chapter6_mwty

__**[[image:http://www.rockwood.k12.mo.us/eurekahs/academics/mcilwee/VSEPR.jpg width="451" height="262" align="right"]]6-Section 1** **Introduction to Chemical Bonding**__
Most atoms are at relatively high potential energy. Atoms are less stable by themselves than combined with other atoms. By bonding with other atoms it decrease the potential energy, thus makes it more stable. Atoms’ valence electrons are redistributed when they bond, in order to make it more stable. Main-group metals lose electrons to form cations, and nonmetals gain electrons to form anions. By calculating the differences in the elements’ electronegativities, we can determine the bonding of atoms.
 * Chemical Bond-** //is a mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together.//
 * Types of Chemical Bonding**
 * Ionic bonding** is //a chemical bonding that result from the electrical attraction between cations and anions//. In ionic bonding, atoms completely give up electrons to other atoms in order to be stabilized. Another chemical bonding is covalent bonding.
 * Covalent bonding** //results from the sharing of electron pairs between two atoms//. In covalent bond, electrons are shared by atoms.
 * Nonpolar-covalent bond** //is a covalent bond in which the bonding electrons are shared equally by the bonded atoms, resulting in a balanced distribution of electrical charge.// Nonpolar-covalent bond’s electronegativity differences are between 0 and 0.3. Bonding that have electrons attracted by the ore-electronegative atom are polar. **Polar** //means that they have an uneven distribution of charge//. A **polar-covalent bond** //is a covalent bond in which the bonded atoms have an unequal attraction for the shared electrons.// Their electronegativity differences fall from 0.3 to 0.7. Atoms in polar-covalent that are partial positive charged uses d+and partial negative chargedd–.

__**6-Section 2** **Covalent Bonding and Molecular Compounds**__
A **molecule** //is a neutral group of atoms that are held together by covalent bonds.// A molecule of a chemical compound may consist of two or more atoms of the same element or two or more different atoms. **Molecular compound** //is a chemical compound whose simplest units are molecules//. **chemical formula** //indicates the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscript//s. A **molecular formula** //shows the types and numbers of atoms combined in a single molecule of molecular compound.// A //diatomic molecule is a molecule containing only two atoms//. Consider two atoms approaching each other. Each atom’s particles begin to interact. The nuclei and electrons are attracted to each other, thus decrease in the total potential energy of the atoms. But, the nuclei and electrons repel each other, and an increase in potential energy. The attractive force continues until a final distance is reached. //The distance between two bonded atoms at their minimum potential energy, that is, the average distance between two bonded atoms, is the bond length.// Atoms release energy as they change from individual atoms to parts of a molecule. This same amount of energy must require separating the bonded atoms. Noble-gas atoms are very stable because their outer s and p orbitals are filled by a total of eight electrons. By sharing electrons with other, atoms can fill their outermost s and p orbitals through bonding. The bonding follows a certain rule, the octet rule, which states that; Chemical compounds tend to form so that teach atom, by gaining, or losing, or sharing electrons, has an octet of electrons in its highest occupied energy level. The Lewis structure of ozone is a resonance structure. **Resonance** //refers to bonding in molecules or ions that cannot be correctly represented by a single Lewis structure.//
 * Formation of a Covalent Bond**
 * Characteristics of the Covalent Bond:**
 * Bond energy** //is the energy required to break a chemical bond and form neutral isolated atom.//
 * The Octet Rule:**
 * Electron-Dot Notation** //is an electron-configuration notation in which only the valence electrons of an atom of a particular element are show, indicated by dots placed around the element’s symbol.//
 * Lewis structures** //are formulas in which atomic symbols represent nuclei and inner-shell electron pairs in covalent bonds, and dots adjacent to only on atomic symbol represent unshared electrons.// A **structural formula** indicates the kind, number, arrangement, and bonds but not the unshared pairs of the atoms in a molecule. **Single bond** //is a covalent bond in which one pair of electrons is shared between two atoms.//

Lewis Structure: 1)Determine the total number of valence electrons 2)Determine the number of bonds each atom will form 3)Determine the central atom - if you only have one of an element, it is almost always the central atom -if Carbon is present, it is the central atom 4)Perpose a structure which allows for each atom to have its desired number of bonds. (COUNT ELECTRONS) 5)Distribute remaining electrons as lone pairs to satisfy octets. Count electrons again.

Example: H2O: Step 1: Arrange the atoms: H O H Step 2: Count up total valence electrons: 2(1) + 6 = 8 e- total Step 3: Draw single bonds between central atom and surrounding atoms: H - O - H Step 4: Place remaining electrons, in pairs, around appropriate atoms; start with outer atoms. 8 - 4 = 4 e- left; therefore 2 lone pairs to add: -In this example, the 2 lone pairs must go around O, since H never gets lone pairs à H-O-H (lone pairs go on top and below the 'O'.)

