Chapter 3 Structure And Properties Of Ionic And Covalent Compounds PDF

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® Because learning changes everything. GENERAL, ORGANIC, AND 11TH Edition BIOCHEMISTRY Katherine J. Denniston Danaè R. Quirk Joseph J. Topping Robert L. Caret Chapter 3 Structure and Properties of Ionic and Covalent Compounds © McGraw Hill LLC. All rights reserved. No reproduction or distribution without the prior written consent of McGraw Hill LLC.  3.1 Chemical Bonding Chemical bond - the force of attraction between any two atoms in a compound. This attractive force overcomes the repulsion of the positively charged nuclei of the two atoms participating in the bond. Interactions involving valence electrons are responsible for the chemical bond. © McGraw Hill LLC 2  Lewis Symbols Lewis symbol - a way to represent atoms using the element symbol and valence electrons as dots. As only valence electrons participate in bonding, this makes it much easier to work with the octet rule. The number of dots used corresponds directly to the number of valence electrons located in the outermost shell of the atoms of the element. © McGraw Hill LLC 3  Writing Lewis Symbols The four sides around the atomic symbol can each have two dots for a maximum of eight (octet of electrons). Writing Lewis symbols: Place one dot on each side until there are four dots around the symbol. Then add a second dot to each side in turn. The number of valence electrons limits the number of dots placed. Each unpaired dot (unpaired valence electron) is available to form a chemical bond. © McGraw Hill LLC 4  Lewis Symbols for Representative Elements © McGraw Hill LLC 5  Principal Types of Chemical Bonds: Ionic and Covalent Ionic bond - attractive force due to the transfer of one or more electrons from one atom to another: The attraction is due to the opposite charges of the ions. Covalent bond - attractive force due to the sharing of electrons between atoms. Some bonds have characteristics of both types and are not easily identified as one or the other. © McGraw Hill LLC 6  Ionic Bonding Representative elements form ions that obey the octet rule. Electrons are lost by a metal and they are gained by a nonmetal: Each atom achieves a “noble gas” configuration. 2 ions are formed - a cation and anion, which are attracted to each other. Ions of opposite charge attract each other creating the ionic bond. © McGraw Hill LLC 7  Ionic Bonding: NaCl Consider the formation of NaCl Na + Cl → NaCl Chlorine has a high electron affinity. Sodium has a low ionization energy (it readily loses its When chlorine gains an electron): electron, it gains the Ar configuration: When sodium loses the electron, it gains the Ne configuration. © McGraw Hill LLC 8  Essential Features of Ionic Bonding Metals tend to form cations because they have low ionization energies and low electron affinities. Nonmetals tend to form anions because they have high ionization energies and high electron affinities. Ions are formed by the transfer of electrons. The oppositely charged ions formed are held together by an electrostatic force. Reactions between metals and nonmetals tend to form ionic compounds. © McGraw Hill LLC 9  Ion Arrangement in a Crystal As a sodium atom loses one electron, it becomes a smaller sodium ion. When a chlorine atom gains that electron, it becomes a larger chloride ion. Attraction of the Na cation with the Cl anion forms NaCl ion pairs that aggregate into a crystal lattice. © McGraw Hill LLC 10  Covalent Bonding Consider the formation of H2: H + H  H2 Each hydrogen has one electron in its valance shell If it were an ionic bond it would look like this: However, both hydrogen atoms have an equal tendency to gain or lose electrons. Electron transfer from one H to another usually will not occur under normal conditions. © McGraw Hill LLC 11  Covalent Bond Instead, each atom attains a noble gas configuration by sharing