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1 Chapter 5 Compounds and Their Bonds 5.1 Octet Rule and Ions.

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1 1 Chapter 5 Compounds and Their Bonds 5.1 Octet Rule and Ions

2 2 An octet  is 8 valence electrons  is associated with the stability of the noble gases  does not occur with He; He is stable with 2 valence electrons (duet) Valence Electrons He 1s 2 2 Ne 1s 2 2s 2 2p 6 8 Ar 1s 2 2s 2 2p 6 3s 2 3p 6 8 Kr 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 8 Octet Rule

3 3 Ionic and Covalent Bonds Atoms form octets  to become more stable  by losing, gaining, or sharing valence electrons  by forming ionic or covalent bonds

4 4 Metals Form Positive Ions Metals form positive ions  by a loss of their valence electrons  with the electron configuration of the nearest noble gas  that have fewer electrons than protons Group 1A(1) metals ion 1+ Group 2A(2) metals ion 2+ Group 3A(3) metalsion 3+

5 5 Formation of a Sodium Ion, Na + Sodium achieves an octet by losing its one valence electron. With the loss of its valence electron, a sodium ion has a 1 + charge.Sodium atom Sodium ion 11p + 11e – 10e – 0 1 +

6 6 Formation of Magnesium Ion, Mg 2+ Magnesium achieves an octet by losing its two valence electrons. With the loss of two valence electrons magnesium forms a positive ion with a 2 + charge Mg atom Mg 2+ ion 12p + 12p + 12e – 10e – 0 2 +

7 7 Formation of Negative Ions In ionic compounds, nonmetals  achieve an octet arrangement  gain electrons  form negatively charged ions with 3 –, 2 –, or 1 – charges

8 8 Formation of a Chloride Ion, Cl – Chlorine achieves an octet by adding an electron to its valence electrons.

9 9 Ionic Charge from Group Numbers  The charge of a positive ion is equal to its Group number. Group 1A(1) = 1 + Group 2A(2) = 2 + Group 3A(3) = 3 +  The charge of a negative ion is obtained by subtracting 8 or 18 from its Group number. Group 6A(16) = 6 – 8 = 2 – or 16 – 18= 2 –

10 Group Number and Ionic Charge Ions  achieve the electron configuration of their nearest noble gas  of metals in Groups 1A(1), 2A(2), or 3A(13) have positive 1 +, 2 +, or 3 + charge.  Of nonmetals in Groups 5A(15), 6A(16), or 7A(17) have negative 3 –, 2 –, or 1 – charge. 10

11 Groups Numbers for Some Positive and Negative Ions 11

12 12 Chapter 5 Compounds and Their Bonds 5.2 Ionic Compounds

13 13 Ionic compounds  consist of positive and negative ions  have attractions called ionic bonds between positively and negatively charged ions  have high melting and boiling points  are solid at room temperature Ionic Compounds

14 14 Salt Is an Ionic Compound Sodium chloride, or “table salt,” is an example of an ionic compound.

15 15 An ionic formula  consists of positively and negatively charged ions  is neutral  has charge balance total positive charge = total negative charge The symbol of the metal is written first, followed by the symbol of the nonmetal. Ionic Formulas

16 16 Charge Balance for NaCl, “Salt” In NaCl,  a Na atom loses its valence electron  a Cl atom gains an electron  the symbol of the metal is written first, followed by the symbol of the nonmetal.

17 17 Charge Balance In MgCl 2 In MgCl 2,  a Mg atom loses two valence electrons  two Cl atoms each gain one electron  subscripts indicate the number of ions needed to give charge balance

18 18 Writing Ionic Formulas from Charges Charge balance is used to write the formula for sodium nitride, a compound containing Na + and N 3−. Na + 3Na + + N 3− = Na 3 N Na + 3(1+) + 1(3–) = 0

19 19 Chapter 5 Compounds and Their Bonds 5.3 Naming and Writing Ionic Formulas

20 20 Naming Ionic Compounds with Two Elements To name a compound with two elements,  identify the cation and anion  name the cation first, followed by the name of the anion

21 21 Formula IonsName Cation Anion NaClNa + Cl – sodium chloride K 2 SK + S 2– potassium sulfide MgOMg 2+ O 2– magnesium oxide CaI 2 Ca 2+ I – calcium iodide Al 2 O 3 Al 3+ S 2– aluminum sulfide Examples of Ionic Compounds with Two Elements

22 22 Transition Metals Form Positive Ions Most transition metals and Group 4(14) metals,  Form 2 or more positive ions  Zn 2+, Ag +, and Cd 2+ form only one ion.

