Chapter 6 Ellen Duong Jake Macneal Mikelanxhelo Novruzaj

Chapter 6 Ellen Duong, Jake Macneal, Mikelanxhelo Novruzaj, Enxhi Rrapi, Weijia Wang

Vocabulary Chemical bond: mutual electrical attraction between nuclei and valence electrons of different atoms that binds the atoms together Ionic Bonding: chemical bonding that results from the electrical attraction between large numbers of cations and anions Covalent Bonding: results from the sharing of electron pairs between two atoms Nonpolar-covalent bond: covalent bond in which bonding electrons are shared equally by bonded atoms resulting in a balanced distribution of electrical charge Polar: have an uneven distribution of charge Polar-covalent bond: covalent bond in which bonded atoms have an unequal attraction for the shared electrons Molecule: a neutral group of atoms held together by covalent bonds

Molecular compound: A chemical compound whose simplest units are molecules Chemical Formula: indicates relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts Molecular Formula: shows types and numbers of atoms combined in a single molecule of molecular compound Diatomic Molecule: molecule containing only two atoms Bond length: distance between two bonded atoms at their minimum potential energy, that is, the avg distance between two bonded atoms Bond energy: the energy required to break a bond and form neutral isolated atoms Octet rule: chemical compounds tend to form so that each atom has an octet (8) of electrons in its highest occupied energy level Unshared/ Lone Pair: a pair of electrons not involved in bonding and belongs exclusively to one atom

Lewis Structures: 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. Structural Formula: indicates kind, number, arrangement, and bonds, but not the unshared pairs of atoms in a molecule Single Bond: covalent bond produced by sharing of one pair of electrons between two atoms Double Bond: a covalent bond produced by the sharing of two pairs of electrons between two atoms Resonance: bonding in molecules or ions that cannot be correctly represented by a single Lewis structure Ionic Compound: composed of positive and negative ions that are combined so that the number of positive and negative charges are equal Formula Unit: simplest collection of atoms from which an ionic compound’s formula can be established

Lattice energy: energy released when one mole of an ionic crystalline compound is formed from gaseous ions Polyatomic Ion: charged group of covalently bonded atoms Metallic Bonding: chemical bonding resulting from the attraction between metal atoms and surrounding sea of electrons Malleability: ability of a substance to be hammered or beaten into thin sheets Ductility: ability of a substance to be drawn, pulled or extruded through a small opening to produce a wire Molecular Polarity: uneven distribution of molecular charge VSPER theory (Valence shell electron repulsion): electrons want to be as far away from each other as possible Hybridization: mixing of two or more atomic orbitals of similar energies on the same atom to produce new orbitals of equal energies

Hybrid Orbitals: orbitals of equal energy produced by combination of two or more orbitals in the same atom Intermolecular Forces: forces of attraction between molecules Dipole: created by equal but opposite charges that are separated by a short distance Dipole-dipole forces: forces of attraction between polar molecules Hydrogen Bonding: Intermolecular force in which a hydrogen atom that is bonded to a highly electronegative atom is attracted to an unshared pair or electrons of an electronegative atom in a nearby molecule London Dispersion Forces: intermolecular attractions resulting from the constant motion of electrons and creation of instantaneous dipole.

Section 1 Introduction to Chemical Bonding

When atoms bond valence electrons are redistributed and the way that they are redistributed determines the type of bond they have. a. In pure ionic bonding atoms completely give up electrons to other atoms. b. In pure covalent bonds the shared electrons are owned equally by each atom.

Ionic or Covalent a. Whether or not a bond is ionic or covalent can be calculated by the difference in the elements’ electronegativities. a. A bond is rarely pure and it falls between ionic or covalent so it depends on how strong the atom of the element attracts the electrons. b. ii. Covalent a. Bonding between atoms with an EN difference of 1. 7 or less has an ionic character of 50% or less which means it is a covalent bond. b. Bonding between atoms of the same element is covalent c. Bonds with 0 -5% ionic character or 0 -. 3 EN difference are considered non-polar covalent bonds. d. Bonds with 5 -50% ionic character or. 3 -1. 7 EN difference are considered polar.

Section 2 Covalent Bonding and Molecular Compounds

The composition of a compound is given by its chemical formula.

Formation of covalent bonding a. Nature favors chemical bonding because most atoms are at lower potential energy when bonding than when independent. b. When two hydrogen atoms are away from each other their potential energy is 0 when they get closer their charged particles began to interact and they become attracted which lowers the total potential energy of the atoms and the two nuclei also repel each other, while the electrons repel each other resulting in an increase of potential energy. c. Relative strength of attraction and repulsion between charged particles depend on distance separating atoms a. When atoms first “See” each other the electron-proton attraction is stronger than electron-electron, and proton repulsions.

