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 created or destroyed - they may only be converted from one form to the other. Energy is the ability to do work and it occurs in a number of 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 are capable of independent, prolonged existence. These are called nuclides. If nuclides are grouped by chemical characteristics, about 100 sets result and these are referred to as elements. The modern Periodic Table was devised in 1869 by Julius Meyer and Dmitri Mendeleev. It organizes the elements into groups and families with similar chemical and physical properties.
Periods Groups or Families Periodic Table
Crustal Abundance Crustal volume:<1% of Earth Mantle volume: 83% of Earth Core volume: 16% of Earth Crustal mass:<1% of Earth Mantle mass: 68% of Earth Core mass: 31% of Earth What element is most abundant for the entire Earth?
Atomic Particles: Basics Atoms are composed of electrons and two large nuclear particles called protons and neutrons. Protons and neutrons are approximately equal in mass and are ~1800 times more massive than the electron. Both nuclear particles are composed of quarks, smaller fundamental particles. Protons have unit positive charge (+1), while electrons have unit negative charge (-1). Neutrons carry no charge. Atoms are electrically neutral and thus the number of electrons must equal the number of protons.
Basic Terminology Atomic number (Z): The atomic number represents the number of unit positive charges on the nucleus and is equal to the number of protons within the nucleus, since each proton carries unit positive charge. In electrically neutral atoms, it also represents the number of electrons, which carry unit negative charge. Mass number (A): The mass number is equal to the total number of nucleons, which is the sum of the number of protons and neutrons. A does not equal the total mass of the atom; rather, it represents a whole number approximation 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, with the A, the mass number, placed in the upper left and Z, the atomic number, placed in the lower left corner. For example, 23 Na 11, has a mass number of 23 and an atomic number of 11. Isotopes are atoms of the same element that differ in mass. For example, 87 Sr and 86 Sr or 238 U and 235 U. Isotopes have similar chemical characteristics and are studied using a mass spectrograph or spectrometer. Most elements have at least two naturally occurring isotopes. Isobars are nuclides that have the same mass number but different atomic numbers. For example, 36 S and 36 Ar are isobars; they both contain a total of 36 nucleons (protons plus neutrons), but the sulfur isotope has 16 protons and 20 neutrons, while the argon isotope has 18 protons and 18 neutrons. Isobars do not have similar chemical characteristics!
Atomic Weight Atomic weight is the weighted average of the atomic masses of the naturally occurring isotopes. For example, a natural sample of the element chlorine contains a mixture of 75.53% 35 Cl and 24.47% 37 Cl. Thus the atomic weight is obtained 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 = ν ν = 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 e ch/"kt $ 1 d" where h is Planck s constant and k is Boltzman s constant! m T = ch 4.97k = 2.90x10"3 mk
Rutherford s Scattering Experiment
Balmer Series: Visible spectrum for Hydrogen
Bohr Atomic Model Oxygen Atom The Bohr model for the atom envisioned these electrons in stable orbits of specified radius and energy, where we could exactly pinpoint the position of any individual electron. Each energy level was permitted to have a specified number of electrons, and was called a shell. We know now that this simple view is not correct; it is impossible to exactly determine the position of an electron in space.
Complete emission spectrum and energy levels 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 that specified regions near the nucleus have higher probabilities for finding an electron than others. Each electron does, however, have a specific energy. Must solve wave equation for specific states! The combination of the energy and probability gives rise to the current understanding for electron distributions, which are referred to as electron 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 additional electron also added to it extra-nuclear cloud. From theory and experiment, we know that these electrons are added in a systematic fashion, 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 orbitals S, P, and D orbital probability density functions P orbitals 1S orbital D orbitals Pauli Exclusion Principle: no two electrons may have the same quantum number Hunds Rule: electrons prefer parallel spins before pairing up in subshell orbital
Quantum Numbers n = 6; l = 4; m = 1 n = principal l = orbital (0 -> (n-1)) m l = magnetic (± l) m s = spin (±1/2) l = 0 -> s orbital = 1 -> p orbital = 2 -> d orbital = 3 -> f orbital From: http://www.daugerresearch.com/orbitals/
Animated Orbital Simulations Superposition of States 4, 3, 3 and 4, 1, 0 3, 2, 2 and 3, 1, -1 From: http://www.daugerresearch.com/orbitals/
H atom quantum numbers and X-rays