23 Metals with Variable Charge The names of transition metals with two or more positive ions (cations) use a Roman numeral after the name of the metal to identify ionic charge. 23

24 24 Naming Ionic Compounds with Variable Charge Metals

25 25 Naming FeCl 2 STEP 1 Determine the charge of the cation from the anion. Fe ion + 2 Cl – = Fe ion + 2– = 0 Fe ion = 2+ = Fe 2+ STEP 2 Name the cation by the element name, and use a Roman numeral to show its charge. Fe 2+ = iron(II) STEP 3 Write the anion with an ide ending. chloride STEP 4 Name the cation first, then the anion. iron(II) chloride

26 26 Naming Cr 2 O 3 STEP 1 Determine the charge of cation from the anion. 2Cr ions + 3O 2– = 2Cr ions + 3(2–) = 2Cr ions + 6– = 0 2Cr ions = 6+ Cr ion = 3+ = Cr 3+ STEP 2 Name the cation by the element name, and use a Roman numeral to show its charge. Cr 3+ = chromium(III) STEP 3 Write the anion with an ide ending. oxide STEP 4 Name the cation first, then the anion. chromium (III) oxide

27 27 Guide to Writing Formulas from the Name

28 28 Writing Formulas Write a formula for potassium sulfide. STEP 1 Identify the cation and anion. potassium= K + sulfide = S 2− STEP 2 Balance the charges. K + S 2− K + 2(1+) + 1(2–) = 0 STEP 3 Write the cation first. 2K + and 1S 2− = K 2 S 1 = K 2 S

29 29 Writing Formulas Write a formula for iron(III) chloride. STEP 1 Identify the cation and anion. iron (III) = Fe 3+ (III = charge of 3+) chloride = Cl − STEP 2 Balance the charges. Fe 3+ Cl − Cl − 1(3+) + 3(1–) = 0 STEP 3 Write the cation first. 1Fe 3+ and 3Cl − = FeCl 3

30 30 Chapter 5 Compounds and Their Bonds 5.4 Polyatomic Ions

31 31 A polyatomic ion  is a group of atoms  has an overall ionic charge Examples: NH 4 + ammoniumOH − hydroxide NO 3 − nitrateNO 2 − nitrite CO 3 2− carbonatePO 4 3− phosphate HCO 3 − hydrogen carbonate (bicarbonate) Polyatomic Ions

32 32 The names of common polyatomic anions  end in ate NO 3 − nitratePO 4 3− phosphate  with one oxygen less end in ite NO 2 − nitritePO 3 3− phosphite  with hydrogen attached use prefix hydrogen (or bi) HCO 3 − hydrogen carbonate(bicarbonate) HSO 3 − hydrogen sulfite (bisulfite) Some Names of Polyatomic Ions

33 Guide to Naming Compounds with Polyatomic Ions 33

34 34  The positive ion is named first followed by the name of the polyatomic ion. NaNO 3 sodium nitrate K 2 SO 4 potassium sulfate Fe(HCO 3 ) 3 iron(III) bicarbonate or iron(III) hydrogen carbonate (NH 4 ) 3 PO 3 ammonium phosphite Examples of Names of Compounds with Polyatomic Ions

35 35 Writing Formulas with Polyatomic Ions The formula of an ionic compound  containing a polyatomic ion must have a charge balance that equals zero (0) Na + and NO 3 − NaNO 3  with two or more polyatomic ions put the polyatomic ions in parentheses. Mg 2+ and 2NO 3 − Mg(NO 3 ) 2 subscript 2 for charge balance

36 36 Chapter 5 Compounds and Their Bonds 5.5 Covalent Compounds

37 37 Forming Octets in Molecules In a fluorine (F 2 ) molecule, each F atom  shares one electron  acquires an octet

38 38 Diatomic Elements  These elements share electrons to form diatomic, covalent molecules.

39 Guide to Writing Electron-Dot Formulas 39

40 40 Single and Multiple Bonds  In a single bond, one pair of electrons is shared.  In a double bond, two pairs of electrons are shared.  In a triple bond, three pairs of electrons are shared.

