10.1 History of the Atom Dalton’s Model of the Atom

Slides:



Advertisements
Similar presentations
Modern Atomic Theory Chapter 10
Advertisements

The Modern Atom Figure: 05-00CO Caption:
Renee Y. Becker CHM 1025 Valencia Community College
Chapter 11 Modern Atomic Theory. Copyright © Houghton Mifflin Company. All rights reserved. 11 | 2 Rutherford’s Atom The concept of a nuclear atom (charged.
Chapter 5 Models of the Atom by Christopher Hamaker
Electron Configuration
Chapter 4: Arrangement of Electrons in Atoms
Chapter 4 Models of the Atom by Christopher G. Hamaker
Electronic Structure of Atoms © 2009, Prentice-Hall, Inc. Chapter 6 Electronic Structure of Atoms Chemistry, The Central Science, 11th edition Theodore.
Dalton Model of the Atom
Introductory Chemistry, 2nd Edition Nivaldo Tro
Chapter 10: Modern atomic theory Chemistry 1020: Interpretive chemistry Andy Aspaas, Instructor.
Chapter 5 Electrons in Atoms Wave and Particle Models of Light
Concept #4 “Electrons in the Atom” Honors Chemistry 1.
1 Ch 4 Electron Energies. 2 Electromagnetic Spectrum Electromagnetic radiation is a form of energy that exhibits wave-like behavior as it travels though.
The Wave Nature of Light. Waves To understand the electronic structure of atoms, one must understand the nature of electromagnetic radiation. The distance.
Chapter 6 Electronic Structure of Atoms
Chapter 4: Arrangement of Electrons in Atoms Chemistry.
Chapter 6 Electronic Structure of Atoms. Waves To understand the electronic structure of atoms, one must understand the nature of electromagnetic radiation.
Chapter 4 Electron Configurations. Early thoughts Much understanding of electron behavior comes from studies of how light interacts with matter. Early.
Chapter 4 Arrangement of Electrons in Atoms 4.1 The Development of a New Atomic Model.
Atomic Models Scientist studying the atom quickly determined that protons and neutrons are found in the nucleus of an atom. The location and arrangement.
Electronic Structure of Atoms © 2009, Prentice-Hall, Inc. Chapter 7 Electronic Structure of Atoms.
Development of Atomic Models
Electronic Structure of Atoms Electronic Structure of Atoms.
Chapter 4 Arrangement of Electrons in Atoms. 4-1 The Development of the New Atomic Model Rutherford’s atomic model – nucleus surrounded by fast- moving.
Chapter 5 Electrons in Atoms Chemistry Section 5.1 Light and Quantized Energy At this point in history, we are in the early 1900’s. Electrons were the.
1 Electronic Structure of Atoms Chapter 6 2 The Wave Nature of Light All waves have a characteristic wavelength,, and amplitude, A. The frequency,, of.
1 Chapter 7 Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Quantum Theory and the Electronic Structure of.
A TOMIC O RBITALS AND E LECTRON C ONFIGURATIONS. Waves  Electrons behave like waves.  The distance between corresponding points on adjacent waves is.
Chapter 5: Electrons in Atoms
Modern Atomic Theory Mr. Heyroth.
Light and Energy Electromagnetic Radiation is a form of energy that emits wave-like behavior as it travels through space. Examples: Visible Light Microwaves.
Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Table of Contents Chapter 4 Arrangement of Electrons in Atoms Section.
Chapter 5 Electrons in Atoms Honors Chemistry Section 5.1 Light and Quantized Energy At this point in history, we are in the early 1900’s. Electrons.
Electrons in Atoms Chapter Wave Nature of Light  Electromagnetic Radiation is a form of energy that exhibits wavelike behavior as it travels through.
Lesson 3 : The Bohr Model. Bohr Model of an Atom  Electrons orbit the nucleus in fixed energy ranges called orbits (energy levels)  An electron can.
Unit 4 Energy and the Quantum Theory. I.Radiant Energy Light – electrons are understood by comparing to light 1. radiant energy 2. travels through space.
Bohr’s Model Rutherford’s model didn’t explain the arrangement of electrons around the nucleus.
Electron in Atoms Chapter 5. Rutherford’s Atomic Model Discovered dense positive piece at the center of the atom- “nucleus” Electrons would surround.
CHAPTER 11 NOTES MODERN ATOMIC THEORY RUTHERFORD’S MODEL COULD NOT EXPLAIN THE CHEMICAL PROPERTIES OF ELEMENTS.
Chapter 4 © Houghton Mifflin Harcourt Publishing Company Section 1 The Development of a New Atomic Model Properties of Light The Wave Description of Light.
E LECTRONS IN A TOMS Chapter 5. L IGHT AND Q UANTIZED E NERGY Nuclear atom and unanswered questions Scientists found Rutherford’s nuclear atomic model.
Chapter 6 Electronic Structure of Atoms John D. Bookstaver St. Charles Community College Cottleville, MO Lecture Presentation © 2012 Pearson Education,
Light Light is a kind of electromagnetic radiation, which is a from of energy that exhibits wavelike behavior as it travels through space. Other forms.
Chapter 4 1 © 2011 Pearson Education, Inc. Dalton Model of the Atom John Dalton proposed that all matter is made up of tiny particles. These particles.
Keep until June 2011! Unit 2.2: Electrons.
Chapter 6 Electronic Structure of Atoms
Arrangement of Electrons in Atoms
5-1 Quantum Theory of the atom
Electronic Structure of Atoms
Chapter 13 Electrons in Atoms.
Electrons in Atoms.
Quantum Theory.
Electromagnetic spectrum
Electrons in Atoms Chapter 5.
A New Atomic Model Chapter 4 Section 1.
Islamic University - Gaza
Chapter 4 Models of the Atom by Christopher G. Hamaker
Electromagnetic spectrum
Chapter 6 Quantum Mechanical Model & Electron Configurations
Presentation transcript:

10.1 History of the Atom Dalton’s Model of the Atom John Dalton proposed that all matter is made up of tiny particles. These particles are molecules or atoms. Molecules can be broken down into atoms by chemical processes. Atoms cannot be broken down by chemical or physical processes. Chapter 5

Thomson’s Model of the Atom J.J. Thomson proposed a subatomic model of the atom in 1903. Thomson proposed that the electrons were distributed evenly throughout a homogeneous sphere of positive charge. This was called the “plum pudding” model of the atom. Chapter 5

Rutherford’s Gold Foil Experiment Rutherford’s student fired alpha particles at thin gold foils. If the “plum pudding” model was correct, α-particles should pass through undeflected. At the center of an atom is the atomic nucleus, which contains the atom’s protons. Chapter 5

10.2 Radiant Energy Spectrum The complete radiant energy spectrum is an uninterrupted band, or continuous spectrum. The radiant energy spectrum includes most types of radiation, most of which are invisible to the human eye. Chapter 5

Visible Spectrum Light usually refers to radiant energy that is visible to the human eye. The visible spectrum is the range of wavelengths between 400 and 700 nm. Radiant energy that has a wavelength lower than 400 nm and greater than 700 nm cannot be seen by the human eye. Chapter 5

The Wave/Particle Nature of Light In 1900, Max Planck proposed that radiant energy is not continuous, but is emitted in small bundles. This is the quantum concept. Radiant energy has both a wave nature and a particle nature. An individual unit of light energy is a photon. Chapter 5

c= ln Wave Nature of Light Light travels through space as a wave, similar to an ocean wave. Wavelength is the distance light travels in one cycle. Frequency is the number of wave cycles completed each second. Light travels at a constant speed: 3.00 × 108 m/s (given the symbol c). c= ln Chapter 5

Waves The distance between corresponding points on adjacent waves is the wavelength (). The number of waves passing a given point per unit of time is the frequency (). Amplitude

Wavelength vs. Frequency The longer the wavelength of light, the lower the frequency. The shorter the wavelength of light, the higher the frequency. Chapter 5

Sample Calculating Frequency from Wavelength The yellow light given off by a sodium vapor lamp used for public lighting has a wavelength of 589 nm. What is the frequency of this radiation? Solution 589 nm 1 m 5.89 X 10-7m 1 X 109 nm 3.00 X 108 m/s 5.89 X 10-7 m 11

The Nature of Energy Einstein used this assumption to explain the photoelectric effect. (electrons are ejected from metals when light from specific wavelengths are applied) He concluded that energy is proportional to frequency: E = h where h is Planck’s constant, 6.63  10−34 J-s.

The Nature of Energy Therefore, if one knows the wavelength of light, one can calculate the energy in one photon, or packet, of that light: c =  E = h

Sample Energy of a Photon Calculate the energy of one photon of yellow light that has a wavelength of 589 nm (n= 5.09 X 1014 s-) Solution The value of Planck’s constant, h, is given both in the text and in the table of physical constants on the inside back cover of the text, and so we can easily calculate E: 14

The Wave Nature of Matter Louis de Broglie proposed that if light can have material properties, matter should exhibit wave properties. He demonstrated that the relationship between mass and wavelength was  = h mv

h mv  = What is the wavelength of an electron moving with a speed of Sample Matter Waves What is the wavelength of an electron moving with a speed of 5.97 × 106 m/s? The mass of the electron is 9.11 × 10-31 kg. Solution  = h mv 6.40 X 1034 J*s (9.11 X10-31 kg)(5.97 X106 m/s) mass Kg Velocity m/s 1.22 X 10-10 m 16

Emission Line Spectra When an electrical voltage is passed across a gas in a sealed tube, a series of narrow lines is seen. These lines are the emission line spectrum. The emission line spectrum for hydrogen gas shows three lines: 434 nm, 486 nm, and 656 nm. Chapter 5

“Atomic Fingerprints” The emission line spectrum of each element is unique. We can use the line spectrum to identify elements using their “atomic fingerprint.” Chapter 5

