Electrons in atoms
This presentation is the property of its rightful owner.
Sponsored Links
1 / 65

Electrons in Atoms PowerPoint PPT Presentation


  • 130 Views
  • Uploaded on
  • Presentation posted in: General

Electrons in Atoms. Chapter 5. 5.1. Models of the Atom.

Download Presentation

Electrons in Atoms

An Image/Link below is provided (as is) to download presentation

Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author.While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server.


- - - - - - - - - - - - - - - - - - - - - - - - - - E N D - - - - - - - - - - - - - - - - - - - - - - - - - -

Presentation Transcript


Electrons in atoms

Electrons in Atoms

Chapter 5


Models of the atom

5.1

Models of the Atom

  • The scale model shown is a physical model. However, not all models are physical. In fact, several theoretical models of the atom have been developed over the last few hundred years. You will learn about the currently accepted model of how electrons behave in atoms.


The development of atomic models

5.1

The Development of Atomic Models

  • The Development of Atomic Models

    • What was inadequate about Rutherford’s atomic model?


The development of atomic models1

5.1

The Development of Atomic Models

  • Rutherford’s atomic model could not explain the chemical properties of elements.

    • Rutherford’s atomic model could not explain why objects change color when heated.


The development of atomic models2

5.1

The Development of Atomic Models

  • The timeline shows the development of atomic models from 1803 to 1911.


The development of atomic models3

5.1

The Development of Atomic Models

  • The timeline shows the development of atomic models from 1913 to 1932.


The bohr model

5.1

The Bohr Model

  • The Bohr Model

    • What was the new proposal in the Bohr model of the atom?


The bohr model1

5.1

The Bohr Model

  • Bohr proposed that an electron is found only in specific circular paths, or orbits, around the nucleus.


The bohr model2

5.1

The Bohr Model

  • Each possible electron orbit in Bohr’s model has a fixed energy.

    • The fixed energies an electron can have are called energy levels or shells. See periodic table

    • A quantum of energy is the amount of energy required to move an electron from one energy level (shell) to another energy level.


The bohr model3

5.1

The Bohr Model

  • Like the rungs of the strange ladder, the energy levels in an atom are not equally spaced.

  • The higher the energy level occupied by an electron, the less energy it takes to move from that energy level to the next higher energy level.


The quantum mechanical model

5.1

The Quantum Mechanical Model

  • The Quantum Mechanical Model

    • What does the quantum mechanical (electron cloud) model determine about the electrons in an atom?


The quantum mechanical model1

5.1

The Quantum Mechanical Model

  • The quantum mechanical model determines the allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus.

    • The exact location of an electron cannot be known with certainty because to do so you would have to interfere with it (Heisenberg Uncertainty Principle).


The quantum mechanical model2

5.1

The Quantum Mechanical Model

  • Austrian physicist Erwin Schrödinger (1887–1961) used new theoretical calculations and results to devise and solve a mathematical equation describing the behavior of the electron in a hydrogen atom.

  • The modern description of the electrons in atoms, the quantum mechanical model, comes from the mathematical solutions to the Schrödinger equation.


The quantum mechanical model3

5.1

The Quantum Mechanical Model

  • The propeller blade has the same probability of being anywhere in the blurry region, but you cannot tell its location at any instant. The electron cloud of an atom can be compared to a spinning airplane propeller.


The quantum mechanical model4

5.1

The Quantum Mechanical Model

  • In the quantum mechanical (electron cloud) model, the probability of finding an electron within a certain volume of space surrounding the nucleus can be represented as a fuzzy cloud. The cloud is more dense where the probability of finding the electron is high.


Atomic orbitals

5.1

Atomic Orbitals

  • Atomic Orbitals

    • How do sublevels of principal energy levels differ?


Atomic orbitals1

5.1

Atomic Orbitals

  • An atomic orbital is often thought of as a region of space in which there is a high probability of finding an electron.

    • Each energy sublevel corresponds to an orbital of a different shape, which describes where the electron is likely to be found.


Atomic orbitals2

5.1

Atomic Orbitals

  • Different atomic orbitals are denoted by letters. The s orbitals are spherical, and p orbitals are dumbbell-shaped.


Atomic orbitals3

5.1

Atomic Orbitals

  • Four of the five dorbitals have the same shape but different orientations in space.


