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Before you begin this chapter, review the following concepts and skills: Atoms and electrons ■ modern model of the atom. You will also learn about the model itself and how it relates to periodic trends and the periodic table. identifying subatomic particles and their properties (from previous studies) green colours in the northernit lights comeresult of streams of The northern lights, also called theThe aurora borealis, the from the interaction of accelerated electrons withWhen oxygen atoms.the The redelectrons colours usually protons and electrons from the Sun. interact with come from nitrogen atoms. How do these describing the structure interactions lead to atoms of different elements gaseous atoms and in organization Earth’sof the upper atmosphere, the atoms emit colored light. emitting light with characteristic wavelengths? ■ periodic table (from previous studies) ■ explaining periodic trends for properties such as atomic radius, first ionization energy, and electron affinity (from previous studies) 118 MHR • Unit 2 Structure and Properties In the twentieth century scientists developed a revolutionary new model of the atom. This model helped to explain many phenomena that the existing theories had failed to explain. n modified form, both of these inventions are in every chemistry course around the world. ed the periodic table, he was well-acquainted He knew nothing, however, about subatomic electron, which is the foundation for the nctive shape. Because the original periodic imental observations, chemists did not need structure to develop it. (As you will see in iodic table easily accommodates details about will learn that the modern periodic table’s l consequence of atomic structure.) spectrum, quantum, photons The atomic model - Dalton In your notebook, list the main ideas in Dalton’s atomic theory. Explain how this theory enabled chemists to explain the three mass laws: the law of conservation of mass, the law of definite proportions, and the law of multiple proportions. John Dalton’s atomic theory (1809) suggested that the atom was a tiny, solid uncuttable sphere. e Modern Atomic Model omic theory to advance their understanding uring chemical reactions. His atomic model, explaining the behaviour of substances. d a system of symbols to show how atoms ances. Figure 3.2 on the next page shows you will no doubt notice, Dalton correctly rbon dioxide and sulfur trioxide, but ran into mmonia, and methane. Dalton’s attempt at hts a crucial limitation with his atomic se it to explain why atoms of elements ch they do. This inability did not prevent studies. It did, however, suggest the need omic model. The model of the atom in 1809. The atom, as Dalton pictured it, was a tiny, solid, indestructible sphere. Figure 3.1 Dalton’s atomic theory no longer applies in its original form Chapter 3 Atoms, Electrons, and Periodic Trends • MHR 119 The atomic model - Dalton Dalton’s atomic model, however, was inadequate for explaining the behavior of substances. He designed a system of symbols to show how atoms combine to form other substances. hydrogen sulfur nitrogen water carbon dioxide sulfur trioxide carbon ammonia oxygen methane Examples of Dalton’s system of symbols for atoms and molecules Figure 3.2 Examples of Dalton’s system of symbols for atoms and molecules Dalton correctly predicted the formulas for carbon dioxide and sulfur The Discovery of the Electron Requires a New Atomic Model trioxide,Inbut ran into serious trouble with water, ammonia, and methane 1897, Dalton’s idea of an indivisible atom was shattered with a startling announcement. A British scientist, Joseph John Thomson, had discovered 1 the existence of a negatively charged particle with mass less than 1000 that of a hydrogen atom. This particle was, of course, the electron. 1 (Later calculations showed that the mass of an electron is 1837 that The atomic model - Thomson oxygen methane Figure 3.2 Examples of Dalton’s system of symbols for atoms and molecules The Discovery of the Electron Requires a New Atomic Model In 1897, Dalton’sIn 1897, ideaDalton’s of an shattered with a startling ideaindivisible of an indivisibleatom atom waswas shattered with a startling A British scientist, JosephJohn John Thomson, had discovered announcement. Aannouncement. British scientist, Joseph Thomson, had discovered the 1 the existence of a negatively charged particle with mass less than 1000 existence of a negatively charged with mass1000 that of a hydrogen atom. Thisparticle particle was, of course, the electron. times smaller than 1 (Later calculations showed that the mass of an electron is 1837 that the mass of a hydrogen atom. This particle was the electron. of a hydrogen atom.) It took several years for chemists to consider the consequences of this discovery. They realized that if atoms contain electrons, atoms must also contain a positive charge of some kind to balance the negative charge. The atomic model that Thomson eventually proposed is shown in Figure 3.3. Keep in mind that scientists had not yet discovered the proton. Therefore, in this model, the entire sphere carries a uniform, positive charge. Chemists realized that if atoms contain electrons, atoms must also contain a positive charge of some kind to balance the negative charge. This was the atomic model