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IV.  Lesson Essential Questions: › Why do atoms form chemical bonds? › How is the type of chemical bond determined? Vocabulary: chemical bond, ionic.

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Presentation on theme: "IV.  Lesson Essential Questions: › Why do atoms form chemical bonds? › How is the type of chemical bond determined? Vocabulary: chemical bond, ionic."— Presentation transcript:

1 IV

2  Lesson Essential Questions: › Why do atoms form chemical bonds? › How is the type of chemical bond determined? Vocabulary: chemical bond, ionic bonding, covalent bonding, nonpolar-covalent bonding, polar, polar-covalent bonding

3  Chemical Bond › attractive force between atoms or ions that binds them together as a unit › bonds form in order to…  decrease potential energy (PE)  increase stability

4 ION Polyatomic Ion Monatomic Ion 1 atom 2 or more atoms NO 3 - Na +

5 IONIC COVALENT Bond Formation Type of Structure Solubility in Water Electrical Conductivity Other Properties e - are transferred from metal to nonmetal high yes (solution or liquid) yes e - are shared between two nonmetals low no usually not Melting Point crystal lattice true molecules Physical State solid Solid, liquid, or gas odorous

6 “electron sea” METALLIC Bond Formation Type of Structure Solubility in Water Electrical Conductivity Other Properties Melting Point Physical State e - are delocalized among metal atoms very high yes (any form) no malleable, ductile, lustrous solid

7 Ionic Bonding - Crystal Lattice RETURN

8 Covalent Bonding - True Molecules RETURN Diatomic Molecule

9 Metallic Bonding - “Electron Sea” RETURN

10  Most bonds are a blend of ionic and covalent characteristics.  Difference in electronegativity determines bond type.

11  Electronegativity › Attraction an atom has for a shared pair of electrons. › higher e - neg atom   - › lower e - neg atom   +

12  Electronegativity Trend › Increases up and to the right.

13  Nonpolar Covalent Bond › e - are shared equally › symmetrical e - density › usually identical atoms

14 ++ --  Polar Covalent Bond › e - are shared unequally › asymmetrical e - density › results in partial charges (dipole)

15 zNonpolar zPolar zIonic View Bonding Animations.Bonding Animations

16 Examples:  Cl 2  HCl  NaCl =0.0 Nonpolar =0.9 Polar =2.1 Ionic

17  What type of bonding would be expected between the following atoms? › Li and Cl › Ca and Ga › I and Cl › K and Na

18  Lesson Essential Questions: › How is a molecular compound formed? › What are some of the characteristics of a covalent bond? Vocabulary: molecule, chemical formula, molecular formula, bond energy, electron-dot, Lewis structure, structural formula, single bond, multiple bonds, resonance

19  Covalent bond – bond that is created by the sharing of electrons  Molecule – neutral group of atoms held together by covalent bonds  Molecular compound – chemical compound made of molecules

20 CHEMICAL FORMULA Molecular Formula Unit IONICCOVALENT CO 2 NaCl

21  Potential Energy › based on position of an object › low PE = high stability

22 no interaction attraction vs. repulsion increased attraction  Potential Energy Diagram

23 balanced attraction & repulsion increased repulsion attraction vs. repulsion  Potential Energy Diagram

24 Bond Energy  Bond Energy › Energy required to break a bond Bond Length

25  Bond Energy › Short bond = high bond energy

26  Electron Dot Diagrams 1. Pick the central atom 2. Count the valence electrons (they are what electron dot diagrams show) 3. Place electrons around the atom

27  Octet Rule › Most atoms form bonds in order to obtain 8 valence e - › Full energy level stability ~ Noble Gases Ne

28 ++ -- ++  Nonpolar Covalent - no charges  Polar Covalent - partial charges

29  On page 186 in your text book do practice problems # Draw the Lewis structure of ammonia, NH 3 2. Draw the Lewis structure for hydrogen sulfide, H 2 S 3. Draw the Lewis structure for silane, SiH 4 4. Draw the Lewis structure for phosphorus trifluoride, PF 3

30  Some elements can share more than one electron pair. › Double bond (two pairs of electrons are shared) › Triple bond (three pairs of electrons are shared)

31  Draw Lewis structures for each of the following molecules: › O 2 › CO 2 › N 3 › N 2

32 Occurs when more than one valid Lewis structure can be written for a particular molecule (due to position of double bond) These are resonance structures of benzene. The actual structure is an average (or hybrid) of these structures.

33 yNeither structure is correct, it is actually a hybrid of the two. To show it, draw all varieties possible, and join them with a double-headed arrow. yNote the different location of the double bond

34 yResonance in a carbonate ion (CO 3 2- ): yResonance in an acetate ion (C 2 H 3 O 2 1- ):

35  Prefix System (binary compounds) 1.Less e - neg atom comes first. 2.Add prefixes to indicate # of atoms. Omit mono- prefix on first element. 3.Change the ending of the second element to -ide.

