Chemical Bonding

VSEPR

"Valence-Shell, electron-pair repulsion", referring to the repulsion between pairs of valence electrons of the atoms in a molecule. The theory states that repulsion between the sets of valence-level electrons surrounding an atom causes these sets to be oriented as far apart as possible.

Bond order

# of bonds involved for that atom / # of R.S.

Single, Double, Triple Bonds (bond length vs. strength)

1. Single bond is a covalent bond in which one pair of electrons is shared between two atoms. Longest and weakest 2. A double bond is a covalent bond in which two pairs of electrons are shared between two atoms. All four electrons in a double bond "belong" to both atoms. Medium length and medium strength 3. A triple bond is a covalent bond in which three pairs of electrons are shared between atoms . Shortest and strongest

Bent

2 bonded atoms and 1 lone pair or 2 bonded atoms and 2 lone pairs

Trigonal pyramidal

3 bonded atoms and 1 lone pair

T-shape

3 bonded atoms and 2 lone pairs

Trigonal planar

3 bonded atoms and no lone pairs (flat)

See-saw

4 bonded atoms and 1 lone pair

Square Planar

4 bonded atoms and 2 lone pairs

Tetrahedral

4 bonded atoms and no lone pairs

Square Pyramidal

5 bonded atoms and 1 lone pair

Trigonal bipyramidal

5 bonded atoms and no lone pairs

Octahedral

6 bonded atoms and no lone pairs

Polyatomic Ions

A charged group of covalently bonded atoms. Certain atoms bond covalently with each other to form a group of atoms that has both molecular and ionic characteristics. Polyatomic ions combine with ions of opposite charge to form ionic compounds. The charge of a polyatomic ion results from an excess of electrons (negative charge) or a shortage of electrons (positive charge). For example, an ammonium ion, a common positively charged polyatomic ion, contains one nitrogen atom and four hydrogen atoms and has a single positive charge. Its formula is sometimes written as [NH4]+ to show that the group of atoms as a whole has a charge of 1+. It behaves as a unit

Molecular Compound

A chemical compound whose simplest units are molecules.

Polar-covalent Bond

A covalent bond in which the bonded atoms have an unequal attraction for the shared electrons.

Nonpolar-covalent Bond

A covalent bond in which the electrons are shared equally by the bonded atoms, resulting in a balanced distribution of electrical charge.

Chemical Bond

A mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together. As independent particles, most atoms are at relatively high potential energy. Nature, however, favors arrangements in which potential energy is minimized. This means that most atoms are less stable existing by themselves than when they are combined. By bonding with each other, atoms decrease in potential energy, thereby creating more stable arrangements of matter.

Molecule

A neutral group of atoms that are held together by covalent bonds. A single molecule of a chemical compound is an individual unit capable of existing on its own. It may consist of two or more atoms of the same element, or of two or more different atoms.

Lone pairs

An unshaired pair of electrons that is not involved in bonding and that belongs exclusively to one atom

Electronegativity in Bonding

Bonding between atoms of different elements is rarely purely ionic or purely covalent. It usually falls somewhere between these two extremes, depending on how strongly the atoms of each element attract electrons (electronegativity). The degree to which bonding between atoms of two elements is ionic or covalent can be estimated by calculating the difference in electronegativity. Almost none: Metallic 0-0.3: Nonpolar-covalent 0.3-1.7: Polar-covalent 1.7-33: Ionic

Electron Deficient

Boron has just three valence electrons. Because electron pairs are shared in covalent bonds, boron tends to form bonds in which it is surrounded by six electrons. Beryllium has only 2 valence electrons and 3 electrons total, therefore it is unable to be surrounded by eight electrons and have a full octet.

