Periodic Properties of the Elements

Presentación del tema: "Periodic Properties of the Elements"— Transcripción de la presentación:

1 Periodic Properties of the Elements
Metallic character corresponds to conductance of heat and electricity.

2

3 Contenidos 2.- Enlace iónico
1.- ¿Por qué se unen los átomos?. Tipos de enlace. 2.- Enlace iónico 3.- Propiedades de los compuestos iónicos. 4.- El enlace covalente. 4.1. Teoría de Lewis. 4.2. Resonancia. 4.3. Repulsion de pares electronicos.RPECV . Geometría. 4.4.  Polaridad en los enlaces y moléculas. Momento dipolar.

4 Metals Tend to Lose Electrons

5 Nonmetals Tend to Gain Electrons

6 Properties of Metal, Nonmetals, and Metalloids

7 Atoms try to attain noble-gas electronic configurations Li  Li+ Ca  Ca2+ Al  Al3+

8 C  C4- P  P3- O  O2- Br  Br-

9 Chemical Bonding I: Basic Concepts
Copyright © The McGraw-Hill Companies, Inc.  Permission required for reproduction or display.

10 ¿Por qué se unen los átomos?
Los átomos, moléculas e iones y se unen entre sí porque al hacerlo se llega a una situación de mínima energía, lo que equivale a decir de máxima estabilidad. Son los electrones más externos, los también llamados electrones de valencia los responsables de esta unión, al igual que de la estequiometría y geometría de las sustancias químicas.

11 Enlace iónico Se da entre un metal que pierde uno o varios electrones y un no metal que los captura Resultan iones positivos y negativos que se mantienen unidos por atracciones electrostáticas, formando redes cristalinas. Las reacciones de pérdida o ganancia de e– se llaman reacciones de ionización: Ejemplo: Na – 1 e–  Na+ O + 2e–  O2– Reac. global: O + 2 Na  O2– + 2 Na+ Formula del compuesto (empírica): Na2O

12 Propiedades de los compuestos iónicos
Puntos de fusión y ebullición elevados ya que para fundirlos es necesario romper la red cristalina tan estable por la cantidad de uniones atracciones electrostáticas entre iones de distinto signo. Son sólidos a temperatura ambiente. Gran dureza.(por la misma razón). Solubilidad en disolventes polares e insolubilidad en disolventes apolares. Conductividad en estado disuelto o fundido. Sin embargo, en estado sólido no conducen la electricidad. Son frágiles.

13 Enlace covalente Dos átomos unidos mediante enlace covalente tienen menos energía que los dos átomos aislados. Al igual que en el enlace iónico la formación de un enlace covalente va acompañada de un desprendimiento de energía. Se llama energía de enlace a la energía necesaria para romper 1 mol de un determinado tipo de enlace. Es siempre endotérmica (positiva). Ejemplo: para romper 1 mol de H2 (g) en 2 moles de H (g) se precisan 436 kJ,  Eenlace(H–H) = kJ La distancia a la que se consigue mayor estabilidad se llama “distancia de enlace”.

14 Enlace Covalente Covalent Bonding
Covalent bonds are formed when atoms share electrons. If the atoms share 2 electrons a single covalent bond is formed. If the atoms share 4 electrons a double covalent bond is formed. If the atoms share 6 electrons a triple covalent bond is formed.

15 This figure shows the potential energy of an H2 molecule as a function of the distance between the two H atoms. Where is H2 most stable?

16 Representation of the formation of an H2 molecule from
two H atoms.

17 Covalent Bonding

18 Formation of Covalent Bonds
H molecule formation representation. HCl molecule formation

19 Teoría de Lewis Se basa en las siguientes hipótesis:
Los átomos para conseguir 8 e– en su última capa comparten tantos electrones como le falten para completar su capa (regla del octete). Cada pareja de e– compartidos forma un enlace. Se pueden formar enlaces sencillos, dobles y triples con el mismo átomo.

