Chemistry: Atoms First Third Edition Julia Burdge and Jason Overby Chapter 2 Atoms and the Periodic Table Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Chapter 2 Atoms and the Periodic Table 2.1 Atoms First 2.2 Subatomic Particles and Atomic Structure Discovery of the Electron Radioactivity The Proton and the Nuclear Model of the Atom The Neutron 2.3 Atomic Number, Mass Number, and Isotopes 2.4 Nuclear Stability Patterns of Nuclear Stability 2.5 Average Atomic Mass 2.6 The Periodic Table 2.7 The Mole and Molar Mass The Mole Molar Mass Interconverting Mass, Moles, and Numbers of Atoms
2.1 Atoms First An atom is the smallest quantity of matter that still retains the properties of matter. An element is a substance that cannot be broken down into two or more simpler substances by any means. Ø Examples: gold, oxygen, helium A DVD collection can be separated into smaller numbers until you have just one DVD left. But a single DVD cannot be separated into smaller pieces that are still DVDs. Atoms can also be divided smaller and smaller until eventually only a single atom remains. Dividing it any smaller would give pieces that are no longer an atom.
Atoms First Dalton said that atoms, of which all matter consists, are tiny, indivisible particles. Once a single atom has been obtained, dividing it smaller produces subatomic particles. The nature, number, and arrangement of subatomic particles determine the properties of atoms, which in turn determine the properties of all things material. We will start by examining the structure of atoms and the tiny subatomic particles that atoms contain.
2.2 Subatomic Particles and Atomic Structure In the late 1800s, many scientists were doing research involving radiation, the emission and transmission of energy in the form of waves. They commonly used a cathode ray tube, which consists of two metal plates sealed inside a glass tube from which most of the air has been evacuated.
Subatomic Particles and Atomic Structure (2) When metal plates are connected to a high-voltage source, the negatively charged plate, or cathode, emits an invisible ray. The cathode ray is drawn to the anode where it passes through a small hole. Although invisible, the path is revealed when the ray strikes a phosphor-coated surface producing a bright light.
Subatomic Particles and Atomic Structure (3) Researches discovered that like charges repel each other, and opposite charges attract one another. J. J. Thomson (1856 1940) noted the rays were repelled by a plate bearing a negative charge, and attracted to a plate bearing a positive charge.
Subatomic Particles and Atomic Structure (4) Thomson s contributions: He proposed the rays were actually a stream of negatively charged particles. These negatively charged particles are called electrons. By varying the electric field and measuring the degree of deflection of cathode rays, Thomson determined the chargeto-mass ratio of electrons to be: 1.76 10 8 C/g. (C is coulomb, the derived SI unit of electric charge.)
Subatomic Particles and Atomic Structure (5) R. A. Millikan (1868 1953) determined the charge on an electron by examining the motion of tiny oil drops. The charge was determined to be 1.6022 10 19 C.
Subatomic Particles and Atomic Structure (6) Knowing the charge, he was then able to use Thomson s chargeto-mass ratio to determine the mass of an electron. -19 charge - 1.6022 10 C -28 mass of an electron = = = 9.10 10 g 8 charge / mass - 1.76 10 C/g
Subatomic Particles and Atomic Structure (7) Wilhelm Rontgen (1845 1923) discovered X-rays. They were not deflected by magnetic or electric fields, so they could not consist of charged particles. Antoine Becquerel (1852 1908) discovered radioactivity, the spontaneous emission of radiation. Radioactive substances, such as uranium, can produce three types of radiation.
Subatomic Particles and Atomic Structure (8) Alpha (α) rays consist of positively charged particles, called α particles. Beta (β) rays, or β particles, are electrons so they are deflected away from the negatively charged plate. Gamma (γ) rays, like X-rays, have no charge and are unaffected by external electric or magnetic fields.
Subatomic Particles and Atomic Structure (9) Ernest Rutherford used α particles to prove the structure of atoms. The majority of particles penetrated the gold foil undeflected. Sometimes, α particles were deflected at a large angle. Sometimes, α particles bounced back in the direction from which they had come.
