(See following explanations)

Covalent crystal: A crystal that consists of only one molecule. All atoms are joined to others with covalent bonds. Also called netword crystal.
Crystal lattice structure: The arrangement of atoms, ions, or molecules in a crystal structure.


Diatomic: A term describing a molecule that contains only two atoms, e.g., HCl or hydrogen gas.
Electron deficient: A term describing a Lewis structure that has fewer than an octet of electrons around one or more of its atoms, except for H.
Electronic configuration: A listing of the electrons within an atom, based on the sublevels that are filled and the relative energies of these sublevels.
Element: Any one of the 109 distinct particles, know as atoms, that are currently known. Each has distinct chemical and physcial properties.
Group (family): A column in the periodic table.
Halide: An organic compound with a halogen in its structure.
Hund's Law: The rule that every orbital in a sublevel must fill with one electron before a second electron of opposite spin can be added to any orbital in that                       sublevel.
Isoelectronic: A term describing any two atoms that have identical electronic configurations. These atoms may be ions or elements.
Isomers: Distinctly different compounds that have the same elemental compositions.
Isotope: A form of an element with a specified number of protons, neutrons, and electrons.
Lewis structure: A molecular structure based on the concept that all atoms try to achieve the noble gas configuration by sharing electrons.
Linear: A term referring to atoms aligned in a straight line; a three-atom arrangement with a 180 degree bond angle.
Magnetic quantum number, m: The quantum number that specifies the orbital in which an electron is located and the orientation of the orbital in space.
Metallic crystal: A crystal formed from a metal in the periodic table. Metallic crystals are malleable, ductile, and conduct electricity.
Molecular crystal: A crystal formed from a molecule. The attractive forces that hole molecular crystals together are London forces, dipole-dipole attractions,                              hydrogen bonding, or a combination of these.
Molecular orbital: An orbital created by the pairing of electrons from different atoms. This orbital encircles the atoms that are bonded together.
Molecule: A group of atoms bound together by covalent bonds with 0 total charge.
Noble gas: An elecment in the last group in the periodic table. They are unusually stable elements and have full valence shells.
Valence electrons: The outermost s and p electrions in an atom. The number and the arrangement of valence electrons define chemicl and physical             
                              properties.
Valence sheel electron-pair repulsion (VSEPR) theory: A method evaluating molecular structure be relating the number of bonding and nonbonding    
                                                                                       electron pairs on an atom to its geometrical strucuture.
Unit cell: The fundamental building blocks of crystals. An entire crystal is formed by repetitive stacking of the unit cells.
Triangular bipyramid: A geometric structure with five atoms covalently bound to a central atom. Three atoms in the equatorial positions are 120 degrees
                                   from each other. Two additional atoms in the axial positions are 90 degrees from the equatorial atoms.
Tetrahedron: A geometric structure with four atoms bound to a central atom by covalent bonds. Each bond is equidistant from any other with a bond angle
                     of 109 degrees.
Symmetrical: A term describing a geometrical property whereby a structure may be rotated some angle less than 360 degrees and after rotation the
                      molecule has the same configuration as before.
Sublevel: A subdivision of an energy level. Electrons in each principal energy level are localized in sublevels. Each sublevel has a distinct shape associated
               with it. Sublevels are numbered from zero up to one less than the number of the principal energy level. These subelevel numbers are the azimuthal                quantum numbers, l. Sublevels are also designated by the letters s, p, d, f.
Structural isomers: Compounds with the same formula but with the atoms bonded in different arrangements.
Spin quantum number, m: The quantum number that specifies the spin of an electron as either +1/2 or -1/2. Two electrons in the same orbital must have                                               opposite spins.
Simple cubic: A term describing a cubic structure with one atom in each of the eight corners of a unit cell.
Resonance structure: A Lewis structure that can be drawn in more than one equally probable way. The actual structure is a mixture of all possible                                                   resonance structures.
Radiosotope: A radioactive isotope of an element.
Radioactivity: The property that some unstable nuclei have of decaying spontaneously with the emission of a small particle and/or energy.
Quantum number: One of four numbers used in a wave-mechanical model of the atom to describe an electron in an atom.
Proton: One of three particles, along with the electrion and neutron, that make up an atom. The proton has a positive charge, equal in magnitude (but with                  the opposite sign) to the charge of the electron. The number of protons is equal to the atomic number, Z, of an element. Protons and neutrons make              up the bulk of the mass of an atom.
Principal quantum number, n: The quantum number that specifies the energy level of the atom in which an electron is located; n may have any integer from                                                 1 to infinity.
Planar triangle: A geometric strucutre of four atoms, three bonded to a central atom with 120 degree angles between the atoms, which are all in the same                              plane.
Pauli exclusion principle: The requirement that no two electrons in a n atom have the same set of four quantum numbers.
Octet rule: A simple but effective rule stating that covalent molecules tend to have octets of electrons atround each atoms in their structures. These octets                       simulate the electron configurations of the noble gases.
Octahedron: A geometric structure of six atoms covalently bound to a central atom. Each atom is 90 degrees from any other.
Nucleus: The center of an atom, which contrains the protons and neutrons. The nucleus is extremely dense and comprises a very small fraction of of the                    atom's volume; the rest of the atom is empty space.
Nucleon: Either a proton or a neutron, both of which are fundamental particles of the nucleus.
Nonbonding electron pair: A pair of electrons in a Lewis structure that is not shared with any other atoms.
Natural adundance: The percentage of an isotope of an element found in nature.