20 Estructuras de Lewis Dibuje la estructura del compuesto mostrando qué átomos están conectados con otros. Coloque el elemento menos electronegativo al centro. Calcule el número total de electrones. Agregue 1 por cada carga negativa y elimine 1 por cada carga positiva. Complete los octetos de electrones para todos los elementos, excepto para el hidrógeno. Si la estructura tiene demasiados electrones, forme enlaces dobles o triples en el átomo central. 9.6

21 Lewis Formulas Hydrogen molecule, H2. Fluorine, F2 Nitrogen, N2

22 NH4+

23 Writing Lewis Formulas: The Octet Rule
The octet rule states that representative elements usually attain stable noble gas electron configurations in most of their compounds. Arrangement of bonding (or shared) electrons vs. nonbonding (or unshared or lone pairs) of electrons

24 Electronegativity

25 Percent Ionic Character

26 Teoría de Lewis Se basa en las siguientes hipótesis:
Los átomos para conseguir 8 e– en su última capa comparten tantos electrones como le falten para completar su capa (regla del octete). Cada pareja de e– compartidos forma un enlace. Se pueden formar enlaces sencillos, dobles y triples con el mismo átomo.

27 Enlace Covalente Símbolos de Lewis
Un símbolo de Lewis representa el núcleo y los electrones internos de un átomo. Los puntos alrededor del símbolo representan a los electrones de de valencia.

28 Enlace Covalente Teoría de Lewis
Los electrones de la capa más externa (de valencia) se transfieren (iónico) o se comparten (covalente) de modo que los átomos adquieren una configuración electrónica estable. De gas noble. Octeto. En esta caso los electrones de valencia de cada átomo se representan por medio de puntos, cruces o círculos. Cada par de electrones compartidos pueden representarse con una línea y si hay dobles o triples enlaces se representan con dos o tres líneas.

29 Enlace Covalente Las estructuras de Lewis no explican
La forma o la geometría de una molécula. La información de los orbitales donde proceden los electrones o de donde se alojan definitivamente estos. Basta con contar los electrones de valencia y distribuirlos correctamente alrededor del átomo. Por ejemplo no explica la diferencia para estos compuestos de azufre

30 Enlace Covalente Reglas para las estructuras de Lewis
El H sólo puede adquirir 2e. Los elementos del 2º período: 8e y lo del 3ª y siguiente pueden ampliar el octeto. Escribir una fórmula con el elemento menos electronegativo en el centro, enlazado por enlaces sigma a los átomos periféricos. Si la molécula es iónica sumar o restar su carga. Para que se cumpla la regla del octecto: sumar los electrones de valencia más los electrones compartidos. Asignar pares solitarios preferentemente a los átomos periféricos.

31 Writing Lewis Structures
All the valence e- of atoms must appear. Usually, the e- are paired. Usually, each atom requires an octet. H only requires 2 e-. Multiple bonds may be needed. Readily formed by C, N, O, S, and P.

32 Lewis Theory: An Overview
Valence e- play a fundamental role in chemical bonding. e- transfer leads to ionic bonds. Sharing of e- leads to covalent bonds. e- are transferred of shared to give each atom a noble gas configuration the octet.

33 Valence electrons are the outer shell electrons of an
atom. The valence electrons are the electrons that particpate in chemical bonding. Group # of valence e- e- configuration 1A 1 ns1 2A 2 ns2 3A 3 ns2np1 4A 4 ns2np2 5A 5 ns2np3 6A 6 ns2np4 7A 7 ns2np5 9.1

34 9.1

35 Lewis Structures Table 1.4 Lewis Structures

36 Lewis Symbols A chemical symbol represents the nucleus and the core e-. Dots around the symbol represent valence e-. • Si • • • N •• • P As Sb Bi •• Al • Se Ar I

37 Estructuras esqueleto
Atomos de hidrógeno son siempre terminales. Atomos centrales son generalmente aquellos con la menor electronegatividad Atomos de carbono son siempre atomos centrales. Estructuras son generalmente compactas y simétricas. H can only accommodate two electrons H and O are common exceptions to rule 2 Organic compounds are not compact nor symmetrical.

38 Reglas para escribir estructuras de Lewis

39

40 Lewis Structures for Ionic Compounds
Ba • O •• •• O Ba 2+ 2- BaO Mg • Cl •• •• Cl Mg 2+ - 2 MgCl2 Binary ionic compounds. Note the types of arrows used to move electrons – fishhooks for single e-. Write the Lewis symbol for each atom Determine how many e- each atom must gain or lose. Use multiples of one or both ions to balance the number of electrons.

