Chemical Bonds Compounds Alloys.pdf
Bonding (and Nomenclature) The bonds & properties of compounds & alloys along with compound nomenclature The textbook chapters for this powerpoint slide show are 6, and 7.1. (7.2 through 7.4 will be covered as a separate mini-math unit as these sections deal with the fundamental math around compounds.) Chapter 6 provides detailed descriptions of bonding with discussion of changes in electrons and various theories. Chapter 7, section 1 explains how to name and write formulas of compounds. Our classroom practice and this powerpoint slide show deal with ionic bonding and naming (chapters 6.3 & part of 7.1), metallic bonding (chapter 6.4) and covalent bonding and naming (chapters 6.2, 6.5 and part of 7.1). Read this material carefully for this unit.
Bonding Some general concepts:
• Compounds form because of resultant lower overall energy within the atoms’ electron configurations along with electrical attraction and repulsion. (Read section 6.1) • The following slides and our general approach to bonding in our Periodic Trends discussion make a clear distinction between ionic and covalent bonding. • In reality, the nature of a chemical bond is more of a continuum between fully ionic where the electrons are transferred and fully covalent where the electrons are equally distributed around both (all) of the nuclei involved in a bond. (See Figure 2 in Ch. 6 page 176)
IONIC COMPOUNDS & IONIC BONDING
• Flip through the chapter 6 pp 190 – 194, examining subsection titles and figures. (Read Chapter 6.3) • Ionic compounds: A compound in which the particles are held together by the attraction between cations (positive ions) and anions (negative ions). • Ionic bond: The attraction between cations and anions
Ion Formation & Ionic Bonding The sodium atom loses its valence electron forming the sodium ion (isoelectronic with neon). The chlorine atom gains the electron sodium lost becoming the chloride ion (isoelectronic with argon). The positive sodium ion is attracted to the negative chloride ion. Electron transfer has occured in all ionic bonding & ionic compounds. (p191) Ionic Bond
• NaCl is the chemical formula for sodium chloride. • The compound is NEUTRAL all the positive charges are balanced by all the negative charges. Crystal Lattice
Crystal Lattice: The regular ordered arrangement of cations and anions in an ionic compound.
Formula Units • A chemical formula shows the kinds and numbers of atoms in the smallest representative unit of a substance. • A formula unit is the lowest whole-number ratio of ions in an ionic compound. • We don’t talk about “molecules” of ionic compounds; a molecule is a unique chemical concept that occurs when there is a specific one-to-one relationship between covalently bonded atoms. This specific one-to-one relationship does not exist within ionic compounds. It is impossible to “observe” a single formula unit of an ionic compound; they don’t really exist without being part of a crystal lattice. • The formula unit is used to count particles with Avogadro’s number.
Crystal Lattice and Formula Unit • Notice that each chloride ion is equally bonded (attracted to) six sodium ions. But each sodium ion is equally bonded to six Crystal chloride ions. lattice
This results in a one-to-one ratio of cations to anions; the formula unit then is NaCl.
Coordination Numbers and Formulas • The coordination number of an ion is the number of ions of opposite charge that surround the ion in a crystal. • In NaCl, each ion has a coordination number of 6.
Coordination Numbers and Formulas • In CsCl, each ion has a coordination number of 8. • In TiO2, each Ti4+ ion has a coordination number of 6, while each O2- ion has a coordination number of 3.
Coordination Numbers and Formulas • The ratio of the ions’ coordination numbers is the inverse of the ratio of the ions in the formula: – Na+ (coordination number 6); Cl- (coordination number 6); 6:6 = 1:1 Na1Cl1 – Ti4+ (coordination number 6); O2- (coordination number 3); 6:3 = 2:1 so Ti1O2
• Coordination Numbers are determined experimentally and are particular to each ionic compound. The size and charge of the ions both impact the coordination numbers in a compound.
• This information will not be tested; this is to provide thoroughness.