electrons. Each hydrogen atom The shared now has two Electron electrons around it pair is called a and attained a He covalent bond. configuration. © McGraw Hill LLC 12  Covalent Bonding in Hydrogen © McGraw Hill LLC 13  Features of Covalent Bonds Covalent bonds form between atoms with similar tendencies to gain or lose electrons. Compounds containing covalent bonds are called covalent compounds or molecules. The diatomic elements have completely covalent bonds (totally equal sharing): H2, N2, O2, F2, Cl2, Br2, I2 Each fluorine is surrounded by 8 electrons – Ne configuration: © McGraw Hill LLC 14  Examples of Covalent Bonding 2e  from 2H 2e  for H 6e  from O 8e  for O 4e  from 4H 2e  for H 4e  from C 8e  for C © McGraw Hill LLC 15  Polar Covalent Bonding and Electronegativity The Polar Covalent Bond: Ionic bonding involves the transfer of electrons. Covalent bonding involves the sharing of electrons. Polar covalent bonding - bonds made up of unequally shared electron pairs. © McGraw Hill LLC 16  Polar Covalent Bond Hydrogen is somewhat Fluorine is somewhat positively charged negatively charged The two electrons between H and F are not shared equally. The electrons spend more time with fluorine. This sets up a polar covalent bond. A truly covalent bond can only occur when both atoms are identical. © McGraw Hill LLC 17  Polar Covalent Bonding in HF Fluorine is electron rich δ  Hydrogen is electron deficient δ  This results in unequal sharing of electrons in the pairs = polar covalent bonds © McGraw Hill LLC 18  Electronegativity Electronegativity - a measure of the ability of an atom to attract electrons in a chemical bond. Elements with high electronegativity have a greater ability to attract electrons than do elements with low electronegativity. Consider the covalent bond as competition for electrons between 2 positive centers: The difference in electronegativity determines the extent of bond polarity. © McGraw Hill LLC 19  Electronegativities of Selected Elements The most electronegative elements are found in the upper right corner of the periodic table. The least electronegative elements are found in the lower left corner of the periodic table. © McGraw Hill LLC 20  Electronegativity Calculations The greater the difference in electronegativity between two atoms, the greater the polarity of their bond. Which would be more polar, a H-F bond or H-Cl bond? H-F... 4.0 − 2.2 = 1.8 H-Cl... 3.2 − 2.2 = 1.0 The HF bond is more polar than the HCl bond. © McGraw Hill LLC 21  3.2 Naming Compounds and Writing Formulas of Compounds Nomenclature - the assignment of a correct and unambiguous name to each and every chemical compound: Two naming systems: ionic compounds. covalent compounds. © McGraw Hill LLC 22  Formulas of Compounds A formula is the representation of the fundamental compound using chemical symbols and numerical subscripts: The formula identifies the number and type of the various atoms that make up the compound unit. The number of like atoms in the unit is shown by the use of a subscript. Presence of only one atom is understood when no subscript is present. © McGraw Hill LLC 23  Ionic Compounds Metals and nonmetals usually react to form ionic compounds. The metals are cations and the nonmetals are anions. The cations and anions arrange themselves in a regular three-dimensional repeating array called a crystal lattice. Formula of an ionic compound is the smallest whole- number ratio of ions in the substance. © McGraw Hill LLC 24  Writing Formulas of Ionic Compounds from the Identities of the Component Ions Determine the charge of each ion: Metals have a charge equal to group number. Nonmetals have a charge equal to the group number minus eight. Cations and anions must combine to give a formula with a net charge of zero: It must have the same number of positive charges as negative charges. © McGraw Hill LLC 25  Predict Formulas Predict the formula of the ionic compounds formed from