36 PREFIX mono- di- tri- tetra- penta- hexa- hepta- octa- nona- deca- NUMBER

37 zCCl 4 zN 2 O zSF 6 ycarbon tetrachloride ydinitrogen monoxide ysulfur hexafluoride

38 zarsenic trichloride zdinitrogen pentoxide ztetraphosphorus decoxide yAsCl 3 yN2O5yN2O5 yP 4 O 10

39  Lesson Essential Questions: › How is an ionic bond formed? › What are some of the characteristics of an ionic bond? Vocabulary: ionic compound, formula unit, lattice energy, polyatomic ion

40  Ionic compound – composed of positive and negative ions that are combined so that the charges are equal.

41 CHEMICAL FORMULA Molecular Formula Unit IONICCOVALENT CO 2 NaCl

42  Electron dot notation is used to note changes.  Form to create an atmosphere of stability

43  Covalent – show sharing of e -  Ionic – show transfer of e -

44  Ions minimize potential energy in crystals by forming a crystal lattice.  Distance between all ions represent a balance of attraction between oppositely charged particles and repulsion between like charged particles

45  Lattice Energy › Energy released when one mole of an ionic crystalline compound is formed from gaseous ions

46 Ionic  High melting temperature  High boiling point  Hard  Brittle, because slight shift of crystal can cause it to break  Conduct electricity when dissolved in water Covalent  Low melting temperature  Low boiling point  Do not conduct electricity  Not as brittle

47 Ionic Formulas  Write each ion, cation first. Don’t show charges in the final formula.  Overall charge must equal zero. › If charges cancel, just write symbols. › If not, use subscripts to balance charges.  Use parentheses to show more than one polyatomic ion.  Stock System - Roman numerals indicate the ion’s charge.

48 Ionic Names  Write the names of both ions, cation first.  Change ending of monatomic ions to -ide.  Polyatomic ions have special names.  Stock System - Use Roman numerals to show the ion’s charge if more than one is possible. Overall charge must equal zero.

49  Consider the following: › Does it contain a polyatomic ion?  -ide, 2 elements  no  -ate, -ite, 3+ elements  yes › Does it contain a Roman numeral?  Check the table for metals not in Groups 1 or 2. › No prefixes!

50 Common Ion Charges NA

51 zpotassium chloride zmagnesium nitrate zcopper(II) chloride  K + Cl   Mg 2+ NO 3   Cu 2+ Cl   KCl  Mg(NO 3 ) 2  CuCl 2

52 zNaBr zNa 2 CO 3 zFeCl 3 ysodium bromide ysodium carbonate yiron(III) chloride

53  Lesson Essential Questions: › How is a metallic bond formed? › What are some of the characteristics of a metallic bond? Vocabulary: metallic bond, alloy

54  Metal ions held together by attraction to free floating electrons. (Sea of electrons)  Good conductors of electricity – Why?

55  Malleable  Ductile  Bond strength – related to enthalpy of vaporization › The more energy required to vaporize, the stronger the bond. › See table on page 196.

56  A mixture of two or more substances, one of which must be a metal.  Common alloys include steel, 14K gold, 18K gold, cast iron, sterling silver, and bronze.  Within different alloys, there can be different types of mixtures – ex. Steel  Where do we find alloys?

57  Use the 3 circle Venn diagram to compare and contrast ionic, metallic, and covalent bonding.

58  Lesson Essential Questions: › How is the VSEPR Theory useful? › What are the different forces present in bonding? Vocabulary: VSEPR theory, hybridization, dipole, hydrogen bonding, London dispersion forces

59  V alence Shell Electron Pair Repulsion Theory  Electron pairs orient themselves in order to minimize repulsive forces.

60  Types of e - Pairs › Bonding pairs - form bonds › Lone pairs - nonbonding e - yLone pairs repel more strongly than bonding pairs!!!

61  Lone pairs reduce the bond angle between atoms. yBond Angle

62  Draw the Lewis Diagram.  Tally up e - pairs on central atom. › double/triple bonds = ONE pair  Shape is determined by the # of bonding pairs and lone pairs. Know the 8 common shapes & their bond angles!

63

64 2 total 2 bond 0 lone LINEAR 180° BeH 2

65 3 total 3 bond 0 lone TRIGONAL PLANAR 120° BF 3

66 3 total 2 bond 1 lone BENT <120° SO 2

67 4 total 4 bond 0 lone TETRAHEDRAL 109.5° CH 4

68 4 total 3 bond 1 lone TRIGONAL PYRAMIDAL 107° NH 3

69 4 total 2 bond 2 lone BENT 104.5° H2OH2O

70 5 total 5 bond 0 lone TRIGONAL BIPYRAMIDAL 120°/90° PCl 5

71 6 total 6 bond 0 lone OCTAHEDRAL 90° SF 6

72  PF 3 4 total 3 bond 1 lone TRIGONAL PYRAMIDAL 107°

73  CO 2 2 total 2 bond 0 lone LINEAR 180°

74  Identify the molecular geometry for the following molecules: › HI › CBr 4 › CH 2 Cl 2

75  Intermolecular forces = forces between molecules. › The boiling point of a liquid is a good measure of the intermolecular forces between its molecules: the higher the boiling point, the stronger the forces between the molecules.  Types of intermolecular forces › Dipole-dipole forces › Hydrogen bonding › London dispersion forces

76  Dipole – created by equal but opposite charges that are separated by a short distance.  A dipole is represented by an arrow with its head pointing toward the negative pole and a crossed tail at the positive pole. The dipole created by a hydrogen chloride molecule is indicated as follows:  Dipole-dipole forces are the forces of attraction between polar molecules.

77  The negative region in one polar molecule attracts the positive region in adjacent molecules. So the molecules all attract each other from opposite sides.  Dipole-dipole forces act at short range, only between nearby molecules.

78  Hydrogen bonding = intermolecular force in which a hydrogen atom that is bonded to a highly electronegative atom is attracted to an unshared pair of electrons in a nearby molecule.

79  London Dispersion Forces = intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles.

80  Modern Chemistry Textbook    m/chemnotes1.htm m/chemnotes1.htm 


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