chapter6_soag

=//Chapter 6//=

//Section 1//
A **chemical bond** is a mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together.
 * **Ionic bonds** are chemical bonds that result from the electrical attraction between cations(positive) and anions(Inegative). In these bonds, atoms completely give or lose electrons to other atoms.(Electronegativity difference is greater than or equal to 2.0)
 * **Covalent bonding** results from the sharing of electron pairs between two atoms.
 * **Non-polar covalent (pure covalent)** - the bonding electrons are shared equally by the bonded atoms, resulting in a balanced distribution of electrical charge. (Electronegativity difference is less than or equal to 0.4)
 * **Polar covalent** - unequal sharing of electrons, gives rise to partial changes. (Electronegativity difference is between 0.4 and 2.0)
 * The higher the electronegativity difference then the higher the percentage that the compound becomes ionic.
 * When thinking of a Nonpolar-covalent bond think of two equal strength people holding hands, it would take the same amount of energy to shifts either one. But for a Polar-covalent bond think of a father holding his young child's hand.  It would be much eaiser to shift the young child because he or she would be much weaker than the father.

===//Section 2//===

A **molecule** is a neutral group of atoms that are held together by covalent bonds. ·Molecules are able to exist on there own · They may consist of two or more atoms of the same element, or of different elements · A **diatomic molecule** consists of only two atoms of only one type of element

A **molecular compound** is a chemical compound whose simplest units are molecules. · The composition of a compound is given by its **chemical formula**, which indicates the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts. ü The chemical formula of a molecular compound is called the **molecular formula**, which shows the types and numbers of atoms combined in a single molecule of a molecular compound.

· Most atoms have lower potential energy when they are bonded to other atoms than they have as they are independent particles. · As atoms move towards each other, the **nuclei** and electrons are attracted to each other. The two nuclei repel each other and the two electrons repel each other. Attraction between particles corresponds to a decrease in potential energy of the atoms, while repulsion corresponds to an increase in potential energy. · When the atoms first notice each other’s presence, the electron-proton attraction is stronger than the electron-electron and proton-proton repulsions, so the atoms are drawn together and their potential energy is lowered. The total energy continues to be lowered as they approach each other until a distance is reached at which the repulsion between the like charges equals the attraction of the opposite charges. At this point, a stable molecule forms. Diagram of a Polar-Covalent Bond · The **bond length** is the average distance between the nuclei of two covalently bonded atoms ü Vary with the types of atoms being combined · While forming a covalent bond, atoms release energy as they change from isolated individual atoms to parts of a molecule (the amount of potential energy released equals the difference between the potential energy at the zero level and that of bonded atoms). The same amount of energy must be added to separate the bonded atoms. · **Bond energy** is the energy required to break a chemical bond and form neutral isolated atoms. Measured in kJ/mol (indicates the energy required to break one mole of bonds in isolated molecules) Vary with the types of atoms being combined
 * Formation of** **Covalent Bonds**
 * Characteristics of Covalent Bonds**

· 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. Exception: some elements can be surrounded by less than 8 electrons. Exception: some elements can be surrounded by more than 8 electrons when they combine with the highly electronegative elements fluorine, oxygen and chlorine. In these bonds, electrons in //d// orbitals as well as in //s// and //p// orbitals are involved.
 * The Octet Rule**

· 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.
 * Electron-Dot Notation**

For example: **F** Also known as **Lewis Dot Structure**

· Electron-dot notation used to represent molecules instead of individual atoms. · 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. For example: H**:**H ßcalled a **single bond** (a covalent bond in which one pair of electrons is shared between two atoms) · A **lone pair** is a pair of electrons that belongs to one atom and does not bond at all with the other atom · A **structural formula** indicates the kind, number, arrangement, and bonds but not the unshared pairs of the atoms in a molecule. For example: F—F The amount bonds depends on which column the element is in. Ex. Carbon is in the fourth collumn so Carbon would have four bonds.
 * Lewis Structures**

· 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. ü Carbon, Nitrogen and Oxygen commonly form them ü Have greater bond energies and are shorter than single bonds · **A triple bond** is a covalent bond in which three pairs of electrons are shared between two atoms. ü Even stronger and shorter than double bonds
 * Multiple Covalent Bonds**

· **Resonance** refers to bonding in molecules or ions that cannot be correctly represented by a single Lewis structure. · To indicate resonance, a double-headed arrow is placed between a molecule’s resonance structures.
 * Resonance Structures (or resonance hybrids)**

===Section 3===

An **ionic compound** is composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal. · Ionic compounds exist as **crystalline solids** which are three-dimensional networks of positive and negatives ions mutually attracted to one another Chemical formulas for ionic compounds show the ratio of ions present in a sample. · A **formula unit** is the simplest collection of atoms from which an ionic compound’s formula can be established.

