1. Ionic Compounds and Covalent Compounds

2. Ionic and Metallic Bonding

3. Valence electrons Scientists learn that all elements within each group behave similarly because they have same number of valence electrons Valence electrons – electrons in highest occupied energy level of element’s atoms - determines chemical properties - look at group number to determine number of valence electrons Example: Hydrogen, Lithium, Sodium, Potassium, etc. have 1 valence electron being they are in group 1A

4. Electron Dot Structure Electron Dot Structure – diagrams that show valence electrons as dots

5. Octet Rule Noble gases, such as neon and argon, are unreactive in chemical reactions Chemist Gilbert Lewis used this fact to explain why atoms form certain kinds of ions and molecules - called his explanation the octet rule: - In forming compounds, atoms tend to achieve electron configuration of a noble gas

6. Octet Rule –Atoms of metals tend to lose valence electrons - leaves a complete octet in next- lowest energy level –Atoms of some non-metals tend to gain electrons -some share electrons with another nonmetal to achieve complete octet

7. Formation of Cations Cation – positively charged ion - due to atom’s loss of valence electron(s) - most common are produced by loss of valence electrons from metal atoms Example:

8. Formation of Cations electron configuration of the sodium ion is the same as that of a neon atom

9. Formation of Cations Using electron dot structures, you can show the ionization more simply

10. Formation of Cations magnesium atom attains electron configuration of neon by losing both valence electrons - loss of valence electrons produces magnesium cation with a charge of 2+

11. Formation of Cations Cations of Group 1A elements always have a charge of 1+ Cations of group 2A elements always have a charge of 2+

12. Formation of Cations Copper atom can ionize to form a 1+ cation (Cu + ) By losing its lone 4s electron, copper attains noble-gas electron configuration

13. Formation of Anions Anion – negatively charged ion - name usually ends in “-ide”

14. Formation of Anions Gain of one electron gives chlorine an octet = converts chlorine atom into chloride ion - It has same electron configuration as the noble gas argon

15. Formation of Anions Each dot in electron dot structure represents an electron in the valence shell in the electron configuration diagram

16. Halide Ions Halide ions- ions produced when atoms of chlorine and other halogens gain electrons – All halogen atoms have 7 valence electrons – All halogen atoms need to gain only one electron to achieve electron configuration of a noble gas

17. Common Anions

18. Formation of Ionic Compounds Ionic Compounds – compounds composed of cations and anions - electrically neutral - total positive charge = total negative charge Ionic Bonds – electrostatic forces that hold ions together in ionic compounds

19. Ionic Bonds

20. Formula Units Chemical Formula - shows kinds and numbers of atoms in smallest representative unit of a substance Formula Unit - lowest whole-number ratio of ions in an ionic compound

21. Formation of Ionic Compounds NaCl is the chemical formula for sodium chloride

22. Properties of Ionic Compounds Most ionic compounds crystalline solids at room temp - generally have high melting points

23. Properties of Ionic Compounds Coordination number - number of ions of opposite charge that surround the ion in a crystal In NaCl, each ion has a coordination number of 6

24. Properties of Ionic Compounds In CsCl, each ion has a coordination number of 8 In TiO 2, each Ti 4+ ion has a coordination number of 6, while each O 2- ion has a coordination number of 3

25. Properties of Ionic Compounds – Ionic compounds can conduct an electric current when melted or dissolved in water

26. Metallic Bonds and Metallic Properties Valence electrons of metal atoms can be modeled as a sea of electrons – Valence electrons are mobile and can drift freely from one part of the metal to another –Metallic bonds- consist of the attraction of the free-floating valence electrons for the positively charged metal ions

27. Metallic Bonds and Metallic Properties Metals are ductile— they can be drawn into wires

28. Metallic Bonds and Metallic Properties Force can change shape of a metal Force can shatter an ionic crystal

29. Crystalline Structure of Metals – Metal atoms are arranged in very compact and orderly patterns

30. Alloys – Alloys - mixtures composed of two or more elements - at least one of which is a metal - properties are often superior to those of their component elements

31. Covalent Bonding

32. Molecules and Molecular Compounds In nature, matter takes many forms - The noble gases, including helium and neon, are monatomic -That means they exist as single atoms

33. Molecules and Molecular Compounds Some compounds are so different from ionic compounds -attractions between ions fail to explain their bonding covalent bond – bonds holding atoms together by sharing electrons

34. Molecules and Molecular Compounds Molecule - a neutral group of atoms joined together by covalent bonds Diatomic molecule - a molecule consisting of two atoms - Example: O 2

