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2 Periodic Trends in Atomic Properties

3 Characteristic properties and trends of the elements are the basis of the periodic table’s design.

4 These trends allow us to use the periodic table to accurately predict properties and reactions of a wide variety of substances.

5 Metals and Nonmetals

6 Chemical Properties of Metals metals tend to lose electrons and form positive ions (cations). nonmetals tend to gain electrons and form negative ions (anions). Chemical Properties of Nonmetals When metals react with nonmetals electrons are usually transferred from the metal to the nonmetal.

7 Physical Properties of Metals lustrous malleable good conductors of heat good conductors of electricity nonlustrous brittle poor conductors of heat poor conductors of electricity Physical Properties of Nonmetals

8 Metalloids have properties that are intermediate between metals and nonmetals

9 The Metalloids 1.boron 2.silicon 3.germanium 4.arsenic 5.antimony 6.tellurium 7.polonium

10 Metals are found to the left of the metalloids Nonmetals are found to the right of the metalloids. 11.1

11 Atomic Radius

12 Radii of atoms increase down a group. For each step down a group, electrons enter the next higher energy level.

13 Radii of atoms tend to decrease from left to right across a period. For representative elements within the same period the energy level remains constant as electrons are added. Each time an electron is added a proton is a added to the nucleus. This increase in positive nuclear charge pulls all electrons closer to the nucleus.

14 Ionization Energy

15 The ionization energy of an atom is the energy required to remove the outermost electron from an atom. Na + ionization energy → Na + + e -

16 The first ionization energy is the amount of energy required to remove the first electron from an atom. He + first ionization energy → He + + e - He + 2,372 kJ/mol → He + + e - The second ionization energy is the amount of energy required to remove the second electron from an atom. He+ + 5,247 kJ/mol → He ++ + e- He + second ionization energy → He + + e -

17 As each succeeding electron is removed from an atom ever higher energies are required.

18 Periodic relationship of the first ionization energy for representative elements in the first four periods. Ionization energies gradually increase from left to right across a period. IA IIA IIIA IVA VA VIA VIIA Noble Gases

19 Periodic relationship of the first ionization energy for representative elements in the first four periods Ionization energies of Group A elements decrease from top to bottom in a group. IA IIA IIIA IVA VA VIA VIIA Noble Gases Distance of Outer Shell Electrons From Nucleus nonmetalsmetals nonmetals have higher ionization potentials than metals

20 Lewis “Dot” Structures of Atoms

21 Metals form cations and nonmetals form anions to attain a stable valence electron structure.valence

22 This stable structure often consists of two s and six p electrons. These rearrangements occur by losing, gaining, or sharing electrons.

23 Na with the electron structure 1s 2 2s 2 2p 6 3s 1 has 1 valence electron. The Lewis structure of an atom is a representation that shows the valence electrons for that atom.valence Fluorine with the electron structure 1s 2 2s 2 2p 5 has 7 valence electrons

24 valence electrons: the electrons that occupy the outermost energy level of an atom. valence electrons are responsible for the electron activity that occurs to form chemical bonds.

25 The Lewis structure of an atom uses dots to show the valence electrons of atoms. The number of dots equals the number of s and p electrons in the atom’s outermost shell. B Paired electrons Unpaired electron Symbol of the element 2s22p12s22p1

26 The number of dots equals the number of s and p electrons in the atom’s outermost shell. S 3s23p43s23p4 The Lewis structure of an atom uses dots to show the valence electrons of atoms.

27 Lewis Structures of the first 20 elements.

28 The chemistry of many elements, especially the representative ones, is to attain the same outer electron structure as one of the noble gases.

29 With the exception of helium, this structure consists of eight electrons in the outermost energy level.

30 The Covalent Bond: Sharing Electrons

31 A covalent bond consists of a pair of electrons shared between two atoms. In the millions of chemical compounds that exist, the covalent bond is the predominant chemical bond.

32 Substances which covalently bond exist as molecules. Carbon dioxide bonds covalently. It exists as individually bonded covalent molecules containing one carbon and two oxygen atoms.

33 The term molecule is not used when referring to ionic substances. Sodium chloride bonds ionically. It consists of a large aggregate of positive and negative ions. No molecules of NaCl exist.

34 Covalent bonding in the hydrogen molecule Two 1s orbitals from each of two hydrogen atoms overlap. Each 1s orbital contains 1 electron. The orbital of the electrons includes both hydrogen nuclei. The most likely region to find the two electrons is between the two nuclei. The two nuclei are shielded from each other by the electron pair. This allows the two nuclei to draw close together. Two 1s orbitals from each of two hydrogen atoms overlap.

