Bond energy- the energy required to break a chemical bond and form neutral atoms Bond length- the average distance between two bonded atoms Chemical bond- a link between atoms that results from the mutual attraction of their nuclei for electrons Chemical formula- represents the relative numbers of atoms of each kind in a chemical compound by sing atomic symbols and numerical subscripts Covalent bond- a chemical bond resulting from the sharing of electrons between two atoms Diatomic molecule- a molecule containing two atoms Dipole- created by equal but opposite charges separated by a short distance Dipole-dipole forces- forces of attraction between polar molecules Double bond- a covalent bond between two atoms produced by the sharing of two pairs of electrons Ductility- the state of being able to be drawn, pulled, or extruded through a small opening to produce a wire Formula unit- the simplest collection of atoms from which a compound[s formula can be established Hybrid orbitals- orbitals of equal energy produced by the combination of two or more orbitals on the same atom Hybridization- he mixing of two or more atomic orbitals of similar energies on the same atom to give new orbitals of equal energies Hydrogen bonding- intermolecular attraction between a hydrogen atom and an unshared pair of electrons on a strongly electronegative atom (such as fluorine, oxygen, or nitrogen) in another molecule Intermolecular forces- the forces of attraction between molecules Ionic bond- the chemical bond resulting from electrostatic attraction between positive and negative ions Ionic compound- is composed of positive and negative ions combined so that the positive and negative charges are equal Lattice energy- the energy released when one mole of an ionic crystalline compound is formed from gaseous ions Lewis structures- formulas in which atomic symbols represent nuclei and inner-shell electrons, dot-pairs or dashes between two atomic symbols represent electron pairs in covalent bonds, and dots adjacent to only one atomic symbol represent unshared electrons London dispersion forces- intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles and induced dipoles Malleability- the state of being able to be shaped or extended by beating with a hammer, rolling, or otherwise exerting physical pressure that results in a change in contour Metallic bond- a chemical bond resulting from the attraction between positive ions and surrounding mobile electrons Molecular compound- a chemical compound whose simplest formula units are molecules Molecular formula- shows the types and numbers of atoms combined in a single molecule Molecule- the smallest unit quantity of matter which can exist by itself and retain all the properties of the original substance Multiple bonds- double and triple bonds are referred, or multiple covalent bonds Nonpolar-covalent bond- a covalent bond which the bonding electrons are shared equally by the bonded atoms, with a resulting balanced distribution of electrical charge Octet rule- chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level Polar-covalent bond- a covalent bond in which the united atoms have an unequal attraction for the shared electrons Polyatomic ion- a charged group of covalently bonded atoms Single bond- a covalent bond produced by the sharing of one pair of electrons between two atoms Structural formula- indicates the kind, number, arrangement, and bonds of the atoms in a molecule Triple bond- a covalent bond produced by the sharing of three pairs of electrons between two atoms Unshared pair- also called a lone pair, is a pair of electrons that is not involved in bonding but instead belongs exclusively to one atom VSEPR theory- states that electrostatic repulsion between the valence-level electron pairs surrounding an atom causes these pairs to be oriented as far apart as possible
Types of Chemical Bonds ¨C Chemical bonds are subdivided into four general classes based on (a) the nature of the electron sharing and (b) types of atomic orbitals involved in the sharing. The four classes of bonds are (1) ionic; (2) covalent; (3) metallic; and (4) coordinate Ionic bonding- chemical bonding that results from the attraction between large numbers of cations and anions. Ions are mutually attracted because of their net electrical charges. The electrons are localized on individual atoms in atomic orbitals rather than shared. This type of bonding involves s- and p-orbitals. Covalent bonding- Chemical bonding that results from the sharing of electron pairs between two atoms. Sharing of electrons between atoms results in the creation of a new class of orbitals. Where atomic orbitals are centered about the atom nucleus, covalent orbitals are multi-polar (involving two or more