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Chapter 4 “Atomic Structure” Defining the Atom OBJECTIVES: Democritus’s ideas about atoms. Describe Defining the Atom OBJECTIVES: Explain theory. Dalton’s atomic Defining the Atom OBJECTIVES: Identify what instrument is used to observe individual atoms. Democritus’s Atomic Philosophy The Greek philosopher Democritus (460 B.C. – 370 B.C.) was among the first to suggest the existence of atoms (from the Greek word “atomos”) He believed that atoms were indivisible and indestructible His ideas did agree with later scientific theory, but did not explain chemical behavior, and was not based on the scientific method – but just philosophy Dalton’s Atomic Theory (experiment based!) John Dalton (1766 – 1844) 1) All elements are composed of tiny indivisible particles called atoms 2) Atoms of the same element are identical. Atoms of any one element are different from those of any other element. 3) Atoms of different elements combine in simple whole-number ratios to form chemical compounds 4) In chemical reactions, atoms are combined, separated, or rearranged – but never changed into atoms of another element. 4.1 Defining the Atom > Early Models of the Atom All elements are composed of tiny indivisible particles called atoms. Slide 7 of 18 © Copyright Pearson Prentice Hall 4.1 Defining the Atom > Early Models of the Atom Atoms of the same element are identical. The atoms of any one element are different from those of any other element. Slide 8 of 18 © Copyright Pearson Prentice Hall 4.1 Defining the Atom > Early Models of the Atom Atoms of different elements can physically mix together or can chemically combine in simple whole-number ratios to form compounds. Slide 9 of 18 © Copyright Pearson Prentice Hall 4.1 Defining the Atom > Early Models of the Atom Chemical reactions occur when atoms are separated, joined, or rearranged. Atoms of one element are never changed into atoms of another element in a chemical reaction. Slide 10 of 18 © Copyright Pearson Prentice Hall Sizing up the Atom Elements are able to be subdivided into smaller and smaller particles – these are the atoms, and they still have properties of that element If you could line up 100,000,000 copper atoms in a single file, they would be approximately 1 cm long Despite their small size, individual atoms are observable with instruments such as scanning tunneling (electron) microscopes 4.1 Defining the Atom > Sizing up the Atom Iron Atoms Seen Through a Scanning Tunneling Microscope Slide 12 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Chapter 4 Atomic Structure 4.1 Defining the Atom 4.2 Structure of the Nuclear Atom 4.3 Distinguishing Among Atoms Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 13 of 18 © Copyright Pearson Prentice Hall Defining the Atom > CHEMISTRY & YOU How did scientists determine the structures that are inside an atom? Doctors often use X-rays to see bones and other structures that cannot be seen through your skin. Scientists use many methods to “see” inside an atom. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 14 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Subatomic Particles Much of Dalton’s atomic theory is accepted today. • One important change, however, is that atoms are now known to be divisible. • They can be broken down into even smaller, more fundamental particles, called subatomic particles. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 15 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Subatomic Particles Subatomic Particles What are three kinds of subatomic particles? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 16 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Subatomic Particles Three kinds of subatomic particles are electrons, protons, and neutrons. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 17 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Subatomic Particles Electrons In 1897, the English physicist J. J. Thomson (1856–1940) discovered the electron. • Electrons are negatively charged subatomic particles. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 18 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Subatomic Particles Electrons Thomson performed experiments that involved passing electric current through gases at low pressure. • He sealed the gases in glass tubes fitted at both ends with metal disks called electrodes. • The electrodes were connected to a source of electricity. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. © Copyright Pearson Prentice Hall Slide 19 of 18 Defining the Atom > Subatomic Particles Electrons • One electrode, the anode became positively charged. • The other electrode, the cathode, became negatively charged. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. © Copyright Pearson Prentice Hall Slide 20 of 18 Defining the Atom > Subatomic Particles Electrons The result was a glowing beam, or cathode ray, that traveled from the cathode to the anode. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 21 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Subatomic Particles Electrons Thomson found that a cathode ray is deflected by electrically charged metal plates. • A positively charged plate attracts the cathode ray, while a negatively charged plate repels it. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. © Copyright Pearson Prentice Hall Slide 22 of 18 Defining the Atom > Subatomic Particles Electrons Thomson knew that opposite charges attract and like charges repel, so he hypothesized that a cathode ray is a stream of tiny negatively charged particles moving at high speed. • Thomson called these particles corpuscles. • Later they were named electrons. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 23 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Subatomic Particles Electrons The U.S. physicist Robert A. Millikan (1868–1953) carried out experiments to find the quantity of an electron’s charge. • In his oil-drop experiment, Millikan suspended negatively charged oil droplets between two charged plates. • He then changed the voltage on the plates to see how this affected the droplets’ rate of Slide 24 of 18 fall. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. © Copyright Pearson Prentice Hall Defining the Atom > Subatomic Particles Electrons The U.S. physicist Robert A. Millikan (1868–1953) carried out experiments to find the quantity of an electron’s charge. • From his data, he found that the charge on each oil droplet was a multiple of 1.60 10–19 coulomb, meaning this must be the charge of an electron. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 25 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Subatomic Particles Electrons The U.S. physicist Robert A. Millikan (1868–1953) carried out experiments to find the quantity of an electron’s charge. • Using this charge and Thomson’s chargeto-mass ratio of an electron, Millikan calculated an electron’s mass. • Millikan’s values for electron charge and mass are similar to those accepted today. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. © Copyright Pearson Prentice Hall Slide 26 of 18 Defining the Atom > Subatomic Particles Electrons An electron has one unit of negative charge, and its mass is 1/1840 the mass of a hydrogen atom. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 27 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Subatomic Particles Protons and Neutrons If cathode rays are electrons given off by atoms, what remains of the atoms that have lost the electrons? • For example, after a hydrogen atom (the lightest kind of atom) loses an electron, what is left? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 28 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Subatomic Particles Protons and Neutrons You can think through this problem using four simple ideas about matter and electric charges. 1. Atoms have no net electric charge; they are electrically neutral. 2. Electric charges are carried by particles of matter. 3. Electric charges always exist in whole-number multiples of a single basic unit; that is, there are no fractions of charges. 4. When a given number of negatively charged particles combines with an equal number of positively charged Slide 29 of 18 particles, an electrically neutral particle is formed. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. © Copyright Pearson Prentice Hall Defining the Atom > Subatomic Particles Protons and Neutrons It follows that a particle with one unit of positive charge should remain when a typical hydrogen atom loses an electron. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 30 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Subatomic Particles Protons and Neutrons In 1886, Eugen Goldstein (1850–1930) observed a cathode-ray tube and found rays traveling in the direction opposite to that of the cathode rays. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 31 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Subatomic Particles Protons and Neutrons In 1886, Eugen Goldstein (1850–1930) observed a cathode-ray tube and found rays traveling in the direction opposite to that of the cathode rays. • He concluded that they were composed of positive particles. • Such positively charged subatomic particles are called protons. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 32 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Subatomic Particles Protons and Neutrons In 1932, the English physicist James Chadwick (1891–1974) confirmed the existence of yet another subatomic particle: the neutron. • Neutrons are subatomic particles with no charge but with a mass nearly equal to that of a proton. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 33 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Interpret Data The table below summarizes the properties of these subatomic particles. Properties of Subatomic Particles Particle Symbol Relative charge Relative mass (mass of proton = 1) Actual mass (g) Electron e– 1– 1/1840 9.11 10–28 Proton p+ 1+ 1 1.67 10–24 Neutron n0 0 1 1.67 10–24 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. © Copyright Pearson Prentice Hall Slide 34 of 18 Defining the Atom > Subatomic Particles Although protons and neutrons are extremely small, theoretical physicists believe that they are composed of yet smaller subnuclear particles called quarks. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 35 of 18 © Copyright Pearson Prentice Hall Defining the Atom > The Atomic Nucleus When subatomic particles were discovered, scientists wondered how the particles were put together in an atom. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 36 of 18 © Copyright Pearson Prentice Hall Defining the Atom > When subatomic particles were discovered, scientists wondered how the particles were put together in an atom. • Most scientists—including J. J. Thompson— thought it likely that the electrons were evenly distributed throughout an atom filled uniformly with positively charged material. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 37 of 18 © Copyright Pearson Prentice Hall Defining the Atom > The Atomic Nucleus When subatomic particles were discovered, scientists wondered how the particles were put together in an atom. • Most scientists—including J. J. Thompson— thought it likely that the electrons were evenly distributed throughout an atom filled uniformly with positively charged material. – In Thomson’s atomic model, known as the “plumpudding model,” electrons were stuck into a lump of positive charge, similar to raisins stuck in dough. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. © Copyright Pearson Prentice Hall Slide 38 of 18 Defining the Atom > The Atomic Nucleus This model of the atom turned out to be short-lived, however, due to the work of a former student of Thomson, Ernest Rutherford (1871–1937). • Born in New Zealand, Rutherford was awarded the Nobel Prize for Chemistry in 1908. His portrait appears on the New Zealand $100 bill. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. © Copyright Pearson Prentice Hall Slide 39 of 18 Defining the Atom > The Atomic Nucleus Rutherford’s Gold-Foil Experiment In 1911, Rutherford and his co-workers wanted to test the existing plum-pudding model of atomic structure. • They devised the gold-foil experiment. • Their test used alpha particles, which are helium atoms that have lost their two electrons and have a double positive charge because of the two remaining protons. Slide Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 40 of 18 © Copyright Pearson Prentice Hall Defining the Atom > The Atomic Nucleus Rutherford’s Gold-Foil Experiment In the experiment, a narrow beam of alpha particles was directed at a very thin sheet of gold. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 41 of 18 © Copyright Pearson Prentice Hall Defining the Atom > The Atomic Nucleus Rutherford’s Gold-Foil Experiment In the experiment, a narrow beam of alpha particles was directed at a very thin sheet of gold. • According to the prevailing theory, the alpha particles should have passed easily through the gold, with only a slight deflection due to the positive charge thought to be spread out in the gold atoms. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. © Copyright Pearson Prentice Hall Slide 42 of 18 Defining the Atom > The Atomic Nucleus Rutherford’s Gold-Foil Experiment Rutherford’s results were that most alpha particles went straight through, or were slightly deflected. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 43 of 18 © Copyright Pearson Prentice Hall Defining the Atom > The Atomic Nucleus Rutherford’s Gold-Foil Experiment Rutherford’s results were that most alpha particles went straight through, or were slightly deflected. • What was surprising is that a small fraction of the alpha particles bounced off the gold foil at very large angles. • Some even bounced straight back toward the source. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. © Copyright Pearson Prentice Hall Slide 44 of 18 Defining the Atom > The Atomic Nucleus The Rutherford Atomic Model Based on his experimental results, Rutherford suggested a new theory of the atom. • He proposed that the atom is mostly empty space. – Thus explaining the lack of deflection of most of the alpha particles. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 45 of 18 © Copyright Pearson Prentice Hall Defining the Atom > The Atomic Nucleus The Rutherford Atomic Model Based on his experimental results, Rutherford suggested a new theory of the atom. • He concluded that all the positive charge and almost all of the mass are concentrated in a small region that has enough positive charge to account for the great deflection of some of the alpha particles. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 46 of 18 © Copyright Pearson Prentice Hall Defining the Atom > The Atomic Nucleus The Rutherford Atomic Model The Rutherford atomic model is known as the nuclear atom. In the nuclear atom, the protons and neutrons are located in the positively charged nucleus. The electrons are distributed around the nucleus and occupy almost all the volume of the atom. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 47 of 18 © Copyright Pearson Prentice Hall Defining the Atom > The Atomic Nucleus The Rutherford Atomic Model According to this model, the nucleus is tiny and densely packed compared with the atom as a whole. • If an atom were the size of a football stadium, the nucleus would be about the size of a marble. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 48 of 18 © Copyright Pearson Prentice Hall Defining the Atom > The Atomic Nucleus The Rutherford Atomic Model Rutherford’s model turned out to be incomplete. • The Rutherford atomic model had to be revised in order to explain the chemical properties of elements. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 49 of 18 © Copyright Pearson Prentice Hall Defining the Atom > What evidence from Rutherford’s Gold-Foil experiment disproves J.J. Thompson’s “plum-pudding model”? Rutherford observed that most of the particles passed through the foil with no deflection, and a small fraction were deflected at large angles or reflected directly back. If the plum-pudding model was true, most alpha particles would have been deflected at small angles by the evenlySlide spaced electrons. 50 of 18 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. © Copyright Pearson Prentice Hall Defining the Atom > Key Concepts • Three kinds of subatomic particles are electrons, protons, and neutrons. • In the nuclear atom, the protons and neutrons are located in the nucleus. The electrons are distributed around the nucleus and occupy almost all the volume of the atom. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 51 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Glossary Terms • electron: a negatively charged subatomic particle • cathode ray: a stream of electrons produced at the negative electrode (cathode) of a tube containing a gas at low pressure • proton: a positively charged subatomic particle found in the nucleus of an atom • neutron: a subatomic particle with no charge and a mass of 1 amu; found in the nucleus of an atom • nucleus: the tiny, dense central portion of an atom, composed of protons and neutrons Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. © Copyright Pearson Prentice Hall Slide 52 of 18 Defining the Atom > BIG IDEA Electrons and the Structure of Atoms Atoms have positively-charged protons and neutral neutrons inside a nucleus, and negatively-charged electrons outside the nucleus. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 53 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Chapter 4 Atomic Structure 4.1 Defining the Atom 4.2 Structure of the Nuclear Atom 4.3 Distinguishing Among Atoms Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 54 of 18 © Copyright Pearson Prentice Hall Defining the Atom > CHEMISTRY & YOU How can there be different varieties of atoms? Just as there are many types of dogs, atoms come in different varieties too. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 55 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Atomic Number and Mass Number Atomic Number and Mass Number What makes one element different from another? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 56 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Atomic Number and Mass Number Atomic Number Elements are different because they contain different numbers of protons. • An element’s atomic number is the number of protons in the nucleus of an atom of that element. • The atomic number identifies an element. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. © Copyright Pearson Prentice Hall Slide 57 of 18 Defining the Atom > Interpret Data For each element listed in the table below, the number of protons equals the number of electrons. Atoms of the First Ten Elements Name Symbol Atomic number Protons Neutrons Mass number Electrons Hydrogen H 1 1 0 1 1 Helium He 2 2 2 4 2 Lithium Li 3 3 4 7 3 Beryllium Be 4 4 5 9 4 Boron B 5 5 6 11 5 Carbon C 6 6 6 12 6 Nitrogen N 7 7 7 14 7 Oxygen O 8 8 8 16 8 Fluorine F 9 9 10 19 9 20 10 Neon Ne 10 10 Copyright © Pearson Education, Inc., or its 10 affiliates. All Rights Reserved. © Copyright Pearson Prentice Hall Slide 58 of 18 Defining the Atom > Atomic Number and Mass Number Atomic Number Remember that atoms are electrically neutral. • Thus, the number of electrons (negatively charged particles) must equal the number of protons (positively charged particles). Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 59 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Sample Problem 4.1 Understanding Atomic Number The element nitrogen (N) has an atomic number of 7. How many protons and electrons are in a neutral nitrogen atom? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 60 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Sample Problem 4.1 1 Analyze Identify the relevant concepts. The atomic number gives the number of protons, which in a neutral atom equals the number of electrons. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 61 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Sample Problem 4.1 2 Solve Apply the concepts to this problem. • Identify the atomic number. • Then use the atomic number to find the number of protons and electrons. The atomic number of nitrogen is 7. So, a neutral nitrogen atom has 7 protons and 7 electrons. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 62 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Atomic Number and Mass Number Mass Number The total number of protons and neutrons in an atom is called the mass number. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 63 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Atomic Number and Mass Number Mass Number If you know the atomic number and mass number of an atom of any element, you can determine the atom’s composition. The number of neutrons in an atom is the difference between the mass number and atomic number. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 64 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Atomic Number and Mass Number Mass Number If you know the atomic number and mass number of an atom of any element, you can determine the atom’s composition. Number of neutrons = mass number – atomic number Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 65 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Atomic Number and Mass Number Mass Number The composition of any atom can be represented in shorthand notation using atomic number and mass number. Au is the chemical symbol for gold. • The atomic number is the subscript. • The mass number is the superscript. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. © Copyright Pearson Prentice Hall Slide 66 of 18 Defining the Atom > Atomic Number and Mass Number Mass Number You can also refer to atoms by using the mass number and the name of the element. • 197 79 Au may be written as gold-197. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Au is the chemical symbol for gold. Slide 67 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Sample Problem 4.2 Determining the Composition of an Atom How many protons, electrons, and neutrons are in each atom? 9 4 a. Be Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 20 10 b. Ne 23 11 c. Na Slide 68 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Sample Problem 4.2 1 Analyze List the knowns and the unknowns. Use the definitions of atomic number and mass number to calculate the numbers of protons, electrons, and neutrons. KNOWNS Beryllium (Be) atomic number = 4 mass number = 9 Neon (Ne) atomic number = 10 mass number = 20 Sodium (Na) atomic number = 11 mass number = 23 UNKNOWNS protons = ? electrons = ? neutrons = ? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 69 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Sample Problem 4.2 2 Calculate Solve for the unknowns. Use the atomic number to find the number of protons. atomic number = number of protons a. 4 b. 10 c. 11 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 70 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Sample Problem 4.2 2 Calculate Solve for the unknowns. Use the atomic number to find the number of electrons. atomic number = number of electrons a. 4 b. 10 c. 11 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 71 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Sample Problem 4.2 2 Calculate Solve for the unknowns. Use the mass number and atomic number to find the number of neutrons. number of neutrons = mass number – atomic number a. number of neutrons = 9 – 4 = 5 b. number of neutrons = 20 – 10 = 10 c. number of neutrons = 23 – 11 = 12 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 72 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Sample Problem 4.2 3 Evaluate Do the results make sense? • For each atom, the mass number equals the number of protons plus the number of neutrons. • The results make sense. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 73 of 18 © Copyright Pearson Prentice Hall Defining the Atom > What information is needed to determine the composition of a neutral atom of any element? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 74 of 18 © Copyright Pearson Prentice Hall Defining the Atom > What information is needed to determine the composition of a neutral atom of any element? The atomic number and mass number are needed to determine an atom’s composition. The atomic number gives the number of protons, which equals the number of electrons. The number of neutrons is the difference between the mass number and the atomic number. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 75 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Isotopes There are three different kinds of neon atoms. • How do these atoms differ? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 76 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Isotopes • All have the same number of protons (10). • All have the same number of electrons (10). • But they each have different numbers of neutrons. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 77 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Isotopes Isotopes are atoms that have the same number of protons but different numbers of neutrons. • Neon-20, neon-21, and neon 22 are three isotopes of neon. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 78 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Isotopes Because isotopes of an element have different numbers of neutrons, they also have different mass numbers. • Despite these differences, isotopes are chemically alike because they have identical numbers of protons and electrons, which are the subatomic particles responsible for chemical Slide behavior. 79 of 18 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. © Copyright Pearson Prentice Hall Defining the Atom > Isotopes Remember the dogs at the beginning of the lesson. • Their color or size doesn’t change the fact that they are all dogs. • Similarly, the number of neutrons in isotopes of an element does not change which element it is because the atomic number does not change. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. © Copyright Pearson Prentice Hall Slide 80 of 18 Defining the Atom > CHEMISTRY & YOU How are the atoms of one element different from the atoms of another element? How are isotopes of the same element different? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 81 of 18 © Copyright Pearson Prentice Hall Defining the Atom > CHEMISTRY & YOU How are the atoms of one element different from the atoms of another element? How are isotopes of the same element different? Atoms of different elements are different because they contain different numbers of protons. Isotopes of the same element are different because they have different numbers of neutrons, and thus different mass numbers. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 82 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Sample Problem 4.3 Writing Chemical Symbols of Isotopes Diamonds are a naturally occurring form of elemental carbon. Two stable isotopes of carbon are carbon-12 and carbon-13. Write the symbol for each isotope using superscripts and subscripts to represent the mass number and the atomic number. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 83 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Sample Problem 4.3 1 Analyze Identify the relevant concepts. Isotopes are atoms that have the same number of protons but different numbers of neutrons. The composition of an atom can be expressed by writing the chemical symbol, with the atomic number as a subscript and the mass number as a superscript. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 84 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Sample Problem 4.3 2 Solve Apply the concepts to this problem. Use Table 4.2 to identify the symbol and the atomic number for carbon. The symbol for carbon is C. The atomic number of carbon is 6. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 85 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Sample Problem 4.3 2 Solve Apply the concepts to this problem. Look at the name of the isotope to find the mass number. For carbon-12, the mass number is 12. For carbon-13, the mass number is 13. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 86 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Sample Problem 4.3 2 Solve Apply the concepts to this problem. Use the symbol, atomic number, and mass number to write the symbol of the isotope. For carbon-12, the symbol is 12 6 C. For carbon-13, the symbol is 13 6 C. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 87 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Atomic Mass Atomic Mass How do you calculate the atomic mass of an element? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 88 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Atomic Mass The mass of even the largest atom is incredibly small. • Since the 1920s, it has been possible to determine the tiny masses of atoms by using a mass spectrometer. • The mass of a fluorine atom was found to be 3.155 x 10–23 g. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 89 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Atomic Mass Such data about the actual masses of individual atoms can provide useful information, but in general these values are inconveniently small and impractical to work with. • Instead, it is more useful to compare the relative masses of atoms using a reference isotope as a standard. • The reference isotope chosen is carbon-12. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. © Copyright Pearson Prentice Hall Slide 90 of 18 Defining the Atom > Atomic Mass This isotope of carbon has been assigned a mass of exactly 12 atomic mass units. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 91 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Atomic Mass In nature, most elements occur as a mixture of two or more isotopes. • Each isotope of an element has a fixed mass and a natural percent abundance. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 92 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Interpret Data Natural Percent Abundance of Stable Isotopes of Some Elements Name Hydrogen Symbol 1 1 2 1 H H 3 1H Helium 3 2 4 2 Carbon 12 6 13 6 He He C C 16 8 Oxygen O 17 8O 18 O 8 Chlorine 35 17 37 17 Cl Cl Natural percent abundance Mass (amu) Atomic mass 99.985 0.015 negligible 1.0078 2.0141 3.0160 1.0079 0.0001 99.9999 3.0160 4.0026 4.0026 98.89 1.11 12.000 13.003 12.011 99.759 0.037 0.204 15.995 16.995 17.999 15.999 75.77 34.969 24.23 Copyright © Pearson Education, Inc., or its affiliates. 36.966 35.453 All Rights Reserved. © Copyright Pearson Prentice Hall Slide 93 of 18 Defining the Atom > Atomic Mass Chlorine occurs as two isotopes: chlorine35 and chlorine-37. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 94 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Atomic Mass Chlorine occurs as two isotopes: chlorine35 and chlorine-37. • If you calculate the arithmetic mean of these two masses ((34.968 amu + 36.966 amu)/2), you get an average atomic mass of 35.986. • However, this value is higher than the actual value of 35.453. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. © Copyright Pearson Prentice Hall Slide 95 of 18 Defining the Atom > Atomic Mass Chlorine occurs as two isotopes: chlorine35 and chlorine-37. • To explain this difference, you need to know the natural percent abundance of the isotopes of chlorine. • Chlorine-35 accounts for 75 percent of the naturally occurring chlorine atoms; chlorine-37 Slide accounts for only 24 percent. 96 of 18 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. © Copyright Pearson Prentice Hall Defining the Atom > Atomic Mass Because there is more chlorine-35 than chlorine37 in nature, the atomic mass of chlorine, 35.453 amu, is closer to 35 than to 37. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 97 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Atomic Mass The atomic mass of an element is a weighted average mass of the atoms in a naturally occurring sample of the element. • A weighted average mass reflects both the mass and the relative abundance of the isotopes as they occur in nature. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 98 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Sample Problem 4.4 Understanding Relative Abundance of Isotopes The atomic mass of copper is 63.546 amu. Which of copper’s two isotopes is more abundant: copper-63 or copper-65? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 99 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Atomic Mass • To calculate the atomic mass of an element, multiply the mass of each isotope by its natural abundance, expressed as a decimal, and then add the products. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 100 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Atomic Mass Carbon has two stable isotopes: carbon-12, which has a natural abundance of 98.89 percent, and carbon-13, which has a natural abundance of 1.11 percent. • The mass of carbon-12 is 12.000 amu; the mass of carbon-13 is 13.003 amu. • The atomic mass of carbon is calculated as follows: Atomic mass of carbon = (12.000 amu x 0.9889) + 13.003 amu x 0.0111) = (11.867 amu) + (0.144 amu) = 12.011 amu Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. © Copyright Pearson Prentice Hall Slide 101 of 18 Defining the Atom > Sample Problem 4.5 Calculating Atomic Mass Element X has two naturally occurring isotopes. The isotope with a mass of 10.012 amu (10X) has a relative abundance of 19.91 percent. The isotope with a mass of 11.009 amu (11X) has a relative abundance of 80.09 percent. Calculate the atomic mass of element X. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 102 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Sample Problem 4.5 2 Calculate Solve for the unknowns. Use the atomic mass and the decimal form of the percent abundance to find the mass contributed by each isotope. for 10X: 10.012 amu x 0.1991 = 1.993 amu for 11X: 11.009 amu x 0.8009 = 8.817 amu Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 103 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Sample Problem 4.5 2 Calculate Solve for the unknowns. Add the atomic mass contributions for all the isotopes. For element X, atomic mass = 1.953 amu + 8.817 amu = 10.810 amu Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 104 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Sample Problem 4.5 3 Evaluate Does the result make sense? The calculated value is closer to the mass of the more abundant isotope, as would be expected. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 105 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Why is the atomic mass of an element usually not a whole number? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 106 of 18 © Copyright Pearson Prentice Hall Defining the Atom > Why is the atomic mass of an element usually not a whole number? The atomic mass of an element is usually not a whole number because it is a weighted average of the masses of the naturally occurring isotopes of the element. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Slide 107 of 18 © Copyright Pearson Prentice Hall