41 41 Electron-Dot Formula of CS 2 Write the electron-dot formula for CS 2. STEP 1 Determine the atom arrangement. The C atom is the central atom. S C S STEP 2 Determine the total number of valence electrons for 1C and 2S. 1 C(4e – ) + 2 S(6e – ) = 16e – STEP 3 Attach each S atom to the central C atom using one electron pair. S : C : S 16e – – 4e – = 12e – remaining STEP 4 Attach 12 electrons as 6 lone pairs..... : S : C : S :

42 42 Electron-Dot Formula of CS 2 (continued) To complete octets, form one or more multiple bonds. Convert two lone pairs to bonding pairs between C and S atoms to make two double bonds.

43 43 A Nitrogen Molecule has A Triple Bond In a nitrogen molecule, N 2,  each N atom shares 3 electrons  each N atom attains an octet  the sharing of 3 sets of electrons is a multiple bond called a triple bond

44 44 Resonance structures are  two or more electron-dot formulas for the same arrangement of atoms  related by a double-headed arrow ( )  written by changing the location of a double bond between the central atom and a different attached atom Resonance Structures

45 45 Sulfur dioxide has two resonance structures. STEP 1 Write the arrangement of atoms. O S O STEP 2 Determine the total number of valence electrons. 1 S(6e − ) + 2 O(6e − ) = 18e − STEP 3 Connect bonded atoms by single electron pairs. O : S : O 4e − used 18e − – 4e − = 14e − remaining Writing Resonance Structures

46 46 STEP 4 Add 14 remaining electrons as 7 lone pairs. STEP 5 Form a double bond to complete octets. Two resonance structures are possible. Writing Resonance Structures (continued)

47 47 Write the ionic formula of the compound containing Ba 2+ and Cl .  Write the symbols of the ions. Ba 2+ Cl   Balance the charges. Ba 2+ Cl  two Cl  needed Cl   Write the ionic formula using a subscript 2 for two chloride ions that give charge balance. BaCl 2 Formula from Ionic Charges

48 48 Chapter 5 Compounds and Their Bonds 5.6 Naming and Writing Covalent Formulas NO nitrogen oxide NO 2 nitrogen dioxide N 2 O 4 dinitrogen tetroxide

49 49 Names of Covalent Compounds Prefixes are used  in the names of covalent compounds  because two nonmetals can form two or more different compounds Examples of compounds of N and O: NO nitrogen oxide NO 2 nitrogen dioxide N 2 O dinitrogen oxide N 2 O 4 dinitrogen tetroxide N 2 O 5 dinitrogen pentoxide

50 50 Naming Covalent Compounds STEP 1 Name the first nonmetal by its element name. STEP 2 Name the second nonmetal with an ide ending. STEP 3 prefixes to indicate the number (from subscripts) of atoms of each nonmetal. Mono is usually omitted.

51 51 What is the name of SO 3 ? STEP 1 The first nonmetal is S sulfur. STEP 2 The second nonmetal is O, named oxide. STEP 3 The subscript 3 of O is shown as the prefix tri. SO 3 → sulfur trioxide The subscript 1(for S) or mono is understood. Naming Covalent Compounds (Continued)

52 52 Name P 4 S 3 STEP 1 The first nonmetal P is phosphorus. STEP 2 The second nonmetal S is sulfide. STEP 3 The subscript 4 of P is shown as tetra. The subscript 3 of O is shown as tri. P 4 S 3 → tetraphosphorus trisulfide Naming Covalent Compounds (Continued)

53 Guide to Writing Formulas for Covalent Compounds 53

54 54 Write the formula for carbon disulfide. STEP 1 Elements are C and S STEP 2 No prefix for carbon means 1 C Prefix di = 2 Formula: CS 2 Writing Formulas of Covalent Compounds

55 55 Chapter 5 Compounds and Their Bonds 5.7 Electronegativity and Bond Polarity

56 56 The electronegativity value  indicates the attraction of an atom for shared electrons  increases from left to right going across a period on the periodic table  decreases going down a group on the periodic table  is high for the nonmetals, with fluorine as the highest  is low for the metals Electronegativity

57 57 Some Electronegativity Values Low values High values

58 58 A nonpolar covalent bond  occurs between nonmetals  has an equal or almost equal sharing of electrons  has almost no electronegativity difference (0.0 to 0.4) Examples: Atoms Electronegativity Type of Bond Difference_______________ N–N 3.0 – 3.0 = 0.0 Nonpolar covalent Cl–Br 3.0 – 2.8 = 0.2 Nonpolar covalent H–Si 2.1 – 1.8 = 0.3 Nonpolar covalent Nonpolar Covalent Bonds

