Chapter 6 Characteristics of Atoms Department of Chemistry and Biochemistry Seton Hall University.

Chapter 6Characteristics of Atoms

Department of Chemistry and Biochemistry

Seton Hall University

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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

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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

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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

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Connections between wavelength and

frequency• c = 3108 m/s in a vacuum• make sure the units all agree!

c

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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-

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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

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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

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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

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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

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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

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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

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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

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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

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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)

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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

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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

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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