2.1 Atomic Structure and Subatomic Particles (p. 40) There are two types of charge. 1. _ 2. Electrons In 1897 J. J. Thomson, using a cathode ray tube, measured the ratio of the charge-to-mass of the electron: -5.6 x 10-9 grams/coulomb. (The coulomb, C, is the SI unit of charge.) Thomson is credited with the discovery of the electron. George Stoney coined the term electron in 1891. Its origin is the Greek word for amber, ήλεκτρον. In 1909, R. A. Millikan, using his oil droplet experiment, determined the unit charge of an electron to be -1.60 x 10-19 C. 2.2 The Nuclear Atom (p. 43) In 1911, Ernest Rutherford, using his gold foil experiment, proposed that the atom has a dense, positively charged core which he called the nucleus. The electrons were envisioned to be someehow located in the space around the nucleus. The term proton seems to have been coined by Rutherford, and first appears in print in 1920. It was at this time that he proposed the possible existence of this type of particle within the atom. Neutrons (p. 43) In 1932, James Chadwick carried out experiments that detected neutral particles within the atom. (The American chemist H.D. Harkins, coined the term "neutron" in 1921.) SUMMARY (p. 44) The atom is % empty space and is composed of three particles. a._ b._ c._ The atom is electrically. This means that the number of is equal to the number of. The electrons, composing less than % of the atom s mass, _ the nucleus. In some atoms, the electrons are _ to move from one atom to another. Protons and neutrons are located in the. The mass of the atom comes predominantly from the _. The nucleus contributes less than _% of the size of the atom, but % of the mass. The electrons occupy most of the volume of the atom, but contribute very little mass. Protons When atoms lose electrons, they develop a charge. The mass of the proton is 1.67262158 x 10-24 g. The charge of the proton is but in sign to that of the electron. 1
Particle Symbol Relative Charge Charge (C) Mass (kg) Relative Mass Mass (amu) Electron -1.60 x 10-19 9.11 x 10-31 0.0005485799 Proton +1.60 x 10-19 1.673x 10-27 1.007276 Neutron - - - - - - - - - - 1.675x 10-27 1.00866490 2.3 The Sizes of Atoms and Units Used to Represent Them (p. 44) Mass is the quantity of matter in an object. The SI unit is the (), a unit about 2.2 pounds. The SI unit of length is the (), a unit about 10% longer than a yard. 1 Angstrom (Å) = 10-10 m Volume is a derived unit. 1 L = dm 3 = m 3 or _ cm 3 = _ ml The SI for time unit is the second. The SI unit for amount of a substance is the (). 1 mole = The student must be able to carry out calculations based on Table 2.1 (p. 45). Convert each of the following so that the power of ten is replaced by a prefix. 1. 3.88 x 10-2 g =? (3.88 cg) 2. 1.72 x 10-9 s =? (1.72 ns) 3. 8.06 x 10 3 L =? (8.06 kl) 4. 6.95 x 10-3 mol =? (6.95 mmol) 2
Use exponential notation to express each of the following in an appropriate SI base unit. 1. 2.33 µg =? (2.33 x 10-9 kg) 2. 4691 ns =? (4.691 x 10-6 s) 3. 55.2 cmol =? (5.52 x 10-1 mol) SI Unit conversions 4 000 m =? km 76,000 kg =? Mg 1,500 mg =? g 2,300 µs =? ms 500 ml =? L 12 pm =? mm 2.4 Uncertainty and Significant Figures (p. 50) There is uncertainty in all measurements. The digits recorded from a measurement will include both the certain and the uncertain ones. These digits are called significant figures. In general, the greater the number of significant figures in a measurement, the greater is the certainty of the measurement. Significant Figures 1. All Non-zero digits are significant. 2. Zeros may or may not be significant. a. Leading zeros (zeros to the left of a digit) are not significant. b. Captive zeros (zeros between non-zero digits) are significant. c. Trailing or terminal zeros (zeros to right of number) are significant only if the number contains a decimal point, otherwise they are not. 3. Exact numbers contain an infinite number of significant figures. Rules for Significant Figures in Calculations 1. When multiplying or dividing, the answer has the same number of significant figures as the value in the operation with the least number of significant figures. 2. When adding and subtracting, the answer has the same number of decimal places as the value in the operation with the least number of decimal places. 3. Exact numbers have no bearing on the number of significant figures and are not considered when determining the number of significant figures in an answer. 3