IONIC COMPOUND NAMES • All ionic compounds are made of a cation and anion, so… • the names of ionic compounds are pretty straight forward, cation ___ anion. • The names of the cations are the same as the metals. (sodium, Na; sodium ion Na+) • Anions of monoatomic non-metals end in ___ide; • Oxygen – oxide • Phosphorous – phosphide
Nitrogen – nitride fluorine -- fluoride
IONIC COMPOUND NAMES • Polyatomic ions have different endings – Most end in ___ ite or ______ ate – The “ite” ions have one less oxygen in them than the “ate” ions – Their names and formulas have to be memorized – Some end in “ide,” so be careful
Table 2, Ch. 7.1, p. 226
• Common Polyatomic Ions • See Handout • Make flashcards • Practice and Memorize
IONIC COMPOUND FORMULAS • Ionic compounds are neutral so the formula must reflect that, so… • The total positive charges from the cation(s) in the formula must be “balanced” (i.e., equal) to the total negative charges from the anion(s) in the formula.
IONIC COMPOUND FORMULAS • The “simplest” method of arriving at a correct ionic compound formula is to “criss-cross” the charge number of the ions, then reduce the subscripts to the “simplest” ratio. • Magnesium Chloride
–Mg2+ Cl(1)–Mg(1)Cl2 MgCl2
IONIC COMPOUND FORMULAS • With polyatomic ions, use a parentheses around any polyatomic ions for which there are more than one of them in the compound. • Ex: Ammonium Carbonate • NH4+ CO32• (NH4)2CO3 • We use parentheses around the ammonium (there are more than one), but not around carbonate (there is only one).
IONIC COMPOUND FORMULAS • Some are “multi-valent.” – This means they can have several different ionic charges (oxidation numbers) – Multi-valent metals are found primarily within the transition metals and the metals of the pblock (especially tin and lead). – Generally, expect multi-valent behavior for all transition metals except zinc and silver. – This initial approach to and understanding of bonding and oxidation states is intentionally simplified
IONIC COMPOUND FORMULAS • There are two ways to name transition metals — The stock method which uses a Roman Numeral after the name of the metal to indicate the charge. For example copper (II) – Cu2+, chromium (VI) Cr6+ — The classical names use _ous and _ic endings. The “ic” endings have higher oxidation numbers than the “ous” endings — And, despite them being technically incorrect, prefix use on ionic compounds can be found throughout chemical literature and catalogs (e.g., TiO2 is often called titanium dioxide, instead of titanium (IV) oxide.)
Multi-Valent Metals See also the bottom of Table 1, Ch. 7, p. 221
You don’t need to memorize the classical names.
IONIC COMPOUND FORMULAS
• The Ionic Bonds & Ionic Compounds Packet is a “paper” lab to help you attain the concepts around ionic bonding and ionic compounds. • This packet is rich with practice in writing formulas and writing names of compounds. • Working through this packet, with the goal of attaining concepts and skills, is necessary for maximum success. • Do not “copy” answers or otherwise let someone else DO the work; if you don’t DO the work you won’t learn the skills.
Metallic Bonding • A metallic bond is the attractive force holding the nuclei of metals together. • There is NO electron transfer; there is NO covalent sharing. • Metals don’t really make chemical compounds with each other, they do make alloys.
Sea of Electrons • The accepted theory behind metallic bonding is the “Sea of Electrons” Theory • In this theory the valence electrons of the metal atoms become “de-localized” so that the valence electrons cannot be “assigned” to any specific nucleus. • The remaining “cations” are then attracted to the “sea of de-localized” electrons.
•This theory explains the generally high melting point of metals; there is great attraction between the cations and the sea of electrons. •This theory explains the maleability of metals; notice that the sea of electrons allows bending rather than breaking (as you would see in most ionic compound crystals). •This theory explains the conductivity of metals: •The nuclei are relatively close together so heat (vibration of particles) can be easily conducted. •The electrons are “free” to move – the very definition of electricity, hence electrical conductivity.
Alloys • An alloy is a solution of metals. • Different metals are mixed to take advantage of their particular properties like corrosion resistance, hardness, electrical conductivity, and high melting point. • In spite of the “random” nature of the sea of electron theory, the nuclei have particular sizes that result in “optimum” distances between them. Changing the speed at which the alloy is heated and cooled can change the nature of the metal “crystals.”