combining ions of the following pairs of elements: 1. sodium and oxygen 2. lithium and bromine 3. aluminum and oxygen 4. barium and fluorine © McGraw Hill LLC 26  Writing Names of Ionic Compounds from the Formula of the Compound 1 Name the cation followed by the name of the anion. A positive ion retains the name of the element; change the anion suffix to -ide. Formula cation and anion stem + ide = Compound name NaCl sodium chlor + ide sodium chloride Na2O sodium ox + ide sodium oxide L i2 S lithium sulf + ide lithium sulfide AlBr3 aluminum brom + ide aluminum bromide CaO calcium ox + ide calcium oxide © McGraw Hill LLC 27  Writing Names of Ionic Compounds from the Formula of the Compound 2 If the cation of an element has several ions of different charges (as with transition metals) use a Roman numeral following the metal name: Roman numerals give the charge of the metal. Examples: FeCl3 is iron(III) chloride FeCl2 is iron(II) chloride CuO is copper(II) oxide © McGraw Hill LLC 28  Common Nomenclature System Use -ic to indicate the higher of the charges that ion might have. Use -ous to indicate the lower of the charges that ion might have. Examples: FeCl2 is ferrous chloride FeCl3 is ferric chloride © McGraw Hill LLC 29  Stock and Common Names for Iron and Copper Ions Table 3.1 Systematic (Stock) Names for Iron and Copper Ions Formula Ion Charge Cation Name Compound Name FeCl2 2 Iron(II) Iron(II) chloride FeCl3 3 Iron(III) Iron(III) chloride Cu 2 O 1 Copper(I) Copper(I) oxide CuO 2 Copper(II) Copper(II) oxide Table 3.1 Common nomenclature for Iron and Copper Ions Formula Ion Charge Cation Name Compound Name FeCl2 2 Ferrous Ferrous chloride FeCl3 3 Ferric Ferric chloride Cu 2 O 1 Cuprous Cuprous oxide CuO 2 Cupric Cupric oxide © McGraw Hill LLC 30  Common Monatomic Cations and Anions Table 3.2 Common Monatomic Cations and Anions Cation Name Anion Name H Hydrogen ion H Hydride ion Li  Lithium ion F Fluoride ion Na  Sodium ion Cl Chloride ion K Potassium Rows 11 and 12 for Br  name are empty.Bromide ion ionanion and Cs  Cesium ion I Iodide ion Be 2 Beryllium ion O 2 Oxide ion Mg 2 Magnesium ion S2 Sulfide ion Ca 2 Calcium ion N 3 Nitride ion Ba 2 Barium ion P 3 Phosphide ion Al3 Aluminum ion Ag  Silver ion Monatomic ions - ions consisting of a single charged atom. © McGraw Hill LLC 31  Polyatomic Ions Polyatomic ions - ions composed of 2 or more atoms bonded together with an overall positive or negative charge: Within the ion itself, the atoms are bonded using covalent bonds. The positive and negative ions will be bonded to each other with ionic bonds. Examples: NH 4 ammonium ion SO 24 sulfate ion © McGraw Hill LLC 32  Common Polyatomic Cations and Anions Ion Name Ion Name H 3O  Hydronium CO32 Carbonate NH 4 Ammonium HCO3 Bicarbonate NO 2 Nitrite ClO  Hypochlorite NO3 Nitrate ClO 2 Chlorite SO32 Sulfite ClO3 Chlorate SO 24 Sulfate ClO 4 Perchlorate HSO 4 Hydrogen sulfate CH 3COO  or C2 H 3O 2  Acetate OH  Hydroxide MnO 4 Permanganate CN  Cyanide Cr2 O72 Dichromate PO34 Phosphate CrO 24 Chromate HPO 24 Hydrogen phosphate O 22 Peroxide H 2 PO 4 Dihydrogen phosphate © McGraw Hill LLC 33  Name These Compounds 1. NH4Cl 2. BaSO4 3. Fe(NO3)3 4. CuHCO3 5. Ca(OH)2 © McGraw Hill LLC 34  Writing Formulas of Ionic Compounds from the Name of the Compound Determine the charge of each ion. Write the formula so that the resulting compound is neutral. Example: Barium chloride: Barium is 2+, Chloride is 1− Formula is BaCl2 © McGraw Hill LLC 35  Determine the Formulas from Names Write the formula for the following ionic compounds: 1. sodium sulfate 2. ammonium sulfide 3. magnesium phosphate 4. chromium(II) sulfate © McGraw Hill LLC 36  Covalent Compounds Covalent compounds are typically formed from nonmetals. Molecules - compounds characterized by covalent bonding: Not a part of a massive three-dimensional crystal structure. Exist as discrete molecules in the solid, liquid, and gas states. © McGraw Hill LLC 37  Naming Covalent Compounds 1 1. The names of the elements are written in the order in which they appear in the formula. 