41 The Ionic Bond - - - Li+ F Li + F 1s22s1 1s22s22p5 [He] 1s2 1s22s22p6
[Ne] Li Li+ + e- e- + F - F - Li+ + Li+ 9.2

42 Covalent Bonding

43 Multiple Covalent Bonds
• • • C O • •• • • O • C • • • O • • • • • • • • C O •• C O ••

44 Multiple Covalent Bonds
• • N • •• •• N • • N •• • • N • •• N ••

45 Writing Lewis Formulas: More rules….
The central atom in a molecule or polyatomic ion is determined by: The atom that requires the largest number of electrons to complete its octet goes in the center. C vs. N vs. O For two atoms in the same column, the less electronegative (or lowest ionization energy) element goes in the center. C vs. Si Hydrogen never is the central atom.

46

47 Writing Lewis Structures
Draw skeletal structure of compound showing what atoms are bonded to each other. Put least electronegative element in the center. Count total number of valence e-. Add 1 for each negative charge. Subtract 1 for each positive charge. Complete an octet for all atoms except hydrogen If structure contains too many electrons, form double and triple bonds on central atom as needed. 9.6

48 HCN H–CN : H2CO H–C=O : | ·· H
Ejemplo: Escribir las estructuras de Lewis completas para las siguientes especies químicas: CH4, HCN, H2CO, H2SO4, NH4+. H H · ·· | CH4 · C · + 4 · H  H ··C ·· H ; H–C–H · ·· | H H HCN H–CN : H2CO H–C=O : | ·· H ·· ·· : O : : O : ·· ·· ··  H2SO4 H ··O ··S ·· O ·· H ; H–O–S–O–H ·· ·· ··  : O : : O : ·· ·· H | NH H–N+H | H : O : || H–O–S–O–H || : O :

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50

51

52

53

54

55 5 + (3 x 7) = 26 valence electrons
Write the Lewis structure of nitrogen trifluoride (NF3). Step 1 – N is less electronegative than F, put N in center Step 2 – Count valence electrons N - 5 (2s22p3) and F - 7 (2s22p5) 5 + (3 x 7) = 26 valence electrons Step 3 – Draw single bonds between N and F atoms and complete octets on N and F atoms. Step 4 - Check, are # of e- in structure equal to number of valence e- ? 3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons F N 9.6

56 4 + (3 x 6) + 2 = 24 valence electrons
Write the Lewis structure of the carbonate ion (CO32-). Step 1 – C is less electronegative than O, put C in center Step 2 – Count valence electrons C - 4 (2s22p2) and O - 6 (2s22p4) -2 charge – 2e- 4 + (3 x 6) + 2 = 24 valence electrons Step 3 – Draw single bonds between C and O atoms and complete octet on C and O atoms. Step 4 - Check, are # of e- in structure equal to number of valence e- ? 3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons Step 5 - Too many electrons, form double bond and re-check # of e- 2 single bonds (2x2) = 4 1 double bond = 4 8 lone pairs (8x2) = 16 Total = 24 O C 9.6

57 Formal Charge CF = V-(PS + 1/2 PE ) FC = N° valence – N°lone pair e- -
N°bond pair e- 2 Formal charge = V – (L + 0.5S) CF = V-(PS + 1/2 PE ) The formal charge on an atom in a Lewis structure is the number of valence e- in the free atom minus the number of e- assigned to that atom in the Lewis structure.

58 Formal Charges on Atoms:
Comparison of the number of electrons “possessed” by an atom relative to its original number of valence electrons Formal charge = V – (L + 0.5S) V = # valence electrons of atom L = # electrons present as lone pair electrons on the atom in its compound form S = # shared electrons on the atom in its compound form

59 ( ) - - Two possible skeletal structures of formaldehyde (CH2O) H C O
An atom’s formal charge is the difference between the number of valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure. formal charge on an atom in a Lewis structure = total number of valence electrons in the free atom - total number of nonbonding electrons - 1 2 total number of bonding electrons ( ) The sum of the formal charges of the atoms in a molecule or ion must equal the charge on the molecule or ion. 9.7