Alloys • As solutions (homogeneous mixtures) of metals, alloys are mixed in the liquid phase then cooled. • The metal component that is in the greatest proportion is the solvent, and the other metal(s) is(are) solutes. • Two key alloy types are possible: – The substitution alloy in which the solute atoms replace (substitute) the existing solvent atoms – The interstitial alloy in which the solute atoms take positions in the “gaps” between the solvent atoms.
Types of Alloys Most commercial jewelry is made of alloys of silver, gold and platinum with copper, iron, zinc, nickel, and other metals.
Stainless steel is an alloy of iron, chromium, carbon and other elements.
Covalent Bonding Molecules & Molecular Compounds
Covalent Bonding • As the name implies, covalent bonding is when the valence electrons are shared between 2 atoms. • Generally, non-metals are the only elements to make covalent bonds. • Most atoms bond to achieve stable octets of valence electrons. (H generally achieves a stable “doublet.”)
Lewis Dot Diagrams •An element’s valence electrons are arranged around the “sides” of the element’s symbol. •Place one dot on each side working around in a single direction, until all the valence electrons are represented. 1
H 1s2
•A pair of “dots” (i.e., valence electrons) between two atoms is a covalent bond.
H2
H
H
H H
Covalent Bonding •When two electrons are shared between two atoms there is a single covalent bond. •When 4 electrons are shared between two atoms there is a double covalent bond. •When 6 electrons are shared between two atoms there is a triple covalent bond. •Dot diagrams (Lewis Dot Structures or Lewis Dot Diagrams) are used to determine/analyze the bonding in covalent compounds. •A dash is often used to indicate a pair of electrons instead of two dots. Dashes are most frequently used to indicate bonds; one dash for a single bond, two for a double bond and three for a triple bond.
H H Single Bond
O O Double Bond
N N Triple Bond
Bonding and Lewis Dot Diagrams • The unpaired valence electrons are the most likely to bond. • There are some bonds in which both of the bonding electrons are shared by a single atom; this is called a coordinate bond. • Coordinate bonds are indicated by an arrow instead of a dash; the arrow should point from the atom that is “sharing” the electrons (had an unshared pair before the bond).
Exceptions to the Octet Rule • Boron is a semi-metal that often “breaks” the octet rule. When boron makes covalent bonds with its three valence electrons, boron ends up with 6 valence electrons and not 8. • Some molecules have atoms with more than 8 valence electrons. • See page 183
Lewis Dot Structures • Lewis Dot Diagrams take practice • After you have “paired” up single electrons, check for octets (doublets for H) for all atoms in the structure. • If Octets are NOT met, consider moving pairs of electrons around to see if sharing an unshared pair (coordinate bonding) can satisfy octet needs. • As a last resort, accept the structure with exceptions to the octet rule.
Resonance • For some molecules more than one “valid” structure can be worked out. • Most times this involves alternative placements for double and triple bonds. • In these instances, all of the valid structures, together, represent the Lewis Dot Structure for the compound. • See p. 189 Also class notes (ozone structure is incorrect on page 189)
Bonding Theory • The various explanations of covalent bonds can be complex, visual, and overlapping. • Three basic bonding theories are presented in our study of bonding – Molecular Orbitals – VSEPR Theory – Orbital Hybridization Theory
Molecular Orbitals • Your textbook does NOT present very good detail about this theory. • Atomic orbitals overlap within molecules giving rise to molecular orbitals. • There are two basic types of molecular orbitals. – Sigma Bonding Molecular Orbital (sigma bond) – Pi Bonding Molecular Orbital (pi bond)
Sigma Bonds • A sigma bonds is formed (read, a sigma molecular orbital occurs) when the overlapping atomic orbitals are co-linear with the axis of the bonding nuclei.
• These “sausage shaped” clouds are the “most likely” place to find the shared (bonding) electrons.
Sigma Bonds • Sigma Bonds are due to the overlap of: • Two overlapping “s” orbitals • Overlapping “s” and “p” orbitals • Two overlapping “p” orbitals
• A sigma bond represents two shared electrons. • All single covalent bonds are sigma bonds.
Pi Bonds • Pi bonds are formed from the overlap of atomic orbitals that are NOT co-linear with the bonding nuclei.
• A Pi bond (read pi molecular orbital) is a pair of “kidney” shaped clouds above & below (or infront & behind) the bonding nuclei.