nuclei). Each covalent (or molecular) orbital can hole two electrons. This type of bonding involves s- and p-orbitals. Metallic bonding- a form of chemical bonding involving multiple atoms and in which the valance electrons are delocalized throughout a network of overlapping vacant orbitals of roughly equal potential energy. This bond type uses ¡°d-orbitals¡± provide electrons to electron poor groups. Bond type and the difference in electronegativity between atoms in a chemical bond. (Ionic and covalent bond types) Non-polar covalent- a covalent bond in which the bonding electrons are shared equally by the bonded atoms, resulting in a balanced distribution of electrical charge Ionic character <5% \ Electronegativity difference roughly <.3 Polar-covalent- a bond with an uneven distribution of charge due to an unequal attraction for the charged electrons by the bonded atoms. 5% ¡Ü ionic character ¡Ü 50% 0.3 ¡Ü electronegativity difference ¡Ü 1.7 Ionic Ionic character > 50% Electronegativity difference > 1.7 Covalent bonding and molecular compounds Molecule- a neutral group of atoms that are held together by covalent bonds. While the number of atoms in a molecule may be quite large (DNA), the molecule has a fixed, definite number of atoms that constitute one molecule Molecular compound- a chemical compound whose simplest units are molecules Chemical formula- a listing of the relative numbers of atoms of each kind in a chemical compound Chemist represent a chemical formula using symbols for the elements and subscripts for the number of atoms of a given element. Empirical formula- a chemical formula that contains the fewest possible number of atoms to represent the elemental composition of a compound Molecular formula- a listing of the types and numbers of atoms in a single molecule of a molecular compound Diatomic molecule- a molecule containing only two atoms Formation of a covalent bond (energy and geometry considerations) Potential energy balance between free atoms and bonded atoms (How do the energy states ¡°before¡± and ¡°after¡± bonding compare. What energy barriers must be over to form a bond Rearrangement of atomic orbitals to molecular orbitals (conversion from single-center orbitals to multi-center affects the arrangement of atoms in 3D space.) Valence electrons (s- and p-orbitals) Inner-shell electrons (d-orbitals) Characteristics of the covalent bond Bond length- the distance between two boned atoms at the potential energy minimum (the separation distance where attractive and repulsive forces are in balance.) Bond energy- the equivalent of the amount of energy that is required to break a chemical bond and form isolated, neutral atoms The octet rule- the concept that atoms in Chemical compounds tend to act so as to form the equivalent of a noble gas s- and p- orbital electronic configuration Bonding proceeds so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level. Rule applies to the main-group elements (groups 1-2 and 13-17) Exceptions: 1st period elements in groups 1-2 of the second period and transition metals bonded to highly electronegative atoms using d orbital electrons as expanded valence electrons. Symbolic representations of chemical bonds Electron Dot Notation- and electron configuration notation in which only the valence electrons of an atom of a particular element are shown, indicated by dots placed around the element¡¯s symbol(Remember that we cannot distinguish electrons based on the atom with which they were originally associated. Dot notation simply seeks to symbolize the octet rule for each atom. Electrons shown between atoms are considered to be shared and available to complete the octet of all atoms involved in the sharing. Fluorine atom = F Fluorine molecule=F F Lewis Structures- Formulas in which atomic symbols represent nuclei and inner-shell electrons, dot-pairs or dashes between to atomic symbols represent electron pairs in covalent bonds, and dots adjacent to only one atomic symbol represent unshared electrons. Fluorine molecule = F-F Hydrogen Chloride = H-Cl Structural formula- a Lewis structure that shows the kind, number, and arrangement of atoms and location of bonds but does not show unshared electrons (A symbolic representation that attempts to depict the organization of atoms within a molecule these 2D symbolic representations do not specify molecular geometry) Single bond- a covalent bond produced by sharing one pair of electrons between two atoms Multiple bonds- a covalent bond produced by sharing two (double bond) or three (triple bond) pairs of electrons between to atoms. Triple bond example: N=N Resonance structures- bonding in molecules or ions that cannot be correctly represented using a single Lewis structure O=O-O or O-O=O