Chemical Bonding - Collier's Classroom

CHEMICAL BONDING CHEMICAL BOND attractive forces between 2 atoms overcomes the repulsion of 2 positively charged nuclei Lewis Symbols represents atoms using the element symbol and valence electrons as dots only valence electrons participate in

bonding LEWIS SYMBOLS Each side of the symbol represents an atomic orbital, which may hold up to two electrons Place 1 dot on each side until 4 dots are around the symbol Add a second dot to each side # of dots = # of valence electrons Each unpaired dot can form a chemical bond LETS TRY SOME:

Na C Pb Mg F Ne Cl-

Ca2+ N3- BONDING AND NOMENCLATURE IONIC COMPOUNDS Transfer of an electron Metal and nonmetal Cation and anion Electrically neutral Each atom achieves a Noble Gas

configuration FORMULA UNIT Chemical formula: represents the kinds and numbers of elements in a compound NaCl Doesnt represent a single discrete unit Formula Unit: the lowest number ratio of ions in an ionic compound NaCl: 1 Na and 1 Cl

Mg2Cl: 2 Mg and 1 Cl AlBr3: 1 Al and 3 Br PREDICATING FORMULAS FOR IONIC COMPOUNDS Determine the charge of each element Determine the ratio needed to have a neutral atom Write the formula LETS TRY SOME:

1. sodium and oxygen 2. lithium and bromine 3. aluminum and oxygen 4. barium and fluorine LETS TRY SOME: 1. sodium and oxygen: Na2O

2. lithium and bromine: LiBr 3. aluminum and oxygen: Al2O3 4. barium and fluorine: BaF2 NAMING IONIC COMPOUNDS 1st name the cation then the anion Cation: name doesnt change Anion: suffix become ide

Chlorine Chloride IONIC COMPOUNDS: NAME THE FOLLOWING NaCl Na2O Li2S

AlBr3 CaO IONIC COMPOUNDS: NAME THE FOLLOWING NaCl sodium chloride Na2O sodium oxide

Li2S lithium sulfide AlBr3 aluminum bromide CaO calcium oxide TRANSITION METALS

cations have multiple ion charges use a Roman numeral to give the charge following the name Examples: FeCl3 is iron(III) chloride FeCl2 is iron(II) chloride CuO

is copper(II) oxide POLYATOMIC IONS 2 or more atoms bonded together with an overall positive or negative charge Within the ion itself, the atoms are bonded using covalent bonds Examples: NH4+

ammonium ion SO42- sulfate ion NAMING POLYATOMIC COMPOUNDS (P. 257) NH4Cl BaSO4 Fe(NO3)3

Cu(HCO3)3 Ca(OH)2 NAMING POLYATOMIC COMPOUNDS (P. 257) NH4Cl: ammonium chloride BaSO4: barium sulfate

Fe(NO3)3 : iron (III) nitrate Cu(HCO3)3: copper (III) bicarbonate Ca(OH)2: calcium hydroxide WRITING IONIC FORMULAS FROM COMPOUNDS NAME Writing Formulas of Ionic Compounds from the Name of the Compound Determine

the charge of each ion Compound must be neutral Example: barium chloride barium is +2, chloride is -1 Formula is BaCl2 Write the Formulas for the following ionic compounds: sodium sulfate

ammonium sulfide magnesium phosphate chromium (II) sulfate Write the Formulas for the following ionic compounds: sodium sulfate:

ammonium sulfide magnesium phosphate chromium (II) sulfate BONDING IN METALS Metallic Bonds: the attraction of freefloating valence electrons from the positively charge metal ions Forces of attraction that hold metals

together valence electrons of metals sea of electrons; electrons are mobile and drift freely Crystalline Structure of Metals: metal atoms are arranged in very compact, orderly patterns ALLOYS mixtures composed of two or more elements, at least one of which is a metal alloys are often superior to their component elements

brass: zinc and copper sterling silver: silver and copper COVALENT COMPOUNDS COVALENT BONDS Two nonmetal atoms share electrons Form between atoms with similar tendencies to gain or lose electrons Form

bonds to obtain a full octet Diatomic Elements H2, N2, O2, F2, Cl2, Br2, I2 The Magic Seven DIATOMIC ELEMENTS TYPES OF BONDS Single bond: share two electrons or 1 pair H2