Pi Bonds • Pi bonds are only present with double and triple covalent bonds. • Generally, pi bonds are formed by overlapping “p” orbitals. • For a double covalent bond; one sigma and one pi bond are formed. • For a triple covalent bond; one sigma and two pi bonds are formed.
VSEPR Theory • VSEPR theory explains molecular geometries by electron pairs in the valence shell repelling each other. • Valence Shell Electron Pair Repulsion • The molecular geometry bonding lab (comes with packet) will help provide sufficient practice and examples for the “common” geometries found in many molecules.
VSEPR • It is important to remember that both bonding pairs and unshared pairs are all pairs of valence shell electrons, and therefore repel. • According to VSEPR theory, the repulsion between electron pairs causes molecular shapes to adjust so that the valenceelectron pairs stay as far apart as possible. • See the summary of molecular geometry and VSEPR theory (p. 200, Table 5)
Orbital Hybridization • Molecular Orbitals satisfy quantum mechanics. • VSEPR theory satisfactorily explains molecular geometry. • Orbital Hybridization does both. • But….
Orbital Hybridization • . . . It is the most “challenging” theory. • In hybridization, several atomic orbitals mix to form the same total number of equivalent hybrid orbitals. • The atomic orbital types we are familiar with “hybridize” in bonding. The level of hybridization is based upon examining the “central” atom in a molecule. • Counting the number of electron pairs directly bonded to the central atom is the best indicator of the hybridization
Orbital Hybridization • sp (unwritten superscripts, 1 each) 1 + 1 = 2 (two electron pairs bonded with the central atom) • sp2 3 e- pairs • sp3 4 e- pairs • dsp3 5 e-pairs • d2sp3 6 e-pairs • REMEMBER – Unshared electron pairs ARE electron pairs – count them.
Orbital Hybridization • See pp 201 – 202 for good illustrations and discussion. • Be sure to participate in the molecular geometry lab & complete the bonding summary table, which “connects” Lewis Dot Structures, Molecular Geometry, orbital hybridization, polarity of molecules, into “predictable” bonding patterns.
Molecular Polarity • Molecular Geometry & Electronegativity can be used to determine the “polarity” of covalent compounds. • Given: All ionic bonds (by definition) are polar bonds. • Bond Polarity: A bond is polar if there is a difference between the electronegativities of the bonding atoms.
Molecular Polarity • The mathematical difference in the e-Ns of the bonding atoms is called the bond dipole -- δ (lower case greek delta). • The sum of all the bond dipoles in a molecule is the molecular dipole. • If all the dipoles add to zero, the molecule is non-polar.
Attractive Forces • There are 3 basic attractive forces which operate on covalent compounds: – Dispersion Forces (London dispersion forces, or Van der Waals forces) – Dipole-Dipole Attractions – Hydrogen Bonding
Dispersion Forces • Operate on “non-polar” molecules. • Caused by the “instantaneous” dipoles resulting from “shifting” electron clouds. • Larger molecules generally have greater dispersion forces.
Dipole-Dipole Attractions • Operate on polar molecules • The positive end of one molecule is attracted to/by the negative end of another molecule • The greater the molecular dipole, generally, the greater the attractive force.
Hydrogen Bonding • An intermolecular force (i.e., NOT a chemical bond) • Due to the very electronegative nature of O, N, and F and… • The low elelctronegativity of H. • Essentially a VERY strong dipole-dipole attraction between molecules where H is chemically bonded to O, N, or F.
