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Terms in this set (43)
mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together.
chemical bonding that results from the electrical attraction between cations and anions.
results from the sharing of electron pairs between two atoms.
a covalent bond in which the bonding electrons are shared equally by the bonded atoms, resulting in a balanced distribution of electrical charge. Electronegativity difference of about 0-0.3.
Examples: H₂, CH₄
bonds that have an uneven distribution of charge.
a covalent bond in which the bonded atoms have an unequal attraction for the shared electrons. Electronegativity difference of about 0.3-1.7.
Examples: HCl, H₂O
a neutral group of atoms that are held together by covalent bonds.
a chemical compound whose simplest units are molecules
indicates the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts.
shows the types and numbers of atoms combined in a single molecule of a molecular compound.
a molecule containing only two atoms.
the energy required to break a chemical bond and form neutral isolated atoms.
the average distance between two bonded atoms
chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons (8) in its highest occupied energy level.
an 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.
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.
indicates the kind, number, arrangement, and bonds but not the unshared pairs of the atoms in a molecule.
a covalent bond in which one pair of electrons is shared between two atoms
a covalent bond in which two pairs of electrons are shared between two atoms.
a covalent bond in which three pairs of electrons are shared between two atoms.
double and triple bonds.
bonding in molecules or ions that cannot be correctly represented by a single Lewis structure.
composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal. Most of these exist as crystalline solids.
the simplest collection of atoms from which an ionic compound's formula can be established.
the energy released when one mole of an ionic crystalline compound is formed from gaseous ions.
a charged group of covalently bonded atoms. Here is a Quizlet deck of some common ones: "http://quizlet.com/31357022/polyatomic-ions-flash-cards/".
chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons.
the ability of a substance to be hammered or beaten into thin sheets.
the ability of a substance to be drawn, pulled, or extruded through a small opening to produce a wire.
the uneven distribution of molecular charge.
repulsion between the sets of valence-level electrons surrounding an atom causes these sets to be oriented as far as possible. To use the chart, know that A stands for the center and is always there, B is the number of bonds from the central atom, and E is any lone pair of electrons. For B and E, the subscripts represent how many of each there are (# of bonds, # of lone pairs) in the particular molecule.
Chart for VSEPR Theory
the mixing of two or more atomic orbitals of similar energies on the same atom to produce new hybrid atomic orbitals of equal energies.
orbitals of equal energy produced by the combination of two or more orbitals in the same atom.
forces of attraction between molecules.
created by equal but opposite charges that are separated by a short distance.
the intermolecular force in which a hydrogen atom that is bonded to a highly electronegative atom is attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule.
London Dispersion Forces
the intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles.
Naming Ionic Compounds
write the name of the metal first and then the nonmetal, adding -ide to the end of the non-metal. Polyatomics do not need endings added, so you use the metal as is, and then add the polyatomic name.
Example: Sodium Chloride, Potassium Flouride, Sodium Nitrite
Naming Covalent Compounds
The formula is written with the more electropositive element (the one further to the left on the periodic table) placed first, then the more electronegative element (the one further to the right on the periodic table).
[Important exception: when the compound contains oxygen and a halogen, the halogen is placed first. If both elements are in the same group, the one with the higher period number is named first.]
The first element in the formula is given the neutral element name, and the second one is named by replacing the ending of the neutral element name with -ide. A prefix is used in front of each element name to indicate how many atoms of that element are present.
Example: Carbon Dioxide, Dinitrogen tetroxide
Oxidation States (oxidation numbers)
Oxidation state shows the total number of electrons which have been removed from an element (a positive oxidation state-called oxidation) or added to an element (a negative oxidation state-reduction) to get to its present state.
Recognising this simple pattern is the single most important thing about the concept of oxidation states. If you know how the oxidation state of an element changes during a reaction, you can instantly tell whether it is being oxidised or reduced without having to work in terms of electron-half-equations and electron transfers.
involves an increase in oxidation state
involves a decrease in oxidation state