Double bond: share four electrons or 2 pairs O2 Triple bond: share six electrons or 3 pairs N2 Unshared Pair: a lone pair or nonbonding pair BOND ENERGIES AND LENGTH

Bond dissociation energy - the energy required to break a bond triple bond > double bond > single bond Bond length - the distance separating the nuclei of two adjacent atoms single bond > double bond > triple bond Resonance: the structure that occurs when it is possible to draw two or more

valid electron dot structures the actual bonding of resonance structures is a hybrid of the possible structures BOND RESONANCE EXCEPTIONS TO THE OCTET RULE Odd number of electrons NO2 Expanded Octet: a few atoms expand the octet to include ten or twelve electrons PCl3

and PCl5 SF6 Less than an Octet BeH2 BF3 NAMING COVALENT COMPOUNDS Nonmetals Molecules covalently bonded compounds Mono Tri

1 3 Penta 5 Hepta 7 Nona 9 Di -2 Tetra 4 Hexa -6 Octa 8 Deca- 10 NAMING COVALENT COMPOUNDS The names of the elements are written in the order in which they appear in the

formula A prefix indicates the number of each kind of atom 1st element: if only one is present, no prefix is used Last element: use the suffix ide Example: CO is carbon monoxide The final vowel in a prefix is often dropped before a vowel in the stem name Correct: monoxide Not: monooxide

NAME THESE COVALENT COMPOUNDS SiO2: N2O5: CCl4: IF7 : NAME THESE COVALENT

COMPOUNDS SiO2: Silicon Dioxide N2O5: Dinitrogen Pentoxide CCl4: Carbon Tetrachloride IF7 : Iodine Heptafluoride WRITING FORMULAS OF COVALENT COMPOUNDS

Use the prefixes in the names to determine the subscripts for the elements Examples: nitrogen trichloride diphosphorus NCl3 pentoxide P 2 O5

Some common names that are used: H2O water C2H5OH NH3 ammonia ethanol C6H12O6 - glucose PROVIDE FORMULAS FOR THESE COVALENT COMPOUNDS nitrogen monoxide

dinitrogen tetroxide diphosphorus pentoxide nitrogen trifluoride PROVIDE FORMULAS FOR THESE COVALENT COMPOUNDS nitrogen monoxide NO

dinitrogen tetroxide diphosphorus pentoxide P2O5 nitrogen trifluoride N2O4 NF3 PROPERTIES OF IONIC & COVALENT COMPOUNDS PROPERTIES OF IONIC AND

COVALENT COMPOUNDS Physical State Ionic compounds: solids Covalent compounds: solids, liquids, and gases Melting and Boiling Points Ionic compounds: higher melting & boiling points

ionic bonds are stronger more energy is needed to break bond melt at several hundred C PROPERTIES OF IONIC AND COVALENT COMPOUNDS Structure of Compounds in the Solid State Ionic compounds: crystalline Covalent compounds are crystalline or amorphous having no regular structure PROPERTIES OF IONIC AND COVALENT COMPOUNDS

Solutions Ionic compounds dissolve in water Ions dissociate or separate Electrolytes - ions in solution conduct electricity Covalent Molecules: solids do not dissociate

Nonelectrolytes: do not conduct electricity Ionic Covalent Composed of Metal + nonmetal 2 nonmetals Electrons Transferred Shared Physical state

Solid / crystal Any / crystal OR amorphous Dissociation Yes, electrolytes Boiling/Melting High No, nonelectrolytes Low Acids: molecular compound that produces H+ (hydrogen ions) in solution

Memorize these Acids HCl HBr HI H2SO4 HNO3 HClO4 HC2H3O2 Hydrochloric acid Hydrobromic acid Hydroiodic acid Surlfuric acid Nitric acid Perchoric acid Acetic acid NAMING ACIDS