Network Solids • Covalently bonded “crystals” that resemble crystal lattice more than individual molecules. • Very high melting & boiling point. Each atom in the network is bonded to at least 2 other but usually 4 other atoms. • See page 189
Naming Covalent Compounds • Use prefixes on all atoms in the formula – Drop mono, on leading elements in a formula – Generally list elements from least e-negative to most e-negative (for formulas also) – Drop some vowels to make name sound better – See prefixes on page 228
Naming Covalent Compounds
• Network solids are named like other covalent compounds, but • much like ionic compounds, the formulas are in their simplest ratio • There are not “individual” molecules • See examples on page 230
Naming Covalent Compounds • Most common acids can be considered “hydrogen salts.” – HA, where ‘A’ is a common anion – The naming protocol is based on the name of the anion in the acid • If the anion ends in ‘ide,’ the acid is named – Hydro ‘root of anion’ ic acid • If the anion ends in ‘ite’ the acid is named – ‘anion root’ ous acid • If the anion ends in ‘ate’ the acid is named – ‘anion root’ ic acid • See page 230
Bonding Some general concepts:
• Compounds form because of resultant lower overall energy within the atoms’ electron configurations along with electrical attraction and repulsion. (Read section 6.1) • The following slides and our general approach to bonding in our Periodic Trends discussion make a clear distinction between ionic and covalent bonding. • In reality, the nature of a chemical bond is more of a continuum between fully ionic where the electrons are transferred and fully covalent where the electrons are equally distributed around both (all) of the nuclei involved in a bond. (See Figure 2 in Ch. 6 page 176)
IONIC COMPOUNDS & IONIC BONDING
• Flip through the chapter 6 pp 190 – 194, examining subsection titles and figures. (Read Chapter 6.3) • Ionic compounds: A compound in which the particles are held together by the attraction between cations (positive ions) and anions (negative ions). • Ionic bond: The attraction between cations and anions
Ion Formation & Ionic Bonding The sodium atom loses its valence electron forming the sodium ion (isoelectronic with neon). The chlorine atom gains the electron sodium lost becoming the chloride ion (isoelectronic with argon). The positive sodium ion is attracted to the negative chloride ion. Electron transfer has occured in all ionic bonding & ionic compounds. (p191) Ionic Bond
• NaCl is the chemical formula for sodium chloride. • The compound is NEUTRAL all the positive charges are balanced by all the negative charges. Crystal Lattice
Crystal Lattice: The regular ordered arrangement of cations and anions in an ionic compound.
Formula Units • A chemical formula shows the kinds and numbers of atoms in the smallest representative unit of a substance. • A formula unit is the lowest whole-number ratio of ions in an ionic compound. • We don’t talk about “molecules” of ionic compounds; a molecule is a unique chemical concept that occurs when there is a specific one-to-one relationship between covalently bonded atoms. This specific one-to-one relationship does not exist within ionic compounds. It is impossible to “observe” a single formula unit of an ionic compound; they don’t really exist without being part of a crystal lattice. • The formula unit is used to count particles with Avogadro’s number.
Crystal Lattice and Formula Unit • Notice that each chloride ion is equally bonded (attracted to) six sodium ions. But each sodium ion is equally bonded to six Crystal chloride ions. lattice
This results in a one-to-one ratio of cations to anions; the formula unit then is NaCl.
Coordination Numbers and Formulas • The coordination number of an ion is the number of ions of opposite charge that surround the ion in a crystal. • In NaCl, each ion has a coordination number of 6.
Coordination Numbers and Formulas • In CsCl, each ion has a coordination number of 8. • In TiO2, each Ti4+ ion has a coordination number of 6, while each O2- ion has a coordination number of 3.
Coordination Numbers and Formulas • The ratio of the ions’ coordination numbers is the inverse of the ratio of the ions in the formula: – Na+ (coordination number 6); Cl- (coordination number 6); 6:6 = 1:1 Na1Cl1 – Ti4+ (coordination number 6); O2- (coordination number 3); 6:3 = 2:1 so Ti1O2
• Coordination Numbers are determined experimentally and are particular to each ionic compound. The size and charge of the ions both impact the coordination numbers in a compound.
• This information will not be tested; this is to provide thoroughness.
IONIC COMPOUND NAMES • All ionic compounds are made of a cation and anion, so… • the names of ionic compounds are pretty straight forward, cation ___ anion. • The names of the cations are the same as the metals. (sodium, Na; sodium ion Na+) • Anions of monoatomic non-metals end in ___ide; • Oxygen – oxide • Phosphorous – phosphide
Nitrogen – nitride fluorine -- fluoride
IONIC COMPOUND NAMES • Polyatomic ions have different endings – Most end in ___ ite or ______ ate – The “ite” ions have one less oxygen in them than the “ate” ions – Their names and formulas have to be memorized – Some end in “ide,” so be careful
Table 2, Ch. 7.1, p. 226
• Common Polyatomic Ions • See Handout • Make flashcards • Practice and Memorize
IONIC COMPOUND FORMULAS • Ionic compounds are neutral so the formula must reflect that, so… • The total positive charges from the cation(s) in the formula must be “balanced” (i.e., equal) to the total negative charges from the anion(s) in the formula.