35 Covalent bonding in the chlorine molecule Each unpaired 3p orbital on each chlorine atom contains 1 electron. Two 3p orbitals from each of two chlorine atoms overlap. The orbital of the electrons includes both chlorine nuclei. The most likely region to find the two electrons is between the two nuclei. The two nuclei are shielded from each other by the electron pair. This allows the two nuclei to draw close together. Two 3p orbitals from each of two chlorine atoms overlap. Each chlorine now has 8 electrons in its outermost energy level.

36 hydrogen chlorine iodine nitrogen Covalent bonding with equal sharing of electrons occurs in diatomic molecules formed from one element. A dash may replace a pair of dots. H-H

37 Electronegativity Linus Pauling

38 electronegativity The relative attraction that an atom has for a pair of shared electrons in a covalent bond.

39 If the two atoms that constitute a covalent bond are identical then there is equal sharing of electrons. This is called nonpolar covalent bonding. Ionic bonding and nonpolar covalent bonding represent two extremes.

40 If the two atoms that constitute a covalent bond are not identical then there is unequal sharing of electrons. This is called polar covalent bonding. One atom assumes a partial positive charge and the other atom assumes a partial negative charge. –This charge difference is a result of the unequal attractions the atoms have for their shared electron pair.

41 Polar and Non-Polar

42 : HCl ++ -- Shared electron pair. : The shared electron pair is closer to chlorine than to hydrogen. Partial positive charge on hydrogen. Partial negative charge on chlorine. Chlorine has a greater attraction for the shared electron pair than hydrogen. Polar Covalent Bonding in HCl The attractive force that an atom of an element has for shared electrons in a molecule or a polyatomic ion is known as its electronegativity.

43 A scale of relative electronegativities was developed by Linus Pauling.

44 Electronegativity decreases down a group for representative elements. Electronegativity generally increases left to right across a period.

45 The electronegativities of the metals are low. The electronegativities of the nonmetals are high.

46 The polarity of a bond is determined by the difference in electronegativity values of the atoms forming the bond.

47 If the electronegativity difference between two bonded atoms is greater than the bond will be more ionic than covalent. If the electronegativity difference is greater than 2, the bond is strongly ionic. If the electronegativity difference is less than 1.5, the bond is strongly covalent.

48 HH Hydrogen Molecule If the electronegativities are the same, the bond is nonpolar covalent and the electrons are shared equally. The molecule is nonpolar covalent. Electronegativity 2.1 Electronegativity 2.1

49 If the electronegativities are the same, the bond is nonpolar covalent and the electrons are shared equally. Cl Chlorine Molecule Electronegativity 3.0 Electronegativity 3.0 The molecule is nonpolar covalent. Electronegativity Difference = 0.0

50 If the electronegativities are not the same, the bond is polar covalent and the electrons are shared unequally. HCl Hydrogen Chloride Molecule Electronegativity 2.1 Electronegativity 3.0 The molecule is polar covalent. ++ -- Electronegativity Difference = 0.9

51 Sodium Chloride Na + Cl - If the electronegativities are very different, the bond is ionic and the electrons are transferred to the more electronegative atom. Electronegativity 0.9 Electronegativity 3.0 The bond is ionic.No molecule exists. Electronegativity Difference = 2.1

52 A dipole is a molecule that is electrically asymmetrical, causing it to be oppositely charged at two points. A dipole can be written as + -

53 An arrow can be used to indicate a dipole. The arrow points to the negative end of the dipole. HClHBrH O H Molecules of HCl, HBr and H 2 O are polar.

54 A molecule containing different kinds of atoms may or may not be polar depending on its shape. The carbon dioxide molecule is nonpolar because its carbon-oxygen dipoles cancel each other by acting in opposite directions.

55 Relating Bond Type to Electronegativity Difference.

56 Lewis Structures of Compounds

57 In writing Lewis structures, the most important consideration for forming a stable compound is that the atoms attain a noble gas configuration.

58 The most difficult part of writing Lewis structures is determining the arrangement of the atoms in a molecule or an ion. In simple molecules with more than two atoms, one atom will be the central atom surrounded by the other atoms.

59 Cl 2 O has two possible arrangements. Cl-Cl-O The two chlorines can be bonded to each other. Cl-O-Cl The two chlorines can be bonded to oxygen. Usually the single atom will be the central atom.

