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Chapter 8 Chemical Bonding

Chapter 8 Chemical Bonding. Section 1: Types of Chemical Bonds Elements typically exist as compounds in nature. They are not often found alone. Type of bonding influences chemical and physical properties. Example: graphite and diamond.

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Chapter 8 Chemical Bonding

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  1. Chapter 8 Chemical Bonding

  2. Section 1: Types of Chemical Bonds Elements typically exist as compounds in nature. They are not often found alone. Type of bonding influences chemical and physical properties. Example: graphite and diamond. Are silicon dioxide and carbon dioxide similar or different? Why? Bond energy - The energy required to break a bond allows us to obtain information about the strength of a bonding interaction

  3. How do we know that a solution of sodium chloride actually consists of sodium ions and chloride ions? It conducts electricity This indicates to us that electrons are transferred from one substance to another.

  4. Why do atoms transfer electrons from one element to another? • This allows elements to achieve the lowest possible energy by behaving in this way • For sodium and chloride, the electron transfer occurs because chlorine has a strong attraction for an extra electron and the oppositely charged ions have great attraction for each other, providing driving forces for the procedure • Sodium chloride is solid, has a high melting point and has great thermal stability because of the electrostatic attractions of closely packed, oppositely charged ions

  5. Ionic Bonding- when bonding forces occur due to the attraction between oppositely charged ions • creates great thermal stability due to electrostatic attractions. • Electrostatic has to do with electromagnetic forces Think light waves- electron waves • Forms when an atom that loses electrons relatively easily reacts with an atom that has a high affinity for electrons • Metal and nonmetal react to form an ionic compound

  6. Coulomb’s law gives the energy of interaction between a pair of ions • E = 2.31 x 10-19 J*nm (Q1*Q2/ r) • E is energy of interaction between a pair of ions. • R is distance between ion centers in nanometers, • Q1 and Q2 are numerical ion charges. • Negative sign indicates attractive force - ion pair has lower energy than the separated ions. • Repulsive energy is calculated for like charges with a positive sign.

  7. Why will 2 hydrogen atoms form H2 rather than staying as H? • The energy of the system is lower than the energy of 2 H atoms • System acts to minimize the sum of the repulsive energy terms and the attractive energy term • Distance at which energy is minimum is . • H2 is stable because the electrons are between the two hydrogen nuclei, so both are attracted by both protons. The attraction between electrons and protons results in a force that pulls the protons toward each other and balances out the proton-proton and electron-electron repulsions.

  8. Covalent bonding • bonds in which electrons are shared by nuclei • Polar covalent bonding • bonds in which electrons are shared unequally between atoms

  9. Section 8.2: Electronegativity • Electronegativity • the ability of an atom in a molecule to attract shared electrons to itself • Helps us determine bond type

  10. Explain how Linus Pauling determined electronegativities for atoms. • the compared the measured H—X bond energy with the expected H—X bond energy. • The expected H—X bond energy = • H—H bond energy + X—X bond energy /2 • The difference between the actual and expected bond energies is indicated by • ∆ = (H—X)act – (H—X)exp • If H and X had the same electronegativity, then the difference would be zero. If X has a greater electronegativity than H, the shared electrons will be closer to the X atom, so the molecule will be polar

  11. REMEMBER THAT THE ELECTRONEGATIVITY CHART IS ON PAGE 334. • Is there a clear division between covalent and ionic bonds? No. There is a continuum. • From approximately 0.0-0.4 you have a non-polar covalent bond • From approximately 0.4-1.9 you have a polar covalent bond • From approximately 2.0 to 4.0 you have an ionic bond.

  12. Section 8.3: Bond Polarity and Dipole Moments • Dipolar/ Dipole Moment- molecules with a center of positive charge and a center of negative charge • Dipolar character is indicated by an arrow pointing toward the negative charge center.