IONIC COMPOUND FORMULAS • The “simplest” method of arriving at a correct ionic compound formula is to “criss-cross” the charge number of the ions, then reduce the subscripts to the “simplest” ratio. • Magnesium Chloride
–Mg2+ Cl(1)–Mg(1)Cl2 MgCl2
IONIC COMPOUND FORMULAS • With polyatomic ions, use a parentheses around any polyatomic ions for which there are more than one of them in the compound. • Ex: Ammonium Carbonate • NH4+ CO32• (NH4)2CO3 • We use parentheses around the ammonium (there are more than one), but not around carbonate (there is only one).
IONIC COMPOUND FORMULAS • Some are “multi-valent.” – This means they can have several different ionic charges (oxidation numbers) – Multi-valent metals are found primarily within the transition metals and the metals of the pblock (especially tin and lead). – Generally, expect multi-valent behavior for all transition metals except zinc and silver. – This initial approach to and understanding of bonding and oxidation states is intentionally simplified
IONIC COMPOUND FORMULAS • There are two ways to name transition metals — The stock method which uses a Roman Numeral after the name of the metal to indicate the charge. For example copper (II) – Cu2+, chromium (VI) Cr6+ — The classical names use _ous and _ic endings. The “ic” endings have higher oxidation numbers than the “ous” endings — And, despite them being technically incorrect, prefix use on ionic compounds can be found throughout chemical literature and catalogs (e.g., TiO2 is often called titanium dioxide, instead of titanium (IV) oxide.)
Multi-Valent Metals See also the bottom of Table 1, Ch. 7, p. 221
You don’t need to memorize the classical names.
IONIC COMPOUND FORMULAS
• The Ionic Bonds & Ionic Compounds Packet is a “paper” lab to help you attain the concepts around ionic bonding and ionic compounds. • This packet is rich with practice in writing formulas and writing names of compounds. • Working through this packet, with the goal of attaining concepts and skills, is necessary for maximum success. • Do not “copy” answers or otherwise let someone else DO the work; if you don’t DO the work you won’t learn the skills.
Metallic Bonding • A metallic bond is the attractive force holding the nuclei of metals together. • There is NO electron transfer; there is NO covalent sharing. • Metals don’t really make chemical compounds with each other, they do make alloys.
Sea of Electrons • The accepted theory behind metallic bonding is the “Sea of Electrons” Theory • In this theory the valence electrons of the metal atoms become “de-localized” so that the valence electrons cannot be “assigned” to any specific nucleus. • The remaining “cations” are then attracted to the “sea of de-localized” electrons.
•This theory explains the generally high melting point of metals; there is great attraction between the cations and the sea of electrons. •This theory explains the maleability of metals; notice that the sea of electrons allows bending rather than breaking (as you would see in most ionic compound crystals). •This theory explains the conductivity of metals: •The nuclei are relatively close together so heat (vibration of particles) can be easily conducted. •The electrons are “free” to move – the very definition of electricity, hence electrical conductivity.
Alloys • An alloy is a solution of metals. • Different metals are mixed to take advantage of their particular properties like corrosion resistance, hardness, electrical conductivity, and high melting point. • In spite of the “random” nature of the sea of electron theory, the nuclei have particular sizes that result in “optimum” distances between them. Changing the speed at which the alloy is heated and cooled can change the nature of the metal “crystals.”
Alloys • As solutions (homogeneous mixtures) of metals, alloys are mixed in the liquid phase then cooled. • The metal component that is in the greatest proportion is the solvent, and the other metal(s) is(are) solutes. • Two key alloy types are possible: – The substitution alloy in which the solute atoms replace (substitute) the existing solvent atoms – The interstitial alloy in which the solute atoms take positions in the “gaps” between the solvent atoms.
Types of Alloys Most commercial jewelry is made of alloys of silver, gold and platinum with copper, iron, zinc, nickel, and other metals.