1. When the anion ends in ide, the acid name begins with hydro- and the anion has the suffix ic then acid. chloride 2. When the anion ends in ite, the anion name has the suffix ous then acid. sulfite hydrochloric acid sulfurous acid 3. When the anion ends in ate, the

anion name has the suffix ic then acid. nitrate nitric acid MOLECULAR STRUCTURE DRAWING LEWIS STRUCTURES ON MOLECULES AND POLYATOMIC IONS Lewis Structure Guidelines Use chemical symbols for the elements Least electronegative atom: center Hydrogen and halogens: outside Carbon: chains of carbon-carbon covalent bonds Determine the # of valence electrons for

each atom Find the total number of valence electrons Polyatomic cations: subtract one e- for every (+) charge Polyatomic anions: add one e- for every (-) charge Connect the central atom to surrounding atoms using electron pairs Complete octets of the atoms bonded to the central atom Hydrogen: only 2 electrons Electrons not involved in bonding are

represented as lone pairs Count the electrons in the diagram & compare to total from step 2 CARBON DIOXIDE Draw the Lewis structure of carbon dioxide, CO2 Arrange the atoms in their most probable order C-O-O

or O-C-O Place the least electronegative atom, carbon, in the center, O-C-O Find the number of valence electrons for each atom and the total for the compound 1 C atom x 4 valence electrons = 4 e 2 O atoms x 6 valence electrons = 12 e 16 e- total Use electron pairs to connect the C to each O with a single bond O

Place electron pairs around the atoms : :C:O O:C:O: This satisfies the rule for the O atoms, but not for C Redistribute the electrons moving 2 e- from each O, placing them between C:O In this structure, the octet rule is satisfied Four

O::C::O electrons in this arrangement signify a double bond Recheck the electron distribution 8 electron pairs = 16 valence electrons, number counted at start 8 electrons around each atom, octet rule satisfied Using the guidelines presented, write Lewis structures for the following: H2O

NH3 CO2 NH4+ CO32- N2

LEWIS STRUCTURES AND RESONANCE Actually, all bonds are the same length, so theres no true double or triple bonds the actual structure is an average the three Lewis structures Resonance - two or more Lewis structures that contribute to the real structure EXCEPTIONS TO THE OCTET RULE

Incomplete octet - less than 8 earound an atom other than H Lets look at BeH2 1 Be atom x 2 valence electrons = 2 e2 H atoms x 1 valence electrons = 2 e total 4 e Resulting Lewis structure: H : Be : H or H Be H

EXCEPTIONS CONTINUED Odd electron - if there is an odd number of valence electrons, it is not possible to give every atom eight electrons NO, nitric oxide I t is impossible to pair all electrons as the compound contains an ODD number of valence electrons 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 x 5 valence electrons = 5 e 5 F atoms x 7 valence electrons = 35 e 40 e- total BONDING THEORIES Molecular Orbitals: when two atoms combine, their atomic orbitals combine each atomic orbital can contain 2 electrons Sigma Bonds (): 2 s atomic orbitals ): 2 s atomic orbitals combine

Oval or oblong Pi Bond (): 2 p atomic orbitals overlap): 2 p atomic orbitals overlap Mirror image jelly beans split by the bond axis Sigma bonds are stronger than Pi bonds HYBRID ORBITALS Carbon: 2s22p2 Draw the orbital diagram. How many lone

electrons does it have? Draw the Lewis dot structure? How many lone pairs does it have? Why? 1-2s electron is moved up to the 2p orbital so that it can have 4 bonds and a full octet. Are the bonds the same? Yes They hybridize to form 4-sp3 orbitals They extend further into space than s or p orbitals Form 4 C-H sigma bonds, which are unusually strong covalent bonds

HYBRIDIZATION AND DOUBLE BONDS C2H4: ethane Draw the Lewis diagram 4 single bonds and1 double bond The C-H bonds are sp2 from 1-2s and 2-2p atomic orbitals The sp2 orbitals are sigma bonds The 3rd sp2 orbital form a C-C sigma bond The

nonhybridized 2p carbon orbitals form 1-pi bond In chemical reactions, a pi bond is more likely to break than a sigma MOLECULAR GEOMETRY MOLECULAR GEOMETRY VSEPR theory - Valance Shell Electron Pair Repulsion theory covalent bonds predicts a molecules shape

electrons around an atom arrange themselves to maximize their distance from each other to minimize electronic repulsion MOLECULAR GEOMETRY Linear structure 2 single bonds or 2 double bonds; 0 lone pairs bond angles of 180