35. Molecules and Molecular Compounds molecular compound - compound composed of molecules Molecular compounds tend to have relatively lower melting and boiling points than ionic compounds Example:

36. Molecules and Molecular Compounds

37. Molecular Formulas Molecular formula- chemical formula of a molecular compound - shows how many atoms of each element a molecule contains

38. Molecular Formulas

39. Formulas of Some Molecular Compounds

40. The Octet Rule in Covalent Bonding In covalent bonds, electron sharing usually occurs so that atoms attain the electron configurations of noble gases - Example: each hydrogen atom has 1 electron - pair of hydrogen atoms share 2 electrons in a covalent bond H 2

41. Single Covalent Bonds Single covalent bond- two atoms held together by sharing a pair of electrons

42. Single Covalent Bonds An electron dot structure such as H:H represents shared pair of electrons of covalent bond by two dots – Structural formula – represents covalent bonds by dashes and shows arrangement of covalently bonded atoms

43. Single Covalent Bonds Halogens form single covalent bonds in their diatomic molecules - Fluorine is one example:

44. Single Covalent Bonds Unshared pair- pair of valence electrons that is not shared between atoms - also known as a lone pair or a nonbonding pair

45. Double and Triple Covalent Bonds Atoms form double or triple covalent bonds if they can attain a noble gas structure by sharing two pairs or three pairs of electrons Double covalent bond - bond that involves two shared pairs of electrons Triple covalent bond- bond formed by sharing three pairs of electrons

46. Double and Triple Covalent Bonds

47. Carbon dioxide is an example of a triatomic molecule:

48. Coordinate Covalent Bonds In carbon monoxide, oxygen has a stable configuration but the carbon does not:

49. Coordinate Covalent Bonds As shown below, the problem is solved if the oxygen donates one of its unshared pairs of electrons for bonding:

50. Coordinate Covalent Bonds Coordinate covalent bond- a covalent bond in which one atom contributes both bonding electrons In a structural formula, you can show coordinate covalent bonds as arrows that point from the atom donating the pair of electrons to the atom receiving them In a coordinate covalent bond, the shared electron pair comes from one of the bonding atoms

51. Coordinate Covalent Bonds Polyatomic ion - such as NH 4 +, is a tightly bound group of atoms that has a positive or negative charge and behaves as a unit Most plants need nitrogen that is already combined in a compound to grow:

52. Coordinate Covalent Bonds

53. Bond Dissociation Energies Bond dissociation energy - energy required to break the bond between two covalently bonded atoms A large bond dissociation energy corresponds to a strong covalent bond

54. Resonance Ozone in the upper atmosphere blocks harmful ultraviolet radiation from the sun Contributes to smog at lower elevations

55. Resonance The actual bonding of oxygen atoms in ozone is a hybrid, or mixture, of the extremes represented by the resonance forms Resonance structure - a structure that occurs when it is possible to draw two or more valid electron dot structures - have same number of electron pairs for a molecule or ion

56. Exceptions to the Octet Rule Octet rule cannot be satisfied in molecules whose total number of valence electrons is an odd number Also molecules in which an atom has fewer, or more, than a complete octet of valence electrons

57. Exceptions to the Octet Rule Two electron dot structures can be drawn for the NO 2 molecule:

58. Molecular Orbitals Molecular orbitals- orbitals that apply to the entire molecule - Just as an atomic orbital belongs to a particular atom - molecular orbital belongs to a molecule as a whole Bonding orbital- molecular orbital that can be occupied by two electrons of a covalent bond

59. Molecular Orbitals Sigma bond- when two atomic orbitals combine to form a molecular orbital that is symmetrical around the axis connecting two atomic nuclei

60. Molecular Orbitals pi bond - (  ) bonding electrons most likely found in sausage-shaped regions above and below the bond axis of bonded atoms - atomic orbitals overlap less than in sigma bonding = weaker than sigma bonds

61. VSEPR Theory tetrahedral angle- bond angle of 109.5˚ - results when central atom forms 4 bonds directed toward center of regular tetrahedron

62. VSEPR Theory VSEPR theory- valence-shell electron- pair repulsion theory - explains three-dimensional shape of methane -repulsion between electron pairs causes molecular shapes to adjust so that the valence-electron pairs stay as far apart as possible

63. VSEPR Theory Measured H—N—H bond angle is only 107°

64. VSEPR Theory Measured bond angle in water is about 105°

65. VSEPR Theory Carbon dioxide molecule is linear

66. VSEPR Theory Nine Possible Molecular Shapes

67. Hybrid Orbitals Orbital hybridization provides information about both molecular bonding and molecular shape –In hybridization, several atomic orbitals mix to form same total number of equivalent hybrid orbitals

