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Ionic and Covalent Bonding Electron and Lewis Dot Structures James Hutchison.

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Presentation on theme: "Ionic and Covalent Bonding Electron and Lewis Dot Structures James Hutchison."— Presentation transcript:

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2 Ionic and Covalent Bonding Electron and Lewis Dot Structures James Hutchison

3 Valence Electrons are…? The electrons responsible for the chemical properties of atoms, and are those in the outer energy level. Valence electrons - The s and p electrons in the outer energy level – the highest occupied energy level Core electrons – are those in the energy levels below.

4 Keeping Track of Electrons Atoms in the same column... 1)Have the same outer electron configuration. 2)Have the same valence electrons. The number of valence electrons are easily determined. It is the group number for a representative element Group 2A: Be, Mg, Ca, etc. – have 2 valence electrons

5 Electron Dot diagrams are… A way of showing & keeping track of valence electrons. How to write them? Write the symbol - it represents the nucleus and inner (core) electrons Put one dot for each valence electron (8 maximum) They don’t pair up until they have to (Hund’s rule) X

6 The Electron Dot diagram for Nitrogen l Nitrogen has 5 valence electrons to show. l First we write the symbol. N l Then add 1 electron at a time to each side. l Now they are forced to pair up. l We have now written the electron dot diagram for Nitrogen.

7 The Octet Rule l In Chapter 6, we learned that noble gases are unreactive in chemical reactions l In 1916, Gilbert Lewis used this fact to explain why atoms form certain kinds of ions and molecules l The Octet Rule: in forming compounds, atoms tend to achieve a noble gas configuration; 8 in the outer level is stable l Each noble gas (except He, which has 2) has 8 electrons in the outer level

8 Formation of Cations Metals lose electrons to attain a noble gas configuration. They make positive ions (cations) If we look at the electron configuration, it makes sense to lose electrons: Na 1s 2 2s 2 2p 6 3s 1 1 valence electron Na 1+ 1s 2 2s 2 2p 6 This is a noble gas configuration with 8 electrons in the outer level.

9 Electron Dots For Cations Metals will have few valence electrons (usually 3 or less); calcium has only 2 valence electrons Ca

10 Electron Dots For Cations Metals will have few valence electrons Metals will lose the valence electrons Ca

11 Electron Dots For Cations Metals will have few valence electrons Metals will lose the valence electrons Forming positive ions Ca 2+ NO DOTS are now shown for the cation. This is named the “calcium ion”.

12 Electron Dots For Cations Let’s do Scandium, #21 The electron configuration is: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 1 Thus, it can lose 2e - (making it 2+), or lose 3e - (making 3+) Sc = Sc 2+ Scandium (II) ion Scandium (III) ion Sc = Sc 3+

13 Electron Dots For Cations Let’s do Silver, element #47 Predicted configuration is: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 9 Actual configuration is: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 1 4d 10 Ag = Ag 1+ (can’t lose any more, charges of 3+ or greater are uncommon)

14 Electron Dots For Cations Silver did the best job it could, but it did not achieve a true Noble Gas configuration Instead, it is called a “pseudo- noble gas configuration”

15 Electron Configurations: Anions Nonmetals gain electrons to attain noble gas configuration. They make negative ions (anions) S = 1s 2 2s 2 2p 6 3s 2 3p 4 = 6 valence electrons S 2- = 1s 2 2s 2 2p 6 3s 2 3p 6 = noble gas configuration. Halide ions are ions from chlorine or other halogens that gain electrons

16 Electron Dots For Anions Nonmetals will have many valence electrons (usually 5 or more) They will gain electrons to fill outer shell. P 3- (This is called the “phosphide ion”, and should show dots)

17 Stable Electron Configurations All atoms react to try and achieve a noble gas configuration. Noble gases have 2 s and 6 p electrons. 8 valence electrons = already stable! This is the octet rule (8 in the outer level is particularly stable). Ar

18 Ionic Bonding Anions and cations are held together by opposite charges (+ and -) Ionic compounds are called salts. Simplest ratio of elements in an ionic compound is called the formula unit. The bond is formed through the transfer of electrons (lose and gain) Electrons are transferred to achieve noble gas configuration.

