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1 4.1 Defining the Atom > 1 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Chapter 4 Atomic Structure 4.1 Defining the Atom 4.2 Structure of the Nuclear Atom 4.3 Distinguishing Among Atoms

2 4.1 Defining the Atom > 2 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. How do you study something that you cannot see it? CHEMISTRY & YOU Similar to how you might study a gift- wrapped present, scientists often study things that cannot be seen with the naked eye.

3 4.1 Defining the Atom > 3 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Early Models of the Atom How did the concept of the atom change from the time of Democritus to the time of John Dalton?

4 4.1 Defining the Atom > 4 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Early Models of the Atom An atom is the smallest particle of an element that retains its identity in a chemical reaction. Although early philosophers and scientists could not observe individual atoms, they were still able to propose ideas about the structure of atoms.

5 4.1 Defining the Atom > 5 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Early Models of the Atom Democritus’s Atomic Philosophy The Greek philosopher Democritus (460 BC –370 BC ) was among the first to suggest the existence of atoms.

6 4.1 Defining the Atom > 6 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Early Models of the Atom Democritus’s Atomic Philosophy Democritus reasoned that atoms were indivisible and indestructible.

7 4.1 Defining the Atom > 7 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Early Models of the Atom Democritus’s Atomic Philosophy Democritus reasoned that atoms were indivisible and indestructible. Although, Democritus’s ideas agreed with later scientific theory, they did not explain chemical behavior They also lacked experimental support because Democritus’s approach was not based on the scientific method.

8 4.1 Defining the Atom > 8 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Dalton’s Atomic Theory The modern process of discovery regarding atoms began with John Dalton (1766–1864), an English chemist and schoolteacher. Early Models of the Atom

9 4.1 Defining the Atom > 9 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Dalton’s Atomic Theory By using experimental methods, Dalton transformed Democritus’s ideas on atoms into a scientific theory. Early Models of the Atom

10 4.1 Defining the Atom > 10 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Dalton’s Atomic Theory Early Models of the Atom Dalton studied the ratios in which elements combine in chemical reactions. The result of his work is known as Dalton’s atomic theory.

11 4.1 Defining the Atom > 11 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Dalton’s Atomic Theory Atoms of element A Early Models of the Atom 1.All elements are composed of tiny indivisible particles called atoms.

12 4.1 Defining the Atom > 12 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Dalton’s Atomic Theory Atoms of element A Atoms of element B Early Models of the Atom 2.Atoms of the same element are identical. The atoms of any one element are different from those of any other element.

13 4.1 Defining the Atom > 13 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Dalton’s Atomic Theory Mixture of atoms of elements A and B Early Models of the Atom 3.Atoms of different elements can physically mix together or can chemically combine in simple whole- number ratios to form compounds.

14 4.1 Defining the Atom > 14 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Compound made by chemically combining atoms of elements A and B Dalton’s Atomic Theory Early Models of the Atom 4.Chemical reactions occur when atoms are separated from each other, joined, or rearranged in different combinations. Atoms of one element are never changed into atoms of another element as a result of a chemical reaction.

15 4.1 Defining the Atom > 15 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. How was Jon Dalton able to study atoms even though he couldn’t observe them directly? What evidence did he use to formulate his atomic theory? CHEMISTRY & YOU

16 4.1 Defining the Atom > 16 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. How was Jon Dalton able to study atoms even though he couldn’t observe them directly? What evidence did he use to formulate his atomic theory? CHEMISTRY & YOU Dalton studied the ratios in which elements combine in chemical reactions. He observed that when atoms mix, they maintain their own identity unless they combine in a chemical reaction.

17 4.1 Defining the Atom > 17 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. How was Democritus’s idea of the atom different from Dalton’s?

18 4.1 Defining the Atom > 18 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. How was Democritus’s idea of the atom different from Dalton’s? Democritus’s idea did not explain chemical behavior and was not a scientific theory because it lacked experimental support. Using experimental support, Dalton transformed Democritus’s ideas about atoms into a scientific theory that explained chemical behavior.