60 Practice Writing Lewis Structures

61 AtomGroupValence Electrons ClVIIA7 HIA1 CIVA4 NVA5 SVIA6 PVA5 IVIIA7 Valence Electrons of Group A Elements

62 3-Dimensional Shapes Linear 180  Bent 105  Trigonal Planar 120  Tetrahedral  Trigonal Pyramidal 107 

63 Covalent Bonding Structures Molecular Formula Lewis “dot” Structure 3-D Structure Name Bond Angle Polar or Non-polar H2OH2O CO 2 PH 3 NO 3 – CH 4 Bent 105  Polar Linear 180  Non-Polar Trigonal Pyramidal 107  Polar Trigonal Planar 120  Non-Polar Tetrahedral  Non-Polar

64 The Ionic Bond: Transfer of Electrons From One Atom to Another

65 After sodium loses its 3s electron it has attained the same electronic structure as neon.

66 After chlorine gains a 3p electron it has attained the same electronic structure as argon.

67 Formation of NaCl

68 The 3s electron of sodium transfers to the half-filled 3p orbital of chlorine. Lewis representation of sodium chloride formation. A sodium ion (Na+) and a chloride ion (Cl - ) are formed. The force holding Na + and Cl - together is an ionic bond.

69 Formation of MgCl 2

70 Two 3s electrons of magnesium transfer to the half-filled 3p orbitals of two chlorine atoms. A magnesium ion (Mg 2+ ) and two chloride ions (Cl - ) are formed. The forces holding Mg 2+ and two Cl - together are ionic bonds.

71 NaCl is made up of cubic crystals.In the crystal each sodium ion is surrounded by six chloride ions.

72 In the crystal each chloride ion is surrounded by six sodium ions.

73 The ratio of Na + to Cl - is 1:1 There is no molecule of NaCl

74 Relative Size of Sodium Ion to Chloride Ion

75 Sodium ion is smaller than a sodium atom because: (1) the sodium atom has lost its outermost electron. (2) The 10 remaining electrons are now attracted by 11 protons and are drawn closer to the nucleus. 11.6

76 Chloride ion is larger than a chloride atom because: (1) the chlorine atom has gained an electron and now has 18 electrons and 17 protons. (2) The nuclear attraction on each electron is thereby decreased, allowing the chlorine to expand. 11.6

77 Metals usually have one, two or three electrons in their outer shells. When a metal reacts it: –usually loses one two or three electrons –attains the electron structure of a noble gas –becomes a positive ion. The positive ion formed by the loss of electrons is much smaller than the metal atom.

78 Nonmetals are usually only a few electrons short of having a noble gas structure. When a nonmetal reacts it: –usually gains one two or three electrons –attains the electron structure of a noble gas –becomes a negative ion. The negative ion formed by the gain of electrons is much larger than the nonmetal atom.

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80 Predicting Formulas of Ionic Compounds

81 In almost all stable chemical compounds of representative elements, each atom attains a noble gas electron configuration.

82 Ions are always formed by adding or removing electrons from an atom.

83 Most often ions are formed when metals combine with nonmetals. Metals will lose electrons to attain a noble gas configuration. Nonmetals will gain electrons to attain a noble gas configuration.

84 The charge on an ion can be predicted from its position in the periodic table.

85 elements of Group IIA have a +2 charge elements of Group IA have a +1 charge elements of Group VA have a -3 charge elements of Group VIA have a -2 charge elements of Group VIIA have a -1 charge

86 Writing Formulas From Names of Compounds

87 A chemical compound must have a net charge of zero.

88 If the compound contains ions, then the charges on all of the ions must add to zero.

89 Write the formula of calcium chloride. Step 1. Write down the formulas of the ions. Ca 2+ Cl - Step 2. Combine the smallest numbers of Ca 2+ and Cl - so that the sum of the charges equals zero. (2+) + 2(1-) = 0 The correct formula is CaCl 2 The lowest common multiple of +2 and –1 is 2 The cation is written first. The anion is written second. (Ca 2+ ) + 2(Cl - ) = 0

90 Write the formula of barium phosphide. Step 1. Write down the formulas of the ions. Ba 2+ P 3- Step 2. Combine the smallest numbers of Ba 2+ and P 3- so that the sum of the charges equals zero. 3(2+) + 2(3-) = 0 The correct formula is Ba 3 P 2 The lowest common multiple of +2 and –3 is 6 3(Ba 2+ ) + 2(P 3- ) = 0 The cation is written first. The anion is written second.

91 Write the formula of magnesium oxide. Step 1. Write down the formulas of the ions. Mg 2+ O 2- Step 2. Combine the smallest numbers of Mg 2+ and O 2- so that the sum of the charges equals zero. (2+) + (2-) = 0 The correct formula is MgO The lowest common multiple of +2 and –2 is 1 ( Mg 2+ ) + (O 2- ) = 0

92 Write the Formula of Sodium Peroxide givesor Na 2 O 2 NaO

93 Write the Formula of Sodium Peroxide NaO does not contain the peroxide anion gives not Na 2 O 2 NaO Don’t mess with the subscripts of polyatomic ions!!

