2 Section 4.1 The Development of a New Atomic Model Objectives:Discuss the significance of the photoelectric effect and the line-emission spectrum of hydrogen to the development of the atomic model.Describe the Bohr model of the hydrogen atom.
3 Electromagnetic radiation = energy that exhibits wavelike behavior All electromagnetic radiation travels at the same speed. Different types are the result of different wavelengths and frequencies.This spectrum shows the different forms of electromagnetic radiation.
4 Light is a small part of the electromagnetic spectrum.
5 Absorption/Emission Spectra When light strikes a metal, the metal ejects e- from the surface and creates an electric current = Photoelectric effectRemember the cathode-ray tube?When the glass tube was filled with a pure gas and an electric current passed through, the gas will gave off light.Different gases give off different colors of light.If this light is passed through a prism, a series of bright lines is seen (emission spectrum).Every element has a distinct emission spectrum.
6 So, what’s happening to create the emission spectrum?CarbonOxygenIron
7 Rutherford’s model of the atom provided information about the structure of atoms, it did not explain where the electrons were located in the space surrounding the nucleus.
8 Max Planck Albert Einstein In 1900, Max Planck suggested that objects could give off energy in small, specific amounts he called quanta.A quantum (singular of quanta) is the minimum amount of energy that can be lost or gained by an atom.Albert EinsteinEinstein proposed that different elements require different frequencies of energy to eject electrons
9 So scientists agreed…Light is a form of energyDifferent colors of light have different levels of energy on the electromagnetic spectrumAtoms of different elements had different values for a quantum (the minimum amount of energy they can gain and lose)When atoms of different elements absorb their quantum of energy they can temporarily eject electronsDifferent colors of light are created by the different levels of energy being absorbed, and then given off, by electrons.
10 The question still remained: Why would different elements absorb different amounts of energy and then give off different light colors when energy was applied and they ejected electrons?What was it about their structures that allowed this?
11 Neils Bohr Model of the H Atom When an e- is hit by light energy, it absorbs the energy.If the energy is of the correct frequency (quantum), the e- will jump to another energy level (excited state vs ground state).The electron cannot stay in excited state so it falls back to the ground state. It cannot take the energy with it so it releases the energy in the form of LIGHT energy!We see this energy as different colors of light!When the electron returns to groundstate, it emits energyin the form ofdifferent colors of light – depending on frequency of energy gained.
12 Section 4.2 The Quantum Model of the Atom Objectives:Compare and contrast the Bohr model and the quantum model of the atom.List the 4 quantum numbers, and describe their significance.Relate the number of sublevels corresponding to each of an atom’s main energy levels, the number of orbitals per sublevel, and the number of orbitals per main energy level.
13 Today’s Quantum Model of the Atom Orbital -- a 3-d region around the nucleus that indicates the probable location of an e- .Quantum theory -- describes mathematically the wave properties of e- and other very small particlesQuantum numbers -- numbers or letters that specify the properties of atomic orbitals and the properties of e- in orbitals
14 4 Quantum Numbers Principal Quantum Number Angular Momentum Quantum NumberMagnetic Quantum NumberSpin Quantum Number
15 Principal quantum number (n) - indicates the main energy level occupied by the e-n = 1,2,3,4, etc.Representation of Bohr’s proposal of orbitals. Lower numbered levels are closer to nucleus and of lower energy.
16 Angular Momentum Quantum Number – indicates the orbital and its shape = s, p, d, f The maximum number of electrons in an energy level is 2n2, thus a shell with n = 2 may hold a maximum of 8 electrons.
17 Magnetic Quantum Number – orientation of orbital around the nucleus (the axis, or axes, it is on or between);uses x, y, z, xy, yz, etc.
18 These orbitals can overlap, cause interference with each other, and affect over all energy of each other.
19 Spin Quantum Number – spin state of an electron in an orbital – can be +½, -½ - often represented with arrows
20 Review 4 quantum numbers describe each e- . Principal quantum number - indicates the main energy level occupied by the e- = 1,2,3,4, etc.Angular Momentum Quantum Number – indicates the orbital and its shape = s, p, d, fMagnetic Quantum Number – orientation of orbital around the nucleus (the axis, or axes, it is on or between) = x, y, z, xy, yz, etc.Spin Quantum Number – spin state of an electron in an orbital – can be +½, -½ - often represented withSo…to accurately describe the energy (location) of electron, all 4 quantum numbers must be used.
21 Section 4.3 Electron Configurations Objectives:List the total number of electrons needed to fully occupy each main energy levelState and explain the Aufbau principle, the Pauli exclusion principle, and Hund’s rule.Write orbital notations, electron-configuration notation and noble-gas notation for atoms.
22 Electron Configuration Electron configuration -- the arrangement of e- in an atomElectron configurations summarize the locations of each e- in atoms.Like all systems in nature, e- tend to assume arrangements that have the lowest possible energiesGround-state configuration -- lowest energy arrangement of the e- for each element
23 3 basic rules help govern these ground-state configurations. Aufbau principle -- an e- occupies the lowest-energy orbital possiblePauli exclusion principle -- no 2 e- in the same atom can have the same set of 4 quantum numbersHund’s rule -- orbitals of equal energy are each occupied by one e- before any orbital is occupied by a second e- , and all e- in singly occupied orbitals must have the same spin
25 Types of Electron Configurations Orbital Notation – includes all 4 quantumnumbers (spin is indicated with an arrow)Fluorine – 9 e-_____ _____ _____ _____ _____1s s px py pz
26 Complete Orbital Notations for Elements 1 – 10 on the Periodic Table HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeon
27 Types of Electron Configurations 2. Electron Configuration Notation – includesprincipal and Angular Momentum (combinesmagnetic and ignores spin)Chromium e-1s2 2s2 2p6 3s2 3p6 4s2 3d4
28 Complete Electron Configurations for Elements 11 – 20 on the Periodic TableSodiumMagnesiumAluminumSiPSClArKCa
29 Types of Electron Configurations Noble Gas Notation – abbreviates part of theelectron configuration by using the NobleGas Symbol just prior to the element andadds the rest of the electron configurationMagnesium 12e-Mg = [Ne] 3s2
31 Complete Noble-Gas Notations for the following elements:SCaRbLiOAl
32 Chapter 4 Study GuideKnow the contributions of Plank and Einstein to the Quantum Theory.Be able to define “ground state” and “excited state” and know what happens to cause an atom to moved from ground state to excited state and what happens when it moves back to the ground state.Know the 4 quantum numbers by name and descriptions.Know the 3 rules: Aufbau, Pauli Exclusion, and Hund and know how they are applied when doing configurations.Know the order of energy levels and orbitals from lowest energy to highest energy.Be able to write configurations: orbital notation, electron configuration, and noble gas notations.Be able to use configurations to identify elements.Know the s, p, d, and f blocks on the Periodic Table so you can double check your work.
33 Class work A quantum of electromagnetic energy is called a(n) . A bright-light spectrum of an atom is caused bythe energy released when electrons .Explain the difference in an atom’s ground stateand excited state.Name and describe the information providedby all four quantum numbers.Describe the shapes of the s, p, and d orbitals.How many possible orientations can eachof the following sublevels have?p sublevel?d sublevel?f sublevel?
34 How many electrons can be held in a p sublevel?d sublevel?f sublevel?How many total electrons can be held in eachenergy level?1st –2nd –3rd –4th –
35 Explain each of the following rules governing electron configurations.Aufbau principle –Pauli exclusion principle –Hund’s rule –Do the following for an atom of Cesium.Orbital NotationElectron ConfigurationNoble Gas Notation