Atomic orbitals4

Atomic Orbitals

F orbitals are weird.


Atomic orbitals5

Atomic Orbitals


Atomic orbitals6

Atomic Orbitals


Atomic orbitals7

5.1

Atomic Orbitals

  • The numbers and kinds of atomic orbitals depend on the energy sublevel.


Atomic orbitals8

Atomic Orbitals

  • The maximum number of electrons per orbital is two. So….

    • n = principal energy level

    • n = the number of sublevels per principal energy level

    • n2 = the number of orbitals per principal energy level

    • 2n2 = the number of electrons per principal energy level

      • Make a chart


Atomic orbitals9

Atomic Orbitals

  • So what is an electron doing in an orbital?

    • Spinning while moving. Either clockwise or counterclockwise.

    • Electrons occupying the same orbital spin in opposite directions.


Quantum numbers

Quantum Numbers

  • Electrons can be described by four quantum numbers:

    • n= principal quantum number (energy level)

    • l= orbital quantum number (sublevel)

    • ml= magnetic quantum number (orbital orientation in space)

    • ms= spin quantum number (+ or -)


Quantum numbers1

Quantum Numbers

  • No two electrons can have the same four quantum numbers (Pauli Exclusion Principle).


Electron arrangement in atoms

5.2

Electron Arrangement in Atoms

  • If this rock were to tumble over, it would end up at a lower height. It would have less energy than before, but its position would be more stable. You will learn that energy and stability play an important role in determining how electrons are configured in an atom.


Electron configurations

5.2

Electron Configurations

  • Electron Configurations

    • What are the three rules for writing the electron configurations of elements?


Electron configurations1

5.2

Electron Configurations

  • The ways in which electrons are arranged in various orbitals around the nuclei of atoms are called electron configurations.

    • Three rules—the aufbau principle, the Pauli exclusion principle, and Hund’s rule—tell you how to find the electron configurations of atoms.


Electron configurations2

5.2

Electron Configurations

  • Aufbau Principle

    • According to the aufbau principle, electrons occupy the orbitals of lowest energy first. In the aufbau diagram below, each box represents an atomic orbital.


Electron configurations3

5.2

Electron Configurations

  • Pauli Exclusion Principle

    • According to the Pauli exclusion principle, an atomic orbital may describe at most two electrons. To occupy the same orbital, two electrons must have opposite spins; that is, the electron spins must be paired.


Electron configurations4

5.2

Electron Configurations

  • Hund’s Rule

    • Hund’s rule states that electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin direction as large as possible.


Writing e lectron c onfigurations

Writing Electron Configurations

  • The principle energy level is the first number, the sublevel is the letter (s,p,d or f) and the number of electrons in the sublevel is the superscript.

Number of electrons in sublevel

3p4

Principal energy level

Type of sublevel


Electron configurations5

5.2

Electron Configurations

  • Orbital Filling Diagram


Pg 135

pg. 135


Exceptional electron configurations

5.2

Exceptional Electron Configurations

  • Exceptional Electron Configurations

    • Why do actual electron configurations for some elements differ from those assigned using the aufbau principle?


Exceptional electron configurations1

5.2

Exceptional Electron Configurations

  • Some actual electron configurations differ from those assigned using the aufbau principle because half-filled sublevels are not as stable as filled sublevels, but they are more stable than other configurations.


Exceptional electron configurations2

5.2

Exceptional Electron Configurations

  • Exceptions to the aufbau principle are due to subtle electron-electron interactions in orbitals with very similar energies.

  • Copper has an electron configuration that is an exception to the aufbau principle.


Physics and the quantum mechanical model

5.3

Physics and the Quantum Mechanical Model

  • Neon advertising signs are formed from glass tubes bent in various shapes. An electric current passing through the gas in each glass tube makes the gas glow with its own characteristic color. You will learn why each gas glows with a specific color of light.


Light

5.3

Light

  • Light

    • How are the wavelength and frequency of light related?


Light1

5.3

Light

  • The amplitude of a wave is the wave’s height from zero to the crest.

  • The wavelength, represented by  (the Greek letter lambda), is the distance between the crests.


