1. Chapter 3 – Atoms and the Periodic Table

2. 3.1 Atomic Structure

3. Atoms Derived from the Greek word meaning “unable to divide”. They are the building blocks of molecules.

4. Every element is made of tiny, unique particles that cannot be subdivided.

5. Atoms of the same element are exactly the same. Atoms of different elements can join to form molecules.

6. Atoms are made of protons, neutrons, and electrons.

7. Nucleus The center of an atom; made up of protons and neutrons.

8. Proton A positively charged subatomic particle in the nucleus of the atom

9. Neutron A neutral subatomic particle in the nucleus of an atom.

10. Electron A tiny negatively charged subatomic particle moving around outside of an atom.

11. Atoms have no overall charge. Even though atoms have charged protons and electrons. They have an equal number of each. So, they cancel each other out.

12. Bohr’s Model In 1913, Neils Bohr suggested that electrons moves about a set path about the nucleus. The path defines the electron’s energy level.

13. Energy Level Any of the possible energies of an electron may have in an atom.

14. Modern theory states that electrons behave more like waves. In 1925, Bohr proposed a new model, that the electrons do not follow a set path.

15. Although we cannot know how the electron travels around the nucleus we can know where it spends the majority of its time (thus, we can know position but not trajectory). The “probability” of finding an electron around a nucleus can be calculated. Relative probability is indicated by a series of dots, indicating the “electron cloud”.

16. Electrons are found in orbitals within energy levels.

17. Orbital A region in an atom where there is a high probability of finding electrons.

18. Electrons may occupy four different orbitals.

19. s - orbital - Simplest orbital. - Can only have one orientation in space, because its shaped like a sphere.

20. -Its shape enables it to surround the nucleus. -90% electron probability/cloud for 1s orbital (notice higher probability toward the center) -It can hold a maximum of 2 electrons.

21. p - orbital -it is a dumbbell- shaped and can be oriented in 3 different ways in space. (3-Dimensions)

23. d - orbital -a more complex orbital. -There are a possible of 5 orientations.

24. -All orbitals are very different in shape, each can contain a maximum 2 electrons. -Can contains a total of 10 electrons in all.

25. Four of the d orbital’s resemble two dumbbells in a clover shape. The last d orbital resembles a p orbital with a donut wrapped around the middle.

26. f - orbital -a more complex orbital. -There are a possible of 7 orbitals.

27. -All orbitals are very different in shape, each can contain a maximum 2 electrons. -Can contains a total of 14 electrons in all.

28. Electrons usually occupy the lowest energy levels available.

29. And within each energy level, electrons occupy orbitals with the lowest energy.

30. An s orbital has the lowest energy. A p orbital has slightly more energy, followed by a d orbital. An f orbital has the greatest energy.

31. Every atom has one or more valence electron.

32. Valence Electron An electron in the outermost energy level of an atom.

33. 3.2 A Guided Tour of the Periodic Table.

35. Periodic Law Properties of elements tend to change in a regular pattern when elements are arranged in order of increasing atomic number, or number of protons in their atoms.

36. Period A horizontal row elements in the periodic table.

37. As you move across a row the # of protons increases by 1. As does the # of electrons.

38. Elements in the same group have similar properties.

39. Group (family) – a vertical column of elements in the periodic table

40. Elements in the same group have the same # of valence electrons.

41. These elements are not exactly alike, they have a different # of protons and electrons.

42. Some atoms form ions Some atoms may under go ionization.

43. Ionization The process of adding electrons to or removing electrons from an atom or groups of atoms. These are valence electrons.

44. Ions An atom or a group of atoms that has lost or gained 1 or more electrons and therefore has a net charge.

45. Cation An ion with a positive charge

46. Example: Lithium atom Lithium is so reactive, it reacts with air. It has 1 electron in the outer level of the s orbital. This one electron makes very reactive.

47. Removing this electron forms a positive ion (Li + ) Li + is less reactive, because now its outer energy level is full.

48. Anion An ion with a negative charge.

49. Example: Fluorine atom Fluorine is also very reactive; however it gains an electron to become less reactive.

50. It has 7 valence electrons and needs only 1 to complete its outer energy level. Therefore, easily gaining 1 electron & becoming a negative ion (F - ).

51. Atomic Structure of Atoms All atoms have different structure and different properties; therefore, different structure.

