chapter 6 characteristics of atoms department of chemistry and biochemistry seton hall university
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Chapter 6Characteristics of Atoms
Department of Chemistry and Biochemistry
Seton Hall University
2
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
3
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
4
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
5
Connections between wavelength and
frequency• c = 3108 m/s in a vacuum• make sure the units all agree!
c
6
Characterization of Radiation
υhcλ
hchυ)moleculeΔE(erg
λ(cm)
1υ
λ(cm)
)secc(cm)υ(sec
molecule
sec erg106.626h
E
hcλor
λ
hcE
energyor υ,υλ,
1
11-
27
7
Wavelength and Energy Units
• Wavelength– 1 cm = 108 Å = 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
8
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
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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
10
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
11
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
12
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
13
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
14
Beer’s Law
lcAI
I 010log
• Io = Intensity of incident light
• I = Intensity of transmitted light = molar extinction coefficient• l = path length of cell• c = concentration of sample
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UV Spectral Nomenclature
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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
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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
18
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
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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
20
Energy level diagrams
• Wiggle lines indicate radiative processes
• Straight lines indicate nonradiative processes
• Each energy level represents an arrangement of electrons in the atom
h
h'
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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)
22
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
23
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
24
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
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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)
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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
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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)
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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
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Waves and nodes
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A variety of radial projections
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Radial depictions
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The p-orbitals
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The d-orbitals
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d-orbital radial projection