19 4.1 Defining the Atom > 19 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Sizing up the Atom What instruments are used to observe individual atoms? Sizing up the Atom

20 4.1 Defining the Atom > 20 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Sizing up the Atom This liquid mercury illustrates Dalton’s concept of the atom. Every drop, no matter its size has the same properties. Even if you could make a drop the size of one atom, it would still have the chemical properties of mercury.

21 4.1 Defining the Atom > 21 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. If you could continue to make the copper dust smaller, you would eventually come upon a particle of copper that could no longer be divided and still have the chemical properties of copper. This final particle is an atom. Sizing up the Atom If you were to grind a copper coin into a fine dust, each speck in the small pile of shiny red dust would still have the properties of copper.

22 4.1 Defining the Atom > 22 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Sizing up the Atom A pure copper coin the size of a penny contains about 2.4  atoms. By comparison, Earth’s population is only about 7  10 9 people. If you could line up 100,000,000 copper atoms side by side, they would produce a line only 1 cm long! Atoms are very small.

23 4.1 Defining the Atom > 23 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Despite their small size, individual atoms are observable with instruments such as scanning electron microscopes. Sizing up the Atom In scanning electron microscopes, a beam of electrons is focused on the sample. Electron microscopes are capable of much higher magnifications than light microscopes.

24 4.1 Defining the Atom > 24 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. The ability to move individual atoms holds future promise for the creation of atomic- sized electronic devices, such as circuits and computer chips. Sizing up the Atom With the help of electron microscopes, individual atoms can even be moved around and arranged in patterns.

25 4.1 Defining the Atom > 25 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. This atomic-scale, or “nanoscale,” technology could become essential to future applications in medicine, communications, solar energy, and space exploration. Sizing up the Atom An example of a device made from individual atoms is the nanocar shown here.

26 4.1 Defining the Atom > 26 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. If an atom has a radius of 1  m, how many of these atoms must be lined up in a row to produce a line 1 m long? 1  (10,000,000,000) atoms of radius 1  m would need to be lined up in a row to produce a line 1 m long.

27 4.1 Defining the Atom > 27 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Key Concepts Democritus reasoned that atoms were indivisible and indestructible. By using experimental methods, Dalton transformed Democritus’s ideas on atoms into a scientific theory. Scientists can observe individual atoms by using instruments such as scanning electron microscopes.

28 4.1 Defining the Atom > 28 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Glossary Definitions atom: the smallest particle of an element that retains its identity in a chemical reaction Dalton’s atomic theory: the first theory to relate chemical changes to events at the atomic level

29 4.1 Defining the Atom > 29 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Electrons and the Structure of Atoms BIG IDEA Atoms are the smallest particles of an element that still have the chemical properties of that element.

30 4.1 Defining the Atom > 30 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. END OF 4.1

31 4.1 Defining the Atom > 31 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Chapter 4 Atomic Structure 4.1 Defining the Atom 4.2 Structure of the Nuclear Atom 4.3 Distinguishing Among Atoms

32 4.1 Defining the Atom > 32 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 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. CHEMISTRY & YOU How did scientists determine the structures that are inside an atom?

33 4.1 Defining the Atom > 33 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Subatomic Particles –What are three kinds of subatomic particles and how they were discovered? Subatomic Particles

34 4.1 Defining the Atom > 34 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 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.

35 4.1 Defining the Atom > 35 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Subatomic Particles Three kinds of subatomic particles are electrons, protons, and neutrons.

36 4.1 Defining the Atom > 36 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Subatomic Particles In 1897, the English physicist J. J. Thomson (1856–1940) discovered the electron. Electrons are negatively charged subatomic particles. Electrons

37 4.1 Defining the Atom > 37 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Electrons Subatomic Particles 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.

38 4.1 Defining the Atom > 38 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Electrons Subatomic Particles One electrode, the anode became positively charged. The other electrode, the cathode, became negatively charged.

39 4.1 Defining the Atom > 39 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Electrons Subatomic Particles The result was a glowing beam, or cathode ray, that traveled from the cathode to the anode.

40 4.1 Defining the Atom > 40 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Electrons Subatomic Particles 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.

