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Electrons in Atoms Chapter 5. Chapter Big Idea The atoms of each element have a unique arrangement of electrons.

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Presentation on theme: "Electrons in Atoms Chapter 5. Chapter Big Idea The atoms of each element have a unique arrangement of electrons."— Presentation transcript:

1 Electrons in Atoms Chapter 5

2 Chapter Big Idea The atoms of each element have a unique arrangement of electrons.

3 Section 1: Light and Quantized Energy

4 Essential Questions & Vocabulary How do the wave and particle natures of light compare? What is a quantum of energy and how is it related to an energy change of matter? How do continuous electromagnetic spectra and atomic emission spectra compare and contrast? Electromagnetic Radiation Wavelength Frequency Amplitude Electromagnetic Spectrum Quantum Planck’s constant Photoelectric effect Photon Atomic emission spectrum Vocabulary

5 Unanswered Questions In Rutherford’s model, the atom’s mass is concentrated in the nucleus and electrons move around it.

6 Unanswered Questions This model doesn’t explain how the electrons were arranged around the nucleus. This model also doesn’t explain why negatively charged electrons aren’t pulled into the positively charged nucleus.

7 The Wave Nature of Light In the early 1900s, scientists observed that certain elements emitted visible light when heated in a flame. Analysis of this light revealed that an element’s chemical behavior is related to the arrangement of the electrons in its atoms.

8 The Wave Nature of Light Visible light is a type of electromagnetic radiation. Electromagnetic radiation: a form of energy that exhibits wave-like behavior as it travels through space. Light is a type of energy that travels through space at a constant speed of 3.0 × 10 8 m/s (186,000 mi/s).

9 The Wave Nature of Light The electromagnetic spectrum includes all forms of electromagnetic radiation:

10 The Wave Nature of Light All waves can be described by several characteristics: The wavelength (λ) is the shortest distance between adjacent wave crests.

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13 The Wave Nature of Light The frequency (v) is the number of waves that pass a given point per second. Wavelength and frequency are inversely related—the shorter the wavelength, the higher the frequency.

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15 The Wave Nature of Light The amplitude of a wave is the wave’s height from the origin to the crest, or from the origin to a trough. Wavelength and frequency DO NOT affect the amplitude of a wave.

16 The Wave Nature of Light The amplitude is the wave’s height from the origin to a crest.

17 Wavelength & Frequency c = 3.0 x 10 8 m/s Speed of ALL light waves ν = frequency Measured in hertz (Hz) or inverse seconds (s -1 ) λ = wavelength Measured in meters If nm, multiply by 10 -9

18 The Particle Nature of Light The wave model of light CANNOT explain all of light’s characteristics. Example: Why heated objects emit only certain frequencies of light at a given temperature.

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20 The Particle Nature of Light In 1900, German physicist Max Planck (1858-1947) began searching for an explanation by studying the light emitted by heated objects.

21 The Particle Nature of Light Planck’s study led him to a startling conclusion: Matter can gain or lose energy only in small, specific amounts called quanta. A quantum is the minimum amount of energy that can be gained or lost by an atom.

22 The Particle Nature of Light

23 Prior belief is that energy could be absorbed and emitted in continually varying quantities Heating Water: The water’s temperature increases by infinitesimal steps as its molecules absorb quanta of energy Because the steps are so small, the temperature seems to rise in continuous rather than a stepwise manner.

24 The Particle Nature of Light The energy of a quantum is given by the product of Planck’s constant and the frequency. According to Planck’s theory, for a given frequency, v, matter can emit or absorb energy only in whole number multiples of hv

25 The Particle Nature of Light The photoelectric effect is when electrons are emitted from a metal’s surface when light of a certain frequency shines on it.

26 The Dual Nature of Light In 1905, Albert Einstein proposed that light has a dual nature A beam of light has wavelike and particlelike properties. A photon is a particle of electromagnetic radiation with no mass that carries a quantum of energy.

27 The Dual Nature of Light A photon of red light (relatively long wavelength) carries less energy than a photon of blue light (relatively short wavelength) does.

28 The Dual Nature of Light Every object gets its color by reflecting a certain portion of incident light. The color is determined by the wavelength of the reflected photons, thus by their energy. What is the energy of a photon from the violet portion of the Sun’s light if it has a frequency of 7.230 x 10 -14 Hz?

29 Atomic Emission Spectra The atomic emission spectrum of an element is the set of frequencies of the electromagnetic waves emitted by the electrons as they come back down to the ground state. The energy of the photon being released corresponds to different wavelengths of light. A higher energy photon might be blue in color while a lower energy photon might be red in color.

30 Atomic Emission Spectra Light in a neon sign is produced when electricity is passed through a tube filled with neon gas. The neon atoms become excited. The excited atoms return to their stable state by emitting light to release energy.

31 Atomic Emission Spectra A white-light spectrum is continuous, with some radiation emitted at every wavelength. The emission spectrum of an individual element includes only certain specific wavelengths.

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33 Atomic Emission Spectra Each element’s atomic emission spectrum is unique.

34 Section 2: Quantum Theory and the Atom

35 Essential Questions & Vocabulary How do the Bohr and quantum mechanical models of the atom compare? What is the impact of de Broglie’s wave-particle duality and the Heisenberg uncertainty principle on the current view of electrons in atoms? What are the relationships among a hydrogen atom’s energy levels, sublevels, and atomic orbitals? Ground state Quantum number De Broglie Equation Heisenberg uncertainty principle Quantum mechanical model of the atom Atomic orbital Principal quantum number Principal energy level Energy sublevel Vocabulary

36 Unanswered Questions Einstein’s theory of light’s dual nature accounted for several unexplainable phenomena by not why atomic emission spectra of elements were discontinuous rather than continuous.