Light2

5.3

Light

  • The frequency, represented by  (the Greek letter nu), is the number of wave cycles to pass a given point per unit of time.

  • The SI unit of cycles per second is called a hertz (Hz).


Light3

5.3

Light

  • The product of the frequency and wavelength always equals a constant (c), the speed of light.


Light4

5.3

Light

  • The wavelength and frequency of light are inversely proportional to each other.


Light5

5.3

Light

  • According to the wave model, light consists of electromagnetic waves.

    • Electromagnetic radiation includes radio waves, microwaves, infrared waves, visible light, ultraviolet waves, X-rays, and gamma rays.

    • All electromagnetic waves travel in a vacuum at a speed of 2.998  108 m/s.


Light6

5.3

Light

  • Sunlight consists of light with a continuous range of wavelengths and frequencies.

    • When sunlight passes through a prism, the different frequencies separate into a spectrum of colors.

    • In the visible spectrum, red light has the longest wavelength and the lowest frequency.


Light7

5.3

Light

  • The electromagnetic spectrum consists of radiation over a broad band of wavelengths.


Atomic spectra

5.3

Atomic Spectra

  • Atomic Spectra

    • What causes atomic emission spectra?


Atomic spectra1

5.3

Atomic Spectra

  • When atoms absorb energy, electrons move into higher energy levels. These electrons then lose energy by emitting light when they return to lower energy levels.


Atomic spectra2

5.3

Atomic Spectra

  • A prism separates light into the colors it contains. When white light passes through a prism, it produces a rainbow of colors.


Atomic spectra3

5.3

Atomic Spectra

  • When light from a helium lamp passes through a prism, discrete lines are produced.


Atomic spectra4

5.3

Atomic Spectra

  • The frequencies of light emitted by an element separate into discrete lines to give the atomic emission spectrum of the element.

Mercury

Nitrogen


An explanation of atomic spectra

5.3

An Explanation of Atomic Spectra

  • An Explanation of Atomic Spectra

    • How are the frequencies of light an atom emits related to changes of electron energies?


An explanation of atomic spectra1

5.3

An Explanation of Atomic Spectra

  • In the Bohr model, the lone electron in the hydrogen atom can have only certain specific energies.

    • When the electron has its lowest possible energy, the atom is in its ground state.

    • Excitation of the electron by absorbing energy raises the atom from the ground state to an excited state.

    • A quantum of energy in the form of light is emitted when the electron drops back to a lower energy level.


An explanation of atomic spectra2

5.3

An Explanation of Atomic Spectra

  • The light emitted by an electron moving from a higher to a lower energy level has a frequency directly proportional to the energy change of the electron.

    • Therefore each transition produces a line of a specific frequency in the spectrum.


An explanation of atomic spectra3

5.3

An Explanation of Atomic Spectra

  • The three groups of lines in the hydrogen spectrum correspond to the transition of electrons from higher energy levels to lower energy levels.


Quantum mechanics

5.3

Quantum Mechanics

  • Quantum Mechanics

    • How does quantum mechanics differ from classical mechanics?


Quantum mechanics1

5.3

Quantum Mechanics

  • In 1905, Albert Einstein successfully explained experimental data by proposing that light could be described as quanta of energy.

    • The quanta behave as if they were particles.

    • Light quanta are called photons.

  • In 1924, De Broglie developed an equation that predicts that all moving objects have wavelike behavior.


Quantum mechanics2

5.3

Quantum Mechanics

  • Today, the wavelike properties of beams of electrons are useful in magnifying objects. The electrons in an electron microscope have much smaller wavelengths than visible light. This allows a much clearer enlarged image of a very small object, such as this mite.


Quantum mechanics3

5.3

Quantum Mechanics

  • Classical mechanics adequately describes the motions of bodies much larger than atoms, while quantum mechanics describes the motions of subatomic particles and atoms as waves.


Quantum mechanics4

5.3

Quantum Mechanics

  • The Heisenberg uncertainty principle states that it is impossible to know exactly both the velocity and the position of a particle at the same time.

    • This limitation is critical in dealing with small particles such as electrons.

    • This limitation does not matter for ordinary-sized object such as cars or airplanes.


Quantum mechanics5

5.3

Quantum Mechanics

  • The Heisenberg Uncertainty Principle


  • Login