59 59 A polar covalent bond  occurs between nonmetal atoms  has an unequal sharing of electrons  has a moderate electronegativity difference (0.5 to 1.7) Examples: Atoms Electronegativity Type of Bond Difference _________________________________________ O–Cl 3.5 – 3.0 = 0.5 Polar covalent Cl–C 3.0 – 2.5 = 0.5 Polar covalent O–S 3.5 – 2.5 = 1.0 Polar covalent Polar Covalent Bonds

60 60 Comparing Nonpolar and Polar Covalent Bonds

61 61 Chapter 5 Compounds and Their Bonds 5.8 Shapes and Polarity of Molecules

62 62 VSEPR In the valence-shell electron-pair repulsion theory (VSEPR), the electron groups around a central atom  are arranged as far apart from each other as possible  have the least amount of repulsion of the negatively charged electrons  have a geometry around the central atom that determines molecular shape

63 63 Shapes of Molecules The three-dimensional shape of a molecule  is the result of bonded groups and lone pairs of electrons around the central atom  is predicted using the VSEPR theory (valence-shell- electron-pair repulsion)

64 64 Guide to Predicting Molecular Shape (VSEPR Theory)

65 65 Two Electron Groups In a molecule of BeCl 2,  there are two electron groups bonded to the central atom, Be (Be is an exception to the octet rule).... : Cl : Be : Cl :....  to minimize repulsion, the arrangement of two electron groups is 180°, or opposite each other  the shape of the molecule is linear

66 66 Two Electron Groups with Double Bonds In a molecule of CO 2,  there are two electron groups bonded to C (electrons in each double bond are counted as one group)  repulsion is minimized with the double bonds opposite each other at 180°  the shape of the molecule is linear

67 67 Three Electron Groups In a molecule of BF 3,  three electron groups are bonded to the central atom B (B is an exception to the octet rule).. : F: : F : B : F :....  repulsion is minimized with 3 electron groups at angles of 120°  the shape is trigonal planar

68 68 Two Electron Groups and One Lone Pair In a molecule of SO 2,  S has 3 electron groups; 2 electron groups bonded to O atoms and one lone pair :O :: S : O :..  repulsion is minimized with the electron groups at angles of 120°, a trigonal planar arrangement  the shape is bent (120°), with two O atoms bonded to S ●

69 69 Four Electron Groups In a molecule of CH 4,  there are four electron groups around C  repulsion is minimized by placing four electron groups at angles of 109°, which is a tetrahedral arrangement  the four bonded atoms form a tetrahedral shape

70 70 Three Bonding Atoms and One Lone Pair In a molecule of NH 3,  three electron groups bond to H atoms, and the fourth one is a lone (nonbonding) pair  repulsion is minimized with 4 electron groups in a tetrahedral arrangement  the three bonded atoms form a pyramidal (~109°) shape

71 71 Two Bonding Atoms and Two Lone Pairs In a molecule of H 2 O,  two electron groups are bonded to H atoms and two are lone pairs (4 electron groups)  four electron groups minimize repulsion in a tetrahedral arrangement  the shape with two bonded atoms is bent (~109°)

72 Determining Molecular Polarity Determine the polarity of the H 2 O molecule. Solution: The four electron groups of oxygen are bonded to two H atoms. Thus the H 2 O molecule has a net dipole, which makes it a polar molecule. 72

73 73 Chapter 5 Compounds and Their Bonds 5.9 Attractive Forces in Compounds

74 74 Dipole–Dipole Attractions In covalent compounds, polar molecules  exert attractive forces called dipole-dipole attractions  form strong dipole attractions called hydrogen bonds between hydrogen atoms bonded to F, O, or N, and other atoms that are very electronegative

75 75 Dipole–Dipole Attractions (continued)

76 76 Dispersion Forces Dispersion forces are  weak attractions between nonpolar molecules  caused by temporary dipoles that develop when electrons are not distributed equally

77 77 Comparison of Bonding and Attractive Forces

78 78 Melting Points and Attractive Forces  Ionic compounds require large amounts of energy to break apart ionic bonds. Thus, they have high melting points.  Hydrogen bonds are the strongest type of dipole– dipole attractions. They require more energy to break than do other dipole attractions.  Dispersion forces are weak interactions, and very little energy is needed to change state.


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