Step 1: Arrange the atoms: B, Cl Cl, Cl Step 2: Count up total valence electrons: 3(7) + 3 = 24 e- total Step 3: Draw single bonds between central atom and surrounding atoms: -->Cl Cl B Cl
 * Example 2:**

Step 4: Place remaining electrons, in pairs, around appropriate atoms; start with outer atoms. 24 − 6 = 18 e- left; therefore 9 lone pairs to add: (* means 3 surrounding lone pairs)



Step 5: Make sure all atoms that need octets have octets: The Cl’s have octets, so they’re okay; Step 5: Make sure all atoms that need octets have octets: done in Step 4
 * B is special element that can have an incomplete octet, so B only needs 6.**

__**6-Section 3** **Ionic Bonding and Ionic Compounds**__
An **ionic compound** //is composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal.// **Formula unit** //is the simplest collection of atoms from which an ionic compound’s formula can be established//. In an ionic crystal, ions minimize their potential energy by combining in an orderly arrangement known as a crystal lattice. Crystal structures of compounds show how the atoms are arranged. **Lattice energy** //is the energy released when one mole of an ionic crystalline compound is formed from gaseous ion.//
 * Formation of Ionic Compounds**

__**6-Section 4** **Metallic Bonding**__
Due to the unique properties of metals, their chemical bonding is different than in ionic, molecular, or covalent compounds.
 * Metallic Bonding** //is the chemical bonding that result from the attraction between metal atoms and the surrounding sea of electrons//.
 * Malleability** //is the ability of a substance to be hammered or beaten into thin sheets.//
 * Ductility** //is the ability of a substance to be drawn, pulled, or extruded through a small opening to produce a wire.//
 * Dissolve** //becomes part of a solution//
 * Dissociate** //splitting of an ionic compound into its parts (cation and anion)//

__**6-Section 5** **Molecular Geometry**__
VSEPR means “valence-shell, electron-pair repulsion”. **VSEPR theory** //states that repulsion between the states of valence-level electrons surrounding an atom caused these sets to be oriented as far apart as possible. **In essence, the VSEPR theory states that electron pairs of electrons repel each other, therefore for a molecule to be at it’s most stable – electron pairs must be as far from each other as possible.** // AB2 molecule is linear with 180 degree between each atom. AB3 molecule is called trigonal-planar, and each atom bears 120 degree angles. AB4 molecule is called tetrahedral, and the bond angles are 109.5 degree. Hybridization model is used to explain how the orbitals of an atom become rearranged when the atom forms covalent bonds. **Hybridization** //is the mixing of two or more atomic orbitals of similar energies on the same atom to produce new hybrid atomic orbitals of equal energies.// Fore example, taking ethane, CH4. There are only 2 electrons in the 2p orbital, thus not filling in the entire orbital. We can then take 2s and 2p to combine the orbitals together, to create sp3. **Hybrid orbitals** //are orbitals of equal energy produced by the combination of two or more orbitals on the same atom.// Dipole: created by equal but opposite charges that are separated by a short distance. Hydrogen bonding: 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.
 * Hybridization**

· VSEPR theory is used to predict the shapes of molecules based on the fact that electron pairs strongly repel each other. · Hybridization theory is used to predict the shapes of molecules based on the fact that orbitals within an atom can mix to form orbitals of equal energy. · Intermolecular forces include dipole-dipole forces and London diepersion forces. Hydrogen bonding is a special case of dipole-dipole forces.

Practice Questions: True or False? if the underlined word is false, correct the answer. 1. A chemical bond results from the mutual attraction of the nuclei of atoms and __electrons__. 2. The electrons involves in the formation of a chemicals bond are called __valence electrons__. 3. As independent particles, atoms are at relatively __low__ potential energy. 4. Atoms are __more/less__ stable when they are combined. 5. The chemical bond formed when two atoms share electrons is called a __covalent bond__. 6. If the atoms that share electrons have an unequal attraction for the electrons, the bond is called __ionic__. 7. A molecule is a __neutral group of atoms held together by covalent bonds.__ 8. Bond energy is __the energy required to break a chemical bond.__ 9. The elements of the __alkali metal__ group satisfy the octet rule without forming compounds. 10. Bonding in molecules or ions that cannot be correctly represented by a single Lewis structure is __double bonding__. 11. The chemical formula for an ionic compound represents the __simplest ratio of the combined ions that balances total charges.__ 12. The ions in an ionic compound are organized into a __crystal__. 13. Compared with ionic compounds, molecular compounds have __higher__ boiling points. 14. A polar molecule contains __a region of positive charge and a region of negative charge.__ 15. The mixing of two or more atomic orbitals of similar on the same atom to produce new orbitals of equal energies is called __dipole-dipole interaction__.

Answers: 1) True 2) True 3) False, High 4) More 5) True 6) False, Polar 7) True 8) True 9) False, Noble Gas 10) False, Resonance 11) True 12) True 13) Flase, Lower 14) True 15) False, Hybridazation




 * Sources:** Davis, Raymond E., Mickey Sarquis, Regina Frey, and Jerry L. Sarquis. __Modern Chemistry__. Austin, Texas: Holt, Rinehart and Winstron, 2006. ii-949, **[], [|http://www.scottsdalecc.edu/chemistry/docs/lewdstps.pdf,]http://tinyurl.com/mp95bq **