Ionic Bond

Chemical bonding that results from the electrical attraction between cations and anions. They are composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal. Most ionic compounds exist as crystalline solids. A crystal of any ionic compound is a three-dimensional network of positive and negative ions mutually attracted to one another. As a result, in contrast to a molecular compound, an ionic compound is not composed of independent, neutral units that can be isolated and examined. The chemical formula of an ionic compound merely represents the simplest ratio of the compound's combined ions that gives electrical neutrality, How electrons are used: Atoms completely give up electrons to other atoms. Atoms are transferred. Types of elements: Most of the rocks and minerals that make up Earth's crust consist of positive and negative ions held together by ionic bonding. Physical properties: The ions in ionic compounds are held together by strong attractive forces, so ionic compounds generally have higher melting and boiling points than do molecular compounds. Ionic compounds are hard but brittle because in an ionic crystal, even a slight shift of one row of ions relative to another causes a large buildup of repulsive forces. These forces make it difficult for one layer to move relative to another, causing ionic compounds to be hard. If one layer is moved, however, the repulsive forces make the layers part completely, causing ionic compounds to be brittle. In the solid state, the ions cannot move, so the compounds are not electrical conductors. In the molten state, ionic compounds are electrical conductors because the ions can move freely to carry electrical current. Many ionic compounds can dissolve in water. When they dissolve, their ions separate from each other and become surrounded by water molecules. These ions are free to move through the solution, so such solutions are electrical conductors. Other ionic compounds do not dissolve in water because the attractions between the water molecules and the ions cannot overcome the attractions between the ions. Structure: Ionic crystal or crystal lattice. In an ionic crystal, ions minimize their potential energy. The attractive forces include those between oppositely charged ions and those between the nuclei and electrons of adjacent ions. The repulsive forces include those between like-charged ions and those between electrons of adjacent ions. The distances between ions and their arrangement in a crystal represent a balance among all these forces. Attraction between the adjacent oppositely charged ions is much stronger than repulsion by other ions of the same charge, which are farther away. Why these bond types occur, based on electronegativity: There is a very large difference in electronegativity because one element will want to lose electrons to become a cation (metals), while another element will want to gain electrons to become an anion (nonmetals).

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.

Importance of hydrogen bonding for water's properties

High boiling point, solvent properties due to its polarity

Polar

It occurs in bonds with significantly different electronegativities where the electrons are more strongly attracted by the more-electronegative atom. Bonds that have an uneven distribution of charge (ionic bonds are always polar)

Lattice energy

Lattice energy is the energy released when one molecule of an ionic crystalline compounded is formed from gaseous ions (amount of energy released when separated ions in a gas come together to form a crystalline solid). It is used to compare bond strengths in ionic compounds. Negative energy values indicate that energy is released when the crystals are formed.

Linear

Only 2 atoms, 2 bonded atoms and 0 lone pairs, 2 bonded atoms and 4 lone pairs

Covalent Bond

Results from the sharing of electron pairs between two atoms. How electrons are used: Electrons are shared or "owned" equally by the two bonded atoms. Types of elements: nonmetals bonded to nonmetals Physical properties: They have a low melting point because the forces of attraction between individual molecules are not very strong. In fact, many molecular compounds are already completely gaseous at room temp. Structure: Bond energy Why these bond types occur, based on electronegativity: There is a low electronegativity difference and therefore the two atoms benefit by sharing electrons rather than transferring electrons

Molecular Formula

Shows the types and numbers of atoms combined in a single molecule of a molecular compound. It is the chemical formula of a molecular compound.

Expanded Octet

Some elements can be surrounded by more than eight electrons when they combine with the highly electronegative elements fluorine, oxygen, and chlorine. Bonding involves electrons in d orbitals as well as in s and p orbitals.

Resonance structures

Some molecules and ions cannot be represented adequately by a single Lewis structure. For example, in ozone, it has one single and one double bond. Chemists once speculated that ozone split its time existing as one of these two structures, constantly alternating, or "resonating," from one to the other. Now, scientists saw that ozone has a single structure that is the average of these two structures, resonance structures.

Formal Charges

The charge the atom would have if bonding electrons were equally divided between the two atoms that are sharing them (# of valence- # of things surrounding the atom such as lone pairs or bonds). To match the E.N. of the atoms: more EN atoms should have - formal charge, less EN atoms should have + formal charge. The overall formal charge must equal the charge of the compound.

Bond Length

The distance between two bonded atoms at their minimum potential energy, the average distance between two bonded atoms.

