Chemistry: Atoms First Julia Burdge & Jason Overby Chapter 2 Atoms and the Periodic Table Kent L. McCorkle Cosumnes River College Sacramento, CA Copyright (c) The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

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 Average Atomic Mass 2.5 The Periodic Table 2.6 The Mole and Molar Mass The Mole Molar Mass Interconverting Mass, Moles, and Numbers of Atoms

Atoms First 2.1 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. Examples: gold, oxygen, helium 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. 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 First 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.

2.2 Subatomic Particles and Atomic Structure In the late 1800’s, 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 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 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 This prompted him to propose 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 charge-to-mass ratio of electrons to be 1.76×108 C/g. (C is coulomb, the derived SI unit of electric charge.)

Subatomic Particles and Atomic Structure 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 Knowing the charge, he was then able to use Thomson’s charge-to-mass ratio to determine the mass of an electron.

Subatomic Particles and Atomic Structure 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 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 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 Rutherford proposed a new model for the atom: 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. A typical atomic radius is about 100 pm A typical nucleus has a radius of about 5×10–3 pm 1 pm = 1×10–12 m

Subatomic Particles and Atomic Structure Protons are positively charged particles found in the nucleus. Neutrons are electronically neutral particles found in the nucleus. Neutrons are slightly larger than protons.

(number of protons + neutrons) Atomic Number, Mass Number, and Isotopes 2.3 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. Atoms are neutral, so it’s also the number of electrons. Protons determine the identity of an element. 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. Protons and neutrons are collectively referred to as nucleons. Mass number (number of protons + neutrons) Elemental symbol Atomic number (number of protons)

Atomic Number, Mass Number, and Isotopes Most elements have two or more isotopes, atoms that have the same atomic number (Z) but different mass numbers (A). 1 proton 0 neutrons 1 proton 1 neutron 1 proton 2 neutrons Isotopes of the same element exhibit similar chemical properties, forming the same types of compounds and displaying 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. 35 17 37 41 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 Z = 17, so 17 protons A = 35, so 35 - 17 = 18 neutrons # of electrons = # of protons, so 17 electrons 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.) Determine the numbers of protons, neutrons, and electrons in each of the following species: (a) Cl, (b) Cl, (c) K, and (d) carbon-14. 35 17 37 41 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.) Determine the numbers of protons, neutrons, and electrons in each of the following species: (a) Cl, (b) Cl, (c) K, and (d) carbon-14. 35 17 37 41 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 C Carbon’s atomic number is 6, so 6 protons A = 14, so 14 - 6 = 8 neutrons 6 protons, so 6 electrons 14 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.

1 amu = 1/12 the mass of a carbon-12 atom 2.4 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. Average mass (C) = (0.9893)(12.00000 amu) + (0.0107)(13.003355 amu) Isotope Isotopic mass (amu) Natural abundance (%) 12C 12.00000 98.93 13C 13.003355 1.07 = 12.01 amu 21

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. 16 8 17 8 18 8 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).

2.5 The Periodic Table The periodic table is a chart in which elements having similar chemical and physical properties are grouped together.

The Periodic Table Elements are arranged in periods, horizontal rows, in order of increasing atomic number.

The Periodic Table 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 A vertical column is known as a group.

The Periodic Table Group 1A elements (Li, Na, K, Rb, Cs, Fr) are called alkali metals.

The Periodic Table Group 2A elements (Be, Mg, Ca, Sr, Ba, Ra) are called alkaline earth metals.

The Periodic Table Group 6A elements (O, S, Se, Te, Po) are called chalcogens.

The Periodic Table Group 7A elements (F, Cl, Br, I, At) are called halogens.

The Periodic Table Group 8A elements (He, Ne, Ar, Kr, Xe, Rn) are called the noble gases.

The Periodic Table Groups 1B and 3B-8B are called the transition elements or transition metals.

2.6 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 (NA). We normally round this to 6.022×1023. 1 mole = 6.022×1023, just like 1 dozen = 12 or 1 gross = 144. NA = 6.0221415 x 1023

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 1.00×1020 Ca atoms. Strategy Use Avogardo’s constant, 1 mole = 6.022×1023, to convert from moles to atoms and from atoms to moles. Solution 30.00 mol Ca × 1.00×1020 Ca atoms × 6.022×1023 Ca atoms 1 mol Ca = 1.807×1025 Ca atoms 1 mol Ca 6.022×1023 Ca atoms = 1.66×10-4 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 (1×1020) is less than Avogadro’s number, so there is less than a mole of substance.

Salt (Sodium Chloride) The Mole One mole each of some familiar substances: Helium (in balloon) Water Sugar (Sucrose) Aluminum Copper Salt (Sodium Chloride)

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. Mass of 1 carbon-12 atom: exactly 12 amu 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. 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. 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 25.00 g C × 10.50 g He × 15.75 g Na × 1 mol C 12.01 g C = 2.082 mol C 1 mol He 4.003 g He = 2.623 mol He 1 mol Na 22.99 g Na = 0.6851 mol Na

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 6.89×1018 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. 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. Setup (a) The molar mass of carbon is 12.01 g/mol. (b) The molar mass of helium is 4.003 g/mol. NA = 6.022×1023 Solution 0.515 g C × 6.89×1018 He atoms × 1 mol C 12.01 g C 6.022 ×1023 C atoms 1 mol C × = 2.58×1022 C atoms 1 mol He 6.022 ×1023 He atoms 4.003 g He 1 mol He × = 4.58×10-5 g He

2 Chapter Summary: Key Points Atoms Elements The Atomic Theory Discovery of the Electron Radioactivity The Proton and the Nucleus Nuclear Model of the Atom The Neutron Atomic Number Mass Number Average Atomic Mass The Periodic Table Molar Mass

Group Quiz #2: Subatomic Particles Fill in the table below: # protons # neutrons # electrons 23Na 21Na1+ 32S2-