41 4.1 Defining the Atom > 41 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Subatomic Particles Thompson 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. Thompson called these particles corpuscles. Later they were named electrons. Electrons

42 4.1 Defining the Atom > 42 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Subatomic Particles A cathode ray can also be deflected by a magnet. Electrons

43 4.1 Defining the Atom > 43 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Subatomic Particles To test his hypothesis, Thomson set up an experiment to measure the ratio of an electron’s charge to its mass. He found this ratio to be constant. Also, the charge-to-mass ratio of electrons did not depend on the kind of gas in the cathode-ray tube or the type of metal used for the electrodes. Electrons

44 4.1 Defining the Atom > 44 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Subatomic Particles Thompson concluded that electrons are a component of the atoms of all elements. Electrons

45 4.1 Defining the Atom > 45 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Subatomic Particles 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 fall. Electrons

46 4.1 Defining the Atom > 46 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Subatomic Particles 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. Electrons

47 4.1 Defining the Atom > 47 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Subatomic Particles 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 charge- to-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. Electrons

48 4.1 Defining the Atom > 48 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Subatomic Particles An electron has one unit of negative charge, and its mass is 1/1840 the mass of a hydrogen atom. Electrons

49 4.1 Defining the Atom > 49 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Subatomic Particles 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? Protons and Neutrons

50 4.1 Defining the Atom > 50 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 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 particles, an electrically neutral particle is formed.

51 4.1 Defining the Atom > 51 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 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.

52 4.1 Defining the Atom > 52 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 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.

53 4.1 Defining the Atom > 53 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 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.

54 4.1 Defining the Atom > 54 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. The table below summarizes the properties of these subatomic particles. Properties of Subatomic Particles ParticleSymbol Relative charge Relative mass (mass of proton = 1) Actual mass (g) Electrone–e– 1–1–1/  10 –28 Protonp+p  10 –24 Neutronn0n  10 –24 Interpret Data

55 4.1 Defining the Atom > 55 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Subatomic Particles Although protons and neutrons are extremely small, theoretical physicists believe that they are composed of yet smaller sub-nuclear particles called quarks.

56 4.1 Defining the Atom > 56 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. What evidence did Thomson have that led him to conclude the cathode ray was a stream of tiny negatively charged particles? The ray was attracted to a metal plate with positive electric charge, and Thomson knew that opposite charges attract and like charges repel.

57 4.1 Defining the Atom > 57 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. The Atomic Nucleus –How can you describe the structure of the nuclear atom?

58 4.1 Defining the Atom > 58 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. When subatomic particles were discovered, scientists wondered how the particles were put together in an atom. Most scientists—including J. J. Thomson— 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 “plum- pudding model,” electrons were stuck into a lump of positive charge, similar to raisins stuck in dough.

59 4.1 Defining the Atom > 59 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Born in New Zealand, Rutherford was awarded the Nobel Prize for Chemistry in His portrait appears on the New Zealand $100 bill. 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).

60 4.1 Defining the Atom > 60 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 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.

61 4.1 Defining the Atom > 61 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 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.

62 4.1 Defining the Atom > 62 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 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.

63 4.1 Defining the Atom > 63 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Scientists determined the structures inside an atom by observing how the atom interacted with particles with known properties, like negatively charged electrons or positively charged alpha particles. CHEMISTRY & YOU How did scientists “see” inside an atom to determine the structures that are inside an atom?

64 4.1 Defining the Atom > 64 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 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.

65 4.1 Defining the Atom > 65 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 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.

66 4.1 Defining the Atom > 66 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 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.

67 4.1 Defining the Atom > 67 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 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

68 4.1 Defining the Atom > 68 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. The Atomic Nucleus The Rutherford Atomic Model Rutherford’s model turned out to be correct yet incomplete. The Rutherford atomic model had to be revised in order to explain the chemical properties of elements. We will discuss the revision of the model later.

69 4.1 Defining the Atom > 69 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. What evidence from Rutherford’s Gold- Foil experiment disproves J.J. Thompson’s “plum-pudding model”?

70 4.1 Defining the Atom > 70 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 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 evenly- spaced electrons.

71 4.1 Defining the Atom > 71 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 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.

