Modern Chemistry Chapter 6 Notes by xjrx91xy17

VIEWS: 8,443 PAGES: 12

More Info
									Chapter 6: Chemical Bonding
Ch. 6

Section 1: Introduction to Chemical Bonding Atoms seldom exist as independent particles in nature Chemical Bond: A mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together As independent particles, atoms have a relatively high potential energy. Atoms act by bonding to reduce this potential energy in the form of chemical bonds. By reducing their potential energies, atoms and the compounds they form, become more stable. Types of Chemical Bonds When atoms bond, their valence electrons are redistributed in ways that make the atoms more stable. The way in which the electrons are redistributed determines the type of bonding which occurs Ionic Bonding: Chemical bonding that results from the electrical attraction between large numbers of cations and anions In purely Ionic Bonding, electrons are completely given up by the containing atom

Covalent Bonding: Chemical bonding that results from the sharing of electron pairs between two atoms In a purely Covalent Bond, the electrons are owned equally by the two atoms

Ionic or Covalent? Bonding between atoms of different elements is rarely purely ionic or covalent The degree to which bonding between two atoms of two elements is ionic or covalent can be estimated by calculating the difference on electronegativities of the two elements Bond Naming Guidelines Electronegativity Difference 1.7 or less 0 to 0.3 0.3 to 1.7 1.7 to 3.3 Electronegativity Difference Ionic Character 50% or less 0 to 5% 5% to 50% 100% Ionic Character Type of Bond Covalent Non-polar Covalent Polar Ionic Type of Bond

Bonding between two of the same atoms is Covalent Non-Polar Covalent Bond: A covalent bond in which the bonding electrons are shared equally by the bonded atoms, resulting in a balanced distribution or electrical charge

Chapter 6: Chemical Bonding
Ch. 6

Polar: A bond which has an uneven distribution of charge Polar-Covalent Bond: A covalent bond in which the bonded atoms have an unequal attraction for the shared electrons

Chapter 6: Chemical Bonding
Ch. 6

Section 2: Covalent Bonding and Molecular Compounds Many chemical compounds, including most of the chemicals that are in living things and are provided by living things, are composed of molecules 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 of the same atom, or two or more different atoms Molecular Compound: A chemical compound whose simplest units are molecules Chemical Formula: Indicates the relative number of atoms or each kind in a chemical compound by using atomic symbols and numerical subscripts Molecular Formula: Shows the types and numbers of atoms combined in a single molecular compound; chemical formula of a molecular compound Diatomic Molecule: A molecule containing only two atoms Formation of a Covalent Bond Two isolated hydrogen atoms separated at a distance which prohibits interference have a potential energy of zero Each atom approaches the other. As the hydrogen atoms near each other, their charged particles begin to interact. The nuclei and electrons of opposite atoms are attracted to one another, which results in a decrease in potential energy. However, due to the repulsion of the two nuclei and the electrons, the potential energy increases The relative strength of attraction and repulsion depends on the distance between the two atoms As the electron-proton attraction becomes stronger than the electron-electron and proton-proton attractions, the potential energy is lowered The attractive force continues to dominate until a distance is reached under which the attraction and repulsion become balanced At this point, a stable molecule forms An even closer approach results in a rise in potential energy, and in turn, a repulsion which is greater than the attraction Characteristics of a Covalent Bond At the point of balanced attraction and repulsion, a stable covalent bond forms where either shared electron occupies an overlapping orbital

Chapter 6: Chemical Bonding
Ch. 6

Bond Length: The distance between two bonded atoms at their minimum potential energy, that is, the average distance between two bonded atoms Energy is released in the forming of a covalent bond. The amount of energy released equals the difference between the potential energy of the separated atoms and that of the bonded atoms Bond Energy: The energy required to break a chemical bond and for neutral isolated atoms; usually represented as kj/mol Bond lengths and bond energies vary according to the types of atoms bonded The energy between the same two atoms also varies, depending on other formed bonds Bonding Electron Pair In Overlapping Orbitals