19 Ionic Compounds 1)Also called SALTS 2)Made from: a CATION with an ANION (or literally from a metal combining with a nonmetal)

20 Ionic Bonding NaCl The metal (sodium) tends to lose its one electron from the outer level. The nonmetal (chlorine) needs to gain one more to fill its outer level, and will accept the one electron that sodium is going to lose.

21 Ionic Bonding Na + Cl - Note: Remember that NO DOTS are now shown for the cation!

22 Ionic Bonding All the electrons must be accounted for, and each atom will have a noble gas configuration (which is stable). CaP Lets do an example by combining calcium and phosphorus:

23 Ionic Bonding CaP

24 Ionic Bonding Ca 2+ P

25 Ionic Bonding Ca 2+ P Ca

26 Ionic Bonding Ca 2+ P 3- Ca

27 Ionic Bonding Ca 2+ P 3- Ca P

28 Ionic Bonding Ca 2+ P 3- Ca 2+ P

29 Ionic Bonding Ca 2+ P 3- Ca 2+ P Ca

30 Ionic Bonding Ca 2+ P 3- Ca 2+ P Ca

31 Ionic Bonding Ca 2+ P 3- Ca 2+ P 3- Ca 2+

32 Ionic Bonding = Ca 3 P 2 Formula Unit This is a chemical formula, which shows the kinds and numbers of atoms in the smallest representative particle of the substance. For an ionic compound, the smallest representative particle is called a: Formula Unit

33 Properties of Ionic Compounds 1.Crystalline solids - a regular repeating arrangement of ions in the solid: Fig. 7.9, page 197 – Ions are strongly bonded together. – Structure is rigid. 2.High melting points Coordination number- number of ions of opposite charge surrounding it

34 - Page 198 Coordination Numbers: Both the sodium and chlorine have 6 Both the cesium and chlorine have 8 Each titanium has 6, and each oxygen has 3 NaCl CsCl TiO 2

35 Do they Conduct? Conducting electricity means allowing charges to move. In a solid, the ions are locked in place. Ionic solids are insulators. When melted, the ions can move around. 3.Melted ionic compounds conduct. – NaCl: must get to about 800 ºC. – Dissolved in water, they also conduct (free to move in aqueous solutions)

36 - Page 198 The ions are free to move when they are molten (or in aqueous solution), and thus they are able to conduct the electric current.

37 Metallic Bonds are… How metal atoms are held together in the solid. Metals hold on to their valence electrons very weakly. Think of them as positive ions (cations) floating in a sea of electrons: Fig. 7.12, p.201

38 Sea of Electrons Electrons are free to move through the solid. Metals conduct electricity.

39 Metals are Malleable Hammered into shape (bend). Also ductile - drawn into wires. Both malleability and ductility explained in terms of the mobility of the valence electrons

40 - Page 201 1) Ductility2) Malleability Due to the mobility of the valence electrons, metals have: and Notice that the ionic crystal breaks due to ion repulsion!

41 Malleable Force

42 Malleable Mobile electrons allow atoms to slide by, sort of like ball bearings in oil. Force

43 Ionic solids are brittle Force

44 Ionic solids are brittle Strong Repulsion breaks a crystal apart, due to similar ions being next to each other. Force

45 Covalent Bonds The word covalent is a combination of the prefix co- (from Latin com, meaning “with” or “together”), and the verb valere, meaning “to be strong”. Two electrons shared together have the strength to hold two atoms together in a bond.

46 Molecules  Many elements found in nature are in the form of molecules:  a neutral group of atoms joined together by covalent bonds.  For example, air contains oxygen molecules, consisting of two oxygen atoms joined covalently  Called a “diatomic molecule” (O 2 )

47 How does H 2 form? The nuclei repel each other, since they both have a positive charge (like charges repel). ++ (diatomic hydrogen molecule) + +

48 How does H 2 form? ++ But, the nuclei are attracted to the electrons They share the electrons, and this is called a “covalent bond”, and involves only NONMETALS!