Subatomic Particles and Atomic Structure (10) Rutherford proposed a new model for the atom, a nuclear model: Positive charge is concentrated in the nucleus. The nucleus accounts for most of an atom s mass and is an extremely dense central core within the atom. ØA typical atomic radius is about 100 pm ØA typical nucleus has a radius of about 5 10 3 pm Ø1 pm = 1 10 12 m
Subatomic Particles and Atomic Structure (11) Protons are positively charged particles found in the nucleus. Neutrons are electronically neutral particles found in the nucleus. Neutrons are slightly larger than protons. Electrons are negatively charged particles distributed around the nucleus. Table 2.1 Masses and Charges of Subatomic Particles Particle Mass (g) Mass (amu) Charge (C) Charge unit Electron* 9.10938 10 28 5.4858 10 4 1.6022 10 19 1 Proton 1.67262 10 24 1.0073 +1.6022 10 19 +1 Neutron 1.67493 10 24 1.0086 0 0 *More refined measurements have resulted in a small change to Millikan s original value.
2.3 Atomic Number, Mass Number, and Isotopes All atoms can be identified by the number of protons and neutrons they contain. The atomic number (Z) is the number of protons in the nucleus. ØAtoms are neutral, so it s also the number of electrons. ØProtons determine the identity of an element. Ø For example, nitrogen s atomic number is 7, so every nitrogen has 7 protons. The mass number (A) is the total number of protons and neutrons. ØProtons and neutrons are collectively referred to as nucleons.
Atomic Number, Mass Number, and Isotopes (2) Most elements have two or more isotopes, atoms that have the same atomic number (Z) but different mass numbers (A). Isotopes of the same element typically exhibit very similar chemical properties Øsame types of compounds and similar reactivities.
Worked Example 2.1 Determine the numbers of protons, neutrons, and electrons in each of the following species: (a) #$!" Cl, (b) #" Cl, (c) (! K, and (d) carbon-14. Strategy Recall the superscript denotes the mass number (A) and the subscript denotes the atomic number (Z). If no subscript is shown, the atomic number can be deduced from the elemental symbol or name. Atoms are neutral so the number of electrons equals the number of protons. #$ Solution (a) Z = 17, so 17 protons A = 35, so 35 17 = 18 neutrons # of electrons = # of protons, so 17 electrons (b) Element is chlorine again, so Z must be 17; 17 protons A = 37, so 37 17 = 20 neutrons 17 protons, so 17 electrons
Worked Example 2.1 (cont. 1) Determine the numbers of protons, neutrons, and electrons in each of the following species: (a) #$!" Cl, (b) #" Cl, (c) (! K, and (d) carbon-14. Strategy Recall the superscript denotes the mass number (A) and the subscript denotes the atomic number (Z). If no subscript is shown, the atomic number can be deduced from the elemental symbol or name. Atoms are neutral so the number of electrons equals the number of protons. Solution (c) Potassium s atomic number is 19, so 19 protons A = 41, so 41 19 = 22 neutrons # of electrons = # of protons, so 19 electrons
Worked Example 2.1 (cont. 2) Determine the numbers of protons, neutrons, and electrons in each of the following species: (a) #$!" Cl, (b) #" Cl, (c) (! K, and (d) carbon-14. Strategy Recall the superscript denotes the mass number (A) and the subscript denotes the atomic number (Z). If no subscript is shown, the atomic number can be deduced from the elemental symbol or name. Atoms are neutral so the number of electrons equals the number of protons. Solution (d) Carbon-14 can also be represented as 14 C Carbon s atomic number is 6, so 6 protons A = 14, so 14 6 = 8 neutrons 6 protons, so 6 electrons Think About It Verify that the number of protons and the number of neutrons for each example sum to the mass number that is given. In part (a), there are 17 protons and 18 neutrons, which sum to give a mass number of 35, the value given in the problem. In part (b), 17 protons + 20 neutrons = 37. In part (c), 19 protons + 22 neutrons = 41. In part (d), 6 protons + 8 neutrons = 14.
