the following: 2. Shells are divided into subshells or sublevels, identified as s, p, d, f. Sublevel or subshell – region of space within an electron shell that contains electrons that have the same energy. Specifically, the Quantum Mechanical Model states the following: 3. The number of sublevels/subshells in each principal energy level or shell is equal to the number of that energy level.  Specifically, the Quantum Mechanical Model states the following: 4. These subshells consist of orbitals. Orbital – region of space where the probability of finding an electron is high.  Specifically, the Quantum Mechanical Model states the following: 4.1 Each subshell contains a specific number of orbital. 4.2 Each orbital can hold two electrons  Specifically, the Quantum Mechanical Model states the following: 4.1 Each subshell contains a specific number of orbital. 4.2 Each orbital can hold two electrons Shell/Energy Sublevels/ Number of Electrons Level Subshells 1 s 2 2 s,p 2+6=8 3 s,p,d 2 + 6 + 10 = 18 4 s,p,d,f 2 + 6 + 10 + 14 =32 5 s,p,d,f 32 ELECTRONIC STRUCTURE OF THE ATOM To completely describe an electron in an atom, quantum numbers are needed which correspond to orbital size, shape, orientation and spin.  The four quantum numbers are: 1. Principal Quantum Number (n) - describes the size and energy of the orbital and relative distance from the nucleus. The possible values of n are positive integers. The smaller the value of n, the lower the energy, and the closer to the orbital is to the nucleus.  The four quantum numbers are: 2. Angular Momentum Quantum number (l) - describes the shape of the orbitals. Its value is related to the principal quantum number and has allowed value of 0 to (n-1). For example, if n = 4, then the possible values of l would be 0, 1, 2, and 3 (= 4-1).  The four quantum numbers are: Things to remember: If l = 0, then the orbital is called an s-orbital and has a spherical shape with the nucleus at the center of the sphere. The greater the value of n, the larger is the sphere. If l = 1, then the orbital is called a p-orbital with two lobes of high electron density on either side of the nucleus, for an hourglass or dumbbell shape. If l = 2, then the orbital is a d-orbital with a variety of shapes. If l = 3, then the orbital is an f-orbital with more complex shapes  The four quantum numbers are: 3. Magnetic Quantum Number (ml) - describes the orientation of the orbital around the nucleus. The allowed values for ml are -l though 0 to +l. For example, for l = 3, the possible values of ml would be -3, -2, -1, 0, +1, +2, +3. For example, if l = 1 (a p-orbital), there are three p- orbitals (sublevels) corresponding to ml values of -1, 0, +1.  The four quantum numbers are: 4. Spin Quantum Number (ms) – indicates the direction the electron is spinning. There are only two possible values for ms: +1/2 and - 1/2. When two electrons are to occupy the same orbital, then one must have an ms = +1/2 and the other electron must have an ms = -1/2.  Let’s try: 1. If n = 3, what are the possible values of l ? 1. If n = 3 and l = 2, then what are the possible values of ml?  Let’s try: 3. List all the possible combinations of all four quantum numbers when n = 3, l = 2, and ml = 0. ELECTRONIC STRUCTURE OF THE ATOM In order to track where all the electrons are, chemists use notation called electron electron configuration. configuration Example: Neon – 1s 2s 2p6 2 2  Steps in writing the electron configuration of elements in neutral state: 1. Determine the element’s atomic number. Atomic number pertains to the number of protons in an atom. (In a neutral atom, the number of protons is equivalent to the number of electrons.)  Steps in writing the electron configuration of elements in neutral state: 2. Apply the rules in deriving the electron configuration: a. Aufbau Principle - electron's occupy orbitals in order of increasing energy. b. Pauli’s Exclusion Principle - electrons occupying the same orbital must have opposite spin ↑↓. c. Hund’s Rule - when electrons enter a sublevel with more than one orbital, they will spread out to the available orbitals with the same spin before pairing.  Steps in writing the electron configuration of elements in neutral state: a. Aufbau Principle - electron's occupy orbitals in order of increasing energy.  Steps in writing the electron configuration of elements in neutral state: a. Aufbau Principle - electron's occupy orbitals in order of increasing energy.  Steps in writing the electron configuration of elements in neutral state: a. Aufbau Principle - electron's occupy orbitals in order of increasing energy. Example: Lithium  Steps in writing the electron configuration of elements in neutral state: a. Aufbau Principle - electron's occupy orbitals in order of increasing energy. Example: Silicon  Steps in writing the electron configuration of elements in neutral state: a. Aufbau Principle - electron's occupy orbitals in order of increasing energy. Example: Calcium-  Steps in writing the electron configuration of elements in neutral state: a. Aufbau Principle - electron's occupy orbitals in order of increasing energy. Example: Manganese-  Steps in writing the electron configuration of elements in neutral state: b. Pauli’s Exclusion Principle - electrons occupying the same orbital must have opposite spin ↑↓.  Steps in writing the electron configuration of elements in neutral state: b. Pauli’s Exclusion Principle - electrons occupying the same orbital must have opposite spin ↑↓. Example: Lithium – 1s22s1  Steps in writing the electron configuration of elements in neutral state: c. Hund’s Rule - when electrons enter a sublevel with more than one orbital, they will spread out to the available orbitals with the same spin before pairing. Let’s Practice! Write the electron configuration of the elements. Elements Electron Configuration Be4 C6 Mg12 S16 Ti22 Elements Electron Configuration Be4 1s22s2 C6 1s22s22p2 Mg12 1s22s22p63s2 S16 1s22s22p63s23p4 Ti22 1s22s22p63s23p64s23d2  THE END. Thank you! CREDITS: This presentation template was created by Slidesgo, including icons by Flaticon, and infographics & images by Freepik. Please keep this slide for attribution