that Thomson proposed. electrons Negative electrons CHEM FA C T Scientists initially described adioactivity solely in terms of radiation. The idea of adioactive particles first appeared around the turn of he twentieth century. In 1909, Ernest Rutherford reported confidently that the alpha positively Positively charged sphere Figure 3.3 The atomic model in 1903. Thomson viewed the atom as a positively charged sphere embedded with sufficient numbers of electrons to balance (neutralize) the total charge. The atomic model - Rutherford • End of XIX century Marie and Pierre Curie discover “radioactive” elements: element whose atoms emit positively charged particles (α-particles), negatively charged particles (β-particles), and energy (γ-radiation) • From 1898 to 1907 Ernest Rutherford studies the chemistry of radioactive elements • In 1909 Rutherford and two students perform the gold-foil experiment and demolish Thomson’s atomic model rutherford-scattering simulation The atomic model - Rutherford gold-foil experiment, shown in Figure 3.4. On the basis of this experiment, Rutherford suggested that the deflections he and his students observed were caused by an encounter between an alpha particle and an intense electric field at the centre of the atom. The gold-foil experiment A Hypothesis: Expected result based on Thomson’s model incoming α particles B Experiment C 1 Radioactive sample emits beam of α particles 2 Beam of α particles meets gold foil little or no deflection cross-section of gold foil showing Thomson’s atoms gold foil 5 Major deflections are observed rarely 4 Minor deflections (exaggerated here) are observed occasionally Figure 3.4 The hypothesis, experiment, and results of Rutherford’s gold foil experiment. The experimental hypothesis and design owed much to the contributions of Rutherford’s students, Hans Geiger (of Geiger-counter fame) and Ernest Marsden. Rutherford performed several calculations that led him to an inescapable Actual Result incoming α particles major deflection minor deflection cross-section of gold foil composed of atoms with tiny, massive nucleus 3 Flashes of light observed when α particles strike screen show that most are transmited with little or no deflection The atomic model - Rutherford The gold-foil experiment 1. Rutherford aimed α-particles (+) at extremely thin metal foils. 2. A small number of the α-particles were deflected significantly (∡>90°) by the atoms of the metal foils. 3. Following Thomson’s model, Rutherford expected the αparticles to pass through the metal atoms with low deflection. 4. Rutherford suggested that the deflections he observed were caused by an encounter between an α-particles and an intense electric field at the centre of the atom. The atomic model - Rutherford Rutherford performed several calculations that led him to an inescapable conclusion: “the atom is made up mainly of empty space, with a small, massive region of concentrated charge at the centre”. Soon afterward, the charge on this central region was determined to be positive, and was named the atomic nucleus. Because Rutherford’s atomic model pictures electrons in motion around an atomic nucleus, chemists often call this the nuclear model of the planetary atom. The atomic model - Bohr in 1913 Niels Bohr, a Danish physicist and student of Rutherford, proposed a new model for the hydrogen atom. Bohr’s atomic model pictures electrons in discrete orbit around a central nucleus, unlike Rutherford’s model, in which electrons may move anywhere within the space around the nucleus. The Bohr Atomic Model 1. Each atom has specific energy levels, corresponding to the atom’s electrons orbits. 2. While in one of its energy levels, electrons do not emit energy. 3. Electron change energy level by emitting or absorbing a specific quantity of energy. that the frequency symbol, ν, is the Greek letter nu, not the letter v.) λ 8s (8 Hz) The atomic model - The Problem of Atomic Spectra C The amplitude ( wave λ wavelength (nm) 10−2 100 102 104 16 s−1 (16 Hz) 106 108 1010 A Despite differences in wave represent A wave with a h amplitude has a 1012 intensity (is brig wave with a low visible ultra-wavelengths and frequency, gamma Infrared infrared x-ray violet microwave radio frequency ray all electromagnetic radiation rays travelsFigure at the3.6 sameThe speed electromagnetic spectrum and its properties in a vacuum: 3.00 × 10−8 m/s. 12 1010 1020 1018 1016 1014 108 106 104 10 1012 This value, the speed of light, −1 The visible portion is a constant represented by of the electromagnetic spectrum is called a continuous frequency (s ) the symbol c. spectrum, because the component colours are indistinct. They appear 400 500 600 700 750 nm “smeared” together into a continuum of colour. According to nineteenthcentury physics, part of the energy emitted by electrons should be observable as a continuous spectrum. This is not, however, the case. Instead, when atoms absorb energy (for example, when they are exposed visible light region to an electric current), you observe a pattern of discrete (distinct), coloured lines separated by spaces of varying length. See Figure 3.7. You The relationship between B higher lower by When atoms absorb energy, we this canline observe a pattern of distinct, linesinseparated frequency (ν) in s−1and (Hz) forcolored can also observe spectrum for hydrogen, other atoms, frequency (symbol ν, measured amplitude amplitude spaces of varying length. (brighter) (dimmer) 1 second in Hz)Investigation and wavelength (symbol λ, 3-A. usually measured in nm or m, wavelength λ depending As 400 on the energy). 