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Unit 5: Atoms and the Periodic Table Chapter 4 Part 1: Atomic Structure Objectives Explain Dalton’s atomic theory, and describe why it was more successful than Democritus’s theory. State the charge, mass, and location of each part of an atom according to the modern model of the atom. Compare and contrast Bohr’s model with the modern model of the atom. Atomic Models Atom comes from the Greek word that means “unable to be divided.” Democritus (4 b.c.) - came up with the first theory of atomic structure; said that the universe was made of invisible units called atoms, but was unable to provide evidence. QuickTi me™ and a TIFF ( Uncompressed) decompressor are needed to see thi s pi ctur e. John Dalton (1808) - used scientific research to claim that atoms could not be divided, that all atoms of an element were exactly alike, and that atoms of different elements could join to form compounds. QuickTime™ and a TIF F (Uncompressed) decompressor are needed to see this pict ure. Neils Bohr (1913) - said that the electrons in an atom orbit around the nucleus like planets around the sun. The path of an electron is determined by how much energy it has, which puts it in a specific energy level. QuickTi me™ and a TIFF ( Uncompressed) decompressor are needed to see thi s pi ctur e. Modern wave model (1925) - says that electrons behave more like waves on a vibrating string, and move back and forth between energy levels. Thus, an electron’s exact location at any given moment cannot be determined. QuickTime™ and a TIFF (Uncompressed) decompressor are needed to see this picture. An Atom’s Contents Atoms contain smaller pieces called subatomic particles. PARTICLE CHARGE LOCATION MASS proton positive in nucleus 1 amu electron negative around nucleus almost 0 amu neutron neutral in nucleus 1 amu Unreacted atoms (atoms that are not part of a chemical compound) have no overall charge. That means the number of positive charges (protons) must equal the number of negative charges (electrons). Neutrons do not affect the overall charge. Electrons and Energy Levels If an electron does not have much energy, it will be closer to the nucleus in the first or second energy level. The first energy level will hold only 2 electrons. If an atom has more than 2 electrons, the first two will fill the first level, and the rest will begin filling the second energy level. The second level will hold 8 electrons. If the first and second levels are full and there are still leftover electrons, they will go to the third level, which holds 18 electrons. Once the first three levels are full (that’s 28 electrons!), electrons must go to the fourth level, which holds 32 electrons. QuickTi me™ and a TIFF ( Uncompressed) decompr essor are needed to see thi s picture. Orbitals There are four different types of orbitals that can be found within the energy levels. They are s, p, d, and f. An s orbital has the lowest amount of energy and can hold 2 electrons. A p orbital has more energy than s orbitals, and there are 3 of them. Each one can hold 2 electrons, for a total of 6. There are 5 d orbitals and 7 f orbitals. F orbitals have the most energy, and each one holds 2 electrons. QuickTime™ and a TIFF (Uncompressed) decompressor are needed to see this picture. Valence Electrons Electrons found in the outermost energy level of an atom. Each atom contains between 1 and 8 valence electrons. These electrons are the ones that are used to form chemical bonds with other atoms to form molecules or compounds. Part 2: The Periodic Table Objectives Relate the organization of the periodic table to the arrangement of electrons within an atom. Explain why some atoms gain or lost electrons to form ions. Determine how many protons, neutrons, and electrons at atoms has given its symbol, atomic number, and mass number. Describe how the abundance of isotopes affects an element’s average atomic mass. Locate alkali metals, alkaline-earth metals, and transition metals in the periodic table. Locate semiconductors, halogens, and noble gases in the periodic table. Relate an element’s chemical properties to the electron arrangement of its atoms. The Periodic Table The periodic law says that if the elements are arranged in a specific order, similarities in their properties will occur in a regular pattern. Period - horizontal row; the number of protons (and therefore, the number of electrons) increases by one as you move from left to right. Group (or family) - vertical column; elements in the same group have the same number of valence electrons, so they have similar characteristics. The Periodic Table QuickTime™ and a TIFF (Uncompressed) decompressor are needed to see this picture. Groups Group 1 - Alkali metals (1 v.e.) Group 2 - Alkaline-earth metals (2 v.e.) Groups 3 -12 - Transition metals (number of v.e. varies) Group 13 - Boron family (3 v.e.) Group 14 - Carbon family (4 v.e.) Group 15 - Nitrogen family (5 v.e.) Group 16 - Oxygen family (6 v.e.) Group 17 - Halogens (7 v.e.) Group 18 - Noble (or inert) gases (8 v.e.) If an atom has only 1 v.e., it will be very reactive because it will want to stabilize itself by giving away its v.e. The goal of an atom is to become stable by having a totally full or totally empty outer energy level. This will cause it to be an ion with a +1 charge. Likewise, an atom with 7 v.e. will be very reactive because it only needs 1 v.e. to be stable. Where can it find 1 v.e.? Calculating P, N and E Atomic number - the number of protons in an atom (and thus, the number of electrons) Mass number - the number of protons plus neutrons in an atom Isotope - atoms of an element that have the same number of protons, but different number of neutrons. This will not affect the number of electrons, but will affect the mass. Average atomic mass - the average mass of all the isotopes of an element To calculate the number of neutrons in an atom: Mass number - atomic number = neutrons Mass is measured in atomic mass units (amu). 1 amu = 1/12 the mass of a standard C12 atom Part 3: Using Moles to Count Atoms Objectives Explain the relationship between a mole of a substance and Avogadro’s constant. Find the molar mass of an element by using the periodic table. Solve problems converting the amount of an element in moles to its mass in grams, and vice versa. Using Moles to Count Atoms We use moles to count atoms because they are so small and so numerous. If we know the mass of the atom, we can estimate how many atoms are in a sample of a substance by counting them in groups. The mole has a value of 6.022 x 1023 particles in exactly 1 mole of substance. This value is called Avogadro’s constant. Molar Mass The mass in grams of 1 mole of a substance is its molar mass, which is the same as its average atomic mass on the periodic table. Using Conversion Factors A fraction that equals 1. For example: 12 in/1 ft = 1; 5280 ft/ 1 mi = 1; 365.25 d/1 yr = 1 Example: If you have 5.50 mol of Fe, and Fe has a molar mass of 55.85 g/mol, what is its mass in g? 5.50 mol Fe x 55.85 g Fe= 307 g Fe 1 mol Fe Examples If you have 2.50 mol of S, and S has a molar mass of 32.07 g/mol, what is its mass in g? 2.50 mol S x 32.07 g S = 80.18 g S 1 mol S A) 1.80 mol Ca B) 0.50 mol C C) 3.20 mol Cu A) 72.14 g Ca B) 6.01 g C C) 203.36 g Cu How many moles are in 620 g of Hg? 620 g Hg x 1 mol Hg = 3.09 mol Hg 200.59 g Hg A) 352 g Fe B) 11 g Si C) 205 g He A) 6.30 mol He B) 0.39 mol Si C) 51.25 mol He