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

8 Atomic Radius

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

10 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.

11 Ionization Energy

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

13 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

14 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

15 Lewis “Dot” Structures of Atoms

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

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

18 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

19 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

20 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.

21 Lewis Structures of the first 20 elements.

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

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

24 The Covalent Bond: Sharing Electrons

25 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.

26 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.

27 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.

28 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.

29 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

30 Electronegativity Linus Pauling

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

32 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.

33 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.

34 Polar and Non-Polar

35 : 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.

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

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

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

39 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.

40 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.

41 Lewis Structures of Compounds

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

43 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.

44 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.

45 Practice Writing Lewis Structures

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

47 3-Dimensional Shapes Linear 180  Bent 105  Trigonal Planar 120  Tetrahedral  Trigonal Pyramidal 107 

48 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

49 “Define ‘resonance’? Sure, that’s where you live.”

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

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

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

53 Formation of NaCl

54 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.

55 Formation of MgCl 2

56 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.

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

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

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

60 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.

61 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.

62 Predicting Formulas of Ionic Compounds

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

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

65 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.

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

67 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

68 Writing Formulas From Names of Compounds

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

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

71 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

72 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.

73 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

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

75 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!!

76 Combine to Give Compounds (Do Not Name!) ionsBr – O –2 NO 3 – PO 4 –3 CO 3 –2 NH 4 + Sn +2 Al +3 H+H+ NH 4 Br(NH 4 ) 2 ONH 4 NO 3 (NH 4 ) 3 PO 4 (NH 4 ) 2 CO 3 SnBr 2 SnOSn(NO 3 ) 2 Sn 3 (PO 4 ) 2 SnCO 3 AlBr 3 Al 2 O 3 Al(NO 3 ) 3 AlPO 4 Al 2 (CO 3 ) 3 HBrH2OH2OHNO 3 H 3 PO 4 H 2 CO 3

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