H _↑_ → H _↑_ 1S 1S H_↑_ → H_↓_ 1S 1S

The Octet Rule Many atoms can reach a stable state, noble-gas configuration, through covalent bonding 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 ~ See written notes for example of octet rule Exceptions to the Octet Rule Most main-group elements tend to form covalent bonds according to the octet rule, but there are exceptions

Chapter 6: Chemical Bonding
Ch. 6

Some of these exceptions are Hydrogen, and Boron Other elements can be surrounded by more than eight electrons when they combine with highly electronegative elements such as Fluorine, Boron, and Chlorine. This results in expanded valence, in which electrons are in d-orbitals Electron-Dot Notation Electron-Dot Notation: An electron configuration in which only the valence electrons of an atom of a particular element are shown, indicated by dots placed around the element’s symbol Lewis Structures Electron-dot notation can also be used to represent molecules A hydrogen molecule, H₂ is represented as H:H Unshared pair: Also called the lone pair, is a pair of electrons that is not involved in bonding and that belongs exclusively to one atom In a covalent bond, the unshared pair is represented by a dash H―H Lewis Structures: Formulas in which atomic symbols represent nuclei and inner-shell electrons, dotpairs, 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: Indicated the kind, number, arrangement, and bonds but not the unshared pairs of atoms in a molecule Single Bond: A covalent bond produced by the sharing of one pair of electrons between two atoms Multiple Covalent Bonds Double Bond: A covalent bond produced by the sharing of two pairs of electrons between two atoms; represented by two successive dashes or electron dot pairs Triple Bond: A covalent bond produced by the sharing of three pairs of electrons between two atoms; represented by three successive dashes or electron dot pairs Double and triple bonds are referred to as multiple bonds. Although these bonds are shorter, they are stronger than a single bond; Single, Double, Triple – In order of increasing strength Resonance Structures Some molecules and ions cannot be represented adequately by a single Lewis Structure Molecules in this manner alternate between their double and triple bonds

Chapter 6: Chemical Bonding
Ch. 6

These Resonance structures are represented by a Lewis structure contained within brackets and copied to the appropriate amount of times as determined by the molecular charge Resonance: Refers to bonding molecules or ions that cannot be accurately represented by a single Lewis Structure Covalent Network Bonding This bonding refers to covalently bonded compounds that do not contain individual molecules, but instead can be said to be continuous 3-D networks of bonded atoms

Chapter 6: Chemical Bonding
Ch. 6

Section 3: Ionic Bonding and Ionic Compounds Ionic Compound: 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 three-dimensional network of positive nad negative ions In contrast to a molecular compound, an ionic compound is not composed of independent, neutral units that can be isolated and examined Formula Unit: The simplest collection of atoms from which an ionic compound’s formula can be established Formation of an Ionic Compound Electron-dot notation can be used to demonstrate the changes that take place in ionic bonding
See textbook for extended elaboration on this formation

Characteristics of Ionic Bonding In Ionic Bonding, ions minimize their potential energy by combining in an orderly arrangement known as a crystal lattice Attractive Forces: Those between oppositely charged ions, and those between the nuclei and electrons of adjacent ions Repulsive Forces: Those between like-charged ions, and those between electrons of adjacent ions The three-dimensional arrangement s of ions and the strength of attraction between them vary with the sizes and charges of the ions and the numbers of ions of different charges To compare bond strengths, chemists compare the amounts of energy released when separated ions in a gas come together to form a crystalline solid Lattice Energy: The energy released when one mole of an ionic crystalline compound is formed from gaseous ions A Comparison of Ionic and Molecular Compounds The force that holds ions together in ionic compounds is a very strong attraction between positive and negative charges The forces of attraction of molecular compounds is less than those of ionic compounds The difference in strength of attraction between basic units of molecular and ionic compounds gives rise to different properties in the two types of compounds

Chapter 6: Chemical Bonding
Ch. 6

Molecular compounds have a lower melting and boiling point than ionic compounds Ionic compounds are hard, but brittle Ionic compounds are not electrical conductors in the solid state, but are in the liquid sate Polyatomic Ions Polyatomic Ion: A charged group of covalently bonded atoms The charges result from an excess or a shortage of electrons Excess: Negative Charge Shortage: Positive Charge