Electronic Configuration:

Principal Energy Levels (Shells), n

• The principal energy level closest to the nucleus has an n value of 1
• Each level succeeds up to the largest element which has an n level of 7
• Each level can hold 2n² electrons: the 1st level holds 2, t2nd holds 8, 3rd 18, 4th 32, and so on

Sublevels (Subshells), l

• l: azimuthal quantum numbers
• The number of possible sublevels is equal to the value of n
• The value of l can never be greater than n-1:   l< or l= n-1

Sublevel Letter

• S, p, d, f
• Sublevel number indicates sublevel letter
• l=0 then letter=s; 1 indicates p, 2 indicates d, and so on 

Orbitals

• To share an orbital two electrons must have opposite spins
• Orbitals are called s, p, d, or f according to the sublevel in numbers of orbitals in a sublevel=2l+1
• Each orbital is given a number, the magnetic quantum number: m_l
• m_l -l to +l

Shape of Electron Cloud In Orbitals

• Note: look at pictures in Barron’s on page 23
• s sublevels: electron cloud is spherical
• p sublevels: dumbbell shape
• d orbitals: see Barron’s or orbital and links on this webpage
• f orbitals: even stranger. You are not required to know these

Electronic Configurations:

The Aufbau, or energy erder, of the numerical values of the energies of the orbitals is the following:

• 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 4f, 5d, 6p, 7s, 5f, 6d
• However do not memorize! You can get this from the structure of the periodic table.
• See Figure 1.16 on page 23 on Barron’s, or our lovely website which has a wonderful diagram of this. Important to remember is how to fill the orbitals:

                                                                                                           F                         14

There are some exceptions to this rule, however, most atoms follow this rule.
Common exceptions are:
Cu, Ag, Au, Cr, Mo


Write the complete electronic configurations for the following:
Na
Cl
Answers:
Na: 1s2, 2s2, 2p6, 3s1
Cl: 1s2, 2s2, 2p6, 3s2, 3p5

Abbreviated Electronic Configuration:

• The electrons up to the last completely filled 6p sublevel are the inner, less important electrons. They may be replaced by the symbol of a noble gas, in brackets.
• For example: Mg can be abbreviated as: [Ne]2s2

Lewis Structure

• Bonded electrons have to be spread out as far as possible to minimize the electron repulsion that occurs according to the
• VSEPR theory of valence electrons between the non-bonded electron pairs.
• Carbon is always considered to be the central atom.
• Molecules that are single atom or two atoms with no non-bonded electrons are linear in shape. However the two sets of bonded electron pairs yield a minimum repulsion to each other, resulting in a distance of 180 degrees. Examples of these profiles would be
• HCl, CO, BeCl2, CO2.
• When there are three atoms, or three groups of atoms, surrounding a central atom, the structure is referred to as having a Trigonal Planar shape, and would have a 120o separation between the atoms. Examples of this would be BF3 and AlH3.
• Molecules with four atoms, or atom groups surrounding the central atom is referred to having a Tetrahedral shape because the
• four groups would be occupying the corners of a tetrahedron. Examples of these are: CH4, CCl4, CHCl3, BF4- ion, NH4+ ion, and CH3-CH3.

To see more on this subject go to
http://members.aol.com/profchm/zero_nb.html

Isotopes

• An element has a fixed number of protons but may exist with different numbers of neutrons. Isotopes of an element have the same chemical properties but different atomic weights.
• If you want to refer to a certain isotope, you write it like this: ^AX_Z. Here X is the symbol for the element, Z is the atomic number, and A is the total number of neutrons and protons combined, and is called the mass number. For example, "normal"
• hydrogen is written ¹H₁, deuterium is ²H₁, and tritium is ³H₁.

• To see a complete table showing the diversity of isotopes go to: http://chemlab.pc.maricopa.edu/periodic/isotopes.html

Test Your Knowledge...Again!

For a practice test on atomic structure, go to http://www.learnchem.net/quiz/quiz9/ or http://chem.neopages.com/practice/ats.shtml