Stainless steel is an alloy of iron, chromium, carbon and other elements.
Covalent Bonding Molecules & Molecular Compounds
Covalent Bonding • As the name implies, covalent bonding is when the valence electrons are shared between 2 atoms. • Generally, non-metals are the only elements to make covalent bonds. • Most atoms bond to achieve stable octets of valence electrons. (H generally achieves a stable “doublet.”)
Lewis Dot Diagrams •An element’s valence electrons are arranged around the “sides” of the element’s symbol. •Place one dot on each side working around in a single direction, until all the valence electrons are represented. 1
H 1s2
•A pair of “dots” (i.e., valence electrons) between two atoms is a covalent bond.
H2
H
H
H H
Covalent Bonding •When two electrons are shared between two atoms there is a single covalent bond. •When 4 electrons are shared between two atoms there is a double covalent bond. •When 6 electrons are shared between two atoms there is a triple covalent bond. •Dot diagrams (Lewis Dot Structures or Lewis Dot Diagrams) are used to determine/analyze the bonding in covalent compounds. •A dash is often used to indicate a pair of electrons instead of two dots. Dashes are most frequently used to indicate bonds; one dash for a single bond, two for a double bond and three for a triple bond.
H H Single Bond
O O Double Bond
N N Triple Bond
Bonding and Lewis Dot Diagrams • The unpaired valence electrons are the most likely to bond. • There are some bonds in which both of the bonding electrons are shared by a single atom; this is called a coordinate bond. • Coordinate bonds are indicated by an arrow instead of a dash; the arrow should point from the atom that is “sharing” the electrons (had an unshared pair before the bond).
Exceptions to the Octet Rule • Boron is a semi-metal that often “breaks” the octet rule. When boron makes covalent bonds with its three valence electrons, boron ends up with 6 valence electrons and not 8. • Some molecules have atoms with more than 8 valence electrons. • See page 183
Lewis Dot Structures • Lewis Dot Diagrams take practice • After you have “paired” up single electrons, check for octets (doublets for H) for all atoms in the structure. • If Octets are NOT met, consider moving pairs of electrons around to see if sharing an unshared pair (coordinate bonding) can satisfy octet needs. • As a last resort, accept the structure with exceptions to the octet rule.
Resonance • For some molecules more than one “valid” structure can be worked out. • Most times this involves alternative placements for double and triple bonds. • In these instances, all of the valid structures, together, represent the Lewis Dot Structure for the compound. • See p. 189 Also class notes (ozone structure is incorrect on page 189)
Bonding Theory • The various explanations of covalent bonds can be complex, visual, and overlapping. • Three basic bonding theories are presented in our study of bonding – Molecular Orbitals – VSEPR Theory – Orbital Hybridization Theory
Molecular Orbitals • Your textbook does NOT present very good detail about this theory. • Atomic orbitals overlap within molecules giving rise to molecular orbitals. • There are two basic types of molecular orbitals. – Sigma Bonding Molecular Orbital (sigma bond) – Pi Bonding Molecular Orbital (pi bond)
Sigma Bonds • A sigma bonds is formed (read, a sigma molecular orbital occurs) when the overlapping atomic orbitals are co-linear with the axis of the bonding nuclei.
• These “sausage shaped” clouds are the “most likely” place to find the shared (bonding) electrons.
Sigma Bonds • Sigma Bonds are due to the overlap of: • Two overlapping “s” orbitals • Overlapping “s” and “p” orbitals • Two overlapping “p” orbitals
• A sigma bond represents two shared electrons. • All single covalent bonds are sigma bonds.
Pi Bonds • Pi bonds are formed from the overlap of atomic orbitals that are NOT co-linear with the bonding nuclei.
• A Pi bond (read pi molecular orbital) is a pair of “kidney” shaped clouds above & below (or infront & behind) the bonding nuclei.
Pi Bonds • Pi bonds are only present with double and triple covalent bonds. • Generally, pi bonds are formed by overlapping “p” orbitals. • For a double covalent bond; one sigma and one pi bond are formed. • For a triple covalent bond; one sigma and two pi bonds are formed.