2. A prefix indicates the number of each kind of atom. Table 3.4 Prefixes Used to Denote Numbers of Atoms in a Compound Prefix Number of Prefix Number of Atoms Atoms Mono- 1 Hexa- 6 Di- 2 Hepta- 7 Tri- 3 Octa- 8 Tetra- 4 Nona- 9 Penta- 5 Deca- 10 © McGraw Hill LLC 38  Naming Covalent Compounds 2 3. If only one atom of a particular element is present in the molecule, the prefix mono- is usually omitted from the first element. 4. Example: CO is carbon monoxide. 5. The stem of the name of the last element is used with the suffix –ide. 6. The final vowel in a prefix is often dropped before a vowel in the stem name. © McGraw Hill LLC 39  Name These Covalent Compounds 1. SiO2 2. N2O5 3. CCl4 4. IF7 © McGraw Hill LLC 40  Writing Formulas of Covalent Compounds Use the prefixes in the names to determine the subscripts for the elements. Examples: nitrogen trichloride NCl3 diphosphorus pentoxide P2O5 Some common names that are used: H2O water NH3 ammonia C2H5OH ethanol or ethyl alcohol C6H12O6 glucose © McGraw Hill LLC 41  Provide Formulas for These Covalent Compounds 1. nitrogen monoxide 2. dinitrogen tetroxide 3. diphosphorus pentoxide 4. nitrogen trifluoride © McGraw Hill LLC 42  3.3 Properties of Ionic and Covalent Compounds Physical State: Ionic compounds are usually solids at room temperature. Covalent compounds can be solids, liquids, and gases. Melting and Boiling Points: Melting point - the temperature at which a solid is converted to a liquid. Boiling point - the temperature at which a liquid is converted to a gas. © McGraw Hill LLC 43  Physical Properties Melting and Boiling Points Ionic compounds have much higher melting points and boiling points than covalent compounds. A large amount of energy is required to break the electrostatic attractions between ions. Ionic compounds typically melt at several hundred degrees Celsius. Structure of Compounds in the Solid State Ionic compounds are crystalline. Covalent compounds are crystalline or amorphous – having no regular structure. © McGraw Hill LLC 44  Electrolytes and Nonelectrolytes Solutions of Ionic and Covalent Compounds Ionic compounds often dissolve in water, where they dissociate - form positive and negative ions in solution. Electrolytes - ions present in solution allowing the solution to conduct electricity. Covalent solids usually do not dissociate and do not conduct electricity - nonelectrolytes. © McGraw Hill LLC 45  Comparison of Ionic vs. Covalent Compounds Property Ionic Covalent OftenTable composed of Summary Metal + nonmetal 2 nonmetals Electrons are Transferred Shared Physical state is Solid and crystalline Any; crystal or amorphous Dissociation Yes, they are No, they are electrolytes nonelectrolytes Boiling and Melting High Low point © McGraw Hill LLC 46  3.4 Drawing Lewis Structures of Molecules and Polyatomic Ions Lewis Structure Guidelines 1. Use chemical symbols for the various elements to write the skeletal structure of the compound: The least electronegative atom will be placed in the central position. Hydrogen always occupies terminal positions. Halogens occupy terminal positions, except when more electronegative elements are present. Carbon often forms chains of carbon-carbon covalent bonds. © McGraw Hill LLC 47  Lewis Structure Guidelines (2) 2. Determine the number of valence electrons associated with each atom in the compound: Combine these valence electrons to determine the total number of valence electrons in the compound. Polyatomic cations, subtract one electron for every positive charge. Polyatomic anions, add one electron for every negative charge. © McGraw Hill LLC 48  Lewis Structure Guidelines (3 and 4) 3. Connect the central atom to each of the surrounding atoms with single bonds. 