72 4.1 Defining the Atom > 72 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 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

73 4.1 Defining the Atom > 73 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Atoms have positively-charged protons and neutral neutrons inside a nucleus, and negatively-charged electrons outside the nucleus. BIG IDEA Electrons and the Structure of Atoms

74 4.1 Defining the Atom > 74 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. END OF 4.2

75 4.1 Defining the Atom > 75 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Chapter 4 Atomic Structure 4.1 Defining the Atom 4.2 Structure of the Nuclear Atom 4.3 Distinguishing Among Atoms

76 4.1 Defining the Atom > 76 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Just as there are many types of dogs, atoms come in different varieties too. CHEMISTRY & YOU How can there be different varieties of atoms?

77 4.1 Defining the Atom > 77 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Atomic Number and Mass Number –What makes one element different from another? Atomic Number and Mass Number

78 4.1 Defining the Atom > 78 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Atomic Number and Mass 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. Atomic Number

79 4.1 Defining the Atom > 79 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. For each element listed in the table below, the number of protons equals the number of electrons. Interpret Data Atoms of the First Ten Elements NameSymbolAtomic number ProtonsNeutronsMass number Electrons HydrogenH11011 HeliumHe22242 LithiumLi33473 BerylliumBe44594 BoronB CarbonC NitrogenN OxygenO FluorineF NeonNe

80 4.1 Defining the Atom > 80 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Atomic Number and Mass Number Remember that atoms are electrically neutral. Atomic Number Thus, the number of electrons (negatively charged particles) must equal the number of protons (positively charged particles).

81 4.1 Defining the Atom > 81 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Understanding Atomic Number The element nitrogen (N) has an atomic number of 7. How many protons and electrons are in a neutral nitrogen atom? Sample Problem 4.1

82 4.1 Defining the Atom > 82 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Sample Problem 4.1 The atomic number gives the number of protons, which in a neutral atom equals the number of electrons. Analyze Identify the relevant concepts. 1

83 4.1 Defining the Atom > 83 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Sample Problem 4.1 Solve Apply the concepts to this problem. 2 Identify the atomic number, which is 7 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.

84 4.1 Defining the Atom > 84 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Atomic Number and Mass Number Mass Number The total number of protons and neutrons in an atom is called the mass number.

85 4.1 Defining the Atom > 85 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Atomic Number and Mass Number The number of neutrons in an atom is the difference between the mass number and atomic 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

86 4.1 Defining the Atom > 86 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Atomic Number and Mass Number The composition of any atom can be represented in shorthand notation using atomic number and mass number. Mass Number The atomic number is the subscript. The mass number is the superscript. Au is the chemical symbol for gold.

87 4.1 Defining the Atom > 87 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Atomic Number and Mass Number You can also refer to atoms by using the mass number and the name of the element. Mass Number Au may be written as gold Au is the chemical symbol for gold.

88 4.1 Defining the Atom > 88 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Sample Problem 4.2 Determining the Composition of an Atom How many protons, electrons, and neutrons are in each atom? a.Beb.Nec.Na

89 4.1 Defining the Atom > 89 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Sample Problem 4.2 Use the definitions of atomic number and mass number to calculate the numbers of protons, electrons, and neutrons. Analyze List the knowns and the unknowns. 1 protons = ? electrons = ? neutrons = ? KNOWNS UNKNOWNS Beryllium (Be) atomic number = 4 mass number = 9 Neon (Ne) atomic number = 10 mass number = 20 Sodium (Na) atomic number = 11 mass number = 23

90 4.1 Defining the Atom > 90 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Sample Problem 4.2 Use the atomic number to find the number of protons. atomic number = number of protons a.4b.10c.11 Calculate Solve for the unknowns. 2 a.Beb.Nec.Na

91 4.1 Defining the Atom > 91 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Sample Problem 4.2 Use the atomic number to find the number of electrons. atomic number = number of electrons a.4b.10c.11 Calculate Solve for the unknowns. 2 a.Beb.Nec.Na

92 4.1 Defining the Atom > 92 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Sample Problem 4.2 Use the mass number and atomic number to find the number of neutrons. a.number of neutrons = 9 – 4 = 5 b.number of neutrons = 20 – 10 = 10 c.number of neutrons = 23 – 11 = 12 Calculate Solve for the unknowns. 2 number of neutrons = mass number – atomic number a.Beb.Nec.Na

93 4.1 Defining the Atom > 93 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Sample Problem 4.2 Evaluate Do the results make sense? 3 For each atom, the mass number equals the number of protons plus the number of neutrons. The results make sense.

