Chapter 9 and 10 Summaries PDF

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Chapter 9 and 10 Summaries which covers Periodic Properties of the Elements and Chemical Bonding I: The Lewis Model...

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Chapter 9 and 10 Summaries Chapter 9 Chemical Bonding I: Lewis Theory Bonding Theories  explain how and why atoms attach together  explain why some combinations of atoms are stable and others are not  why is water H2O, not HO or H3O  one of the simplest bonding theories was developed by G.N. Lewis and is called Lewis Theory  Lewis Theory emphasizes valence electrons to explain bonding  using Lewis Theory, we can draw models – called Lewis structures – that allow us to predict many  aka Electron Dot Structures  such as molecular shape, size, polarity Section 9.1 Why Do Atoms Bond?  processes are spontaneous if they result in a system with lower potential energy  chemical bonds form because they lower the potential energy between the charged particles that compose atoms  the potential energy between charged particles is directly proportional to the product of the charges  the potential energy between charged particles is inversely proportional to the distance between the charges Potential Energy between charged Particles  0 is a constant  = 8.85 x 10-12 C2/J∙m  for charges with the same sign, Epotential is + and the magnitude gets less positive as the particles get farther apart  for charges with the opposite signs, Epotential is  and the magnitude gets more negative as the particles get closer together  remember: the more negative the potential energy, the more stable the system becomes Bonding  A chemical bond forms when the potential energy of the bonded atoms is less than the potential energy of the separate atoms  Have to consider following interactions:  Nucleus-to-nucleus repulsion  Electron-to-electron repulsion  Nucleus-to-electron attraction

Types of Bonding

Section 9.2 Ionic Bonds  When metals bond to nonmetals, some electrons from the metal atoms are transferred to the nonmetal atoms  Metals have low ionization energy, relatively easy to remove an electron from  Nonmetals have high electron affinities, relatively good to add electrons to Covalent Bonds  Nonmetals have relatively high ionization energies, so it is difficult to remove electrons from them  When nonmetals bond together, it is better in terms of potential energy for the atoms to share valence electrons  potential energy lowest when the electrons are between the nuclei  Shared electrons hold the atoms together by attracting nuclei of both atoms

Lewis Symbols of Atoms  Aka electron dot symbols  Use symbol of element to represent nucleus and inner electrons  Use dots around the symbol to represent valence electrons  pair first two electrons for the s orbital  put one electron on each open side for p electrons  then pair rest of the p electrons

What We Know  The noble gases are the least reactive group of elements  The alkali metals are the most reactive metals and their atoms almost always lose 1 electron when they react  The halogens are the most reactive group of nonmetals and in a lot of reactions they gain 1 electron Stable Electron Arrangements and Ion Charge  Metals form cations by losing enough electrons to get the same electron configuration as the previous noble gas  Nonmetals form anions by gaining enough electrons to get the same electron configuration as the next noble gas  The noble gas electron configuration must be very stable Section 9.3 Octet Rule  When atoms bond, they tend to gain, lose, or share electrons to result in 8 valence electrons  ns2np6  noble gas configuration  many exceptions  H, Li, Be, B attain an electron configuration like He  He = 2 valence electrons  Li loses its one valence electron  H shares or gains one electron  Though it commonly loses its one electron to become H+  Be loses 2 electrons to become Be2+  B loses 3 electrons to become B3+  Though it commonly shares its three electrons in covalent bonds, resulting in 6 valence electrons Lewis Theory  The basis of Lewis Theory is that there are certain electron arrangements in the atom that are more stable  octet rule  Bonding occurs so atoms attain a more stable electron configuration  More stable = lower potential energy  No attempt to quantify the energy as the calculation is extremely complex

Properties of Ionic Compounds

Lewis Theory and Ionic Bonding  Lewis symbols can be used to represent the transfer of electrons from metal atom to nonmetal atom, resulting in ions that are attracted to each other and therefore bond Section 9.4 Predicting Ionic Formulas using Lewis Symbols  Electrons are transferred until the metal loses all its valence electrons and the nonmetal has an octet  Numbers of atoms are adjusted so the electron transfer comes out even Energetics of Ionic Bond Formation  The ionization energy of the metal is endothermic  Na(s) → Na+(g) + 1 e ─ DH° = +603 kJ/mol  the electron affinity of the nonmetal is exothermic  ½Cl2(g) + 1 e ─ → Cl─(g) DH° = ─ 227 kJ/mol  Generally, the ionization energy of the metal is larger than the electron affinity of the nonmetal, therefore the formation of the ionic compound should be endothermic  But the heat of formation of most ionic compounds is exothermic and generally large; Why?  Na(s) + ½Cl2(g) → NaCl(s) DH°f = -410 kJ/mol Lattice Energy  The lattice energy is the energy released when the solid crystal forms from separate ions in the gas state  Always exothermic  Hard to measure directly, but can be calculated from knowledge of other processes  Lattice energy depends directly on size of charges and inversely on distance between ions