94 Metals will lose electrons to attain a noble gas configuration. Nonmetals will gain electrons to attain a noble gas configuration. Barium and Sulfur Combine. –sulfur gains two electrons from barium and attains an argon configuration. –barium loses two electrons to sulfur and attains a xenon configuration. S [Ne]3s 2 3p 4 Ba [Xe]6s 2 Ba → [Xe] + 2e - S + 2e - → [Ar] Ba + S → BaS

95 Because of similar electron structures, the elements of a family generally form compounds with the same atomic ratios.

96

97 The elements of a family have the same outermost electron configuration except that the electrons are in different energy levels.

98 The atomic ratio of the alkali metal sodium to chlorine is 1:1 in NaCl. The atomic ratios of the other alkali metal chlorides can be predicted to also be 1:1. LiCl, KCl, CsCl, FrCl

99 The atomic ratio of hydrogen to nitrogen is 3:1 in ammonia (NH 3 ). Nitrogen is the first member of group VA. The atomic ratio of hydrogen when combined with other group VA elements can be predicted to also be 3:1. PH 3, AsH 3, SbH 3, BiH 3

100

101 The Following Slides are for Extra Practice/ Explanation

102 Procedures for Writing Lewis Structures

103 Step 1 Obtain the total number of valence electrons to be used in the structure by adding the number of valence electrons in all the atoms in the molecule or ion. –If you are writing the structure of an ion, add one electron for each negative charge or subtract one electron for each positive charge on the ion.

104 Step 1. The total number of valence electrons is eight, two from the two hydrogen atoms and six from the oxygen atom. Write the Lewis structure for H 2 O.

105 Step 2. Write the skeletal arrangement of the atoms and connect them with a single covalent bond (two dots or one dash). –Hydrogen, which contains only one bonding electron, can form only one covalent bond. –Oxygen atoms normally have a maximum of two covalent bonds (two single bonds, or one double bond).

106 Step 2. The two hydrogen atoms are connected to the oxygen atom. Write the skeletal structure: Write the Lewis structure for H 2 O. Place two dots between the hydrogen and oxygen atoms to form the covalent bonds. H O H or H O H : : ::

107 Step 3. Subtract two electrons for each single bond you used in Step 2 from the total number of electrons calculated in Step 1. –This gives you the net number of electrons available for completing the structure.

108 Step 3. Subtract the four electrons used in Step 2 from eight to obtain four electrons yet to be used. Write the Lewis structure for H 2 O. H O H ::

109 Step 4. Distribute pairs of electrons (pairs of dots) around each atom (except hydrogen) to give each atom a noble gas configuration.

110 Step 4. Distribute the four remaining electrons in pairs around the oxygen atom. Hydrogen atoms cannot accommodate any more electrons. Write the Lewis structure for H 2 O. These arrangements are Lewis structures because each atom has a noble gas electron structure. H O H or H O H : : :: : : : : The shape of the molecule is not shown by the Lewis structure.

111 Step 1. The total number of valence electrons is 16, four from the C atom and six from each O atom. Write a Lewis structure for CO 2.

112 Step 2. The two O atoms are bonded to a central C atom. Write the skeletal structure and place two electrons between the C and each oxygen. O C O :: Write a Lewis structure for CO 2.

113 Write a Lewis structure for CO 2. Step 3. Subtract the four electrons used in Step 2 from 16 (the total number of valence electrons) to obtain 12 electrons yet to be used. O C O ::

114 O C O :: Step 4. Distribute the 12 electrons (6 pairs) around the carbon and oxygen atoms. Three possibilities exist. Many of the atoms in these structures do not have eight electrons around them. Write a Lewis structure for CO 2. O C O :: :: :: : : : : : : : : : : 4 electrons 6 electrons 6 electrons 6 electrons : : : : : : 6 electrons IIIIII

115 Write a Lewis structure for CO 2. O C O :::: : : : : Step 5. Remove one pair of unbonded electrons from each O atom in structure I and place one pair between each O and the C atom forming two double bonds. O C O : : :: : : : : : : :: : : : : Each atom now has 8 electrons around it. Carbon is sharing 4 electron pairs. double bond

116 Complex Lewis Structures

117 There are some molecules and polyatomic ions for which no single Lewis structure consistent with all characteristics and bonding information can be written.

