Chemical Bonding • Chemical Bond- A mutual attraction between nuclei and the valence electrons of different atoms that binds the atoms together. • Most atoms are found in compounds. Atoms are rarely found alone in nature.

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Types of Bonding • • •

Ionic Bonding Transfer of electrons Atoms with opposite charges attract

Covalent Bonding Sharing of electrons. The protons of one atom are attracted to the electrons of a neighboring atom, but repelled by the protons of the other atom. When this attraction is stronger than the repulsive forces, a bond is formed.

Types of Bonding • Metallic Bonding • Bonding between the delocalized electrons between metal atoms.

Determining Types of Bonds • The difference in electronegativities in bonded atoms determines the type of bond involved. • 0-0.3 Nonpolar Covalent (evenly shared electrons) • 0.3-1.7 Polar Covalent (unevenly shared electrons) • 1.7-3.3 Ionic

Determining Types of Bonds • NaCl

• O2

• HBr

Molecular Compounds • • • • • • •

Molecular compounds are held together by covalent bonds Electrons are evenly shared in nonpolar covalent bonds Nonpolar molecules have identical charges on either end of the molecule Electrons are unevenly shared in polar covalent bonds Polar molecules have different charges on opposite ends of the molecule Percent Ionic Character is a measure of how unevenly shared the electrons are in a bond Few bonds are totally ionic or totally covalent

Molecular Compounds • • • •

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Molecular Compounds Molecules are the smallest units Water, H2O, is an example A chemical formula shows the number and types of atoms in a compound A molecular formula shows the number and types of atoms in a single molecule of a compound Diatomic molecules Two atoms of the same element covalently bonded Only 7 elements are diatomic H, N,O,F,Cl,Br,I

Energy • Chemical potential energy • Repulsion and attraction between electrons of one atom and protons of another • Bond Energy • The energy required to break a bond in kJ/mol

Energy

• Bond Length-Average distance between two bonded atoms • High energy=strong bond=short bond=stable • Low energy=weak bond=long bond=unstable

Single, Double, and Triple Bonds • •

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A single bond is a sigma bond A double bond has a sigma bond in the center and a pi bond that wraps around the sigma bond Double bonds are stronger and shorter than single bonds A triple bond has a sigma bond in the center and two pi bonds that wrap around the sigma bond Triple bonds are stronger and shorter than single or double bonds

Covalent Networks •

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Covalent networks lack a single cohesive molecule The entire structure is a single molecule Diamond is an example of carbon atoms in a covalent network Allotropes are structures of the same element whose structures have different properties Allotropes of carbon are graphite, diamond, buckminsterfullerene, carbon nanotubes, and graphene

The Octet Rule • Compounds form when atoms gain, lose, or share electrons so that each bonded atom has 8 valence electrons • Exceptions to the Octet Rule • Some atoms require fewer than 8 valence electrons • Some can bond with more than 8 electrons by using d-block electrons

Lewis Structures • Lewis Structures are 2-D representations of the bonding structures of compounds • In order to draw a Lewis Structure, draw electron dot structures for each type of atom involved

Lewis Structure (Covalent) • • •

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Determine the central atom. It is the least electronegative atom A good rule of thumb is that it will be the atom with the most empty slots on a dot diagram Write the symbol for the central atom with its dots and draw a circle around them Surround it with symbols of the other atoms in the compound Move electrons within their circles into overlapping regions when an atom does not have an octet of electrons

Lewis Structure (Covalent) • • •

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Electrons must be shared in pairs 2, 4, or 6 electrons can be shared between two atoms Although each atom in a bond often donates an electron, sometimes, one atom can donate all of the electrons in a bond Only pairs of electrons can be left over on atoms (unless the atom is hydrogen) These are called lone pairs of electrons If more than one correct Lewis structure is possible, the molecule has resonance

Ionic Compounds • Ionic compounds are formed due to electrical attraction from a transfer of electrons • Most are crystalline solids • Chemical formulasimplest ratio of atoms • Formula Unit- smallest collection of atoms from which a formula can be established

Ionic Compounds • Lewis Structures can also be drawn for ionic compounds

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Ionic compounds have a crystal lattice structure A crystal lattice is the bonding framework that determines the arrangement of atoms in a crystal Lattice Energy- the energy released when one mole of an ionic crystalline compound is formed from gaseous ions A negative value indicates a release of energy

Polyatomic Ions •

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Polyatomic ions are covalently bonded groups of atoms that act as a single ion that will bond ionically When drawing Lewis structures for polyatomic ions, add a number of electrons equal to the charge to the central atom for anions and remove a number of electrons equal to the charge for cations Lewis structures for polyatomic ions are drawn inside of brackets with the charge at outside of the upper right bracket

Ionic vs. Covalent Properties • • • • • • •

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Ionic Properties High melting point High boiling point Mostly solids Hard and brittle Poor conductors when solid Decent conductors when molten Soluble in water Have crystalline structures

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Often have low melting points due to weak forces between molecules Often have low boiling points due to weak forces between molecules Some are gases Some are liquids Some are solids Softer than ionic compounds Poor conductors Sometimes form crystals Can form covalent networks

Metallic Bonding • Outer energy levels overlap and loosely held electrons move fairly freely between atoms • Delocalized electrons form a “sea of electrons” • This allows for conductivity of heat and electricity • Electrons absorb and reemit energy (light), making metals shiny

Metallic Bonding • Malleability and ductility work because metallic bonding is the same in all directions in the solid • Strength of metallic bonds is measured by observing the heat of vaporization • The heat of vaporization is the amount of heat needed to vaporize the metal and is measured in kiloJoules/mole.

Molecular Geometry • Properties of molecules depend partially on their shape, or molecular geometry • Molecular polarity, or the distribution of electrons (even/uneven sharing) affects forces between molecules

Molecular Geometry • •

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Atoms and valence electrons in compounds will repel each other as much as is possible VSEPR (valence shell electron pair repulsion) VSEPR diagrams show 3-D representations of molecules Unshared electrons repel each other more than bonding pairs do and alter the angles of bonds more than atoms alone In the diagram here, the lone pairs cause angles smaller than the 109.5 angle found in a true tetrahedral compound Double and triple bonds are treated in the same manner as single bonds

Molecular Geometry •

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A Lewis Structure is needed to draw a VSEPR diagram VSEPR diagrams do not show resonance of molecules VSEPR diagrams can be thought of in terms of an ABE format A=central atom B=atoms bonded to central atom E=lone pairs of electrons These can be found by looking at the Lewis Structure

VSEPR

VSEPR

Hybridization • Hybridization is a mixing of orbitals of different energies to create orbitals of equal energies • A similar idea follows mixing hot and cold water to produce warm water

Intermolecular Forces •

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Forces acting between molecules Generally weaker than the forces making up true bonds Boiling point is a measure of the attractive force between molecules The strongest forces exist between polar molecules Molecules with uneven electron distribution create dipoles The direction of a dipole is from the positive end toward the negative end

Intermolecular Forces • Dipole-dipole forces • Forces of attraction between polar molecules • Uneven distribution of electrons among atoms causes one end to be more negative than the other

Intermolecular Forces • Hydrogen bonding • Not a “real” type of bonding • H atoms bonded to a very electronegative atom are attracted to the lone pair of electrons of a nearby molecule • Responsible for some of the properties of water, such as surface tension and boiling point

Intermolecular Forces • •

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London Dispersion Forces Forces of attraction between atoms because of the constant motion of electrons As electrons move within an atom, they cause regions of uneven charge within the atom, creating instantaneous, temporary dipoles These dipoles attract dipoles created in other atoms and molecules