94 4.1 Defining the Atom > 94 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 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.

95 4.1 Defining the Atom > 95 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Atomic number of the elements can be found in the periodic table Rows are called periods, and columns are called groups.

96 4.1 Defining the Atom > 96 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.Isotopes Isotopes –How do isotopes of an element differ?

97 4.1 Defining the Atom > 97 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.Isotopes There are three different kinds of neon atoms. How do these atoms differ?

98 4.1 Defining the Atom > 98 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.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.

99 4.1 Defining the Atom > 99 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.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.

100 4.1 Defining the Atom > 100 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Because isotopes of an element have different numbers of neutrons, they also have different mass numbers.Isotopes Despite these differences, isotopes are chemically alike because they have identical numbers of protons and electrons, which are the subatomic particles responsible for chemical behavior.

101 4.1 Defining the Atom > 101 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.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.

102 4.1 Defining the Atom > 102 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. How are the atoms of one element different from the atoms of another element? How are isotopes of the same element different? CHEMISTRY & YOU

103 4.1 Defining the Atom > 103 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. How are the atoms of one element different from the atoms of another element? How are isotopes of the same element different? CHEMISTRY & YOU 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.

104 4.1 Defining the Atom > 104 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Writing Chemical Symbols of Isotopes Sample Problem 4.3 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.

105 4.1 Defining the Atom > 105 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 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. Sample Problem 4.3 Analyze Identify the relevant concepts. 1

106 4.1 Defining the Atom > 106 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 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. Sample Problem 4.3 Solve Apply the concepts to this problem. 2

107 4.1 Defining the Atom > 107 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 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. Sample Problem 4.3 Solve Apply the concepts to this problem. 2

108 4.1 Defining the Atom > 108 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Use the symbol, atomic number, and mass number to write the symbol of the isotope. Sample Problem 4.3 Solve Apply the concepts to this problem. 2 For carbon-12, the symbol is C. For carbon-13, the symbol is C

109 4.1 Defining the Atom > 109 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Why are atoms with different numbers of neutrons still considered to be the same element? Despite differences in the number of neutrons, isotopes are chemically alike. They have identical numbers of protons and electrons, which determine chemical behavior.

110 4.1 Defining the Atom > 110 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Atomic Mass –How do you calculate the atomic mass of an element?

111 4.1 Defining the Atom > 111 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 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 x 10 –23 g.

112 4.1 Defining the Atom > 112 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 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.

113 4.1 Defining the Atom > 113 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Atomic Mass An atomic mass unit (amu) is defined as one-twelfth of the mass of a carbon-12 atom. This isotope of carbon has been assigned a mass of exactly 12 atomic mass units.

114 4.1 Defining the Atom > 114 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Atomic Mass The six protons and six neutrons account for nearly all of this mass. Therefore, the mass of a single proton or a single neutron is about one-twelfth of 12 amu, or about 1 amu. A carbon-12 atom has six protons and six neutrons in its nucleus, and its mass is set at 12 amu.

115 4.1 Defining the Atom > 115 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 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 (relative abundance).

116 4.1 Defining the Atom > 116 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Interpret Data Natural Percent Abundance of Stable Isotopes of Some Elements NameSymbol Natural percent abundance Mass (amu)Atomic mass Hydrogen HHHHHH negligible Helium He Carbon CCCC Oxygen OOOOOO Chlorine Cl

117 4.1 Defining the Atom > 117 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Atomic Mass Chlorine occurs as two isotopes: chlorine-35 and chlorine-37. If you calculate the arithmetic mean of these two masses (( amu amu)/2), you get an average atomic mass of However, this value is higher than the actual value of

118 4.1 Defining the Atom > 118 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Atomic Mass Chlorine occurs as two isotopes: chlorine-35 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 accounts for only 24 percent.

119 4.1 Defining the Atom > 119 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Atomic Mass Because there is more chlorine-35 than chlorine- 37 in nature, the atomic mass of chlorine, amu, is closer to 35 than to 37.

120 4.1 Defining the Atom > 120 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 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.