60 ( ) - -1 +1 H C O = 1 2 = 4 - 2 - ½ x 6 = -1 = 6 - 2 - ½ x 6 = +1
C – 4 e- O – 6 e- 2H – 2x1 e- 12 e- 2 single bonds (2x2) = 4 1 double bond = 4 2 lone pairs (2x2) = 4 Total = 12 H C O formal charge on an atom in a Lewis structure = 1 2 total number of bonding electrons ( ) total number of valence electrons in the free atom - total number of nonbonding electrons formal charge on C = 4 - 2 - ½ x 6 = -1 formal charge on O = 6 - 2 - ½ x 6 = +1 9.7

61 ( ) - H C O = 1 2 = 4 - 0 - ½ x 8 = 0 = 6 - 4 - ½ x 4 = 0 C – 4 e-
H C O C – 4 e- O – 6 e- 2H – 2x1 e- 12 e- 2 single bonds (2x2) = 4 1 double bond = 4 2 lone pairs (2x2) = 4 Total = 12 formal charge on an atom in a Lewis structure = 1 2 total number of bonding electrons ( ) total number of valence electrons in the free atom - total number of nonbonding electrons formal charge on C = 4 - 0 - ½ x 8 = 0 formal charge on O = 6 - 4 - ½ x 4 = 0 9.7

62 Formal Charge and Lewis Structures
For neutral molecules, a Lewis structure in which there are no formal charges is preferable to one in which formal charges are present. Lewis structures with large formal charges are less plausible than those with small formal charges. Among Lewis structures having similar distributions of formal charges, the most plausible structure is the one in which negative formal charges are placed on the more electronegative atoms. Which is the most likely Lewis structure for CH2O? H C O -1 +1 H C O 9.7

63 Example Writing a lewis Structure for a Polyatomic Ion.
Write the Lewis structure for the nitronium ion, NO2+. Step 1: Total valence e- = – 1 = 16 e- Step 2: Plausible structure: O—N—O Step 3: Add e- to terminal atoms: O—N—O •• Step 4: Determine e- left over: 16 – 4 – 12 = 0

64 Example Step 5: Use multiple bonds to satisfy octets. O=N=O O—N—O
•• •• •• O=N=O + O—N—O •• •• •• •• Step 6: Determine formal charges: FC(O) = – (4) = 0 2 1 FC(N) = – (8) = +1 2 1

65 Alternative Lewis Structure
•• •• •• O N O + + - O—N—O •• •• •• •• FC(O≡) = – (6) = +1 2 1 FC(N) = – (8) = +1 2 1 FC(O—) = – (2) = -1 2 1

66 Alternative Lewis Structures
Sum of FC is the overall charge. FC should be as small as possible. Negative FC usually on most electronegative elements. FC of same sign on adjacent atoms is unlikely. + •• O≡N—O -

67 Formal charge = V – (L + 0.5S)
For C: 4 – ( x 8) = 0 Each O: 6 – ( x 4) = 0

68 Formal charge = V – (L + 0.5S)

69 Formal charge = V – (L + 0.5S)
What are the formal charges of C and O in carbon monoxide?

70 Example Using the Formal Charge Concept in Writing Lewis Structures.
Write the most plausible Lewis structure of nitrosyl chloride, NOCl, one of the oxidizing agents present in aqua regia. 2+ 2- - +

71 Resonancia. No siempre existe una única estructura de Lewis que pueda explicarlas propiedades de una molécula o ion. Por ejemplo, en el ion carbonato CO32– el C debería formar un doble enlace con uno de los O y sendos enlaces sencillos con los dos O– . Esto conllevaría a que las distancias C–O y C=O deberían ser distintas y ángulos de enlace distintos. Por difracción de rayos X se sabe que tanto distancias como los ángulos O–C–O son iguales.

72 Resonancia. Para explicar tales datos, se supone que los e– de enlace así como los pares electrónicos sin compartir, pueden desplazarse a lo largo de la molécula o ion, pudiendo formar más de una estructura de Lewis distinta. En el caso del ion CO32–, se podrían formar tres estructuras de Lewis en las que el doble enlace se formara con cada uno de los átomos de oxigeno, siendo las tres válidas. Cada una de estas formas contribuye por igual al la estructura del ion CO32–, siendo la verdadera estructura una mezcla de las tres.

73 What are the resonance structures of the carbonate (CO32-) ion?
A resonance structure is one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure. O + - O + - What are the resonance structures of the carbonate (CO32-) ion? O C - O C - O C - 9.8

74 Resonancia. Los tres enlaces C–O tienen 1/3 de doble enlace, por lo que la distancia es intermedia. Los tres átomos de oxígeno participan de 2/3 de carga negativa. Se utiliza el símbolo  entre las distintas formas resonantes.