  13. Is it possible for a molecule to have polar bonds but not have a dipole moment? • Yes- the individual bond polarities are arranged in such a way that they cancel each other out. • There is no way that these molecules will line up in an electric field

  14. IN CLASS PRACTICE/ DISCUSSION: p. 387 # 114

  15. Section 8.4: Ions: Electron Configurations and Sizes • What is true about the electron configurations of atoms in a stable compound? • They have a noble gas arrangement of electrons • How do nonmetallic elements achieve a noble gas electron configuration? • By sharing electrons with other nonmetals to form covalent bonds or by taking electrons from metals to form ions Examples: NaCl, CaO

  16. Ionic compound usually refers to what state of matter? solid • Describe ionic compounds: • The ions are close together to minimize the repulsions among like charges and maximize the attractions between opposite charges. • In gases, the ions are not close enough to interact, so when we discuss stability of ionic compounds, we refer to the solid state • You should be able to predict the charges of ions that form based upon electron configurations • How do you predict the formula of the ionic compounds? Remember that compounds are neutral or have equal numbers of positive and negative charges.

  17. What are some of the exceptions to the rule that ions form noble gas configurations? • Tin (II) and Tin (IV), Lead (II) and Lead (IV), Tl (I) and Tl (III) • Reason for exceptions will not be discussed currently • What factors influence ionic size? • positive ions are smaller than parent atoms because electrons are removed. • Addition of electrons to neutral atoms produces anions larger than parent atoms. • Ion size increases down a group because more energy levels of electrons are added. • Across a row, sizes vary because of cation and anion differences

  18. What are isoelectronic ions? • Ions with the same number of electrons • What are the 2 factors you need to consider when trying to determine sizes of ions? • # of electrons and # of protons

  19. Section 8.5: Formation of Binary Ionic Compounds • Lattice energy- the change in energy that takes place when separated gaseous ions are packed together to form an ionic solid. • indicates how strongly the ions attract each other in the solid state • The sign of lattice energy is negative because energy is released when the gases combine to form a solid • Equations: • Lattice energy = k (Q1Q2/r) • k is a proportionality constant dependent upon the structure of the solid and the electron configurations of the ions Practice: p. 383-384 #’s 45 - 52

  20. Section 8.6: Partial Ionic character of Covalent Bonds • Ionic character increases with: • electronegativity difference • None of the bonds is completely ionic. • Compounds with more than 50% ionic character are considered ionic

  21. Represents the bonds of elements in the gaseous state and do not necessarily apply to the solid state.

  22. Why is the existence of ions favored by the solid state? Many ions can interact and stabilize each other. What is the “official” definition of ionic compounds? Any compound that conducts an electric current when melted. Equations: % ionic character = measured dipole moment of X—Y / calculated dipole moment of X+Y-

  23. Section 8.7: The Covalent Chemical Bond: A Model • Reminders: • Chemical bonds- forces that cause groups of atoms to behave as a unit. • Chemical bonds occur because: bonds result from the tendency of a system to seek its lowest possible energy or to become more stable • Bonds represent: • a quantity of energy obtained from the overall molecular energy of stabilization in an arbitrary way. • Bonds are a model to describe what is happening • It simplifies what is happening and makes it understandable to us.

  24. Why is it useful to think of a protein as a group of C—C, C—H, C—N, C—O, and N—H bonds? This makes the molecules more understandable. We expect certain bonds to behave in certain ways What is the delocalization of electrons? The ability of electrons to move through the entire molecule- one of the flaws of the bond model See page 350 for information on models

  25. Section 8.8: Covalent Bond Energies and Chemical Reactions • Discuss AVERAGE vs REAL bond dissociation energies (CH4) • Single bond- type of bond in which one pair of electrons is shared • Double bond- atoms sharing 2 pairs of electrons • Triple bond- atoms sharing 3 pairs of electrons • How is bond length related to the number of shared electrons? • More shared electrons = shorter bond length