VSEPR Theory • VSEPR theory explains molecular geometries by electron pairs in the valence shell repelling each other. • Valence Shell Electron Pair Repulsion • The molecular geometry bonding lab (comes with packet) will help provide sufficient practice and examples for the “common” geometries found in many molecules.
VSEPR • It is important to remember that both bonding pairs and unshared pairs are all pairs of valence shell electrons, and therefore repel. • According to VSEPR theory, the repulsion between electron pairs causes molecular shapes to adjust so that the valenceelectron pairs stay as far apart as possible. • See the summary of molecular geometry and VSEPR theory (p. 200, Table 5)
Orbital Hybridization • Molecular Orbitals satisfy quantum mechanics. • VSEPR theory satisfactorily explains molecular geometry. • Orbital Hybridization does both. • But….
Orbital Hybridization • . . . It is the most “challenging” theory. • In hybridization, several atomic orbitals mix to form the same total number of equivalent hybrid orbitals. • The atomic orbital types we are familiar with “hybridize” in bonding. The level of hybridization is based upon examining the “central” atom in a molecule. • Counting the number of electron pairs directly bonded to the central atom is the best indicator of the hybridization
Orbital Hybridization • sp (unwritten superscripts, 1 each) 1 + 1 = 2 (two electron pairs bonded with the central atom) • sp2 3 e- pairs • sp3 4 e- pairs • dsp3 5 e-pairs • d2sp3 6 e-pairs • REMEMBER – Unshared electron pairs ARE electron pairs – count them.
Orbital Hybridization • See pp 201 – 202 for good illustrations and discussion. • Be sure to participate in the molecular geometry lab & complete the bonding summary table, which “connects” Lewis Dot Structures, Molecular Geometry, orbital hybridization, polarity of molecules, into “predictable” bonding patterns.
Molecular Polarity • Molecular Geometry & Electronegativity can be used to determine the “polarity” of covalent compounds. • Given: All ionic bonds (by definition) are polar bonds. • Bond Polarity: A bond is polar if there is a difference between the electronegativities of the bonding atoms.
Molecular Polarity • The mathematical difference in the e-Ns of the bonding atoms is called the bond dipole -- δ (lower case greek delta). • The sum of all the bond dipoles in a molecule is the molecular dipole. • If all the dipoles add to zero, the molecule is non-polar.
Attractive Forces • There are 3 basic attractive forces which operate on covalent compounds: – Dispersion Forces (London dispersion forces, or Van der Waals forces) – Dipole-Dipole Attractions – Hydrogen Bonding
Dispersion Forces • Operate on “non-polar” molecules. • Caused by the “instantaneous” dipoles resulting from “shifting” electron clouds. • Larger molecules generally have greater dispersion forces.
Dipole-Dipole Attractions • Operate on polar molecules • The positive end of one molecule is attracted to/by the negative end of another molecule • The greater the molecular dipole, generally, the greater the attractive force.
Hydrogen Bonding • An intermolecular force (i.e., NOT a chemical bond) • Due to the very electronegative nature of O, N, and F and… • The low elelctronegativity of H. • Essentially a VERY strong dipole-dipole attraction between molecules where H is chemically bonded to O, N, or F.
Network Solids • Covalently bonded “crystals” that resemble crystal lattice more than individual molecules. • Very high melting & boiling point. Each atom in the network is bonded to at least 2 other but usually 4 other atoms. • See page 189
Naming Covalent Compounds • Use prefixes on all atoms in the formula – Drop mono, on leading elements in a formula – Generally list elements from least e-negative to most e-negative (for formulas also) – Drop some vowels to make name sound better – See prefixes on page 228
Naming Covalent Compounds
• Network solids are named like other covalent compounds, but • much like ionic compounds, the formulas are in their simplest ratio • There are not “individual” molecules • See examples on page 230
Naming Covalent Compounds • Most common acids can be considered “hydrogen salts.” – HA, where ‘A’ is a common anion – The naming protocol is based on the name of the anion in the acid • If the anion ends in ‘ide,’ the acid is named – Hydro ‘root of anion’ ic acid • If the anion ends in ‘ite’ the acid is named – ‘anion root’ ous acid • If the anion ends in ‘ate’ the acid is named – ‘anion root’ ic acid • See page 230