75 Resonance O O O O O O O + - - + + -½ •• •• •• •• •• •• •• •• •• •• ••
Many Lewis structures may be written for a given structure.. Ozone has two good possibilities, but neither gives the correct structure that has two equivalent O-O bonds.

76 Resonance There are three possible structures for CO3-2 or SO3.
The double bond can be placed in one of three places. When two or more Lewis formulas are necessary to show the bonding in a molecule, we must use equivalent resonance structures to show the molecule’s structure. Double-headed arrows are used to indicate resonance formulas.

77 The best Lewis formula of SO3 is:
There are really no single or double bonds in CO32- or SO3! The best Lewis formula of SO3 is:

78

79 Benzene, C6H6, is another compound famous for resonance.

80 Benzene

81 Ejercicio A: Escribir las distintas formas resonantes del ácido nítrico. HNO3
+ N O H : O · · – N O H : O + – · ·  O también N O H · · : O – +

82 Resonance (Lewis) Structures:
ozone = O3 1. Assume a structure and draw one bond between the bonded atoms. 2. Add all valence electrons to each atom. 3. Rearrange electrons to maximize the number of bonds without exceeding the noble gas configuration for each element. 4. Assign formal charge to each atom.

83 Resonance (Lewis) Structures:
. .. . .. .. .. .. ozone = O3 1. Assume a structure and draw one bond between the bonded atoms. 2. Add all valence electrons to each atom. 3. Rearrange electrons to maximize the number of bonds without exceeding the noble gas configuration for each element. 4. Assign formal charge to each atom.

84 Resonance (Lewis) Structures:
.. .. . . .. . .. .. .. .. .. 1. Assume a structure and draw one bond between the bonded atoms. 2. Add all valence electrons to each atom. 3. Rearrange electrons to maximize the number of bonds without exceeding the noble gas configuration for each element. 4. Assign formal charge to each atom.

85 Resonance (Lewis) Structures:
.. .. . . .. . .. .. .. .. .. .. .. .. .. .. 1. Assume a structure and draw one bond between the bonded atoms. 2. Add all valence electrons to each atom. 3. Rearrange electrons to maximize the number of bonds without exceeding the noble gas configuration for each element. 4. Assign formal charge to each atom.

86 Resonance (Lewis) Structures:
.. ..     .. .. .. .. .. .. .. .. .. Curved arrow formalism, an important concept. This formalism denotes electron movement. The arrow head identifies where the electron pair is going and the tail identifies its initial position.

87 Resonance (Lewis) Structures:
.. ..     .. .. .. .. .. O .. O .. O .. O .. O .. O Curved arrow formalism, an important concept. This formalism denotes electron movement. The arrow head identifies where the electron pair is going and the tail identifies its initial position.

88 Resonance (Lewis) Structures:
.. ..   .. .. .. O .. O .. O Curved arrow formalism, an important concept. This formalism denotes electron movement. The arrow head identifies where the electron pair is going and the tail identifies its initial position.

89 Resonance (Lewis) Structures:
.. ..     .. .. .. .. O3 ozone All resonance structures taken together provide a better electron description of a molecule than any single structure.

90 Rules for Resonance Structures:
1. Resonance structures exist only on paper. formaldehyde   .. ..

91 Rules for Resonance Structures:
2. In writing resonance structures we are only allowed to move electrons. O3

92 Rules for Resonance Structures:
2. In writing resonance structures we are only allowed to move electrons. O3

93 Rules for Resonance Structures:
3. All of the structures must be proper Lewis structures.

94 Rules for Resonance Structures:
4. The energy of the actual molecule is lower than the energy that might be estimated for any contributed structure. ..   ..

95 Rules for Resonance Structures:
5. Equivalent resonance structures make equal contributions to the hybrid, and a system described by them has a large resonance stabilization. ..     .. .. .. ..

96 .. Rules for Resonance Structures:
6. The more stable a structure is (when taken by itself), the greater its contribution to the hybrid. ..   ..   ..