118 Step 1. The total number of valence electrons is 24, 5 from the nitrogen atom and 6 from each O atom, and 1 from the –1 charge. Write a Lewis structure for NO 2.

119 Since the extra electron present results in nitrate having a –1 charge, the ion is enclosed in brackets with a – charge. Step 2. The three O atoms are bonded to a central N atom. Write the skeletal structure and place two electrons between each pair of atoms. Write a Lewis structure for NO 2. O N O :: O : -

120 Step 3. Subtract the 6 electrons used in Step 2 from 24, the total number of valence electrons, to obtain 18 electrons yet to be placed. Since the extra electron present results in nitrate having a –1 charge, the ion is enclosed in brackets with a – charge. O N O :: O : - Write a Lewis structure for NO 2.

121 O N O O Step 4. Distribute the 18 electrons around the N and O atoms. Write a Lewis structure for NO 2. : ::: :: : : : :: : electron deficient

122 Write a Lewis structure for NO 2. Step 5. One pair of electrons is still needed to give all the N and O atoms a noble gas structure. Move the unbonded pair of electrons from the N atom and place it between the N and the electron-deficient O atom, making a double bond. : ::: :: : : : O N O O :: : -

123 Write a Lewis structure for NO 2. Step 5. One pair of electrons is still needed to give all the N and O atoms a noble gas structure. Move the unbonded pair of electrons from the N atom and place it between the N and the electron-deficient O atom, making a double bond. N O : : O : : : O : : : : -

124 A molecule or ion that shows multiple correct Lewis structures exhibits resonance. Write a Lewis structure for NO 2. Step 5. There are three possible Lewis structures. N O : : O : : : O : : : : - : N O : : O : : O : : : : - Each Lewis structure is called a resonance structure. N O : : O : : : O : : : : - 

125 Compounds Containing Polyatomic Ions

126 A polyatomic ion is a stable group of atoms that has either a positive or negative charge and behaves as a single unit in many chemical reactions.

127 Sodium nitrate, NaNO 3, contains one sodium ion and one nitrate ion. sodium ion Na + nitrate ion N O : : O : : : O : : : : - Na +

128 The nitrate ion is a polyatomic ion composed of one nitrogen atom and three oxygen atoms. N O : : O : : : O : : : : - Na + It has a charge of –1 One nitrogen and three oxygen atoms have a total of 23 valence electrons.

129 The –1 charge on nitrate adds an additional valence electron for a total of 24. N O : : O : : : O : : : : - Na + The additional valence electron comes from a sodium atom which becomes a sodium ion.

130 Sodium nitrate has both ionic and covalent bonds. N O : : O : : : O : : : : - Na + Ionic bonds exist between the sodium ions and the carbonate ions. covalent bond ionic bond Covalent bonds are present between the carbon and oxygen atoms within the carbonate ion.

131 When sodium nitrate is dissolved in water, the ionic bond breaks. N O : : O : : : O : : : : - Na + The sodium ions and nitrate ions separate from each other forming separate sodium and nitrate ions. N O : : O : : : O : : : : - Na + The nitrate ion, which is held together by covalent bonds, remains as a unit.

132 Molecular Shape

133 The 3-dimensional arrangement of the atoms within a molecule is a significant determinant of molecular interactions.

134 (bent)

135 The Valence Shell Electron Pair (VSEPR) Model

136 The VSEPR model is based on the idea that electron pairs will repel each other electrically and will seek to minimize this repulsion. To accomplish this minimization, the electron pairs will be arranged as far apart as possible around a central atom.

137 BeCl 2 is a molecule with only two pairs of electrons around beryllium, its central atom. Its electrons are arranged 180 o apart for maximum separation.

138 BF 3 is a molecule with three pairs of electrons around boron, its central atom. Its electrons are arranged 120 o apart for maximum separation. This arrangement of atoms is called trigonal planar.

139 CH 4 is a molecule with four pairs of electrons around carbon, its central atom. An obvious choice for its atomic arrangement is a 90 o angle between its atoms with all of its atoms in a single plane. However, since the molecule is 3-dimensional the molecular structure is tetrahedral with a bond angle of o.

140 Ball and stick models of methane, CH 4, and carbon tetrachloride, CCl

141 Ammonia, NH 3, has four electron pairs around nitrogen. The arrangement of electron pairs around nitrogen is tetrahedral.

142 The NH 3 molecule is pyramidal. NH 3 has one unbonded pair of electrons.

143 Water has four electron pairs around oxygen. The arrangement of electron pairs around oxygen is tetrahedral.

144 The H 2 O molecule is bent. H 2 O has two unbonded pairs of electrons.

145