52. Atomic Number (Z) The number of protons in the nucleus of an atom.

53. Examples: Helium (He) has 2 protons Z = 2 Cesium (Cs) has 55 protons Z = 55

54. Mass Number (A) Mass number equals the # of protons and neutrons in the nucleus in the atom.

55. Ex: Fluorine (F) has 9 protons and 10 neutrons, A = 19 for fluorine.

56. Although atoms of the same element always have the same Atomic number, they can have a different mass numbers.

57. Isotopes Any atoms have the same # of protons, but a different # of neutrons.

58. Ex: Hydrogen has 3 isotopes. 1. Protium – has 1 proton in its nucleus. (A = 1) Most common form

59. 2. Deuterium – has 1 proton and 1 neutron in its nucleus, call heavy hydrogen. (A = 2)

60. 3. Tritium – has 1 proton and 2 neutrons in its nucleus. (A = 3)

61. Atomic numbers (Z) and mass numbers (A) maybe included with the chemical symbol.

62. U 235 92 Atomic Number Mass Number

63. Calculating the # of neutrons in an atom. If you know the Atomic number and the Mass number, all you have to do is subtract.

64. # of Neutrons = mass # – atomic # (# of Neutrons = A – Z )

65. Ex: For our Uranium - 235 # of Neutrons = A – Z # of Neutrons = 235 – 92 # of Neutrons = 143

66. The mass of an atom. The mass of a single atom is very small. A single Fluorine atom is one trillionth of one billionth of a gram.

67. Because it is very hard to work with atomic mass are expressed in atomic mass units (amu).

68. atomic mass unit (amu) A quantity equal to 1/12 of the mass of a Carbon-12 atom.

69. average atomic mass The weighted average of the masses of all naturally occurring isotopes of an element.

70. EX: If we look back at Hydrogen. There are 3 isotopes. The average atomic mass 1.00794 amu. This means there are more isotopes of Hydrogen-1, than Hydrogen-2 or Hydrogen- 3.

71. 3.3 Families of Elements

72. Elements are either classified as Metals or Nonmetals.

73. Metals The elements that are good conductors for heat and electricity.

74. Most elements are metals. Usually solids and shiny.

75. Nonmetals The elements that are usually poor conductors of heat and electricity.

76. Nonmetals, except Hydrogen, are found on the right side of the periodic table. Can be solids, liquids, or gases. Solids are dull and brittle.

77. There are some nonmetals that can conduct under certain circumstances.

78. Semiconductors The intermediate conductors of heat and electricity. Sometimes called metalloids.

79. Metals can be classified even further into 4 different groups.

80. Alkali Metals Highly reactive metallic elements located in Group 1 of the periodic table.

81. Has 1 valence electron. Can be easily removed to form a +1 ion. Highly reactive.

82. Look on your Periodic table. Lithium Sodium Potassium Rubidium Cesium Francium

83. Since they are so reactive, these elements are not found in nature. They combine with other elements to form compounds.

84. Alkaline Earth Metals The reactive metallic elements located in Group 2 of the Periodic Table.

85. They have 2 valence electrons. They are less reactive than alkali metals. May form an ion of a +2 charge if both valence electrons are removed.

86. Look at the Periodic Table Beryllium Magnesium CalciumStrontium BariumRadium

87. They are combine with other elements to form compounds. Ex: Calcium Compounds – shells of sea life, coral reefs, limestone, or marble.

88. Magnesium Compounds – speeds up the processes of the human body.

89. Transition Metals The metallic elements located in groups 3 – 12 of the periodic table.

90. Much less reactive than Alkali or Alkaline Metals. Can lose electrons to form + ions. Some metals can lose up to 4 electrons.

91. All Metals are good heat and electrical conductors. They can be stretched and shaped without breaking.

92. Mercury is the only metal that is a liquid at room temperature.

93. Technetium and Promethium are synthetic elements.

94. Man made elements that are radioactive. Technetium is used to help diagnose cancer and other medical problems.

95. Promethium is used in ‘glow in the dark’ paint. Elements with atomic # greater than 92 are man made as well.

96. Nonmetals Found on the right side of the periodic table. Some elements in groups13-16 and all in groups 17 & 18.

97. May gain electrons to form – ions

98. Carbon combines with other elements to form millions of carbon containing compounds. Carbon compounds are found in both living and nonliving things.