BeH2 or CO2 Bent or Angular 2 shared electron pairs & 2 lone pairs 104.5 bond angles H2 O LINEAR AND BENT MOLECULES MOLECULAR GEOMETRY Trigonal planar 3

single bonds or 2 single bonds & 1 double; 0 lone pairs bond angles of 120 BF3, SO3 Trigonal Pyramidal 3 shared electron pairs and 1 lone pair 107 angles NH3

MOLECULAR GEOMETRY Tetrahedron 4 shared electron and 0 lone pairs Bond angle: 109.5 CH4 Trigonal Bipyramidal 5 atoms around the central atom

Bond angles: 90 , 120 , 180 MOLECULAR GEOMETRY T-shaped 3 Bond angles: 90 , 120 , <180 Octahedral 6

atoms and 2 lone pairs atoms around the central atom Bond angles: 90 , 180 Square Planar 4 atoms and 2 lone pairs Bond angles: 90, 180 DETERMINE THE MOLECULAR GEOMETRY PCl3

SO2 PH3 SiH4 PCl5 SF6 DETERMINE THE

MOLECULAR GEOMETRY PCl3 Trigonal pyramidal SO2 Bent PH3 Trigonal pyramidal SiH4 tetrahedral PCl5 Trigonal

bipyramidal SF6 octahedral POLARITY POLARITY Polar molecules Molecules that are polar behave as a dipole (having two poles or ends) One end is positively charged the other is negatively charged Atoms in molecules, covalent bonds, share electrons, but dont share them equally

Polar Bonds The more electronegative atom will pull the electrons toward it If the electrons are not shared equally polar bond Positive end of the bond, the less electronegative atom Dipole: a molecule with two ends; positive and negative POLAR BOND NONPOLAR BONDS When electrons are shared equally Bond

between two of the same atom: N2 A central atom with 4 identical species surrounding it: CH4 A central atom with 3 identical species and no lone pairs; BH3 DETERMINE WHETHER THE FOLLOWING MOLECULES ARE POLAR:

NH3 O2 BH3 HF BH2F CH4 CO2 H2O CH3F DETERMINE WHETHER THE FOLLOWING MOLECULES ARE POLAR:

NH3:P O2: NP BH3: NP HF: P BH2F: P CH4: NP CO2 : NP H2O: P CH3F: P INTERMOLECUL AR FORCES

PROPERTIES BASED ON ELECTRONIC STRUCTURE AND MOLECULAR GEOMETRY Intramolecular forces: attractive forces within molecules chemical bonds Intermolecular forces: attractive forces between molecules Solubility - the maximum amount of solute that dissolves in a given amount of solvent at a specific temperature Like dissolves like

Polar dissolves in polar; Nonpolar dissolves in nonpolar INTERMOLECULAR FORCES Intermolecular attractions are weaker than either ionic or covalent bonds. Van der Waals Forces: dipole interaction and dispersion forces Dipole interactions: polar molecules are attracted to one another Much weaker than ionic bonds

DISPERSION FORCES when the moving electrons happen to be momentarily more on the side of the a molecule closest to a neighboring molecule, their electric force influences the neighboring molecules electrons to be momentarily more on the opposite side, this causes an attraction between the two molecules Nonpolar molecules Caused by the motion of electrons Weakest intermolecular force

Increases as the number of electrons increase HYDROGEN BONDS hydrogen that is covalently bonded to a very electronegative atom is also weakly bonded to an unshared electron pair of another electronegative atom H2O, DNA WATER & AMMONIA - end of NH3, N, is attracted to + end

of H2O molecule, H + end of NH3, H, is attracted to - end of H2O molecules, O The attractive forces, hydrogen bonds, distribute NH3 molecules throughout H2O, forming a homogeneous solution FACTORS INFLUENCING BOILING & MELTING POINTS Molecular mass Larger molecules have higher m.p. and b.p. b/c it is more difficult to convert a larger mass to another phase

Polarity Polar molecules have higher m.p. and b.p. than nonpolar molecules of similar molecular mass due to their stronger attractive force

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