Lewis Structures of Polyatomic Ions If the polyatomic ion is negative, add the number of valence electrons to the number of electrons corresponding to the ion’s negative charge If the polyatomic ion is positive, subtract the number of corresponding electrons to the positive charge from the number of valence electrons

Chapter 6: Chemical Bonding
Ch. 6

Section 4: metallic Bonding Chemical bonding is different in metals than other compounds. This is evident in the occurance of 2 to 3 electrons in the outermost energy level This difference is due to a metal’s high electrical conductivity The Metallic-Bond Model The orbitals in a metal atom’s outer energy levels overlap. The electrons are delocalized, which forms a sea of electrons around the metal atoms Metallic Bonding: The chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons Metallic Properties All metals have high electrical and thermal conductivity characteristics. Metals can absorb a wide range of light frequencies due to the extremely small energy differences of metals’ many orbitals Malleability: The ability of a substance to be hammered or beaten into thin sheets Ductility: The ability of a substance to be drawn, pulled, or extruded through a small opening to produce a wire Metallic Bond Strength Metallic bond strength varies with the nuclear charge of the metal atoms and the number of electrons in the metal’s electron sea Heat of Vaporization: The amount of heat required to vaporize a metal; a measure of the strength of the bonds that hold a metal together

Chapter 6: Chemical Bonding
Ch. 6

Section 5: Molecular Geometry The properties of molecules depend not only on the bonding of atoms but also on molecular geometry The polarity and geometry of a molecule determine molecular polarity Molecular Polarity: The uneven distribution of molecular charge VSEPR Theory To determine the geometry of molecules that are not diatomic, more complex, the locations of all electron pairs must be considered VSEPR Theory: “Valence-Shell Electron Pair Repulsion.” States that repulsion between the sets of valence-level electrons surrounding an atom causes these sets to be oriented as far apart as possible Common Geometries and Their Formulas Linear Trigonal-Planar Tetrahedral AB AB₂ AB₃ AB₄ 180˚ 120˚ 90˚

VSEPR and Unshared Electron Pairs VSEPR Theory postulates that a molecule with both shared and unshared electron pairs forms a tetrahedral, but refers to only the positions of the atoms. Molecules such as this are called AB₃E molecules, where E is the unshared electron pair Water is an AB₂E₂ molecule An AB₂E molecule is formed with a central atom with two bonds and one unshared electron pair In VSEPR, double and triple bonds are treated the same as single bonds, and polyatomic ions are treated similarly to molecules Hybridization Hybridization is used to explain how the orbitals of an atom become rearranged when an atom forms a covalent bond Hybridization: The 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 the combination of two or more orbitals on the same atom

Chapter 6: Chemical Bonding
Ch. 6

Hybridization also explains the bonding and geometry of many molecules formed by Group 15 and 16 elements Hybridization of S-Orbital The 1s orbital and 1p orbital, during hybridization, produces the sp orbital The 1s orbital and one occupied p orbital form the sp² orbital The 1s orbital and three p orbitals form the sp³ orbital Intermolecular Forces The higher the boiling point, the stronger the forces between particles Intermolecular Forces: The forces of attraction between molecules Intermolecular forces vary in strength Molecular Polarity and Dipole-Dipole Forces Dipole: Created by equal, but opposite charges that are separated by a short distance
See textbook for example

Dipole-Dipole Forces: The forces of attraction between polar molecules The polarity of diatomic molecules is determined by one bond For molecules containing more than two atoms, polarity is determined by the polarity and orientation of each bond A polar molecule can induce a dipole in a non-polar by temporarily attracting its electrons Hydrogen Bonding Some Hydrogen containing compounds have unusually high boiling points 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 Hydrogen bonds are usually represented by dotted lines connecting the hydrogen bonded hydrogen to the unshared pair of electrons of the electronegative atom to which it is attracted London Dispersion Forces London Dispersion Forces: The intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles

Chapter 6: Chemical Bonding
Ch. 6

London forces act between all atoms and molecules. They are the only intermolecular forces acting among noble-gas and non-polar molecules Due to the dependability of the motion of electrons to London forces, the strength increases with the number of electrons in the interacting atoms or molecules


								
To top