Perform the following operations and give answers in the correct number of significant figures. A. 4.184 x 100.62 x (25.27-24.16) =? E. 0.1654 + 2.07-2.114 =? B. 8.925-8.904 x 100 =? 8.925 F. 8.27(4.987-4.962) =? C. (9.04-8.23 + 21.954 + 81.0) / 3.1416 =? G. 9.5 + 4.1 + 2.8 + 3.175 =? 4 D. 9.2 x 100.65 =? 8.321 + 4.026 H. 9.025-9.024 x 100 =? 9.025 Answers: A. 467.; B. 0.24; C. 33.03; D. 75.; E. 0.12; F. 0.21; G. 4.9; H. 0.01 2.5 Atomic Numbers and Mass Numbers (p. 53) 2.6 Isotopes and Atomic Weight (p. 56) [Power Point presentation] 4
2.5 Atomic Numbers and Mass Numbers The atomic number (Z) = the on the nucleus = the number of in the nucleus = the number of _ in the atom Atoms of the same element have the same number of in the nucleus. The mass number (A) = the sum of the number of protons and in the nucleus A Z = number of 21 11Na protons, electrons, and neutrons Isotopes are atoms whose nuclei have the same atomic number but different mass numbers; that is the same number of protons but different number of neutrons. Species 23 Na Protons Neutrons Lagniappe Atoms are electrically. Ions are _ particles. _ are positively charged ions. _ are negatively charged ions. 22 Na 35 Cl 37 Cl Na 1+ S 2- Z - charge = 2.6 Isotopes and Atomic Weight Atomic weight is related to the standard of carbon-12. By definition, an atom of carbon-12 has exactly 12 amu. Masses of all other atoms are relative to this scale. There are two naturally occurring isotopes of carbon on earth: 12 C and 13 C. 12 C has a mass of 12 amu 13 C has a mass of 13.0034 amu 1
The atomic weight (molar mass) given at the bottom of a symbol, is the weighted average of the masses of the naturally occurring isotopes of that element. The atomic weight is the sum of the isotope contributions. 12 C has a natural abundance of 98.892%, while the natural abundance of 13 C is 1.108%. 12 C has a mass of 12 amu 13 C has a mass of 13.0034 amu The atomic mass of carbon is: Isotope contribution = =( ) x ( ) Naturally occurring chlorine is a mixture of two naturally occurring isotopes: 35 Cl which has an abundance of 75.77% and an isotopic mass of 34.969 amu, and 37 Cl with an abundance of 24.23% and an isotopic mass of 36.966 amu. What is the atomic mass of chlorine? 2
2.7 Amounts of Substances The Mole (p. 59) Mole (mol) The SI unit for amount of substance is the mole. It is a counting quantity for particles such as atoms or molecules, or for other chemical entities. Thus, the amount of water in a beaker is the number of water molecules in the beaker and not the mass or the volume of water. 1 mole () = 602,213,673,600,000,000,000,000 = - A mole is the amount of substance in a sample that contains as many elementary entities as the number of atoms in exactly 12 g of C-12. - The mass of one mole in grams is numerically equal to the molar mass. The mole is so defined that a sample of an element with a mass equal to its atomic mass in grams is one mole of that element. How many years would it take for the world s population to spend a mole of dollars at a rate of $1000 per second? 2.8 Molar mass and Problem Solving (p. 60) 1. How many moles are in 0.100 g of silicon? [3.56 x 10-3 mol] 2. What is the mass of 5.00 moles of silver? [5.39 x 10 2 g] 3. How many atoms are in 5.23 x 10-4 mol of potassium? [3.15 x 10 20 ] 4. How many moles are in 9.83 x 10 22 atoms of mercury? [0.163 mol] 5. How many atoms are in 3.50 g of iron? [3.77 x 10 22 atoms] 6. What is the mass of 8.50 x 10 24 atoms of lead? [2.92 x 10 3 g] 7. How many electrons are in a 1.000 g sample of tungsten? [2.423x10 23 ] 5