68. Hybrid Orbitals Hybridization Involving Single Bonds

69. Hybrid Orbitals Hybridization Involving Double Bonds

70. Hybrid Orbitals Hybridization Involving Triple Bonds

71. Bond Polarity nonpolar covalent bond- when atoms in a bond pull equally - bonding electrons are shared equally

72. Bond Polarity polar covalent bond- (polar bond) is a covalent bond between atoms in which the electrons are shared unequally –The more electronegative atom attracts electrons more strongly and gains a slightly negative charge –The less electronegative atom has a slightly positive charge

73. Bond Polarity Chlorine atom attracts electron cloud more than the hydrogen atom does

74. Bond Polarity

75. Polar Molecules polar molecule- one end of the molecule is slightly negative and the other slightly positive Dipole- molecule that has two poles When polar molecules are placed between oppositely charged plates, they tend to become oriented with respect to the positive and negative plates

76. Polar Molecules A hydrogen chloride molecule is a dipole

77. Attractions Between Molecules Intermolecular attractions are weaker than either ionic or covalent bonds –Responsible for determining whether a molecular compound is a gas, liquid, solid at a given temperature

78. Attractions Between Molecules Van der Waals Forces –The two weakest attractions between molecules are collectively called van der Waals forces – named after the Dutch chemist Johannes van der Waals (1837–1923)

79. Attractions Between Molecules Dipole interactions- occur when polar molecules are attracted to one another

80. Attractions Between Molecules Dispersion forces- weakest of all molecular interactions –caused by the motion of electrons –Strength of dispersion forces generally increases as number of electrons in a molecule increases

81. Attractions Between Molecules –Hydrogen Bonds Hydrogen bonds- attractive forces in which a hydrogen covalently bonded to a very electronegative atom is also weakly bonded to an unshared electron pair of another electronegative atom

82. Attractions Between Molecules Hydrogen Bonding in Water

83. Attractions Between Molecules The relatively strong attractive forces between water molecules cause water to form small drops on a waxy surface

84. Intermolecular Attractions and Molecular Properties Network solids - (or network crystals) solids in which all atoms are covalently bonded to each other Network solids consist of molecules that do not melt until the temperature reaches 1000°C or higher, or they decompose without melting at all – Melting a network solid would require breaking covalent bonds throughout the solid

85. Intermolecular Attractions and Molecular Properties Diamond is an example of a network solid – Vaporizes to a gas at 3500°C or above

86. Intermolecular Attractions and Molecular Properties Silicon Carbide is a network solid - Has a melting point of about 2700°C

87. Intermolecular Attractions and Molecular Properties

88. Chemical Names and Formulas

89. Monatomic Ions Monatomic ions- consist of a single atom with a positive or negative charge resulting from the loss or gain of one or more valence electrons, respectively When metals in Groups 1A, 2A, and 3A lose electrons, they form cations with positive charges equal to their group number Names of cations of Group 1A, Group 2A, and Group 3A metals are the same as the name of the metal, followed by “ion” or “cation”

90. Monatomic Ions These elements have ionic charges that can be obtained from their group numbers

91. Monatomic Ions –Anions –The charge of any ion of a Group A nonmetal is determined by subtracting 8 from the group number – Start with stem of the element name and end in “–ide”

92. Monatomic Ions These Group A elements form anions

93. Monatomic Ions

94. -Charges of the cations of many transition metal ions must be determined from number of electrons lost –These colorful solutions contain the transition metal ions Co 3+, Cr 3+, Fe 3+, Ni 2+, and Mn 2+ –Many transition metals used as pigments

95. Monatomic Ions Two methods used to name the ions of transition metals: –The Stock system –The classical method Roman numeral in parentheses is placed after name of element to indicate numerical value of the charge (stock) Classical name of element is used to form the root name for the element (classical)

96. Monatomic Ions

97. Polyatomic Ions Polyatomic ions- ions composed of more than one atom –Names of most polyatomic anions end in “–ite” or “-ate”

98. Polyatomic Ions Names and Formulas of Some Common Polyatomic Ions

99. Binary Ionic Compounds Naming Binary Ionic Compounds: –Binary compound- composed of two elements and can be either ionic or molecular Place cation name first followed by anion name when naming binary compounds

100. Binary Ionic Compounds Writing Formulas for Binary Ionic Compounds: Write symbol of cation and then anion Add subscripts needed to balance charges Example: potassium nitride K + and N 3- K3NK3N

