Atomic Structure and the Composition of Matter

• The atom is a basic building block of minerals.• Matter is a special form of energy; it has mass and

occupies space. Neither matter nor energy may be createdor destroyed - they may only be converted from one formto the other.

• Energy is the ability to do work and it occurs in a numberof forms, including:– Potential; Kinetic; Electrical; Heat; Chemical; Nuclear– Radiant (the only form in which there is an absence of

matter)• Atoms are the smallest division of matter that retain the

characteristics of the elements.

Elements and the Periodic Table

• Approximately 300 different kinds of atoms that arecapable of independent, prolonged existence. Theseare called nuclides.

• If nuclides are grouped by chemical characteristics,about 100 sets result and these are referred to aselements.

• The modern Periodic Table was devised in 1869 byJulius Meyer and Dmitri Mendeleev. It organizesthe elements into groups and families with similarchemical and physical properties.

Periodic Table

Groups or Families

Periods

Crustal Abundance

Crustal volume:<1% of EarthMantle volume: 83% of EarthCore volume: 16% of Earth

Crustal mass:<1% of EarthMantle mass: 68% of EarthCore mass: 31% of Earth

What element is most abundant for the entire Earth?

Atomic Particles: Basics

• Atoms are composed of electrons and two largenuclear particles called protons and neutrons.

• Protons and neutrons are approximately equal inmass and are ~1800 times more massive than theelectron. Both nuclear particles are composed ofquarks, smaller fundamental particles.

• Protons have unit positive charge (+1), whileelectrons have unit negative charge (-1). Neutronscarry no charge.

• Atoms are electrically neutral and thus thenumber of electrons must equal the number ofprotons.

Basic Terminology

• Atomic number (Z): The atomic number represents thenumber of unit positive charges on the nucleus and isequal to the number of protons within the nucleus, sinceeach proton carries unit positive charge. In electricallyneutral atoms, it also represents the number of electrons,which carry unit negative charge.

• Mass number (A): The mass number is equal to the totalnumber of nucleons, which is the sum of the number ofprotons and neutrons. A does not equal the total mass ofthe atom; rather, it represents a whole numberapproximation of the mass, as expressed in amu.

• The number of neutrons is simply defined as the A - Z.

Isotopes and Isobars• A specific type of atom is designated by using its chemical symbol,

which is an abbreviation of its name in German, Latin, or English, withthe A, the mass number, placed in the upper left and Z, the atomicnumber, placed in the lower left corner. For example, 23Na11, has amass number of 23 and an atomic number of 11.

• Isotopes are atoms of the same element that differ in mass. Forexample, 87Sr and 86Sr or 238U and 235U. Isotopes have similarchemical characteristics and are studied using a mass spectrographor spectrometer. Most elements have at least two naturally occurringisotopes.

• Isobars are nuclides that have the same mass number but differentatomic numbers. For example, 36S and 36Ar are isobars; they bothcontain a total of 36 nucleons (protons plus neutrons), but the sulfurisotope has 16 protons and 20 neutrons, while the argon isotope has 18protons and 18 neutrons. Isobars do not have similar chemicalcharacteristics!

Atomic Weight

• Atomic weight is the weighted average of the atomicmasses of the naturally occurring isotopes. For example,a natural sample of the element chlorine contains a mixtureof 75.53% 35Cl and 24.47% 37Cl. Thus the atomic weight isobtained by multiplying the mass of each isotope (in amu)times its fractional abundance:

• 0.7553 (34.97 amu) + 0.2447 (36.95 amu) = 35.45 amu

Atomic Models

• Bohr Model– Electron shells

• Quantum Mechanics– Orbitals– Afbau Filling Order– Quantum numbers and superposition of states

EM propagation and spectrum

λ is wavelength (m)ν is frequency (cycles/s= s-1)c is velocity of EM rad (ms-1)

λ = ν/ c and cλ = νν = 1/λ is wavenumber (m-1)

Note that in 1900 Planck Determined that the energy of a photon is quantized:

E = hυ

where h is Planck’s constant

Blackbody radiation Planck’s Law

!"d" =

8#ch"$5

ech /"kT

$1d"

where h is Planck’s constantand k is Boltzman’s constant

!mT =

ch

4.97k= 2.90x10

"3mK

Rutherford’s Scattering Experiment

Balmer Series: Visible spectrum for Hydrogen

Bohr Atomic Model

The Bohr model for the atom envisioned these electrons in stable orbits of specifiedradius and energy, where we could exactly pinpoint the position of any individualelectron. Each energy level was permitted to have a specified number of electrons, andwas called a shell. We know now that this simple view is not correct; it is impossibleto exactly determine the position of an electron in space.

Oxygen Atom

Complete emissionspectrum and energylevels for H atom

The Quantum Mechanical View• Using the theory of quantum or wave mechanics we can calculate the

probabilities of various electron configurations, and thus show thatspecified regions near the nucleus have higher probabilities for findingan electron than others. Each electron does, however, have a specificenergy. Must solve wave equation for specific states!

• The combination of the energy and probability gives rise to the currentunderstanding for electron distributions, which are referred to aselectron orbitals; these orbitals are referred to as s (sharp), p(principal), d (diffuse), and f (fundamental).

• With increasing atomic number, each new element has an additionalelectron also added to it extra-nuclear cloud. From theory andexperiment, we know that these electrons are added in a systematicfashion, with the lowest energy orbitals being filled first.

• This process is called aufbau filling (1s -> 2s -> 2p -> 3s -> 3p -> 4s -> 3d -> 4p -> 5s, etc. ).

S, P, and D orbital probability density functions S orbitals

P orbitals

D orbitals

1S orbital

Pauli Exclusion Principle:no two electrons may havethe same quantum number

Hunds Rule: electrons preferparallel spins before pairingup in subshell orbital

Quantum Numbers

• n = principal• l = orbital (0 -> (n-1))• ml = magnetic (± l)• ms = spin (±1/2)

– l = 0 -> s orbital = 1 -> p orbital = 2 -> d orbital = 3 -> f orbitaln = 6; l = 4; m = 1

From: http://www.daugerresearch.com/orbitals/

Animated Orbital Simulations

| 4, 3, 3 | and | 4, 1, 0 | | 3, 2, 2 | and | 3, 1, -1 |Superposition of States

From: http://www.daugerresearch.com/orbitals/

H atom quantum numbers and X-rays