4. Next, place electrons as lone pairs around the terminal atoms to complete the octet for each: Hydrogen needs only two electrons. Electrons not involved in bonding are represented as lone pairs. After the terminal atoms have an octet, provide the central atom with an octet if valence electrons are still available. © McGraw Hill LLC 49  Lewis Structure Guidelines (5 and 6) 5. If the octet rule is not satisfied for all the atoms, move one or more lone pairs on a terminal atom in to make a bond with the central atom: Continue to shift the electrons until all atoms have an octet. 6. After you are satisfied with the Lewis structure that you have constructed, perform a final electron count. © McGraw Hill LLC 50  Drawing Lewis Structures of Covalent Compounds Draw the Lewis structure of carbon dioxide, CO2. 1. Draw a skeletal structure of the molecule Arrange the atoms in their most probable order. C-O-O and/or O-C-O Find the electronegativity of O and C. O=3.5 & C=2.5 Place the least electronegative atom as the central atom, here carbon is the central atom. Result is the O-C-O structure from above. © McGraw Hill LLC 51  Drawing Lewis Structures of Covalent Compounds 1 Draw the Lewis structure of carbon dioxide, CO2. 2. Find the number of valence electrons for each atom and the total for the compound. 1 C atom 4 valence electrons 4 e  2 O atoms 6 valence electrons 12 e  16 e  total 3. Use electron pairs to connect the C to each O with a single bond: O C O 4. Place electron pairs around the atoms: This satisfies the rule for the O atoms, but not for C. © McGraw Hill LLC 52  Drawing Lewis Structures of Covalent Compounds 2 Draw the Lewis structure of carbon dioxide, CO2. 5. Redistribute the electrons moving 2 e- from each O, placing them between C  O In this structure, the octet rule is satisfied. This is the most probable structure. Four electrons are between C and O. These electrons are shared in covalent bonds. Four electrons in this arrangement signify a double bond. 6. Recheck the electron distribution 8 electron pairs = 16 valence electrons, number counted at start. 8 electrons around each atom, octet rule satisfied. © McGraw Hill LLC 53  Lewis Structure of Polyatomic Anions 1 Draw the Lewis structure of carbonate ion, CO32 1. Draw a skeletal structure of the molecule. Carbon is less electronegative than oxygen: This makes carbon the central atom. Skeletal structure and charge. 2. The total number of valence electrons is determined by adding one electron for each unit of negative charge: 1 C atom 4 valence electrons 4 e  3 O atoms 6 valence electron 18 e  + 2 negative charges 2 e  24 e  total 3. Distribute these e  around the skeletal structure. © McGraw Hill LLC 54  Lewis Structure of Polyatomic Anions 2 Draw the Lewis structure of carbonate ion, CO32 4. Distributing the electrons around the central carbon atom (4 bonds) and around the surrounding O atoms attempting to satisfy the octet rule results in: 5. This satisfies the octet rule for the 3 oxygen, but not for the carbon. 6. Move a lone pair from one of the O atoms to form another bond with C: © McGraw Hill LLC 55  Lewis Structures Practice Using the guidelines presented, write Lewis structures for the following: 1. H2O 2. NH3 3. CO2 4. NH  4 5. N2 © McGraw Hill LLC 56  Lewis Structure, Stability, Multiple Bonds, and Bond Energies Single bond - one pair of electrons are shared between two atoms. Double bond - two pairs of electrons are shared between two atoms. Triple bond - three pairs of electrons are shared between two atoms. © McGraw Hill LLC 57  Bond Energy and Bond Length Bond energy - the amount of energy required to break a bond holding two atoms together. triple bond > double bond > single bond Bond length - the distance separating the nuclei of two adjacent atoms. single bond > double bond > triple bond © McGraw Hill LLC 58  Lewis Structures and Resonance Write the Lewis structure of CO32 © McGraw Hill LLC 59  Carbonate Ion Resonance In some cases it is possible to write more than one Lewis structure that satisfies the octet rule for a particular compound. Experimental evidence shows all bonds are the same length, meaning there is not really any double bond in this ion: None of theses three Lewis structures exist, but the actual structure is an average or hybrid of these three Lewis structures. Resonance - two or more Lewis structures that contribute to the real structure. © McGraw Hill LLC 60  Lewis Structures and Exceptions to the Octet Rule 1. Incomplete octet - less then eight electrons around an atom other than H: Let’s look at BeH2. 