2.4 Nuclear Stability The nucleus is a small portion of the total volume of an atom ØThe nucleus contains most of the atom s mass ØStability of the atomic nucleus can be related to density Example of a nucleus with 30 protons and 30 neutrons r = (5 10 density = -3 æ 1 10 pm) ç è 1 pm mass volume -12 1 10 = 4 pr 3 m öæ100 cm ö ç = 1 m øè ø -22 3 g = 1 10 4 p 3 5 10-22 g -13-13 ( 5 10 cm) cm 3 = 2 10 14 g / cm 3 ØThe highest know element density is 22.6 g/cm 3 ØThe average atomic nucleus is roughly 9 10 12 (or 9 trillion) times as dense as the densest element known!
Nuclear Stability (2) Principle factor for nuclear stability is neutron-to-proton ratio (n/p) Ø Ø Ø Ø There are more stabile nuclei with 2, 8, 20, 50, 82, or 126 protons or neutrons More with even # s All with atomic number > 83 are radioactive All isotopes of Tc and Pm are radioactive Table 2.2 Number of Stable Isotopes with Even and Odd Numbers of Proton and Neutrons Protons Neutrons Number of stable isotopes Odd Odd 4 Odd Even 50 Even Odd 53 Even Even 164
Nuclear Stability (3) The figure shows the number of neutrons versus the number of protons in various isotopes. Ø Stable nuclei are located in an area of the graph known as the belt of stability Ø Most radioactive nuclei lie outside the belt Ø Above the belt of stability, the nuclei have higher neutron-toproton ratio
2.5 Average Atomic Mass Atomic mass is the mass of an atom in atomic mass units (amu). 1 amu = 1/12 the mass of a carbon-12 atom The average atomic mass on the periodic table represents the average mass of the naturally occurring mixture of isotopes. Isotope Isotopic mass (amu) Natural abundance (%) 12 C 12.00000 98.93 13 C 13.003355 1.07 Average mass (C) = (0.9893)(12.00000 amu) + (0.0107)(13.003355 amu) = 12.01 amu
Average Atomic Mass (2) Measuring Atomic Mass Ø The most direct and most accurate method for determining atomic and molecular masses is mass spectrometry, using a mass spectrometer. Ø The mass spectrum of neon is shown here.
Worked Example 2.2 Oxygen is the most abundant element in both Earth s crust and the human body. The atomic masses of its three stable isotopes, "#! O (99.757 percent), "%! O (0.038 percent), "!! O (0.205 percent), are 15.9949, 16.9991, and 17.9992 amu, respectively. Calculate the average atomic mass of oxygen using the relative abundances given in parentheses. Strategy Each isotope contributes to the average atomic mass based on its relative abundance. Multiplying the mass of each isotope by its fractional abundance (percent value divided by 100) will give its contribution to the average atomic mass. Solution (0.99757)(15.9949 amu) + (0.00038)(16.9991 amu) + (0.00205)(17.992 amu) = 15.9994 amu Think About It The average atomic mass should be closest to the atomic mass of the most abundant isotope (in this case, oxygen-16) and, to four significant figures, should be the same number that appears in the periodic table on the inside front cover of your textbook (in this case, 16.00 amu).
2.6 The Periodic Table The periodic table is a chart in which elements having similar chemical and physical properties are grouped together.
The Periodic Table (2) Elements are arranged in periods, horizontal rows, in order of increasing atomic number.
The Periodic Table (3) Elements can be categorized as metals, nonmetals, or metalloids. Metals are good conductors of heat and electricity. Nonmetals are poor conductors of heat or electricity. Metalloids have intermediate properties.
The Periodic Table (4) A vertical column is known as a group.
The Periodic Table (5) Group 1A elements (Li, Na, K, Rb, Cs, Fr) are called alkali metals.
The Periodic Table (6) Group 2A elements (Be, Mg, Ca, Sr, Ba, Ra) are called alkaline earth metals.
The Periodic Table (7) Group 6A elements (O, S, Se, Te, Po) are called chalcogens.
The Periodic Table (8) Group 7A elements (F, Cl, Br, I, At) are called halogens.
The Periodic Table (9) Group 8A elements (He, Ne, Ar, Kr, Xe, Rn) are called the noble gases.
The Periodic Table (10) Groups 1B and 3B 8B are called the transition elements or transition metals.
2.7 The Mole and Molar Mass The mole is defined as the amount of a substance that contains as many elementary entities as there are atoms in exactly 12 g of carbon-12. This experimentally determined number is called Avogadro s number (N A ). N A = 6.0221415 x 10 23 We normally round this to 6.022 10 23. 1 mole = 6.022 10 23, just like 1 dozen = 12 or 1 gross = 144.