10.5 Bohr Model of the Atom Niels Bohr speculated that electrons orbit about the nucleus in fixed energy levels. Electrons are found only in specific energy levels, and nowhere else. The electron energy levels are quantized. Chapter 5

The Quantum Concept The quantum concept states that energy is present in small, discrete bundles. For example: A tennis ball that rolls down a ramp loses potential energy continuously. A tennis ball that rolls down a staircase loses potential energy in small bundles. The loss is quantized. Chapter 5

Evidence for Energy Levels Bohr realized that this was the evidence he needed to prove his theory. The electric charge temporarily excites an electron to a higher orbit. When the electron drops back down, a photon is given off. The red line is the least energetic and corresponds to an electron dropping from energy level 3 to energy level 2. Chapter 5

10.7 Quantum Mechanics Erwin Schrödinger developed a mathematical treatment into which both the wave and particle nature of matter could be incorporated. This is known as quantum mechanics.

Quantum Model It was later shown that electrons occupy energy sublevels within each level. These sublevels are given the designations s, p, d, and f. These designations are in reference to the sharp, principal, diffuse, and fine lines in emission spectra. The number of sublevels in each level is the same as the number of the main level. Chapter 5

Energy Levels and Sublevels The first energy level has 1 sublevel: 1s The second energy level has 2 sublevels: 2s and 2p The third energy level has 3 sublevels: 3s, 3p, and 3d Chapter 5

Electron Occupancy in Sublevels The maximum number of electrons in each of the energy sublevels depends on the sublevel: The s sublevel holds a maximum of 2 electrons. The p sublevel holds a maximum of 6 electrons. The d sublevel holds a maximum of 10 electrons. The f sublevel holds a maximum of 14 electrons. The maximum electrons per level is obtained by adding the maximum number of electrons in each sublevel. Chapter 5

Electrons per Energy Level Chapter 5

Quantum Mechanical Model An orbital is the region of space where there is a high probability of finding an electron. In the quantum mechanical atom, orbitals are arranged according to their size and shape. The higher the energy of an orbital, the larger its size. s-orbitals have a spherical shape Chapter 5

Shapes of p-Orbitals Recall that there are three different p sublevels. p-orbitals have a dumbbell shape. Each of the p-orbitals has the same shape, but each is oriented along a different axis in space. Chapter 5

d- orbitals

F-orbitals

10.9 Electron Configurations Electrons are arranged about the nucleus in a regular manner. The first electrons fill the energy sublevel closest to the nucleus. Electrons continue filling each sublevel until it is full, and then start filling the next closest sublevel. A partial list of sublevels in order of increasing energy is: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d … Chapter 5

Filling Diagram for Sublevels The order does not strictly follow 1, 2, 3, etc. For now, use this figure to predict the order of sublevel filling. Chapter 5

Orbital Diagrams Each box in the diagram represents one orbital. Half-arrows represent the electrons. The direction of the arrow represents the relative spin of the electron.

Electron Configurations The electron configuration of an atom is a shorthand method of writing the location of electrons by sublevel. The sublevel is written followed by a superscript with the number of electrons in the sublevel. If the 2p sublevel contains 2 electrons, it is written 2p2. The electron sublevels are arranged according to increasing energy. Chapter 5

Electron Configurations follow 3 rules Rules for e- Configs Electron Configurations follow 3 rules Aufbau- electrons fill in lowest energy first (start at the bottom) Pauli Exclusion- 2 electrons maximum in an orbital, with opposite spins to reduce repulsion Hund’s- everyone (box) gets one e- before anyone gets seconds- (in degenerate orbitals)

Writing Electron Configurations First, determine how many electrons are in the atom. Bromine has 35 electrons. Arrange the energy sublevels according to increasing energy: 1s 2s 2p 3s 3p 4s 3d … Fill each sublevel with electrons until you have used all the electrons in the atom: Br: 1s2 2s2 2p6 3s2 3p6 4s2 3d 10 4p5 The sum of the superscripts equals the atomic number of bromine (35). Chapter 5

Valence Electrons When an atom undergoes a chemical reaction, only the outermost electrons are involved. These electrons are of the highest energy and are furthest away from the nucleus. These are the valence electrons. The valence electrons are the s and p electrons beyond the noble gas core. For our purposes we will include ALL valence electrons past the noble gas core.

10.10 e- Configs using the Periodic Table We fill orbitals in increasing order of energy. Different blocks on the periodic table (shaded in different colors in this chart) correspond to different types of orbitals.

Blocks and Sublevels We can use the periodic table to predict which sublevel is being filled by a particular element.

Noble Gas Core Electron Configuration Recall, the electron configuration for Na is: Na: 1s2 2s2 2p6 3s1 We can abbreviate the e- config by indicating the innermost electrons with the symbol of the preceding noble gas. The preceding noble gas before sodium is neon, Ne. We rewrite the electron configuration: Na: [Ne] 3s1