Chemical Bonding Chapter 6

Chemical bond A mutual electrical attraction between the nuclei and valence electrons of different atoms that binds them together Chemistry chapter 6

Types of bonds Ionic electrons are transferred from one atom to another, forming ions Electrical attraction of large numbers of cations and anions holds the compound together Covalent – electron pairs are shared between the atoms Chemistry chapter 6

Ionic Bonds Ions get very close together and are attracted to each other. Example: NaCl Cl- Na+ Chemistry chapter 6

Covalent Bonds The atoms share electrons, and their electron clouds overlap. Example: carbon monoxide O C Chemistry chapter 6

Electronegativity difference Can be used to predict how two atoms will react with each other. Has no units, because it is a comparison. When it is high, the bond is ionic. When it is low, the bond is covalent. Chemistry chapter 6

Ionic – covalent continuum Real bonds are usually sort of ionic and sort of covalent. See figure 6-2 on page 162. Difference of 1.7 or less is considered covalent 50% ionic character or less Difference of more than 1.7 is considered ionic More than 50% ionic character Chemistry chapter 6

Nonpolar-covalent bond Bonding electrons are shared equally Balanced distribution of electrical charge 0 to 5% ionic character 0 to 0.3 electronegativity difference Chemistry chapter 6

Polar Having uneven charge distribution Chemistry chapter 6

Polar-covalent bond The bonded atoms have unequal attraction for the shared electrons 5 to 50% ionic character 0.3 to 1.7 electronegativity difference Chemistry chapter 6

Polar and nonpolar bonds Electron density is greater around Cl than H because it has greater electronegativity d+ and d- indicate partial charge Electron cloud drawings Chemistry chapter 6

Comparing bond types Comparison between ionic and covalent bonding Chemistry chapter 6

Discuss Page 163 Sample problem Practice problem Section review 1 and 2 Chemistry chapter 6

Molecule A neutral group of atoms that are held together by covalent bonds A single molecule is capable of existing on its own May consist of atoms of the same element or different elements Chemistry chapter 6

Molecular compound Compound whose simplest units are molecules Chemistry chapter 6

Chemical formula Uses atomic symbols and subscripts to indicate the relative numbers of atoms of each kind in a chemical compound. Chemistry chapter 6

Molecular formula Shows the types and numbers of atoms in a single molecule of a molecular compound H2O O2 C12H22O11 Chemistry chapter 6

Diatomic molecule Contains only two atoms Chemistry chapter 6

Bonding Bonding makes atoms be at lower potential energy levels, which is favored by nature. Attractive forces (nucleus-electron) are balanced by repulsive forces (nucleus-nucleus and electron-electron). Chemistry chapter 6

Bond length Average distance between bonded nuclei at their lowest potential energy Chemistry chapter 6

Bond energy The amount of energy that must be added to break a chemical bond and form neutral isolated atoms. The same amount of energy that the atoms had to release when they bonded. Chemistry chapter 6

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 The outermost s and p sublevels want to be full Chemistry chapter 6

Exceptions to octet rule Hydrogen only needs two electrons, not eight Boron tends to only be surrounded by six electrons. Some atoms can have more than eight when combining with fluorine, oxygen, and chlorine. This expanded valence also includes the d orbitals. Chemistry chapter 6

Electron dot diagrams Usually only the valence electrons are involved in chemical reactions. Electron dot diagrams allow us to draw these electrons around the symbol for an element. Chemistry chapter 6

Drawing electron dot diagrams The element’s symbol represents the nucleus and all the electrons in the inner energy levels. Write the noble gas configuration of the element. Add the superscripts of the s and p orbitals to get the number of valence electrons. Chemistry chapter 6

Draw dots on the sides to represent electrons. Chemistry chapter 6

Examples Examples Chemistry chapter 6

Molecules Can be represented by dot diagrams Example: HCl Shared pairs can be replaced by long dashes Chemistry chapter 6