500 600 750 nm 4 s−1 (4 Hz) wavelength increases, frequency decreases. As wavelength decreases, frequency increases. 3.7= 1The This1spectrum characteristic of lines H atoms. Nospectrum other atoms display this pattern coloured of this are characteristic of (Note Figure that Hz s−1is .discrete, Note also −1 (8 Hz) 8 s λ that the frequency symbol, ν, of coloured lines. hydrogen atoms. No other atoms display this pattern of coloured lines. wavelength, λ is the Greek letter nu, not the letter v.) C The amplitude (height) of a wave represents its intensity. 16 s−1 (16 Hz) λ A wave with a higher amplitude has a greater intensity (is brighter) than a The atomic model - The Problem of quanta In 1900, a physicist named Max Planck suggested that matter, at the atomic level, can absorb or emit only discrete quantities of energy. Each of these specific quantities is called a quantum of energy. The era of quantum physics was beginning. A quantum is an extremely small “packet” of energy. quantum n=4 n=3 n=2 n=1 Energy levels The atomic model - The Problem of photons region In 1905, Albert Einstein theorized that light comesvisible in energylight packets (quanta) which are called photons that have particle-like properties. Light is emitted as photons of energy. ured mbol λ, m, As uency ases. e also ν, he frequency (ν) in s−1 (Hz) 1 second wavelength λ 4 s−1 (4 Hz) ph λ 8 s−1 (8 Hz) How Bohr’s Atomic Model Explains the Spectrum Bohr proposed that the energy that is emitted and absorbed by an atom must have specific values. He was applying the new ideas of quanta in his model. The Bohr Atomic Model 1. Each atom has specific energy levels, corresponding to the atom’s electrons orbits. 2. While in one of its energy levels, electrons do not emit energy. 3. Electron change energy level by emitting or absorbing a specific quantity of energy (quanta) century physics, part of the energy emitted by electrons should be observable as a continuous spectrum. This is not, however, the case. Instead, when atoms absorb energy (for example, when they are exposed againyou theobserve spectrum H, the(distinct), energy that is associated with toAnalyzing an electric current), a patternfor of discrete coloured lines separated by spaces of varying length. See Figure 3.7. You colored lines in this spectrum corresponds to the change in energy of can also observe this line spectrum for hydrogen, and for other atoms, in electron 3-A. as it moves to higher or lower energy levels. Investigation How Bohr’s Atomic Model Explains the Spectrum 400 500 600 the an 750 nm The discrete, coloured lines of this spectrum are characteristic of hydrogen atoms. No other atoms display this pattern of coloured lines. Figure 3.7 When a H atom is exposed to an electromagnetic radiation, its electron absorbs photons. en=2 n=3 n=4 n=5 n=6 The atom is now said to be in an excited state. Chapter 3 Atoms, Electrons, and Periodic Trends • MHR 123 When an electron is excited to a higher energy level, it falls back to the original energy level and emits light of certain energy. E.g., an electron that falls from the third energy to the second energy level emits a photon of red light with a λ of 656 nm. spectra are called emission spectra. century physics, partalso of the energy emitted by electrons should be 3.9 shows the energy transitions that are the responsible observable asFigure a continuous spectrum. This is not, however, case. for the coloured lines in hydrogen’s emission spectrum. Notice use of the Instead, when atoms absorb energy (for example, when they are the exposed symbol n to designate the allowed energy levels for lines the hydrogen atom: emission spectrum. energy transitions responsible for colored in hydrogen’s toThe an electric current), youare observe a pattern ofthe discrete (distinct), n = 1, n = 2, and so on. This symbol, n, represents a positive integer (such coloured lines “n” separated by spaces of varying length. Seefor Figure 3.7. You The symbol the possible energy levels an hydrogen as 1, 2, ordesignate 3), and is called a quantum number. You will learn moreatom: aboutn=1, n=2, and so on. can also observe this line spectrum for hydrogen, and for other atoms, in the significance of quantum numbers in section 3.2. Investigation 3-A. This symbol, “n” is called a quantum number. How Bohr’s Atomic Model Explains the Spectrum 400 400 500 600 600 500 750 nm 750 nm n=6 n=5 n=4 The discrete, coloured lines of this spectrum are characteristic of hydrogen atoms. No other atoms display this pattern of coloured lines. Figure 3.7 n=3 e-n = 1 n=2 n=3 n=4 n=2 n=3 n=4 n=5 n=6 A Note: Orbits are not n=5drawn to scale. n=6 n=2 123 energy Chapter 3 Atoms, Electrons, and Periodic Trends • MHR n=1 InBitsNote: unexcited hydrogen‘s electron lays inelectron the first is in In itsstate, unexcited state, hydrogen‘s energy Thislevel is theclosest lowest-possible energy n level, thelevel. energy to the nucleus: = 1. This is representing a state of greatest stability for the hydrogen the lowest-possible energy level, representing a atom state of greatest stability for the hydrogen atom.