  26. Bond energies can be used to calculate what? approximate energies for reactions To break bonds, what must happen? energy must be added to the system endothermic- energy has + sign To form bonds, what must happen? energy is released exothermic- energy has – sign

  27. Equation: ∆H = ∑D(bonds broken) - ∑D(bonds formed) Energy of reaction = energy required to break bonds – energy required to form bonds D always has a positive sign and is bond energy PER MOLE

  28. Section 8.9 The Localized Electron Bonding Model • Localized electron (LE) Model • a molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms • Electron pairs are assumed to be localized on a particular atom or in the space between two atoms • Electron pairs on atoms are called lone pairs • Electrons between atoms are bonding pairs

  29. What are the 3 parts of the LE model? • Description of the valence electron arrangement in a molecule using Lewis Structures • Prediction of the geometry of the molecule using the VSEPR model • Description of the type of atomic orbitals used by the atoms to share electrons or hold lone pairs

  30. Section 8.10: Lewis Structures • Lewis Structures • show how the valence electrons are arranged among the atoms in the molecule • The most important requirement for the formation of a stable compound is: • that atoms achieve noble gas electron configurations • What electrons are included in Lewis structures? • just the valence • Duet rule- when H bonds to share 2 electrons • Octet Rule- when 8 electrons are required to fill the orbitals in order to make stable ions

  31. What are the rules for writing Lewis Structures? • Add all the valence electrons from all atoms • Use a pair of electrons to form a bond between each pair of bound atoms • Arrange remaining electrons to satisfy the duet rule for hydrogen and the octet rule for the other elements • Always start with single bonds. • Double bonds may be used later and triple bonds if the others don’t work. • C, N, O, S can form double bonds • N and C can form triple bonds

  32. Section 8.11: Exceptions to the Octet Rule • List and describe the exceptions to the Octet Rule • Boron tends to have fewer than 8 electrons. Usually it has only 6 electrons around it. • This results in very reactive compounds because the compounds are electron-deficient • Be also has fewer than 8 electrons • Elements in period 3 of the periodic table and beyond may have more than 8 electrons • 3) C, N, O, and F always obey the octet rule • 4)Second row elements never exceed the octet rule because orbitals 2s and 2p can hold only 8 electrons

  33. Summarize any rules about Lewis Structures that you have not previously included: When writing Lewis structures, satisfy the octet rule 1st and if electrons remain after the octet rule has been satisfied, place them on elements having available d orbitals How do you determine where extra electrons go when it is not immediately clear? assume extra electrons should be placed on the central atom

  34. Section 8.12: Resonance • Resonance: • occurs when more than one valid Lewis structure can be written for a particular molecule • All bonds in the resonance structure are equivalent. There are not separate double and single bonds. All are a hybrid of the bonds • Resonance structures: • Lewis structures of all possible forms of a molecule separated by a double headed arrow • indicates the actual structure is an average of all of the structures • Necessary because electrons are delocalized- able to move around the entire molecule

  35. Why are oxidation states useful? Why are they flawed? • Useful because it helps for bookkeeping in redox reactions. • Flawed because they are not realistic estimates of the actual charges on individual atoms in a molecule- not good for judging the appropriateness of Lewis structures • Formal Charge • difference between number of valence electrons on the free atom and the number of valence electrons assigned to the atom in the molecule • What must you know to determine formal charge? • Number of valence electrons on the free neutral atom • Number of valence electrons belonging to the atom in a molecule • Compare the numbers

  36. What assumptions must be made to compute formal charge? • Lone pair electrons belong to the atom in question • Shared electrons are divided equally between the 2 sharing atoms • What 2 assumptions do we use to evaluate Lewis structures? • Atoms in molecules try to achieve formal charges as close to zero as possible • Any negative formal charges are expected to reside on the most electronegative atoms.

  37. Equations: Formal charge = number of valence electrons on free atom – number of valence electrons assigned to the atom in the molecule Valence electrons assigned = number of lone pair electrons + ½ number of shared electrons

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