Chapter 7. Ionic & Covalent Bonds

Transcription

1 Chapter 7 Ionic & Covalent Bonds

2 Ionic Compounds Covalent Compounds

3 7.1 EN difference and bond character >1.7 = ionic = polar covalent <0.4 = nonpolar covalent

4 Ionic >1.7 Electrons not shared at all Atom with higher electronegativity takes electrons from atom with lower electronegativity NaCl K 2 O

5 Polar Covalent Electrons are shared by not equally HCl More electronegative atom pulls electrons closer

6 Nonpolar Covalent <0.4 Electrons equally shared H 2 Directly in the middle of the two atoms O 2

7 Molecule EN of 1 st atom EN of 2 nd atom EN difference Bond Type O 2 HCl NaCl NO 3 NH 4

8 7.2 Ion Formation Ions are formed when atoms gain or lose valence electrons to achieve a stable octet electron configuration

9 Valence Electrons and Chemical Bonds Recall: All elements in the same group have the same number of valence electrons and therefore

10 Valence electrons are involved in the formation of chemical bonds The force that holds two atoms together Attraction between positive and negative ions

11 Recall: Dot structures Show only valence electrons Carbon: 1s 2 2s 2 2p 2 Bromine: [Ar] 4s 2 3d 10 4p 5

12 Recall: octet rule Atoms will gain, lose, or share electrons to acquire the stable electron configuration of a noble gas Metals Nonmetals -

13 Name: Formed by: Group 1: Group 2: Group13: Positive Ion formation

14 Transition metal ions Form cations only Charges may vary in some atoms Fe can lose 2 or 3 electrons Fe 2+ or Fe 3+ Periodic table

15 Name: Formed by: Group 15 Group 16 Group 17 Negative Ion Formation

16 Ionic Bonds & Ionic Compounds Oppositely charged ions attract each other, forming electrically neutral compounds

17 Formation of an Ionic Bond Ionic bond the force of attraction that holds oppositely charged ions together

18 Ionic compounds compounds that contain ionic bonds Cation + anion Metal + nonmetal Also called salts

19 Binary ionic compounds contain two different elements (a metal + a nonmetal) NaCl MgO K 2 S CaI 2

20 Compound formation and charge Ionic compounds are electrically neutral Total positive charge must = total negative charge Net charge of all ionic compounds = 0

21 Formation of Sodium Chloride Na: [Ne]3s 1 + Cl: [Ne]3s 2 3p 5 Dot Structure: Na + Cl [Na] + +[ Cl ] -

22 Na 2 O Na + O 2- Na + Total positive charge = Total negative charge =

23 Al 2 O 3 Al 3+ O 2- Al 3+ O 2- O 2- Total positive charge = Total negative charge =

24 How would an ionic compound form from Na + N each of the following: Li + O Sr + F Group 1 + group 15

25 Formulas for ionic compounds Formula unit = the chemical formula for an ionic compound Simplest ratio of ions involved Mg 6 Cl 12 MgCl 2 Overall charge = 0

26 Monatomic ions one atom ions Ex: Oxidation number - the charge of a monatomic ion Most transition metals have more than one oxidation number

27 What is the oxidation number of the ions FeO in the following compounds? MgCl 2 Cu 3 N Cu 3 N 2 Fe 2 O 3

28 Properties of ionic compounds Physical structure ions are packed into a regular repeating pattern

29 Crystal lattice 3D geometric arrangement of particles in an ionic compound Formed by the strong attractions among positive and negative ions Each positive ion is surrounded by negative ions & each negative ion is surrounded by positive ions

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31 Physical properties ionic bonds are very strong, take a lot of energy to be broken apart High melting point High boiling point Hard, rigid, brittle solids

32 More physical properties Brilliant colors due to transition metals in crystal lattices Electrolytes when dissolved or melted Conducts electricity IONIC SOLIDS DO NOT CONDUCT ELECTRICITY

33 Energy and the Ionic Bond Exothermic reactions Endothermic reactions Formation of ionic compounds always releases energy & therefore is

34 Lattice energy the amount of energy required to separate 1 mol of ions in an ionic compound Greater lattice energy = stronger force of attraction

35 Lattice energy is directly related to size of ions bonded Smaller ions = stronger bond Which is stronger KCl or LiCl?

36 Lattice energy is also related to the charge of the ions Bond formed from attraction of ions with larger charges = stronger Which is stronger MgO or NaF?

37 7.4 What is a covalent bond? Bond in which atoms share electrons Always 2 nonmetals Called a molecule

38 Diatomic molecules H 2 N 2 Exist because they are more stable than individual atoms O 2 F 2 Cl 2 Br 2 I 2

39 Single Covalent Bonds One pair of electrons is shared Cl Cl H - F

40 Sigma bonds (σ) Pair of shared electrons is in an area centered between the two atoms Valence orbitals overlap which concentrates electrons in a bonding orbital between the two atoms s overlaps with s or p or two p overlap

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42 Double Covalent Bond Two pairs of electrons are shared Typically C, N, O, S O = O

43 Triple covalent bond Three pairs of electrons are shared Typically C or N N N

44 The pi Bond Forms when parallel orbitals overlap and share electrons A double bond = 1 sigma + 1 pi bond A triple bond = 1 sigma + 2 pi bonds

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46 If atoms need 1 electron, it will usually form 1 covalent bond. H and Halogens typically only form one bond If atoms need 2 electrons, it will usually form 2 covalent bonds. If atoms needs 3 electrons, it will usually form 3 covalent bonds.

