Chemistry 1411 Exam 4 Lessons 21-22 and 24-27

Valence-shell electron-pair repulsion (VSEPR) model

26. A model that accounts for the geometrical arrangements of shared and unshared electron pairs around a central atom in terms of the repulsions between electron pairs

resonance structure

26. One of two or more ALTERNATIVE Lewis structures for a molecule that cannot be described fully with a single Lewis Structure

bond enthalpy

26. The enthalpy change required to break a bond in 1 mole of gaseous molecules

Valence Shell

26. The outermost electron-occupied shell of an atom, which holds the electrons that are usually involved in bonding.

Resonance

26. The use of two or more Lewis structures to represent a particular molecule

sigma (σ) bonds

27. A covalent bond formed by orbitals overlapping end-to-end; its electron density is concentrated between he nuclei of the bonding atoms

pi (π) bonds

27. A covalent bond formed by sideways overlapping orbitals; its electron density is concentrated above and below the plane of the nuclei of the bonding atoms

homonuclear diatomic molecules

27. A diatomic molecule containing atoms of the same element

pi molecular orbital

27. A molecular orbital in which the electron density is concentrated above and below the plane of the two nuclei of the bonding atoms

sigma molecular orbital

27. A molecular orbital in which the electron density is concentrated around a line between the two nuclei of the bonding atoms

antibonding molecular orbital

27. A molecular orbital that is of higher energy and lower stability than the atomic orbitals from which it was formed

bonding molecular orbital

27. A molecular orbital that is of lower energy and greater stability than the atomic orbitals from which it was formed

diatomic molecules

27. A molecule that consists of two atoms

nonpolar molecules

27. A molecule that does not possess a dipole moment

Polar Molecules

27. A molecule that possess a dipole moment

molecular orbitals

27. An orbital that results from the interaction of the atomic orbitals of the bonding atoms

hybrid orbitals

27. Atomic orbitals obtained when two or more nonequivalent orbitals of the same atom combine

delocalized molecular orbitals

27. Molecular orbits that are not confined between two adjacent bonding atoms but actually extended over three of more atoms

bond order

27. The difference between the numbers of electrons in bonding molecular orbitals and antibonding molecular orbitals, divided by two

hybridization

27. The process of mixing the atomic orbitals in an atom (Usually the central atom) to generate a set of new atomic orbitals

Dipoole Moment (μ)

27. The product of charge and the distance between the charges in a molecule

electronegativity

25. the ability of an atom of an element to attract electrons toward itself in a chemical bond

Formal charge

25. the difference between the valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure

bond length

25. the distance between the nuclei of two bonded atoms in a molecule

single bond

25. two atoms are held together by one electron pair

triple bond

25. two atoms are held together by three pairs of electrons

double bond

25. two atoms are held together by two pairs of electrons

Coordinate covalent bond

26. A bond in which the pair of electrons is supplied by one of the two bonded atoms; also called a dative bond

polar covalent bond

25. in such a bond, the electrons spend more time in the vicinity of one atom than the other

Aufbau principle

21. An electron occupies the lowest-energy orbital that can receive it

paramagnetic

21. Attracted by a magnet. A paramagnetic substance contains one or more unpaired electrons.

representative elements

21. Elements in Groups 1A through 7A, all of which have incompletely filled s or p subshells of the highest principal quantum number

transition metals

21. Elements that have incompletely filled d subshells or readily give rise to cations that have incompletely filled d subshells.

Pauli exclusion principle

21. No two electrons in an atom can have the same four quantum numbers

diamagnetic

21. Repelled by a magnet; a diamagnetic substance contains only paired electrons.

electron configuration

21. The distribution of electrons among the various orbitals in an atom or molecule

noble gas core

21. The electron configuration of the noble gas element that most nearly precedes the element being considered

valence electrons

21. The electrons in the outermost shell (main energy level) of an atom; these are the electrons involved in forming bonds.

Hund's Rule

21. The most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins.

actinide series

21. a group of 14 elements following actinium in the periodic table, in which the 5f orbitals are being filled

core electrons

21. all nonvalence electrons in an atom

lanthanides (rare earth series)

21. have incompletely filled 4f subshells or readily give rise to cations that have incompletely filled 4f subshells

Isoelectronic

22. Ions, or atoms and ions, that possess the same number of electrons, and hence the same ground-state electron configuration, are said to be isoelectronic.

atomic radius

22. One-half the distance between the two nuclei in two adjacent atoms of the same element in a metal. For elements that exist as diatomic units, the atomic radius is one-half the distance between the nuclei of the two atoms in a particular molecule.

ionization energy (IE)

22. The minimum energy required to remove an electron from an isolated atom (or an ion) in its ground state

effective nuclear charge (Zeff)

22. The nuclear charge felt by an electron when both the actual charge (Z) and the repulsive effect (shielding) of the other electrons are taken into account

ionic radius

22. The radius of a cation or an anion as measured in an ionic compound (Distance from the center of an ion's nucleus to its outermost electron)

Born-Haber cycle

24. The cycle that relates lattice energies of ionic compounds to ionization energies, electron affinities, heats of sublimation and formation, and bond enthalpies

ionic bond

24. The electrostatic force that holds ions together in an ionic compound

electron affinity (EA)

24. The negative of the enthalpy change when an electron is accepted by an atom in the gaseous state to form an anion

Coulomb's law

24. The potential energy between two ions is directly proportional to the product of their charges and inversely proportional to the distance between them

Lewis dot symbol

24. The symbol of an element with one or more dots that represent the number of valence electrons in an atom of the element.

amphoteric

24. a substance that can act as both an acid and a base

diagonal relationships

24. similarities between pairs of elements in different groups and periods of the periodic table

Covalent Bond

25. A bond in which two electrons are shared by two atoms

Lewis Structure

25. A representation of covalent bonding using Lewis symbols; shared electron pairs are shown either as lines or as pairs of dots between two atoms, and lone pairs are shown as pairs of dots on individual atoms

Octet rule

25. An atom other than Hydrogen, tends to form bonds until it is surrounded by eight valence electrons

Covalent Compounds

25. Compounds containing only covalent bonds

Lone Pairs

25. Valence electrons that are not involved in covalent bond formations

multiple bonds

25. bonds formed when two atoms share two or more pairs of electrons