49 Covalent bonds Nonmetals hold on to their valence electrons. They can’t give away electrons to bond. – But still want noble gas configuration. Get it by sharing valence electrons with each other = covalent bonding By sharing, both atoms get to count the electrons toward a noble gas configuration.

50 Covalent bonding Fluorine has seven valence electrons (but would like to have 8) F

51 Covalent bonding Fluorine has seven valence electrons A second atom also has seven FF

52 Covalent bonding l Fluorine has seven valence electrons l A second atom also has seven l By sharing electrons… FF

53 Covalent bonding l Fluorine has seven valence electrons l A second atom also has seven l By sharing electrons… FF

54 Covalent bonding l Fluorine has seven valence electrons l A second atom also has seven l By sharing electrons… FF

55 Covalent bonding l Fluorine has seven valence electrons l A second atom also has seven l By sharing electrons… FF

56 Covalent bonding l Fluorine has seven valence electrons l A second atom also has seven l By sharing electrons… FF

57 Covalent bonding l Fluorine has seven valence electrons l A second atom also has seven l By sharing electrons… l …both end with full orbitals FF

58 Covalent bonding l Fluorine has seven valence electrons l A second atom also has seven l By sharing electrons… l …both end with full orbitals FF 8 Valence electrons

59 Covalent bonding l Fluorine has seven valence electrons l A second atom also has seven l By sharing electrons… l …both end with full orbitals FF 8 Valence electrons

60 Molecular Compounds Compounds that are bonded covalently (like in water, or carbon dioxide) are called molecular compounds Molecular compounds tend to have relatively lower melting and boiling points than ionic compounds – this is not as strong a bond as ionic

61 Molecular Compounds Thus, molecular compounds tend to be gases or liquids at room temperature – Ionic compounds were solids A molecular compound has a molecular formula: – Shows how many atoms of each element a molecule contains

62 Molecular Compounds The formula for water is written as H 2 O – The subscript “2” behind hydrogen means there are 2 atoms of hydrogen; if there is only one atom, the subscript 1 is omitted Molecular formulas do not tell any information about the structure (the arrangement of the various atoms).

63 A Single Covalent Bond is... A sharing of two valence electrons. Only nonmetals and hydrogen. Different from an ionic bond because they actually form molecules. Two specific atoms are joined. In an ionic solid, you can’t tell which atom the electrons moved from or to

64 Water H O Each hydrogen has 1 valence electron - Each hydrogen wants 1 more The oxygen has 6 valence electrons - The oxygen wants 2 more They share to make each other complete

65 Water Put the pieces together The first hydrogen is happy The oxygen still needs one more H O

66 Water So, a second hydrogen attaches Every atom has full energy levels H O H Note the two “unshared” pairs of electrons

67 Multiple Bonds Sometimes atoms share more than one pair of valence electrons. A double bond is when atoms share two pairs of electrons (4 total) A triple bond is when atoms share three pairs of electrons (6 total) Table 8.1, p Know these 7 elements as diatomic: Br 2 I 2 N 2 Cl 2 H 2 O 2 F 2 What’s the deal with the oxygen dot diagram?

68 Dot diagram for Carbon dioxide CO 2 - Carbon is central atom ( more metallic ) Carbon has 4 valence electrons Wants 4 more Oxygen has 6 valence electrons Wants 2 more O C

69 Carbon dioxide Attaching 1 oxygen leaves the oxygen 1 short, and the carbon 3 short O C

70 Carbon dioxide l Attaching the second oxygen leaves both of the oxygen 1 short, and the carbon 2 short O C O

71 Carbon dioxide l The only solution is to share more O C O

72 Carbon dioxide l The only solution is to share more O C O

73 Carbon dioxide l The only solution is to share more O CO

74 Carbon dioxide l The only solution is to share more O CO

75 Carbon dioxide l The only solution is to share more O CO

76 Carbon dioxide l The only solution is to share more O CO

77 Carbon dioxide l The only solution is to share more l Requires two double bonds l Each atom can count all the electrons in the bond O CO

78 Carbon dioxide l The only solution is to share more l Requires two double bonds l Each atom can count all the electrons in the bond O CO 8 valence electrons

79 Carbon dioxide l The only solution is to share more l Requires two double bonds l Each atom can count all the electrons in the bond O CO 8 valence electrons

80 Carbon dioxide l The only solution is to share more l Requires two double bonds l Each atom can count all the electrons in the bond O CO 8 valence electrons

81 How to draw them? 1)Add up all the valence electrons. 2)Count up the total number of electrons to make all atoms happy. 3)Subtract; then Divide by 2 4)Tells you how many bonds to draw 5)Fill in the rest of the valence electrons to fill atoms up.