Lone pair Unshared pair of electrons Not involved in bonding Chemistry chapter 6

Discuss Read sample problem 6-2 on page 170 Chemistry chapter 6

Lewis diagram Formulas in which atomic symbols represent the nuclei and inner-shell electrons Dot diagrams with dots or dashes for shared electrons Chemistry chapter 6

Structural diagrams Leaves out the unshared pairs Indicates the kind, number, arrangement, and bonds of atoms Chemistry chapter 6

Single bond Covalent bond produced when one pair of electrons is shared Chemistry chapter 6

Discuss Sample problem 6-3 on pages 171-172 Chemistry chapter 6

Double bond Covalent bond produced when two pairs of electrons are shared Example: CO2 Chemistry chapter 6

Triple bond Covalent bond formed when three pairs of electrons are shared Example: N2 Chemistry chapter 6

Multiple bonds Double or triple bonds Have higher bond energies Have shorter bond lengths Chemistry chapter 6

Resonance structures Bonding in molecules or ions that cannot be correctly represented by a Lewis structure A double-headed arrow is placed between the two possibilities Example: ozone Chemistry chapter 6

Discuss Covalent Bonds More Covalent Bonds Sample problem 6-4 on page 174 Chemistry chapter 6

Examples Draw a Lewis Structure and the structural formula for each of the following: H2O CH4 CH4O O2 C2H2 C2H4 Chemistry chapter 6

Discuss Section Review on page 175 Chemistry chapter 6

Ionic compound Composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal Usually crystalline solids Not composed of independent units like molecules Ionic bonding More ionic bonding Chemistry chapter 6

Chemical formula For ionic compounds Shows the simplest ratio of the ions that gives no net charge Chemistry chapter 6

Formula unit The simplest collection of atoms from which an ionic compound’s formula can be established. For sodium chloride, NaCl The simplest way to show a one to one ratio Chemistry chapter 6

Ratio of ions Depends on the charges of the ions combined. Na+ and Cl- form NaCl Mg2+ and F- form MgF2 Chemistry chapter 6

Forming ionic compounds Examples: NaCl MgF2 Chemistry chapter 6

Crystal lattice Ions minimize their potential energy by forming an orderly arrangement Chemistry chapter 6

Lattice energy The energy released when one mole of an ionic crystalline compound is formed from gaseous ions. Used to compare bond strengths Chemistry chapter 6

Molecular vs. ionic compounds Forces between molecules are weaker than ionic bonding forces Molecular compounds have lower melting points Ionic compounds are hard but brittle. Ionic compounds conduct electricity when melted or dissolved in solution Chemistry chapter 6

Polyatomic ions A charged group of covalently bonded atoms. Form ionic bonds just like other ions. Examples Ammonium ion NH4+ Sulfate ion SO42- Chemistry chapter 6

Discuss Use electron-dot notation to demonstrate the formation of ionic compounds involving Li and Cl Ca and I What basic unit are molecular compounds composed of? Ionic compounds? Chemistry chapter 6

Metals High electrical conductivity. Metals form crystals and their orbitals overlap. All the atoms have the same attraction for their electrons, so the electrons can move easily from one atom to another. Chemistry chapter 6

Valence electrons Metals have very few Their p sublevels are empty Transition metals also have many vacant d sublevels Chemistry chapter 6

Metallic bonding Delocalized electrons – don’t belong to one specific atom Sea of electrons formed around the metal atoms Metallic bonding results from the attraction between metal atoms and the surrounding sea of electrons Chemistry chapter 6

Metallic properties High electrical and thermal conductivity Ability to absorb a wide range of light frequencies Electrons enter excited states and then return to ground states This makes metals look shiny Chemistry chapter 6

Metallic properties Metals bond the same way in all directions Their layers can slide past each other without breaking Malleabiltiy - the ability to be hammered or beaten into thin sheets Ductility – the ability to be drawn, pulled or extruded into a thin wire Chemistry chapter 6

Metallic bond strength When there is a bigger nucleus and more electrons, the bond is stronger. Stronger bonds have higher heats of vaporization Chemistry chapter 6