97 Rules for evaluating the relative energies of resonance structures:
a. The more covalent bonds a structure has the more stable it is. .. more stable   .. .. .. .. more stable

98 Rules for evaluating the relative energies of resonance structures:
b. Structures in which all of the atoms have a complete valence shell of electrons are more stable.   .. more stable

99 Rules for evaluating the relative energies of resonance structures:
c. Opposite charge separation decreases stability. decreasing stability

100 Rules for evaluating the relative energies of resonance structures:
d. Resonance contributors with negative charge on electronegative atoms are more stable. least stable most stable  ..  ..   .. .. .. .. .. ..

101 Rules for evaluating the relative energies of resonance structures:
e. Resonance contributors with negative charge on electronegative atoms are more stable. more stable more stable

102 Excepciones a la teoría de Lewis
Moléculas tipo NO y NO2 que tienen un número impar de electrones. Moléculas tipo BeCl2 o BF3 con marcado carácter covalente en las cuales el átomo de Be o de B no llegan a tener 8 electrones. Moléculas tipo PCl5 o SF6 en las que el átomo central tiene 5 o 6 enlaces (10 o 12 e– ). Sólo en caso de que el no-metal no esté en el segundo periodo, pues a partir del tercero existen orbitales “d” y puede haber más de cuatro enlaces.

103 Enlace Covalente

104 Enlace Covalente Algunos ejemplos a la excepción al Octeto son:
Moléculas deficientes de electrones (octeto incompleto ) Moleculas hipervalentes (expansión del octeto)

105 Exceptions to the Octet Rule
The Incomplete Octet Be – 2e- 2H – 2x1e- 4e- BeH2 H Be B – 3e- 3F – 3x7e- 24e- 3 single bonds (3x2) = 6 9 lone pairs (9x2) = 18 Total = 24 F B BF3 9.9

106 Exceptions to the Octet Rule
The covalent compounds of Be. The covalent compounds of the IIIA Group. Species which contain an odd number of electrons. Species in which the central element must have a share of more than 8 valence electrons to accommodate all of the substituents (expanded valence shell). Compounds of the d- and f-transition metals.

107 Lewis dot formula for BBr3

108 Exceptions to the Octet Rule
Incomplete octets. •• B F - + •• B F - •• + •• F •• B •• F F •• •• •• •• ••

109 Limitations of the Octet Rule
AsF5

110 Expanded valence shell

111 Exceptions to the Octet Rule
Odd-Electron Molecules N – 5e- O – 6e- 11e- NO N O The Expanded Octet (central atom with principal quantum number n > 2) S F S – 6e- 6F – 42e- 48e- 6 single bonds (6x2) = 12 18 lone pairs (18x2) = 36 Total = 48 SF6 9.9

112 Exceptions to the Octet Rule
Odd e- species. N=O • •• •• •• H •• • O—H H—C—H H • ••

113 Exceptions to the Octet Rule
Expanded octets. S F •• •• P Cl •• •• Cl •• P •• •• Cl Cl •• •• •• •• ••

114 Expanded Valence Shell
It is not clear which is the more correct representation.

115 Exceptions to the Octet Rule
Expanded octets. S F •• •• P Cl •• •• Cl •• P •• •• Cl Cl •• •• •• •• ••

116 Polar Covalent Bonds Covalent bonds: electrons in a bond are not shared equally The two atoms involved in the bond must have different electronegativities.

117 Some examples of polar covalent bonds:
HF

118 Shown below is an electron density map of HF.
Blue areas indicate low electron density. Red areas indicate high electron density. Electron-rich end. Is this the F or H?

119 Shown below is an electron density map of HI. HI is only slightly polar.

120

121 Polar molecules can be attracted by magnetic and electric fields.

122 Dipole Moments Consider molecules whose centers of positive and negative charge do not coincide, have an asymmetric charge distribution, and are polar. Ex. HF The dipole moment has the symbol .  is the product of the distance,d, separating charges of equal magnitude and opposite sign, and the magnitude of the charge, q.

123 Dipole Moment

124 Dipole Moments Molecules that have a small separation of charge have a small . Molecules that have a large separation of charge have a large .

125 Polarity of a Molecule Two conditions that must be true for a molecule to be polar: 1) There must be at least one polar bond present or one lone pair of electrons. 2) The polar bonds, if there are more than one, and lone pairs must be arranged so that their dipole moments do not cancel one another.

126 CO2 Polar or non-polar? This is a nonpolar molecule!