101. Compounds With Polyatomic Ions Compounds with Polyatomic Ions: –Write the symbol for the cation followed by the formula for the polyatomic ion and balance the charges Example: calcium nitrate is composed of a calcium cation (Ca 2+ ) and a polyatomic nitrate anion (NO 3 – ) In calcium nitrate, two nitrate anions, each with a 1– charge, are needed to balance the 2+ charge of each calcium cation calcium nitrate = Ca(NO 3 ) 2

102. Prefix in the name of a binary molecular compound tells how many atoms of an element are present in each molecule of the compound Naming Binary Molecular Compounds

103. Some guidelines for naming binary molecular compounds: Name the elements in the order listed in the formula Use prefixes to indicate the number of each kind of atom Omit the prefix mono- when the formula contains only one atom of the first element in the name The suffix of the name of the second element is “–ide” Naming Binary Molecular Compounds

104. Writing Formulas for Binary Molecular Compounds How do you write the formula for a binary molecular compound? –Use prefixes in the name to tell you the subscript of each element in the formula –Write correct symbols for the two elements with the appropriate subscripts

105. Silicon carbide is a hard material like diamond SiC Writing Formulas for Binary Molecular Compounds

106. acid- a compound that contains one or more hydrogen atoms – produces hydrogen ions (H + ) when dissolved in water Naming Acids

107. Three rules can help you name an acid with the general formula H n X –When the name of the anion (X) ends in - ide, the acid name begins with the prefix hydro- The stem of the anion has the suffix -ic and is followed by the word acid Naming Acids

108. When the anion name ends in -ite, the acid name is the stem of the anion with the suffix -ous, followed by the word acid When the anion name ends in -ate, the acid name is the stem of the anion with the suffix -ic followed by the word acid

109. Writing Formulas for Acids Use reverse process to write acid formulas

110. Base – ionic compound that produces hydroxide ions when dissolved in water Bases are named in the same way as other ionic compounds - name of cation is followed by name of anion Example: aluminum hydroxide consists of the aluminum cation (Al 3+ ) and the hydroxide anion (OH – ) Al(OH) 3 Names and Formulas for Bases

111. Sodium hydroxide (NaOH) is a base that is used to make paper Names and Formulas for Bases

112. Rules for naming and writing formulas for compounds are possible only because compounds form from elements in predictable ways –Summed up in two laws: the law of definite proportions the law of multiple proportions The Laws of Definite and Multiple Proportions

113. –The Law of Definite Proportions law of definite proportions- states that in samples of any chemical compound, the masses of the elements are always in the same proportions The Laws of Definite and Multiple Proportions

114. In every sample of water, the mass ratio of oxygen to hydrogen is always 8:1 The Laws of Definite and Multiple Proportions

115. Mass ratio of oxygen to hydrogen is always 16:1 in Hydrogen peroxide The Laws of Definite and Multiple Proportions

116. The Law of Multiple Proportions –law of multiple proportions- When same two elements form more than one compound, the different masses of one element that combine with the same mass of the other element are in small whole number ratios The Laws of Definite and Multiple Proportions

117. A Diagram of the Law of Multiple Proportions: The Laws of Definite and Multiple Proportions

118. Calculating Mass Ratios Example: Carbon reacts with oxygen to form compound A and B. Compound A contains 2.41 g C for each 3.22 g O. Compound B contains 6.71 g C for each 17.9 g O. What is the lowest whole number mass ratio of carbon that combines with a given mass of oxygen?

119. Calculating Mass Ratios Compound A = 2.41 g C and 3.22 g O Compound B = 6.71 g C and 17.9 g O Compound A = 2.41g C = 0.748 g C 3.22 g O 1.00 g O Compound B = 6.71 g C = 0.375 g C 17.9 g O 1.00 g O

120. Calculating Mass Ratios 0.748 g C (in compound A) = 1.99 = 2 0.375 g C (in compound B) 1 1 = 2:1

121. Practicing Skills: Naming Chemical Compounds This flowchart will help you name chemical compounds. Begin with the letters Q and R in the general formula Qx Ry. Q and R can be atoms, monatomic ions, or polyatomic ions.

122. CuSO 4 is an example from the flowchart. The compound will end in -ite or -ate. Cu is not part of Group A, so you must name the ions and use a Roman numeral to identify the charge of the transition metal. The name is copper(II) sulfate. Practicing Skills: Naming Chemical Compounds

123. Writing a chemical formula from a chemical name guidelines: –An -ide ending generally indicates a binary compound –An -ite or -ate ending means a polyatomic ion that includes oxygen is in the formula –Prefixes in a name generally indicate that the compound is molecular –A Roman numeral after the name of a cation shows the ionic charge of the cation Practicing Skills: Naming Chemical Compounds