Discuss Describe the electron-sea model of metallic bonding. What is the relationship between metallic bond strength and heat of vaporization? Explain why most metals are malleable and ductile but ionic crystals are not. Chemistry chapter 6

Molecular polarity Uneven distribution of molecular charge Determined by the polarity of each bond and the geometry of the molecule Chemistry chapter 6

Molecular geometry There are two equally successful theories One accounts for molecular bond angles The other describes the orbitals that contain the valence electrons Chemistry chapter 6

VSEPR theory Valence-shell, electron-pair repulsion The repulsion between the sets of valence-level electrons surrounding an atom causes these sets to be oriented as far apart as possible To predict geometry, one must consider the location of all electron pairs around the bonded atoms. Chemistry chapter 6

Linear molecules A is the central atom B is an atom bonded to it Formulas like AB2 form linear molecules, with the B atoms on opposite sides of the A atom Chemistry chapter 6

Trigonal Planar AB3 A is in the center. The 3 Bs make a triangle around it. The molecule is flat (planar) Chemistry chapter 6

Tetrahedral AB4 A pyramid with A at the center and a B at each corner Chemistry chapter 6

When there are lone pairs The repel the other electrons just like shared pairs do Use an E to represent the unshared pairs Chemistry chapter 6

Bent or angular AB2E Like trigonal planar, but without the top of the triangle Chemistry chapter 6

Trigonal pyramidal AB3E Like tetrahedral, but without the top of the pyramid A is at the center, the Bs are the base, and E is at the top Chemistry chapter 6

Bent or angular AB2E2 Like a tetrahedral, but only the atoms (not the unshared pairs) determine the shape Chemistry chapter 6

Trigonal bipyramidal AB5 Like two pyramids on top of each other Chemistry chapter 6

Octahedral AB6 8 sided shape Chemistry chapter 6

VSEPR Double and triple bonds are treated the same way as single bonds. Use Lewis structures and table 6-5 on page 186 to predict molecule shapes Chemistry chapter 6

Hybridization Explains how orbitals get rearranged when atoms form covalent bonds Describes how atoms are bonded Chemistry chapter 6

Hybrid orbitals Orbitals of equal energy produced by the combining of two or more orbitals on the same atom The number of hybrid orbitals equals the number of orbitals they were made from Chemistry chapter 6

Example: Methane CH4 To get 4 equal orbitals, the s and p orbitals combine. They form 4 sp3 orbitals Chemistry chapter 6

Other hybrid orbitals Chemistry chapter 6

Intermolecular forces The forces of attraction between molecules Usually weaker than covalent, ionic, and metallic bonds Can be measured by boiling point Higher boiling point means stronger intermolecular forces Chemistry chapter 6

Dipole Equal but opposite charges separated by a short distance Direction is from the positive pole to the negative pole Chemistry chapter 6

Dipole-dipole forces Forces of attraction between polar molecules Strongest intermolecular forces Only act over short distances Chemistry chapter 6

More than two atoms Polarity of the molecule depends on the polarity and orientation of each bond Chemistry chapter 6

Chemistry chapter 6

Induced dipoles A polar molecule can induce a dipole in a nonpolar molecule by attracting its electrons. It is temporary Chemistry chapter 6

Hydrogen bonding A hydrogen atom in a compound with a very electronegative atom is attracted to an unshared pair of electrons in a nearby molecule. Examples: Water (H2O) Hydrogen fluoride (HF) Ammonia (NH3) Chemistry chapter 6

Hydrogen bonding Causes higher boiling points Chemistry chapter 6

Polarity and Hydrogen Bonding Chemistry chapter 6

London dispersion forces The intermolecular attractions caused by the constant motion of electrons and the creation of instantaneous dipoles Act between all atoms and molecules The only intermolecular forces among noble gas atoms and nonpolar molecules Low boiling points Chemistry chapter 6

discuss What is the difference between intramolecular forces and intermolecular forces? What is a dipole? In what kind of molecule are dipoles common? What is hydrogen bonding? What are London dispersion forces? Molecular Geometry review Chemistry chapter 6