Covalent Bond Characteristics When forming a covalent bond energy needs to be released. The amount of energy released equals the difference between potential energy (when atoms are separated at 0) and the bonded atoms. The same amount of energy must be added to separate bonded atoms.

The Octet Rule a. Noble gas atoms exist independently in nature. b. Exceptions to octet rule a. Most main group elements form covalent bonds according to octet rule. b. When bonding with fluorine, oxygen, and chlorine elements tend to have more than 8 electrons.

Resonance Structures a. There can be more than one way to depict a molecule with a Lewis Structure. b. Signs to look out for: a. if there is a single bond a double bond in the molecule that connect the same two types of elements. a. ex/ [CO 3] 2 -

Section 3 Ionic Bonding and Ionic Compounds

Ionic Bonds a. Ionic Bonds do not form molecules. a. Chemical formulas for Ionic Compounds represent the simplest ratio of ion types (formula unit) b. the formula unit is made up of positively charged ions (cations) and negatively charged ions (anions)

Ionic Bonds cont. a. Ionic Compounds are arranged in a crystal lattice to minimize their potential energy. a. the distances between ions and their arrangement in a crystal represent a balance among all forces (negative and positive). b. The three dimensional arrangements of ions and the strengths or attraction between them vary with the sizes and charges of the ions and the numbers of ions of different charges. c. Lattice energy is used in comparing bond strengths in ionic compounds. a. The larger amount of energy released the stronger the bond b. Lattice energy is negative since it is released

Ionic vs. Covalent Bonds Ionic Covalent

Section 4 Metallic Bonding

Bonding of Metals a. Within a metal, the vacant orbitals in the atoms’ outer energy levels overlap. a. overlapping allows outer electrons of atoms to roam freely throughout the entire metal. a. creates “sea of electrons” the electrons are said to be delocalized (do not belong to any one atom but move freely about the metal’s network of empty atomic orbitals.

Metal Properties a. malleable, ductile, and luster. b. Bond strength a. depends on nuclear charge of the metal atoms and the number of electrons in the metal’s electron sea. b. heat of vaporization a. amount of heat required to turn solid, bonded metal atoms into gaseous individual atoms.

Section 5 Molecular Geometry

Valence-Shell, Electron-Pair Repulsion (VSEPR) Theory a. basically, electron pairs want to be as far away from each other as possible because they repel each other. a. i. note: when there are unshared electron pairs on the central atom (electrons that aren’t in bonds) they also repel other electron pairs. Unshared electron pairs results in a bent/angular molecule. b. Memorize the chart on pg 186. a. note: the “E” stands for an unshared electron pair. c. Be able to predict the shape of a molecule from the Lewis Structure.

Hybridization a. when s and p orbitals join b. please go to page 187 -188 of the textbook.

Intermolecular Forces a. force of attraction between molecules. b. varies in strength, but typically weaker than intramolecular forces (bonds inside a molecule)

IF: Molecular Polarity and Dipole-Dipole Forces a. Basically, being able to draw a line through a molecule splitting it in half and showing that one side is slightly positive and the other side is slightly negative. b. ex/ a. The “+” side of them arrow shows the positive end and the “>” shows the negative end. c. How do we know which side is positive and which side is negative? a. Compare the electronegativities. The element with the higher electronegativity (-3. 4 as oppose to -1) is the negative end. Thus making the other side positive. d. Which molecules have dipole-dipole moment? a. Diatomic molecules with a polar bond. b. 2. Some polyatomic molecules IF their polar bonds do not cancel each other out.

Dipole-Dipole Cont. Which molecules can’t have a dipole moment? a. Some polyatomic molecules with polar bonds. o molecules with identical bonds. § ex/ linear with 2 identical bonds , CO 2 § O=C=O (+)---(><)---(+) That doesn’t work.

Hydrogen Bonding a. 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 a. Water. H 2 O b. represented by dotted lines that connects the molecule’s unshared electron pair to the hydrogen of a different molecule nearby.

London Dispersion Forces a. temporary dipole moment. b. the intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles. c. act between all atoms and molecules d. the only molecular forces acting among noble-gas atoms and non-polar molecules. e. dependent on motion of electrons. f. strength increases a. with the number of electrons and in interacting atoms are molecules b. with increasing atomic or molar mass.

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