82 Example NH 3, which is ammonia N – central atom; has 5 valence electrons, wants 8 H - has 1 (x3) valence electrons, wants 2 (x3) NH 3 has 5+3 = 8 NH 3 wants 8+6 = 14 (14-8)/2= 3 bonds 4 atoms with 3 bonds N H

83 NHH H Examples Draw in the bonds; start with singles All 8 electrons are accounted for Everything is full – done with this one.

84 Example: HCN HCN: C is central atom N - has 5 valence electrons, wants 8 C - has 4 valence electrons, wants 8 H - has 1 valence electron, wants 2 HCN has = 10 HCN wants = 18 (18-10)/2= 4 bonds 3 atoms with 4 bonds – this will require multiple bonds - not to H however

85 HCN Put single bond between each atom Need to add 2 more bonds Must go between C and N (Hydrogen is full) NHC

86 HCN l Put in single bonds l Needs 2 more bonds l Must go between C and N, not the H l Uses 8 electrons – need 2 more to equal the 10 it has NHC

87 HCN l Put in single bonds l Need 2 more bonds l Must go between C and N l Uses 8 electrons - 2 more to add l Must go on the N to fill its octet NHC

88 Another way of indicating bonds Often use a line to indicate a bond Called a structural formula Each line is 2 valence electrons HHO = HHO

89 Other Structural Examples H CN C O H H

90 A Coordinate Covalent Bond... When one atom donates both electrons in a covalent bond. Carbon monoxide (CO) is a good example: OC Both the carbon and oxygen give another single electron to share

91 Coordinate Covalent Bond l When one atom donates both electrons in a covalent bond. l Carbon monoxide (CO) is a good example: OC Oxygen gives both of these electrons, since it has no more singles to share. This carbon electron moves to make a pair with the other single.

92 Coordinate Covalent Bond l When one atom donates both electrons in a covalent bond. l Carbon monoxide (CO) OC C O The coordinate covalent bond is shown with an arrow as:

93 Molecular Orbitals are... The model for covalent bonding assumes the orbitals are those of the individual atoms = atomic orbital Orbitals that apply to the overall molecule, due to atomic orbital overlap are the molecular orbitals – A bonding orbital is a molecular orbital that can be occupied by two electrons of a covalent bond

94 VSEPR: stands for... V alence S hell E lectron P air R epulsion Predicts the three dimensional shape of molecules. The name tells you the theory: – Valence shell = outside electrons. – Electron Pair repulsion = electron pairs try to get as far away as possible from each other. Can determine the angles of bonds.

95 VSEPR Based on the number of pairs of valence electrons, both bonded and unbonded. Unbonded pair also called lone pair. CH 4 - draw the structural formula Has 4 + 4(1) = 8 wants 8 + 4(2) = 16 (16-8)/2 = 4 bonds

96 VSEPR for methane (a gas): Single bonds fill all atoms. There are 4 pairs of electrons pushing away. The furthest they can get away is 109.5º CHH H H This 2-dimensional drawing does not show a true representation of the chemical arrangement.

97 4 atoms bonded Basic shape is tetrahedral. A pyramid with a triangular base. Same shape for everything with 4 pairs. C HH H H 109.5º

98 - Page 232 Methane has an angle of o, called tetrahedral Ammonia has an angle of 107 o, called pyramidal Note the unshared pair that is repulsion for other electrons.

99 Bond Polarity Covalent bonding means shared electrons – but, do they share equally? Electrons are pulled, as in a tug-of-war, between the atoms nuclei – In equal sharing (such as diatomic molecules), the bond that results is called a nonpolar covalent bond

100 Bond Polarity When two different atoms bond covalently, there is an unequal sharing – the more electronegative atom will have a stronger attraction, and will acquire a slightly negative charge – called a polar covalent bond, or simply polar bond.

