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Chapter 4 Compounds and Their Bonds

Chapter 4 Compounds and Their Bonds. 4.1 Valence Electrons 4.2 Octet Rule and Ions . Valence Electrons. The valence electrons are the electrons in the outer shell.

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Chapter 4 Compounds and Their Bonds

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  1. Chapter 4 Compounds and Their Bonds 4.1 Valence Electrons 4.2 Octet Rule and Ions

  2. Valence Electrons • The valence electrons are the electrons in the outer shell. • The electrons in the outer shell have the most contact with other atoms and strongly influence the chemical properties of atoms.

  3. Number of Valence Electrons • For Group A elements, the number of valence electrons is the number of electrons in the s and p subshells of the outer shell. • In the electron configuration for phosphorus, there are 5 valence electrons in the s and p subshells with the highest number. 5 valence electronsP Group 5A 1s22s22p63s23p3

  4. Valence Electrons for Groups

  5. Electron Dot Structure • An electron-dot structure is a convenient way to represent the valence electrons. • For example, the two valence electrons for magnesium are placed as single dots on any two sides of the Mg symbol.

  6. Electron-Dot Structures • Dot structures are used for Group A elements. • The valence electrons are placed on the sides of the symbol of an element.

  7. Octet Rule • The stability of the noble gases is associated with 8 valence electrons (He has 2). Ne 2, 8Ar 2, 8, 8Kr 2, 8, 18, 8 • Atoms can become more stable by acquiring an octet (8 electrons) in the outer shell. • The process of acquiring an octet involves the loss, gain, or sharing of valence electrons. • Main Group Elements!!

  8. Ionization Energy • Ionization energy is the energy it takes to remove a valence electron. • Metals have lower ionization energies and nonmetals have higher ionization energies.

  9. Metals Form Positive Ions • Metals acquire octets by losingvalenceelectrons. • The loss of electrons converts an atom to an ion that has the electron configuration of the nearest noble gas. • Metals form positive ionsbecause they have fewer electrons than protons. Group 1A metals ion 1+Group 2A metals ion 2+Group 3A metals ion 3+

  10. Formation of a Sodium Ion, Na+ • Sodium forms an octet by losing its one valence electron. Na  – e Na + 1s22s22p63s11s22s22p6 (= Ne) • A positive ion forms with a +1 charge. Sodium atom Sodium ion 11 p+ 11 p+ 11 e-10 e- 0 1 +

  11. Formation of Mg2+ Magnesium forms an octet by losing its two valence electrons. Magnesium atom Magnesium ion  Mg  – 2e Mg2+ 1s22s22p63s21s22s22p6 (= Ne) • A positive ion forms with a +2 charge. 12 p+ 12 p+ 12 e- 10 e- 0 2 +

  12. Formation of Negative Ions • Nonmetals gain electrons to achieve an octet arrangement, they form negative ions. • The ionic charge of a nonmetal is 3-, 2-, or 1-.

  13. Formation of a Fluoride Ion, F- • Fluorine forms an octet by adding an electron to its seven valence electrons.   1- : F  + e:F :  1s22s22p51s22s22p6 (= Ne) • A negative ion forms with a -1 charge. Fluorine atom Fluoride ion 9 p+ 9 p+ 9 e-10 e- 0 1 –

  14. Group Number and Ions • The Group number can be used to determine the charge of an ion. • The charge of a positive ion is equal to its Group number. Group 3A = 3+ • The charge of a negative ion is obtained by subtracting its Group number from 8. Group 6A = - (8-6) = 2-

  15. Examples of Ionic Charges

  16. Some Important Ions in the Body

  17. Chapter 4 Compounds and Their Bonds 4.3 Ionic Compounds 4.4 Naming and Writing Ionic Formulas

  18. Ionic Compounds • Ionic compounds consist of positive and negative ions. • An ionic bond is an attraction between the positive and negative charges. • In an ionic formula, the total charge of the positive ions is equal to the total charge of the negative ions. total positive charge = total negative charge

  19. Ionic Formulas • The formulas of ionic compounds are determined from the charges on the ions. atoms ions  – Na +  F : Na+ : F : NaF  sodium fluorine sodium fluoride • The overall charge of NaF is zero (0). (1+ ) + (1-) = 0

  20. Charge Balance In NaCl • The formula does not show the charges of the ions in the compound. • The symbol of the metal is written first, followed by the symbol of the nonmetal.