· **Crystal lattice** is when ions minimize their potential energy by combining in an orderly arrangement · Attractive forces at work within ionic crystal are those between oppositely charged ions and those between the nuclei and electrons of adjacent ions. · Repulsive forces include those between like-charged ions and those between electrons of adjacent ions · The strengths of attraction and three-dimensional arrangements of ions differ with the number of ions of different charges and the sizes · To compare bond strength, compare the amounts of energy released when separated ions in a gas come together to form a crystalline solid · **Lattice energy** is the energy released when one mole of an ionic crystalline compound is formed from gaseous ions
 * Characteristics of Ionic Bonding**

Here are some common examples of lattice enrgy found in some compounds:
 * Compund   Lattice Energy|
 * NaCl____      -787.5        |
 * NaBr   ____   -751.4        |
 * CaF2 ____     -2634.7       |
 * LiCl ____     -861.3        |
 * LiF  ____     -1032         |
 * MgO  ____     -3760         |
 * KCl  ____     -715          |

Ionic Compounds Molecular Compounds · Strong over-all reaction between positive ● Covalent bond of the atoms making and negative charges up each molecule are strong · Forces of attraction much stronger between ● Forces of attraction much weaker molecules between molecules
 * Comparison of Ionic and Molecular Compounds**

· Ionic compounds are hard but brittle · In the solid state, ions ant move therefore they cannot conduct electricity · In molten state, the ions can move freely to carry electrical current making them **electrical conductors**

· **Polyatomic ions** are a charged group of covalently bonded atoms · Polyatomic ions combine with ions of opposite charge to form ionic compounds · Charge of compounds result form left over electrons (negative charge) or shortage of electrons (positive charge)
 * Polyatomic Ions**

===//Section 4//===

· Valence electrons of the atoms that make up metals move around a lot, so they conduct electricity very well. · **Metallic Bonding** is the chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons.
 * Metallic Bonding**

· High electrical and thermal conductivity · Strong absorbers and reflectors of light because they have many orbitals that are very close together. When the electrons fall from higher to lower energy levels, they emit light. · **Malleability** is the ability of a substance to be hammered into or beaten thin sheets. · **Ductility** is the ability of a substance to be drawn, pulled or extruded though a small opening to produce a wire
 * Metallic Properties**

· Varies with the charge of the metal atoms and the number of electrons in the metal · One way to test bond strength is to see how much heat is required to vaporize the metal · The **enthalpy of vaporization** is the amount of energy absorbed as heat when a specified amount of a substance vaporizes at constant pressure. m
 * Metallic Bond Strength**

===Section 5===

The **molecular Geometry** is the three-dimensional arrangement of a molecule
 * Molecular Geometry**
 * The uneven distribution of molecular charge is the **molecular polarity** which is determined by the polarity of each bond and the geometry of the molecule

ü For example: F:Be:F The fluorine atoms are 180 **۫** apart because that is as far away from each other as possible
 * VSEPR Theory**
 * VSEPR = **V**alence **S**hell **E**lectrons**-P**air **R**epulsion
 * **VSEPR theory** states that repulsion between the sets of valence-level electrons surrounding an atom causes these sets to be oriented as far apart as
 * Shared pairs will be as far away from each other as possible


 * VESPR Theory and Unshared Electron Pairs**
 * Lone pairs occupy space around the central atom just as the bonding pairs
 * For Example: Ammonia (NH3) There are three bonds and 1 lone pair
 * The shape of the molecule is determined by the position of the atoms not the lone pairs
 * Unshared electron pairs repel electrons more strongly than bonding electron pairs


 * Hybridization**
 * **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
 * Orbitals of equal energy produced by the combination of two or more orbitals on the same atom are **hybrid orbitals**
 * The number of hybrid orbitals produced is the number of orbitals combined
 * Each electron pair can forma single bond


 * Intermolecular Forces**
 * The forces of attraction between molecules are **intermolecular forces**
 * The higher the boiling point, the stronger the forces are between molecules


 * Molecular Polarity and Dipole-Dipole Forces**
 * **Dipole** is created by equal but opposite charges that are separated by a short distance (represented by an arrow with its head pointing toward negative pole and crossed tail at positive pole )
 * Stronger intermolecular forces exist between polar molecules
 * **Dipole-dipole forces** are the forces of attraction between polar molecules
 * For molecules with two or more atoms, molecular polarity depends on the polarity and the orientation of each bond


 * Hydrogen Bonding**
 * **Hydrogen bonding** are the intermolecular forces in which a hydrogen atom that is bonded to a highly electronegative atom is attracted to am unshared pair of electrons of an electronegative atom in a nearby molecule
 * In compounds between hydrogen atoms and fluorine, oxygen, or nitrogen atoms there are large electronegativity differences making the bonds connecting them highly polar


 * London** **Dispersion Forces**
 * The intermolecular attractions resulting from constant motion of electrons and the creation of instantaneous dipoles are **London** **dispersion forces.**
 * London dispersion forces are the only intermolecular forces acting among noble gases although they are weak
 * London forces increase with increasing atomic or molar mass
 * Everything has London Dispersion Forces
 * How To Stay Safe in The lab when Working with combining and spliting Molecules**
 * Always wear goggles and aprons
 * Make sure directions are followed
 * Don't Do Dumb Stuff
 * Don't Inhale any mixtures or compounds in the lab
 * Do listen to your instructer when in the lab.
 * If you have long hair, pull it back with a rubber band and keep it away from your experiments.
 * Do not bring food or drink into the lab


 * Practice Your Knowledge**
 * How do hybridization models help explain the way they are formed?
 * Draw the molecular geometry of ONF and SF6.
 * What types of atoms form ionic, covalent, and metallic bonds?
 * Why do most atoms tend to chemically bond to other atoms?

media type="youtube" key="L_zQGk8oBlc" height="344"
 * Here is a Simple Video for You to Understand Bonding A little More**