121 4.1 Defining the Atom > 121 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Understanding Relative Abundance of Isotopes The atomic mass of copper is amu. Which of copper’s two isotopes is more abundant: copper-63 or copper-65? Sample Problem 4.4

122 4.1 Defining the Atom > 122 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. The atomic mass of an element is the weighted average mass of the atoms in a naturally occurring sample of the element. Sample Problem 4.4 Analyze Identify the relevant concepts. 1

123 4.1 Defining the Atom > 123 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Compare the atomic mass to the mass of each isotope. The atomic mass of amu is closer to 63 than it is to 65. Sample Problem 4.4 Solve Apply the concepts to this problem. 2

124 4.1 Defining the Atom > 124 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Determine the most abundant isotope based on which isotope’s mass is closest to the atomic mass. Because the atomic mass is a weighted average of the isotopes, copper-63 must be more abundant than copper-65. Sample Problem 4.4 Solve Apply the concepts to this problem. 2

125 4.1 Defining the Atom > 125 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Atomic Mass You can determine atomic mass based on relative abundance. To do this, you must know three things: the number of stable isotopes of the element, the mass of each isotope, and the natural percent abundance of each isotope.

126 4.1 Defining the Atom > 126 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 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.

127 4.1 Defining the Atom > 127 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Atomic Mass Carbon has two stable isotopes: carbon-12, which has a natural abundance of percent, and carbon-13, which has a natural abundance of 1.11 percent. The mass of carbon-12 is amu; the mass of carbon-13 is amu. The atomic mass of carbon is calculated as follows: Atomic mass of carbon = ( amu x ) amu x ) = ( amu) + (0.144 amu) = amu

128 4.1 Defining the Atom > 128 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Calculating Atomic Mass Element X has two naturally occurring isotopes. The isotope with a mass of amu ( 10 X) has a relative abundance of percent. The isotope with a mass of amu ( 11 X) has a relative abundance of percent. Calculate the atomic mass of element X. Sample Problem 4.5

129 4.1 Defining the Atom > 129 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Sample Problem 4.5 The mass each isotope contributes to the element’s atomic mass can be calculated by multiplying the isotope’s mass by its relative abundance. The atomic mass of the element is the sum of these products. Analyze List the knowns and the unknown. 1 atomic mass of X = ? KNOWNSUNKNOWN Isotope 10 X: mass = amu relative abundance = 19.91% = Isotope 11 X: mass = amu relative abundance = 80.09% =

130 4.1 Defining the Atom > 130 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Use the atomic mass and the decimal form of the percent abundance to find the mass contributed by each isotope. for 10 X: amu x = amu for 11 X: amu x = amu Sample Problem 4.5 Calculate Solve for the unknowns. 2

131 4.1 Defining the Atom > 131 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Add the atomic mass contributions for all the isotopes. Sample Problem 4.5 Calculate Solve for the unknowns. 2 For element X, atomic mass = amu amu = amu

132 4.1 Defining the Atom > 132 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Sample Problem 4.5 Evaluate Does the result make sense? 3 The calculated value is closer to the mass of the more abundant isotope, as would be expected.

133 4.1 Defining the Atom > 133 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 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.

134 4.1 Defining the Atom > 134 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Key Concepts –Elements are different because they contain different numbers of protons. –Because isotopes of an element have different numbers of neutrons, they also have different mass numbers. –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.

135 4.1 Defining the Atom > 135 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Key Equation number of neutrons = mass number – atomic number

136 4.1 Defining the Atom > 136 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Glossary Terms –atomic number: the number of protons in the nucleus of an atom of an element –mass number: the total number of protons and neutrons in the nucleus of an atom –isotopes: atoms of the same element that have the same atomic number but different atomic masses due to a different number of neutrons

137 4.1 Defining the Atom > 137 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Glossary Terms –atomic mass unit (amu): a unit of mass equal to one-twelfth the mass of a carbon- 12 atom –atomic mass: the weighted average of the masses of the isotopes of an element

138 4.1 Defining the Atom > 138 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. –Atoms of the same element have the same number of protons, which is equal to an atom’s atomic number. –But atoms of the same element can have different numbers of neutrons. –Atoms of the same element with different numbers of neutrons are isotopes. BIG IDEA Electrons and the Structure of Atoms

139 4.1 Defining the Atom > 139 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. END OF 4.3


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