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This chapter explores the characteristics of atoms, including their mass, nucleus, electrons, and properties. It also delves into the wave aspects of light and its interactions with atoms through absorption and emission processes. The concepts of photon energy, the photoelectric effect, and the quantization of absorption and emission are discussed. UV and visible spectroscopy techniques are explained.
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Chapter 6Characteristics of Atoms Department of Chemistry and Biochemistry Seton Hall University
Characteristics of Atoms • Atoms posses mass • most of this mass is in the nucleus • Atoms contain positive nuclei • Atoms contain electrons • Atoms occupy volume • electrons repel each other, so no other atom can penetrate the volume occupied by an atom • Atoms have various properties • arises from differing numbers of protons and electrons • Atoms attract one another • they condense into liquids and solids • Atoms can combine with one another
Wave aspects of Light • Most useful tool for studying the structure of atoms is electromagnetic radiation • Light is one form of that radiation • Light is characterized by the following properties: • frequency, , nu • wavelength, , lambda • amplitude
Electric and magnetic field components of plane polarized light • Light travels in z-direction • Electric and magnetic fields travel at 90° to each other at speed of light in particular medium • c (= 3 × 1010 cm s-1) in a vacuum
Connections between wavelength and frequency • c = 3108 m/s in a vacuum • make sure the units all agree!
Wavelength and Energy Units • Wavelength • 1 cm = 108Å = 107 nm = 104 =107 m (millimicrons) • N.B. 1 nm = 1 m (old unit) • Energy • 1 cm-1 = 2.858 cal mol-1 of particles = 1.986 1016 erg molecule-1= 1.24 10-4 eV molecule-1 • E (kcal mol-1) (Å) = 2.858 105 • E(kJ mol-1) = 1.19 105/(nm)297 nm = 400 kJ
The photoelectric effect • A beam of light impacts on a metal surface and causes the release of electrons (the photoelectron) if certain conditions are satisfied • Conditions • light must have a frequency above the threshold, o • number of photoelectrons increases with light intensity, but not the kinetic energy
Explanation of the photoelectric effect • Ephoton = hphoton • h = Planck’s constant = 6.626 10-34 J s • Applying the Law of the Conservation of Energy • energy of the photon is absorbed by the metal surface and is transferred to the photoelectron • the minimum frequency is the binding energy of the electron • the remaining energy shows up as the kinetic energy of the electron
Photoelectric effect • Electron kinetic energy = Photon energy - Binding energy • Ekinetic(electron) = h - ho • Comments • if frequency is too low, the photo energy is insufficient to overcome the binding energy of the electron • energy in excess of the binding energy shows up as the kinetic energy of the electron • increasing the intensity of the light increases the number of photons impacting on the metal
Particle properties of light • Light has a dual nature of acting like a wave and acting like a particle • The photoelectric effect confirmed that light occurs as little packets of energy • Light is still diffracted like a wave, has wavelength and frequency
Light and atoms • When matter absorbs photons of light, the energy of the photon is transferred to the matter • In the case of atoms, the absorption process yields information about the atom • Absorption of a photon transforms the atom to a higher energy state • All higher energy states are referred to as excited states • The most stable state is the ground state
Absorption and Emission • White light (light containing all energies of light) is passed through a sample • Sample absorbs some of the light • Light that passes through the sample is dispersed by a prism or other wavelength selecting device • Photodetector records the intensity of the light passing through the sample, which is then interpreted as absorption of light
Beer’s Law • Io = Intensity of incident light • I = Intensity of transmitted light • = molar extinction coefficient • l = path length of cell • c = concentration of sample
UV and Visible Spectroscopy • Vacuum UV or soft X-rays • 100 - 200 nm • Quartz, O2 and CO2 absorb strongly in this region • N2 purge good down to 180 nm • Quartz region • 200 – 350 nm • Source is D2 lamp • Visible region • 350 – 800 nm • Source is tungsten lamp
Emission • Sample is excited by light • Excited sample emits the light • Emitted light is wavelength selected • The light is detected by a photodetector • Plot of emission intensity vs wavelength is generated
Quantization of absorption and emission • One of the three things that led to quantum theory was that the absorption and emission of light occurred at discrete frequencies, not continua • Interpreted as the energy of the photon must match the difference in energy of two energy levels in the atom or molecule
Molecular process • Absorption and emission of visible and ultraviolet light • Photon is annihilated upon absorption, and the electrons in the molecule are rearranged into the excited state • Emission results from the conversion of excited electron energy being converted to a photon of light • Ephoton = Eatom
Energy level diagrams • Wiggle lines indicate radiative processes • Straight lines indicate nonradiative processes • Each energy level represents an arrangement of electrons in the atom
Properties of electrons • Each electrons have the same mass and charge • Electrons behave like magnets through a property called spin (actually, magnets are magnets because electrons have this property) • Electrons have wave properties (diffract just like photons)
Heisenberg uncertainty principle • A particle has a particular location, but a wave has no exact position • The wave properties of electrons cause them to spread out, hence the position of the electron cannot be precisely defined • They are referred to as being delocalized in a region of space • Heisenberg proposed that the motion and position of the particle-wave cannot be precisely known at the same time
Bound electrons and quantization • The properties of electrons bound to a nucleus can only take on certain specific values (most importantly, energy) • Absorption and emission spectra provide experimental values for the quantized energies of atomic electrons • Theory of quantum mechanics links these data to the wave characteristics of electrons bound to nuclei
The Schrödinger Equation • A second order partial differential equation • The solutions to such equations are other equations • These equations describe three-dimensional waves called orbitals • These solutions have indexes that are integers (the solutions are quantized naturally) • These indexes are called quantum numbers
Quantum numbers • n - principle quantum number • values of the positive integers • n = 1,2,3,… • l - azimuthal quantum number • values correlate with the number of preferred axes of a particular orbital, indicating its shape • l = 0,1,2,…(n - 1) • value of l is often indicated by a letter (s, p, d, f, for l = 0, 1, 2, 3)
Quantum number • ml - magnetic quantum number • directionality of orbital • ml = 0, ±1, ±2, ±l • ms - spin orientation quantum number • ms = ±½ • A complete description of an atomic electron requires a set of four unique quantum number that meet the restrictions of quantum mechanics
Shapes of atomic orbitals • Each atomic energy level can be associated with a specific three-dimensional atomic orbital • Orbitals are maps of the probability of the electron being in a particular location around the nucleus • While there are many representations, the most important to learn are the 90% probability volumes (which I will draw for you)
Depictions of orbitals • electron density plot - electron density plotted against the distance from the nucleus • orbital density plots • electron contour diagrams (90% probability drawings) • All are useful in helping us visualize the orbital