Trends in Lattice Energy Ion Size  The force of attraction between charged particles is inversely proportional to the distance between them  Larger ions mean the center of positive charge (nucleus of the cation) is farther away from negative charge (electrons of the anion)  larger ion = weaker attraction = smaller lattice energy Trends in Lattice Energy Ion Charge  The force of attraction between oppositely charged particles is directly proportional to the product of the charges  Larger charge means the ions are more strongly attracted  larger charge = stronger attraction = larger lattice energy  Of the two factors, ion charge generally more important

Ionic Bonding Model vs. Reality  Ionic compounds have high melting points and boiling points  MP generally > 300°C  all ionic compounds are solids at room temperature  Because the attractions between ions are strong, breaking down the crystal requires a lot of energy  The stronger the attraction (larger the lattice energy), the higher the melting point Section 9.5 Single Covalent Bonds  Two atoms share a pair of electrons  2 electrons  One atom may have more than one single bond Double Covalent Bond  Two atoms sharing two pairs of electrons  4 electrons

Covalent Bonding Predictions from Lewis Theory  Lewis theory allows us to predict the formulas of molecules  Lewis theory predicts that some combinations should be stable, while others should not  Because the stable combinations result in “octets”  Lewis theory predicts in covalent bonding that the attractions between atoms are directional  The shared electrons are most stable between the bonding atoms  Resulting in molecules rather than an array Electronegativity and Bond Polarity  If difference in electronegativity between bonded atoms is 0, the bond is pure covalent  equal sharing  If difference in electronegativity between bonded atoms is 0.1 to 0.4, the bond is nonpolar covalent  If difference in electronegativity between bonded atoms 0.5 to 1.9, the bond is polar covalent  If difference in electronegativity between bonded atoms larger than or equal to 2.0, the bond is ionic Chapter 10 Chemical BONDING Chemical Bond  A bond results from the attraction of nuclei for electrons  All atoms trying to achieve a stable octet  IN OTHER WORDS  The p+ in one nucleus are attracted to the e- of another atom  Electronegativity Section 10.1 Two Major Types of Bonding  Ionic Bonding  forms ionic compounds  transfer of e Covalent Bonding  forms molecules  sharing e-

One minor type of bonding  Metallic bonding  Occurs between like atoms of a metal in the free state  Valence e- are mobile (move freely among all metal atoms)  Positive ions in a sea of electrons  Metallic characteristics  High mp temps, ductile, malleable, shiny  Hard substances  Good conductors of heat and electricity as (s) and (l)

Section 10.2 Ionic Bonding  Electrons are transferred between valence shells of atoms  Ionic compounds are made of ions  ionic compounds are called Salts or Crystals IONic Bonding  Electronegativity difference > 2.0  Look up e-neg of the atoms in the bond and subtract  NaCl  CaCl2  Compounds with polyatomic ions  NaNO3

Properties of Ionic Compounds

   

Hard solid @ 22oC High mp temperatures Nonconductors of electricity in solid phase Good conductors in liquid phase or dissolved in water (aq)

Covalent Bonding  Pairs of e- are shared between non-metal atoms  electronegativity difference < 2.0  forms polyatomic ions Properties of Molecular Substances  Low m.p. temp and b.p. temps  Relatively soft solids as compared to ionic compounds  Nonconductors of electricity in any phase Draw the Lewis Diagrams    

LiF MgO CaCl2 K2S

Draw Lewis Dot Structures  You may represent valence electrons from different atoms with the following symbols x  CO2  NH3

Draw the Lewis Dot Diagram for polyatomic ions  Count all valence e- needed for covalent bonding  Add or subtract other electrons based on the charge Draw Polyatomics  Ammonium  Sulfate Types of Covalent Bonds  NON-Polar bonds  Electrons shared evenly in the bond  E-neg difference is zero Types of Covalent Bonds  Polar bond

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Electrons unevenly shared E-neg difference greater than zero but less than 2.0

Place these molecules in order of increasing bond polarity      

HCl CH4 CO2 NH3 N2 HF

Space filling model “Electron-Cloud” model

Section 10.3 Water is a bent molecule

VSEPR Theory  Valence Shell Electron Pair Repulsion Theory  Electron pairs orient themselves in order to minimize repulsive forces.

 Types of e- Pairs  Bonding pairs - form bonds  Lone pairs - nonbonding e Linear (straight line)

 Space filling model

 Bent

Ball and stick model

Space filling model Trigonal pyramid

Ball and stick model

Space filling model Tetrahedral

Ball and stick model

Space filling model

Hydrogen “Bonding”  Strong polar attraction  Occurs ONLY between H of one molecule and N, O, F of another

Section 10.4 Intermolecular forces dictate chemical properties  Strong intermolecular forces cause high b.p., m.p. and slow evaporation (low vapor pressure) of a substance. Intermolecular attractions  Attractions between molecules  Weak attractive forces between non-polar molecules  Strong attraction between special polar molecules...