 # Chapter 10: Modern atomic theory Chemistry 1020: Interpretive chemistry Andy Aspaas, Instructor.

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Chapter 10: Modern atomic theory Chemistry 1020: Interpretive chemistry Andy Aspaas, Instructor

Rutherford’s atom Recall Rutherford’s atomic theory –Positively charged nucleus –Surrounded by negatively charged electrons Unanswered questions –How are electrons arranged –How do they move?

Electromagnetic radiation Electromagnetic radiation: energy transmitted by waves, “radiant energy” Wavelength: distance between peaks of these waves Different forms of electromagnetic radiation have different wavelengths

Electromagnetic spectrum

Types of electromagnetic radiation Radio waves: low frequency and energy Microwaves Infrared Visible Ultraviolet X-rays Gamma rays: high frequency and energy

Energy and electromagnetic radiation The shorter the wavelength, the higher the energy transmitted –Blue light: shorter wavelength: higher frequency: higher energy –Red light: longer wavelength: lower frequency: lower energy

Wave calculations Velocity = c = speed of light 2.998 x 10 8 m/s All types of light energy travel at same speed Amplitude = A = height of wave, brightness of light Wavelength =  = distance between peaks Frequency = = number of waves that pass a point in a given amount of time –Generally measured in Hertz (Hz) –1 Hz = 1 wave/sec = 1 sec -1 c =  x 

Planck’s nuclear theory Light energy behaves as particles in certain situations Each particle of light (a photon) has a certain fixed amount of energy –Energy of photon is directly proportional to frequency of the light –Higher frequency = more energy in photon

Atomic emission spectra Atoms that gain extra energy will release that energy in the form of light Light is given off in very specific wavelengths Different atoms give off different characteristic wavelengths of light when excited –Line spectrum: shows wavelengths of light that are emitted –Only certain wavelengths are given off, so only specific amounts of energy can be absorbed or given off for any one type of atom –Atoms are “quantized” - only specific energy levels

Bohr’s model of the atom Explains line spectrum of hydrogen Energy of atom is related to distance of electron from nucleus –Electrons can “jump” to different possible orbits around nucleus –Gain in energy: electron jump to higher quantum level - “excited state” –Lines in spectrum correspond to difference in energy levels Ground state: minimum energy level But, only explains hydrogen atom behavior –Plus, electrons do not have simple circular orbits

Wave mechanical model of the atom Electrons can be treated as waves (in the same way that light can also be treated as particles) Mathematics can calculate the probability densities of finding an electron in a particular region of the atom –Schrödinger equation - cannot predict location of any one particle, only probability of it being a certain place

Orbitals Solutions to wave equations give regions of high probability for finding electrons –Called orbitals –90% probability of finding an electron –3-dimensional shape

Orbitals and energy levels Principal energy level (n) = how much energy the electrons in the orbital have –Higher values mean higher energy and farther average distance from nucleus Each principal energy level has n sublevels –Different shape and energy –Named s, p, d, f Each sublevel has 1 or more orbitals –s = 1 orbital, p = 3, d = 5, f = 7

Pauli exclusion principle No orbital may have more than 2 electrons Electrons in same orbital must have opposite spin –s holds 2 electrons –p holds 6 electrons –d holds 10 electrons –f holds 14 electrons

Electron configurations Hydrogen electron configuration: 1s 1 –Superscript indicates number of electrons in orbital Helium: 1s 2 Follow the periodic table: row number = principal energy level (first number in electron configuration) Column and section determine which sublevel (s, p, d, f) is filled

Valence electrons Valence electrons: only those in outermost energy level - determine most of an atom’s reactivity properties Can indicate Na electron configuration as –1s 2 2s 2 2p 6 3s 1 or [Ne]3s 1 (using nearest Noble gas with smaller atomic number than the atom)

Atomic properties and the periodic table Ionization energy: energy required to remove an electron from an atom –Decreases down a group (less energy required to remove electron) –Increases across a period (more energy required to remove an electron) Atomic size (atomic radius) –Increases down group to account for greater mass –But decreases across period because more electrons mean more attraction to the nucleus

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