The Mole One mole each of some familiar substances:
Worked Example 2.3 Calcium is the most abundant metal in the human body. A typical human body contains roughly 30 moles of calcium. Determine (a) the number of Ca atoms in 30.00 moles of calcium and (b) the number of moles of calcium in a sample containing 1.00 10 20 Ca atoms. Strategy Use Avogadro s constant, 1 mole = 6.022 10 23, to convert from moles to atoms and from atoms to moles. Solution (a) 30.00 mol Ca!.#$$ &#'( )* *+,-. & -,/ )* = 1.807 10 $5 Ca atoms (b) 1.00 10 20 Ca atoms & -,/ )*!.#$$ &# '( )* *+,-. = 1.66 10<= mol Ca Think About It Make sure that units cancel properly in each solution and that the result makes sense. In part (a), for example, the number of moles (30) is greater than one, so the number of atoms is greater than Avogadro s number. In part (b), the number of atoms (1 10 20 ) is less than Avogadro s number, so there is less than a mole of substance.
Molar Mass The molar mass of a substance is the mass in grams of one mole of the substance. By definition, the mass of a mole of carbon-12 is exactly 12 g. Ø Mass of 1 carbon-12 atom: exactly 12 amu Ø Mass of 1 mole of carbon-12: exactly 12 g Although molar mass specifies the mass of one mole, making the units (g), we usually express molar masses in units of grams per mole (g/mol) to facilitate cancellation of units in calculations.
Worked Example 2.4 Determine (a) the number of moles of C in 25.00 g of carbon, (b) the number of moles of He in 10.50 g of helium, and (c) the number of moles of Na in 15.75 g of sodium. Strategy Molar mass of an element is numerically equal to its average atomic mass. Use the molar mass for each element to convert from mass to moles. Setup (a) The molar mass of carbon is 12.01 g/mol. (b) The molar mass of helium is 4.003 g/mol. (c) The molar mass of sodium is 22.99 g/mol. Solution (a) 25.00 g C! "#$ %!&.(! ) % = 2.082 mol C (b) 10.50 g He! "#$ *+,.((- ) *+ = 2.623 mol He (c) 15.75 g Na! "#$./ &&.00 )./ = 0.6851 mol Na
Worked Example 2.4 (cont.) Think About It Always double-check unit cancellations, in problems such as these errors are common when molar mass is used as a conversion factor. Also make sure that the results make sense. For example, in the case of part (c), a mass smaller than the molar mass corresponds to less than a mole.
Interconverting Mass, Moles, and Number of Atoms Molar mass is the conversion factor from mass to moles, and vice versa. Avogadro s constant converts from moles to atoms.
Worked Example 2.5 Determine (a) the number of C atoms in 0.515 g of carbon, and (b) the mass of helium that contains 6.89 10 18 He atoms. Strategy Use the conversions depicted in the previous slide to convert (a) from grams to moles to atoms and (b) from atoms to moles to grams. Setup (a) The molar mass of carbon is 12.01 g/mol. (b) The molar mass of helium is 4.003 g/mol. N A = 6.022 10 23 Solution (a) 0.515 g C! "#$ % *.(&&!(,- %./#"0!&.(! ) %! "#$ % = 2.58 10 22 C atoms (b) 6.89 10 18 He atoms! "#$ 12 ) 12 *.(&&!(,- 3.((4 12./#"0! "#$ 12 = 4.58 10 5 g He
Worked Example 2.5 (cont.) Think About It A ballpark estimate can help you prevent common errors. For example, the mass in part (a) is smaller than the molar mass of carbon. Therefore, you should expect a number of atoms smaller than Avogadro s number. Likewise, the number of atoms in part (b) is smaller than Avogadro s number. Therefore, you should expect a mass of helium smaller than the molar mass of helium.
2 Chapter Summary: Key Points Atoms Elements Discovery of the Electron Radioactivity The Proton and the Nucleus Nuclear Model of the Atom The Neutron Atomic Number Mass Number Isotopes Nuclear Stability Average Atomic Mass The Periodic Table The Mole Avogadro s Number Molar Mass