1 Be atom 2 valence electrons 2 e  2 H atoms 1 valence electrons 2 e  total 4 e  Resulting Lewis structure: H  Be  H © McGraw Hill LLC 61  Odd Electron 2. Odd electron - if there is an odd number of valence electrons, it is not possible to give every atom eight electrons: Let’s look at NO, nitric oxide. It is impossible to pair all electrons as the compound contains an ODD number of valence electrons. © McGraw Hill LLC 62  Expanded Octet 3. Expanded octet - an element in the 3rd period or below may have 10 and 12 electrons around it: Expanded octet is the most common exception. Consider the Lewis structure of PF5. Phosphorus is a third period element. 1 P atom 5 valence electrons 5 e  5 F atoms 7 valence electrons 35 e  40 e  total Distributing the electrons results in this Lewis structure. © McGraw Hill LLC 63  Lewis Structures and Molecular Geometry; VSEPR Theory Molecular shape plays a large part in determining properties and shape. VSEPR theory - Valance Shell Electron Pair Repulsion theory: Used to predict the shape of the molecules. All electrons around the central atom arrange themselves so they can be as far away from each other as possible – to minimize electronic repulsion. © McGraw Hill LLC 64  VSEPR Theory In the covalent bond, bonding electrons are localized around the nucleus. The covalent bond is directional, having a specific orientation in space between the bonded atoms. Ionic bonds have electrostatic forces which have no specific orientation in space. © McGraw Hill LLC 65  A Stable Exception to the Octet Rule Consider BeH2: Only 4 electrons surround the beryllium atom. These electrons in the bonds to the two atoms have minimal repulsion when located on opposite sides of the structure. Linear structure having bond angles of 180°. © McGraw Hill LLC 66  Another Stable Exception to the Octet Rule Consider BF3: There are 3 bonded atoms around the central atom. These bonded atoms have minimal repulsion when placed in a plane, forming a triangle. Trigonal planar structure with bond angles of 120°. © McGraw Hill LLC 67  Basic Electron Pair Repulsion of a Full Octet Consider CH4 There are 4 bonded atoms around the central carbon. Minimal electron repulsion when electrons are placed at the four corners of a tetrahedron. Each H-C-H bond angle is 109.5°. Tetrahedral is the primary structure of a full octet. © McGraw Hill LLC 68  Basic Electron Pair Repulsion of a Full Octet with One Lone Pair Consider NH3: There are three bonded atoms and one lone pair (four groups). A lone pair is more electronegative with a greater electron repulsion. The lone pair takes one of the corners of the tetrahedron without being visible, distorting the arrangement of electron pairs. Ammonia has a trigonal pyramidal structure with 107° bond angles. © McGraw Hill LLC 69  Basic Electron Pair Repulsion of a Full Octet with Two Lone Pairs Consider H2O: There are two bonded atoms and two lone pair (four groups). All 4 electron pairs are approximately tetrahedral to each other. The lone pairs take two of the corners of the tetrahedron without being visible, distorting the arrangement of electron pairs. Water has a bent or angular structure with 104.5° bond angles. © McGraw Hill LLC 70  Predicting Geometric Shape Using Electron Pairs Bonded Atoms Nonbonding Bond Angle Molecular Geometry Example Lone Electron Pairs 2 0 180 Linear CO 2 3 0 120 Trigonal planar SO3 2 1