  21. Charge Balance In MgCl2

  22. Writing a Formula from Charges Write the formula of the ionic compound that forms from Ba2+ and Cl. • Write the symbols of the positive ion and the negative ion. Ba2+ Cl • Balance the charges until the positive charge is equal to the negative charge. Ba2+ CltwoCl-neededCl • Write the formula using subscripts for the number of ions for charge balance. BaCl2

  23. Names of Ions • Positive ions are named like the element. • Negative ions are named by changing the end of the element name to –ide.

  24. Naming Ionic Compounds with Two Elements • The name of a binary ionic compound (two elements) gives the name of the metal ion first and the name of the negative ion second. Examples: NaCl sodium chloride K2S potassium sulfide CaI2 calcium iodide Al2O3 aluminum oxide

  25. Ionic Charges of Transition Metals • Most transition elements have two or more positive ions.

  26. Summary of Common Ions • Of the transition metals, silver and zinc are important elements that form only one ion.

  27. Naming Compounds with Transition Metals • Transition metals with two different ions use a Roman numeral following the name of the metal to indicate ionic charge.

  28. Chapter 4 Compounds and Their Bonds 4.5 Covalent Bonds 4.6 Naming and Writing Formulas of Covalent Compounds 4.7 Bond Polarity

  29. Covalent Bonds • Covalent bonds form between two nonmetals from Groups 4A, 5A, 6A, and 7A. • In a covalent bond, electrons are shared to complete octets.

  30. H2, A Covalent Molecule • In hydrogen, two hydrogen atoms share their electrons to form a covalent bond. • Each hydrogen atom acquires a stable outer shell of two (2) electrons like helium (He). H+H H : H = HH = H2 hydrogen molecule

  31. Diatomic Elements • As elements, the following share electrons to form diatomic, covalent molecules.

  32. Covalent Bonds in NH3 • The compound NH3 consists of a N atom and three H atoms.   N and 3 H  • By sharing electrons to form NH3, the electron dot structure is written as H Bonding pairs   H : N : H  Lone pair of electrons

  33. Number of Covalent Bonds • Often, the number of covalent bonds formed by a nonmetal is equal to the number of electrons needed to complete the octet.

  34. Dot Structures and Models of Some Covalent Compounds

  35. Multiple Bonds • Sharing one pair of electrons is a single bond.X : X or X–X • In multiple bonds, two pairs of electrons are shared to form a double bond or three pairs of electrons are shared in a triple bond.X : : X or X =XX ::: X or X ≡X

  36. Multiple Bonds in N2 • In nitrogen, octets are achieved by sharing three pairs of electrons. • When three pairs of electrons are shared, the multiple bond is called a triple bond. octets       N  + N  N:::N  triple bond

  37. Naming Covalent Compounds • In the name of a covalent compound, the first nonmetal is named followed by the name of the second nonmetal ending in –ide. • Prefixes indicate the number of atoms of each element.

  38. Formulas and Names of Some Covalent Compounds

  39. Electronegativity • Electronegativity is the attraction of an atom for shared electrons. • The nonmetals have high electronegativity values with fluorine as the highest. • The metals have low electronegativity values.

  40. Some Electronegativity Values for Group A Elements

  41. Nonpolar Covalent Bonds • The atoms in a nonpolar covalent bond have electronegativity differences of 0.3 or less. • Examples: Atoms Electronegativity Type of Difference Bond N-N3.0 - 3.0 = 0.0 Nonpolar covalentCl-Br3.0 - 2.8 = 0.2 Nonpolar covalentH-Si 2.1 - 1.8 = 0.3 Nonpolar covalent

  42. Polar Covalent Bonds • The atoms in a polar covalent bond have electronegativity differences of 0.4 to 1.6. • Examples: Atoms Electronegativity Type of Difference BondO-Cl3.5 - 3.0 = 0.5 Polar covalentCl-C3.0 - 2.5 = 0.5 Polar covalentO-S 3.5 - 2.5= 1.0 Polar covalent

  43. Comparing Nonpolar and Polar Covalent Bonds

  44. Ionic Bonds • The atoms in an ionic bond have electronegativity differences of 1.7 or more. • Examples: Atoms Electronegativity Type of Difference BondCl-K3.0 – 0.8 = 2.2 IonicN-Na3.0 – 0.9 = 2.1 IonicS-Cs 2.5 – 0.7 = 1.8 Ionic

  45. Range of Bond Types

  46. Predicting Bond Type

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