101 Electronegativity? The ability of an atom in a molecule to attract shared electrons to itself. The ability of an atom in a molecule to attract shared electrons to itself. Linus Pauling

102 Table of Electronegativities

103 Bond Polarity Refer to Handout Consider HCl H = electronegativity of 2.1 Cl = electronegativity of 3.0 – the bond is polar – the chlorine acquires a slight negative charge, and the hydrogen a slight positive charge

104 Bond Polarity Only partial charges, much less than a true 1+ or 1- as in ionic bond Written as: HCl the positive and minus signs (with the lower case delta: ) denote partial charges.    and  

105 Bond Polarity Can also be shown: – the arrow points to the more electronegative atom. Table 8.3, p.238 shows how the electronegativity can also indicate the type of bond that tends to form HCl

106 Polar molecules Sample Problem 8.3, p.239 A polar bond tends to make the entire molecule “polar” – areas of “difference” HCl has polar bonds, thus is a polar molecule. – A molecule that has two poles is called dipole, like HCl

107 Polar molecules The effect of polar bonds on the polarity of the entire molecule depends on the molecule shape – carbon dioxide has two polar bonds, and is linear = nonpolar molecule!

108 Polar molecules The effect of polar bonds on the polarity of the entire molecule depends on the molecule shape – water has two polar bonds and a bent shape; the highly electronegative oxygen pulls the e - away from H = very polar!

109 Polar molecules When polar molecules are placed between oppositely charged plates, they tend to become oriented with respect to the positive and negative plates.

110 Attractions between molecules They are what make solid and liquid molecular compounds possible. The weakest are called van der Waal’s forces - there are two kinds: #1. Dispersion forces weakest of all, caused by motion of e - increases as # e - increases halogens start as gases; bromine is liquid; iodine is solid – all in Group 7A

111 #2. Dipole interactions Occurs when polar molecules are attracted to each other. 2. Dipole interaction happens in water – Figure 8.25, page 240 – positive region of one molecule attracts the negative region of another molecule.

112 #2. Dipole interactions Occur when polar molecules are attracted to each other. Slightly stronger than dispersion forces. Opposites attract, but not completely hooked like in ionic solids. HFHF  HFHF 

113 #2. Dipole Interactions    

114 #3. Hydrogen bonding …is the attractive force caused by hydrogen bonded to N, O, F, or Cl N, O, F, and Cl are very electronegative, so this is a very strong dipole. And, the hydrogen shares with the lone pair in the molecule next to it. This is the strongest of the intermolecular forces.

115 Order of Intermolecular attraction strengths 1)Dispersion forces are the weakest 2)A little stronger are the dipole interactions 3)The strongest is the hydrogen bonding 4)All of these are weaker than ionic bonds

116 #3. Hydrogen bonding defined: When a hydrogen atom is: a) covalently bonded to a highly electronegative atom, AND b) is also weakly bonded to an unshared electron pair of a nearby highly electronegative atom. – The hydrogen is left very electron deficient (it only had 1 to start with!) thus it shares with something nearby – Hydrogen is also the ONLY element with no shielding for its nucleus when involved in a covalent bond!

117 Hydrogen Bonding (Shown in water) H H O ++ -- ++ H H O ++ -- ++ This hydrogen is bonded covalently to: 1) the highly negative oxygen, and 2) a nearby unshared pair.

118 Hydrogen bonding allows H 2 O to be a liquid at room conditions. H H O H H O H H O H H O H H O H H O H H O

119 Attractions and properties Why are some chemicals gases, some liquids, some solids? – Depends on the type of bonding! Network solids – solids in which all the atoms are covalently bonded to each other

120 Attractions and properties Network solids melt at very high temperatures, or not at all (decomposes) – Diamond does not really melt, but vaporizes to a gas at 3500 o C and beyond – SiC, used in grinding, has a melting point of about 2700 o C

121 Covalent Network Compounds Some covalently bonded substances DO NOT form discrete molecules. Diamond, a network of covalently bonded carbon atoms Graphite, a network of covalently bonded carbon atoms


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