Bond Energy

The energy required to break a chemical bond and form neutral isolated atoms. The amount of energy released in the forming of a bond equals the difference between the potential energy at the zero level (separated atoms) and that at the bottom of the valley (bonded atoms). The same amount of energy must be added to separate the bonded atoms. The shorter the bond, the higher the bond energy

Ionic vs. Molecular

The force that holds ions together in ionic compounds is a very strong overall attraction between positive and negative charges. In a molecular compound, the covalent bonds of the atoms making up each molecule are also strong. But the forces of attraction between molecules are much weaker than the forces among formula units in ionic bonding.

Intermolecular Forces

The forces of attraction between molecules. Intermolecular forces vary in strength but are generally weaker than bonds that join atoms in molecules, ions in ionic compounds, or metal atoms in solid metals.

Dipole-dipole

The forces of attraction between polar molecules. A dipole is created by equal but opposite charges that are separated by a short distance. These forces are short-range forces, acting only between nearby molecules. For molecules containing more than two atoms, molecular polarity depends on both the polarity and the orientation of each bond.

Metallic Bond

The highest energy levels of most metal atoms are occupied by very few electrons. Within a metal, the vacant orbitals in the atoms' outer energy levels overlap. This overlapping of orbitals allows the outer electrons of the atoms to roam freely throughout the entire metal. How electrons are used: Electrons are delocalized, which means that they do not belong to any one atom but move freely about the metal's network of empty atomic orbitals. These mobile electrons form a sea of electrons around the metal atoms, which are packed together in a crystal lattice. Types of elements: Physical properties: The freedom of electrons to move in a network of metal atoms accounts for the high electrical and thermal conductivity characteristic of all metals. In addition, metals are both strong absorbers and reflectors of light. Because they contain many orbitals separated by extremely small energy differences, metals can absorb a wide range of light frequencies. This absorption of light results in the excitation of the metal atoms' electrons to higher energy levels. However, in metals the electrons immediately fall back down to lower levels, emitting energy in the form of light as a frequency similar to the absorbed frequency. This re-radiated (or reflected) light is responsible for the metallic appearance or luster of metal surfaces. Most metals are malleable and ductile because metallic bonding is the same in all directions throughout the solid. When struck, one plane of atoms in a metal can slide past another without encountering resistance or breaking bonds. Teh amount of energy as heat required to vaporize the metal is a measure of the strength of the bonds that hold the metal together. Structure: A sea of electrons Why these bond types occur, based on electronegativity: Very low electronegativity

London dispersion

The intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles. In any atom or molecule-polar or nonpolar-the electrons are in continuous motion. As a result, at any instant the electron distribution may be slightly uneven. The momentary, uneven charge creates a positive pole in one part of the atom or molecule and a negative pole in another. This temporary dipole can then induce a dipole in an adjacent atom or molecule. The two are held together for an instant by the weak attraction between the temporary dipoles. London forces act between all atoms and molecules. But they are the only intermolecular forces acting among nonpolar molecules. They have low boiling points because they are very weak attractions. Because London fores are dependent on the motion of electrons, their strength increases with the number of electrons in the interaction atoms or molecules. London forces increase with increasing atomic or molecular mass.

Hydrogen Bonding

The intermolecular force in which a hydrogen atom that is bonded to a highly electronegative atoms is attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule (F, O, N). They are highly polar due to the large electronegativity differences. They have a very high boiling point due to their strong attractive force.

Formula Unit

The simplest collection of atoms from which an ionic compound's formula can be established. For example, one formula unit of sodium chloride, NaCl, is one sodium cation plus one chloride anion. (the ending of the negative element's name is replaced with -ide). The ratio of ions in a formula unit depends on the charges of the ions combined.

Molecular polarity

The strongest intermolecular forces exist between polar molecules. Polar molecules act as tiny dipoles because of their uneven charge distribution. A dipole is created by equal but opposite charges that are separated by a short distance. The direction of a dipole is from the dipole's positive pole to its negative pole.

Ionic Compounds

Writing formulas based on charges: The ratio of ions in a formula unit depends on the charges of the ions combined. For example, to achieve electrical neutrality in the ionic compound calcium fluoride, two fluoride anions, F-, each with a charge of 1-, must balance the 2+ charge of each calcium cation, Ca2+. Therefore, the formula of calcium fluoride is CaF2.