47 Strength of Covalent Bonds Bond length the distance between bonded nuclei Determined by size of bonding atoms & how many electrons are shared As more electrons are shared, bond length decreases Cl 2 = 1.43 x m O 2 = 1.21 x m N 2 = 1.10 x m

48 Shorter bond length = stronger bond Which bond type is weakest? Which bond type is strongest?

49 Bonds and energy Energy released when bonds form Energy absorbed to break bond Bond dissociation energy Smaller bond length = larger bond dissociation energy

50 7.5 Molecular Structures Molecular formula tells the type and number of atoms in a molecule PH 3 Lewis structure uses electron dot structures to show how electrons are arranged in molecules

51 Structural formula uses letter symbols and bonds to show relative positions of atoms (Can be predicted from Lewis structure)

52 Rules for drawing Lewis structures 1. Atoms want 8 electrons in their outer energy levels (octet rule) Exception: 2. Number of dots in the molecule = sum of dots in individual atoms

53 If atoms need 1 electron, it will usually form 1 covalent bond. H and Halogens typically only form one bond If atoms need 2 electrons, it will usually form 2 covalent bonds. If atoms needs 3 electrons, it will usually form 3 covalent bonds.

54 Single Bonds NH 3 N = H = Total = valence electrons valence electron

55 CO 2 Double Bonds

56 Triple bonds N 2 C 2 H 2

57 CH 4 PCl 3 C 2 H 4

58 Lewis Structures for Polyatomic Ions Number of dots in the molecule must take the ions charge into account NH 4 1+ N 5 H 1(4) +1 charge subtract 1 electron

59 PO 4 3- P 5 O 6(4) -3 charge add 3 electrons

60 Resonance structures Resonance occurs when more than one valid Lewis structure can be written for a molecule or ion SO 3 2-

61 O 3 Draw all resonance structures for the following CO 3 2-

62 Exceptions to the octet rule Suboctets stable configurations with fewer than 8 electrons around an atom Boron typically forms a suboctet BH 3

63 Expanded octets central atoms contain more than eight valence electrons Phosphorous and sulfur can form expanded octets PCl 5

64 Coordinate covalent bond forms when one atom donates both of the electrons to be shared CO

65 7.6 Molecular Shapes VSEPR model V alence S hell E lectron P air R epulsion The shape of a molecule is determined by minimizing repulsion between lone pairs

66 Bond angle VSEPR forces cause atoms in a molecule to be positioned at fixed angles relative to one another

67 Linear 180 degrees Binary molecules or molecules with 2 atoms bonded to the central atom and no lone pairs on the central atom CO 2

68 Bent degrees Two atoms bonded to central atom, two lone pairs on central atom H 2 O

69 Trigonal Planar 120 degrees Three atoms bonded to central atom, no lone pairs on central atom BH 3

70 Tetrahedral degrees Pyramid shaped 4 atoms bonded to central atom, no lone pair on central atom CH 4

71 Trigonal Pyramidal degrees A triangle that is not flat Three atoms bonded to central atom, one lone pair on central atom NH 3

72 7.7 Electronegativity & Polarity A chemical bond s character is related to each atoms attraction for the electrons in the bond

73 EN difference and bond character >1.7 = mostly ionic = polar covalent <0.4 = nonpolar covalent

74 Ionic >1.7 Electrons not shared at all Atom with higher electronegativity takes electrons from atom with lower electronegativity NaCl K 2 O

75 Polar Covalent Electrons are shared by not equally HCl More electronegative atom pulls electrons closer

76 Nonpolar Covalent <0.4 Electrons equally shared H 2 Directly in the middle of the two atoms O 2

77 Molecule EN of 1 st atom EN of 2 nd atom EN difference Bond Type O 2 HCl NaCl NO 3 NH 4

78 Polar and Nonpolar Molecules Polar bond = electrons shared unequally (dipole) More electronegative atom pulls electrons closer to it and therefore has a slightly negative charge Less electronegative atom has a slightly positive charge

79 Polar molecules always have: 1. At least one polar covalent bond 2. Asymmetric geometry 1. Lone pairs on the center atom 2. Different atoms bonded to the center atom

80 HCl H 2 O

81 Not every molecule with polar bonds is polar!!! CH 4 CO 2

82 Polar Bond vs. Polar Molecule Polar bond = electrons shared unequally between molecules Polar molecule = entire molecule has different partial charges on opposite sides of the molecule Depends on shape

83 CCl 4 NH 3 H 2 0 Linear vs. bent

84 7.8 Intermolecular Forces Weak attractive forces between different moleucles Weaker than bonds

85 Dispersion forces WEAKEST!!! Caused by attraction that results from temporary closeness of electrons Only last a fraction of a second

86 Dipole Dipole forces Occur when polar molecules are attracted to one another Permanent attraction

87 Hydrogen Bonds STRONGEST!!! Between H of one polar molecule and O, N, or F of another polar molecule Without them we wouldn t exist

88 Forces and properties Weak forces holding molecules together give certain properties Low melting points Low boiling points Exist as gasses, liquids or soft solids at room temperature