99. Nonmetals and there compounds are plentiful on Earth.

100. Halogens The highly reactive elements located in Group 17 of the periodic table.

101. The valence shell is almost full; therefore, these elements easily accept an electron. Creating a (-) ion, anion.

102. Fluorine Chlorine Bromine Iodine Astatine

103. Noble Gases The unreactive gases located in Group 18 of the Periodic Table.

104. These elements exist only as single atoms, instead of molecules. All gases are inert, non-reactive.

105. Outer level is full of electrons. They don’t form with other elemental atoms to form atoms.

106. HeliumNeon ArgonKrypton XenonRadon

107. Semiconductors Elements classified as nonmetals, each one has some properties of metals. Known as ‘metalloids’

108. They are able to conduct heat & electricity. Silicon most familiar, used for computers and other electronic devices.

109. Boron Silicon Germanium Arsenic Antimony Tellurium

110. 3.4 Using Moles to Count Atoms.

111. Counting Things One of the first things we do as child. When we count large #’s of small things, we use counting units.

112. The mole is useful for counting small particles.

113. Mole The SI base unit that describes the amount of a substance.

114. A mole is a collection of a very large number of particles. 602,213,670,000, 000,000,000,000

115. Avogadro’s Constant The # of particles in 1 mol; equals 6.022 x 10 23 /mol.

116. Molar Mass The mass in grams of 1 mol of a substance.

117. Ex: 1 mole of Titanium-22 atoms has a molar mass of 47.87 g

118. Think of it as 47.87 grams per mole of Titanium. 47.87 g/mol

119. Molar mass of an element in grams is the same as average atomic mass in amu.

120. Find the molar mass for the following elements: 1.Gold 2.Einsteinium 3.Gallium 4.Cesium

121. Find the molar mass for the following elements: 1.Gold196.97 g 2.Einsteinium252.08 g 3.Gallium 69.72 g 4.Cesium132.91 g

122. Conversion Factors A ratio equal to one that expresses the same quantity in two ways.

123. Can be anything: 1 ream of = 500 sheets paper of paper 1 inch = 2.54 cm 10 km = 6.2 miles

124. Converting Amount (mol) to Mass (g)

125. 1 st establish the amount of matter in moles from the problem. 2 nd look up the element of the period table and find its molar mass in g for conversion factor

126. Conversion factor will be: Atomic mass(g)=1 mole or 1 mole=Atomic mass(g)

127. Set up for factor label method. Given info

128. Set up for factor label method. Given atomic mass of info element

129. Set up for factor label method. Given atomic mass of info element 1mole of element

130. Ex: How many grams in 23 moles of Silver? 23 mol of Ag

131. Conversion Factor 1 mol of Ag = 107.8682 g of Ag

132. Ex: How many grams in 23 moles of Silver? 23 mol 107.8682 g of Ag of Ag

133. Ex: How many grams in 23 moles of Silver? 23 mol 107.8682 g of Ag of Ag 1 mol of Ag

134. 2480.96 g of Ag - Multiply #’s above the horizontal line - Units of moles of Ag will cancel leaving just grams of Ag.

135. Ex: How many grams in 2.8 moles of Radon? 2.8 mol of Rn

136. Conversion Factor 1 mol of Rn = 222.0176 g of Rn

137. Ex: How many grams in 2.8 moles of Radon? 2.8 mol 222.0176 g of Rn of Rn

138. Ex: How many grams in 2.8 moles of Radon? 2.8 mol 222.0176 g of Rn of Rn 1 mol of Rn

139. 621.64928 g of Rn

140. Converting Mass (g) to Amount (mol) Set up is the same except we switch/flip the conversion factor ratio.

141. Ex: Convert 729 g of Th to moles. 729 g of Th

142. Conversion Factor 1 mol of Th = 232.0381 g of Th

143. Ex: Convert 729 g of Th to moles. 729 g 1 mol of Th of Th

144. Ex: Convert 729 g of Th to moles. 729 g 1 mol of Th of Th 232.0381 g of Th

145. 3.142 mol of Th

146. Ex: Convert 50 g of Sr to moles. 50 g of Sr

147. Conversion Factor 1 mole of Sr = 87.62 g of Sr

148. Ex: Convert 50 g of Sr to moles. 50 g 1 mol of Sr of Sr

149. Ex: Convert 50 g of Sr to moles. 50 g 1 mol of Sr of Sr 87.62 g of Sr

150. 0.571 mol of Sr