37 Bohr’s Model of the Atom In 1913, Neils Bohr, a Danish physicist working in Rutherford’s laboratory, proposed a quantum model for the hydrogen atom that seemed to answer this question. Bohr correctly predicted the frequency lines in hydrogen’s atomic emission spectrum.

38 Bohr’s Model of the Atom Bohr suggested that an electron moves around the nucleus only in certain allowed circular orbits.

39 Ground state: the lowest allowable energy state of an atom Excited State: an atom that has gained energy Bohr’s Model of the Atom

40 Each orbit was given a number, called the quantum number. Hydrogen’s single electron in the n=1 orbit in the ground state. When energy is added, the electron moves to the n=2 orbit.

41 Bohr’s Model of the Atom Ground State: first energy level Atom does NOT radiate energy. When external energy is added, the electron moves to a higher- energy orbit (excited state) The electron eventually drops from the higher energy orbit to a lower- energy orbit. Emits a photon corresponding to the energy difference between the two energy levels

42 Balmer Series

43 Limits of Bohr’s Model It explained hydrogen’s spectral lines but failed to explain any other element’s spectral lines. Bohr’s model did not fully account for chemical behavior of atoms Behavior of electrons is still not fully understood but evidence indicates they do not move around the nucleus in circular orbits

44 De Broglie (1892-1987) If waves can have particlelike behavior, could the opposite also be true?

45 Quantum Mechanical Model Louis de Broglie hypothesized that particles, including electrons, could also have wavelike behaviors. Electrons orbit the nucleus only in whole-number wavelengths.

46 The de Broglie Equation predicts that all moving particles have wave characteristics. Impossible to notice the wavelength of a fast moving car but it can for an electron moving at the same speed. Quantum Mechanical Model

47 Werner Heisenberg (1901 – 1976) Balloon in Dark Room It is impossible to take any measurement of an object without disturbing the object. Heisenberg Uncertainty principle states that it is fundamentally impossible to know precisely both the velocity and position of a particle at the same time.

48 Quantum Mechanical Model When a photon interacts with an electron, both the velocity and the position of the electron are modified. The only quantity that can be known is the probability for an electron to occupy a certain region around the nucleus.

49 Erwin Schrodinger (1897 – 1961) Furthered the wave-particle theory by deriving an equation that treated the hydrogen’s electron as a wave. Unlike Bohr, Schrodinger’s equation also applied to other elements. Quantum Mechanical model of the atom – model in which electrons are treated as waves. Does NOT attempt to describe electron’s path; simply describes the PROBABILITY of finding an electron.

50 Atomic Orbitals A three-dimensional area around the nucleus which describes the electron’s probable location. Each dot represents the electron’s location at an instant in time. The high density of dots near the nucleus indicates the electron’s most probable location.

51 Atomic Orbitals Principal Quantum Number – (n) indicates the relative size and energy of atomic orbitals As n increase, the orbital size and energy increase and electron spends more time away from the nucleus Principal Energy Level – (n) specifies the atom’s major energy levels Energy Sublevels are contained within the principal energy levels.

52 Orbital Shapes Each energy sublevel relates to orbitals of different shapes s (1), p (3), d (5), or f (7) Each orbital can contain, at most, 2 electrons

53 Principal Energy Levels Principal Quantum Number (n) Sublevels (Types of Orbitals) Number of Orbitals related to sublevel Total number of Orbitals related to Principal Level (n 2 ) Total Number of electrons by sublevel Total Number of Electrons in Principal Energy Level (2n 2 ) 1s1122 2 spsp 1313 4 2626 8 3 spdspd 135135 9 2 6 10 18 4 spdFspdF 13571357 16 2 6 10 14 32

54 Section 3: Electron Configuration

55 Essential Questions & Vocabulary How are the Pauli exclusion principle, the aufbau principle, and Hund’s rule used to write electron configurations using orbital diagrams and electron configuration notation? What are valence electrons, and how do electron-dot structures represent an atom’s valence electrons? Electron configuration Aufbau principle Pauli exclusion principle Hund’s Rule Valence electron Electron-dot structure Vocabulary

56 Ground-State Electron Configuration Electron Configuration – the arrangement of electrons in an atom. Low-energy systems are more stable than high- energy systems Electrons in an atom tend to arrange in a manner that gives the atom the LOWEST energy possible.

57 aufbau Principle German roots meaning “to build up” or “arrange” Each electron occupies the lowest energy orbital available

58 Aufbau Diagrams Describes the sequence in which orbitals are filled with electrons, however, atoms are NOT built up electron by electron

59 Pauli Exclusion Principle A maximum of two electrons can occupy a single atomic orbital, but only if the electrons have opposite spins.

60 Hund’s Rule Single electrons with the SAME spin must occupy each equal-energy orbital before additional electrons with opposite spins can occupy the same orbitals. Remember: orbitals within the same sublevel have the same energy.

61 Electron Configuration: Neon Orbital Diagram

62 Electron Configuration Notation Designates the principal energy level and energy sublevel associated with each of the atom’s orbitals and includes a superscript representing the number of electrons in the orbital.

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64 Electron Configuration Noble-Gas Notation A shorthand method that uses bracketed symbols around the noble gas symbol that appears last before the element of interest.

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