Electrons in Atoms Chapter 2 Section 2

Objectives  How does the electron relate to the modern atomic theory?  How do electron energy levels in an atom differ from one another?  How are Lewis electron dot diagrams used to illustrate valence electrons?

Important Vocabulary  Electromagnetic spectrum  Emission spectrum  Energy level  Electron cloud  Valence electrons  Lewis dot diagrams  Wavelength  Frequency

Electrons Motion & Energy  Electrons have enough energy to keep them in constant motion around the nucleus  This enables them to overcome the attraction of the positive nucleus  Electrons occupy orbitals of only certain amounts of energy  For them to move up a level of energy, energy must be added  For them to move down a level of energy, energy must be released in the form of light or heat

Waves Transfer Energy  Energy is the ability to exert a force over a certain distance  It is also the ability to do work  Waves carry energy because they can do work  For example:  Water waves can transfer energy to a leaf, to a boat, or onto a beach  Sound waves can transfer energy to your eardrum  Light waves can transfer energy to your eye  The bigger the wave the _______ energy it carries

Electromagnetic Radiation  Electromagnetic radiation travels in the form of waves that have both electric and magnetic properties  Electromagnetic waves travel through a vacuum at the speed of light (300 million m/s)  Two properties of waves are frequency and wavelength

Wavelength & Frequency Wavelength is the distance from one crest to the next Frequency is the number of waves per second A low frequency results in a long wavelength and a high frequency results in a shorter wavelength

Electromagnetic Spectrum  The electromagnetic spectrum consists of electromagnetic radiation waves at all possible energies, frequencies, and wavelengths  The spectrum ranges from 10 3 m to meters  Each part of the electromagnetic spectrum has unique properties

Radio Waves  Heinrich Hertz proved the existence of radio waves in the late 1880s  Radio waves have wavelengths that range from 200 to 600 m  Have the longest wavelengths and the lowest energy  They are used as TV signals, AM and FM radio signals, and for radar equipment  Radio telescopes view planets, comets, giant clouds of gas and dust, stars, and galaxies

Microwaves  Are low energy, low frequency radiation waves  They are used in Doppler radar for weather forecasting and to cook your food  Microwaves are also used to carry telecommunication signals  Most mobile phones use microwaves to transmit information, and space probes transmit signals back to Earth with microwaves  Different wavelengths of microwaves (grouped into "sub-bands") provide different information to scientists.

Infrared Waves  In 1800, William Herschel discovered them  Have less energy than visible light  Are given off by the human body and other warm objects  We experience infrared rays as heat from fires and electric heaters  Through night-vision goggles and infrared thermal cameras we can see infrared waves

Visible Spectrum  Cone-shaped cells in our eyes act as receivers tuned to the wavelengths in this narrow band of the spectrum  A typical human eye will respond to wavelengths from about 380 to 750 nm  The spectrum does not contain all the colors that the human eye and brain can distinguish  Unsaturated colors such as pink, and purple colors such as magenta, are absent because they can only be made by a mix of multiple wavelengths

 When white light shines through a prism, the white light is broken apart into the colors of the visible light spectrum  Water vapor in the atmosphere can also break apart wavelengths creating a rainbow

Ultraviolet Waves  In 1801, Johann Ritter discovered them  Have shorter wavelengths than visible light  UV waves are invisible to the human eye, but some insects, such as bumblebees, can see them.  The Sun is a source of the full spectrum of ultraviolet radiation, which is commonly subdivided into UV-A, UV-B, and UV-C

X-rays  X-rays were first observed and documented in 1895 by German scientist Wilhelm Conrad Roentgen  X-rays have very small wavelengths, between 0.03 and 3 nanometers  X-rays are used by doctors to see the internal structures of the body  X rays have very high energies, so they may kill living cells or turn them into cancer cells when exposed to too much of this type of radiation

Gamma Rays  Have the highest energy and shortest wavelengths  They are produced by the hottest and most energetic objects in the universe, such as neutron stars and pulsars, supernova explosions, and regions around black holes  On Earth, gamma waves are generated by nuclear explosions, lightning, and the less dramatic activity of radioactive decay  Gamma rays can be used to treat cancer by killing the diseased cells

Electrons and Light  When electrons become excited they give off light  The spectrum of light released is called the emission spectrum  Each element has a different emission spectrum  So it serves as evidence of energy levels within atoms

Energy Levels  There are seven levels of energy available for electrons to occupy  Electrons can move between energy levels like the rungs of a ladder absorbing or releasing energy  The number of filled energy levels depends on the number of electrons  Lower levels are filled first

Energy Levels  1 st energy level = 2 electrons  2 nd energy level = 8 electrons  3 rd energy level = 8 electrons  4 th energy level = 18 electrons  For example: Na has 11 electrons  2 in level 1  8 in level 2  1 in level 3

Valence Electrons  Every atom has between 1 and 8 valence electrons  Valence electrons are electrons in the outermost energy level of an atom  They determine an atom’s chemical properties and its ability for form bonds  For example: Neon Has 10 electrons 2 electrons in the lowest level 8 electrons in the 2 nd level Thus, it has 8 valence electrons!

Lewis Dot Diagrams  Valence electrons are usually the only electrons used in chemical reactions  We represent valence electrons as Lewis dot diagrams, which illustrate the valence electrons of an element as dots  Normally, all the elements within a group have the same electron dot structure with the exception of Helium

Lewis Dot Diagrams