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AP Chemistry NOTES 14-1 TYPES of CHEMICAL BONDS

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AP Chemistry NOTES 14-1 TYPES OF CHEMICAL BONDS Chemical Bond – attractive forces that hold atoms together TYPES OF CHEMICAL BONDS *IONIC BOND – generally involves metals from the left side of the periodic table interacting with nonmetals from the right side of the table through a transfer of electrons This generally results in the formation of a solid, regular array of positive and negative ions called a crystal lattice: The strength of the attraction between these ions is known as the crystal lattice energy. (Related to the enthalpy of dissociation). The likelihood of this type of bond forming can be predicted using the following variation of Coulomb’s Law: where E = crystal lattice energy k = a proportionality constant Q1 = charge of first ion Q2 = charge of second ion d = distance between ion centers A negative crystal lattice energy indicates that the compound is lower in energy than the individual atoms, so it favors the formation of the compound. QUESTION: Discuss the formation of a compound containing potassium and barium QUESTION: Which would have the higher melting point, sodium chloride or potassium chloride? Why? Which would be more soluble in water? Why? QUESTION: Which would have the higher melting point, sodium chloride or magnesium oxide? Why? Which would be more soluble in water? Why? Characteristics of Ionic Compounds 1. They are solids with high melting points (typically >400oC) 2. Many are soluble in polar solvents such as water. 3. Most are insoluble in nonpolar solvents. 4. Molten compounds conduct electricity well because they contain mobile charged particles (ions). 5. Aqueous solutions conduct electricity well because they contain mobile charged particles (ions). 6. They tend to be brittle and shatter when hammered. 7. The formulas are usually given in simplest ratio of elements (empirical) *COVALENT BOND – Involves the sharing of electrons between two atoms, usually between two nonmetals Characteristics of Covalent Compounds: 1. They are gases, liquids, or solids with low melting points (typically <300oC) 2. Many are insoluble in polar solvents. 3. Most are soluble in nonpolar solvents 4. Liquid and molten compounds do not conduct electricity. 5. Aqueous solutions are usually poor conductors of electricity because most do not contain charged particles. 6. The formulas are usually given in the true ratios of atoms (molecular). *Metallic Bond – formed between identical metal atoms when a “sea” of valence electrons surrounds an array of positively charged “core” ions *Characteristics of Metals: 1. They are usually crystalline solids. 2. They have a range of melting points usually dependant upon the number of valence electrons. 3. They are good conductors of electricity in both the solid and molten states. 4. They are malleable and ductile. 5. They have a luster. ALLOYS Larger metals Smaller metals Substitute in the “holes” AP Chemistry NOTES 14-2 MOLECULAR GEOMETRY TYPES OF COVALENT BONDS *Single Bond – consists of 1 pair of shared electrons forming a sigma bond (σ) (2 “p” orbitals overlap “head-to-head”) *Double Bond – consists of 2 pairs of shared electrons *1 sigma bond *1 pi bond (π) – formed when 2 “p” orbitals overlap “side-to-side” *Triple Bond – consists of 3 pairs of shared electrons *1 sigma bond 2 pi bonds RELATED FACTS CONCENRING COVALENT BOND TYPES 1. A single sigma bond is stronger than a pi bond 2. The triple bond is the shortest of the bonds and is stronger than the rest 3. The single bond is the longest of the bonds and is weaker than the rest *A single bond has a bond order of “1” *A double bond has a bond order of “2” *A triple bond has a bond order of “3” 4. Bond Dissociation Energy – the energy required to break a chemical bond (also referred to as “bond energy” for short); increases with bond order The molecular geometry is determined by the number of electron pairs on the central atom, and whether these electron pairs are bonding pairs or lone pairs (non-bonding pairs). The electron pair geometry can be determined using the table below: TOTAL ELECTRON PAIRS ON ELECTRON PAIR OR STRUCTURAL CENTRAL ATOM GEOMETRY HYBRIDIZATION 2 linear sp 3 trigonal planar sp2 4 tetrahedral sp3 5 trigonal bipyramidal sp3d 6 octahedral sp3d2 MOLECULAR GEOMETRIES Linear Trigonal Planar Tetrahedral Pyramidal Trigonal Bipyramidal Bent See-Saw T-Shaped Linear Octahedral Square Pyramidal Square Planar BOND POLARITY A polar bond forms when two atoms of different electronegativity values bond together covalently. This results in pair of electrons that is shared “un-equally” between the two atoms. The diagram below represents a hydrogen atom (on the left) bonded with a fluorine atom (on the right): hydrogen fluorine Since fluorine in the above example has a higher electronegativity value (4.0) than hydrogen (2.1), their shared electron pair is displaced more toward fluorine, resulting in a partial negative “pole” on the fluorine end and a partial positive pole on the hydrogen end – hence, a polar bond. Molecules may be polar or non-polar, depending upon their molecular structure. Water, for instance, has an asymmetrical structure that consists of two hydrogen atoms and two “lone pairs” of electrons in a tetrahedral orientation around a single oxygen atom. This gives the water molecule an uneven distribution of electric charge, resulting in a molecule that has a partial positive pole and a partial negative pole – a polar molecule: Molecules like carbon tetrachloride consist of polar bonds. However, the symmetrical structure of the molecule results in an even distribution of charge – the molecule as a whole is nonpolar in nature: In general, molecules are nonpolar if *there are no lone pairs on the central atom and *all peripheral atoms are the same Websites to view molecular geometries: http://intro.chem.okstate.edu/1314F97/Chapter9/VSEPR.html http://www.chemmybear.com/shapes.html http://people.bath.ac.uk/ch3tam/inorganic/vsepr/html/vseprmolecularshapes.htm EXAMPLES: Lewis Structure Electron Pair Molecular Polar or Drawing Molecule Geometry Geometry Nonpolar? BeCl2 BF3 CH4 NH3 H2O Lewis Structure Electron Pair Molecular Polar or Drawing Molecule Geometry Geometry Nonpolar? PCl5 SeF4 IBr3 - ICl2 SCl6 ClF5 XeF4 RESONANCE STRUCTURES AND BOND ORDER Resonance – when two or more plausible structures can be drawn for one molecule; the shared electrons are delocalized among the bonding sites Bond Order – the bond order of a molecule can be determined using the following equation: EXAMPLE: Draw the possible resonance structures of the carbonate ion and determine the bond order of the molecule. FORMAL CHARGE The use of oxidation numbers is actually artificial – more or less o bookkeeping method for keeping track of the electrons. A more realistic method is that of formal charges. The formal charge is the difference between the number of valence electrons on the free atom and the number of electrons assigned to the atom in the molecule. (A formal charge of + 2 is about as large as it gets.) (Note: The sum of the formal charges in a species must equal the total charge of the species) Formal Charge = (# of val e- on the free atom) – (# of val e- assigned in molecule) EXAMPLE: Determine the formal charge on both oxygen and chlorine in the perchlorate ion. .
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