The Atom & Unanswered Questions:

Transcription

1

2 The Atom & Unanswered Questions: 1) Recall-Rutherford s model, that atom s mass is concentrated in the nucleus & electrons move around it. a) Doesn t explain how the electrons were arranged around the nucleus b) Doesn t explain why negatively charged electrons aren t pulled into the positively charged nucleus

3 The Atom & Unanswered Questions: 2) Early 1900 s, scientists observed certain elements emitted visible light when heated in a flame 3) Analysis of the emitted light revealed that an element s chemical behavior is related to the arrangement of the electrons in its atoms.

4 Note:Light and electrons in atoms have some properties in common so let s learn the basics about light

5 Electromagnetic Radiation (ER): 1) Radiation = a general term for any type of energy that emanates or radiates outward in all directions 2) Electromagnetic radiation (ER) = radiation moving at the speed of light; a form of energy that exhibits wavelike behavior as it travels through space.

6 Electromagnetic Radiation (ER): Electromagnetic spectrum = all the forms of electromagnetic radiation Microwaves

7 Electromagnetic Radiation (ER): 4) Forms of ER included in the electromagnetic spectrum: a) Gamma rays b) X-rays c) Ultraviolet light d) Visible light e) Infrared light f) Microwaves g) Radio waves

8 Electromagnetic Radiation (ER): 5) Properties that all ER has in common: a) Travels at the speed of light ( m/s) b) Emitted by atoms after they are energized or as they decay c) Acts like both a wave and a particle

9 THE WAVE THEORY-The WAVE Nature of Light: 1) All waves can be described by several characteristics: a) Wavelength ( λ ) = the distance between corresponding points on adjacent waves b) Frequency ( ν ) = the number of waves that pass a given point in a specific time, usually one second

10 THE WAVE THEORY-The WAVE Nature of Light: 2) Continuous spectrum = contains every wavelength between the wavelength on which the spectrum starts and the wavelength on which the spectrum ends (no gaps)

11 THE WAVE THEORY-The WAVE Nature of Light: 3) Visible spectrum = all the colors of light we can see (rainbow) a) When white light passes through a prism; it is separated into a continuous spectrum of its different components (red, orange, yellow, green blue, indigo, and violet light).

12 THE WAVE THEORY-The WAVE Nature of Light: b) The frequency of light varies; increasing energy increases the frequency. Ex: violet light (greater frequency) has more energy than red light (lower frequency) Red Violet

13 THE PARTICLE THEORY- The PARTICLE Nature of Light: 1) The wave model of lightcannot explain all of light s characteristics. a) Why do heated objects emit only certain frequencies of light at a given temperature or why some metals emit electrons when light of a specific frequency shines on them? b) A new model was needed to address these phenomena.

14 THE PARTICLE THEORY- The PARTICLE Nature of Light: 2) When objects are heated, they emit glowing light.

15 By studying glowing metals, Max Planck discovered that only certain wavelengths of light are emitted at each specific temperature. THE PARTICLE THEORY- The PARTICLE Nature of Light: By studying glowing metals, Max Planck discovered that only certain wavelengths of light are emitted at each specific temperature.

16 THE PARTICLE THEORY- The PARTICLE Nature of Light: 4) Because temperature is a measure of energy, Planck postulated that each wavelength emitted also represents a certain energy. a) Planck concluded that energy is quantized: that it only comes in discrete packets. b) He empirically found that at each temperature the energy was related to the frequency of light by an equation.

17 THE PARTICLE THEORY- The PARTICLE Nature of Light: 5) Quantum = a discrete packet of energy that can be gained or lost by an atom 6) In the photoelectric effect, when light hits a metal surface, electrons are ejected.

18 THE PARTICLE THEORY- The PARTICLE Nature of Light: a) This only happens when light of a minimum of frequency is used. b) Lower frequency light will not eject electrons. c) Albert Einstein explained the photoelectric effect by saying that light acts like particles which he called photons.

19 THE PARTICLE THEORY- The PARTICLE Nature of Light: 7) Photon = a particle of electromagnetic radiation with no rest mass that carries a quantum of energy.

20 THE PARTICLE THEORY- The PARTICLE Nature of Light: 8) Mystery of the photoelectric effect- When a light s frequency was BELOW a certain minimum (depending on the metal), the photoelectric effect was not observed, no matter how long the light shone on a metal. a) The wave theory of electromagnetic radiation DID NOT explain this!!!!!

21 Dual Wave-Particle Nature of Light: 1) Albert Einstein proposed in 1905 that light (and all ER) has a dual wave-particle nature; sometimes a beam of light exhibits wavelike and particle-like properties

22 Dual Wave-Particle Nature of Light: a) In some experiments, light acts like waves. Ex: Diffraction and interference of light are explained by a wave model of electromagnetic radiation. b) In some experiments, light acts likeparticles. Ex: The photoelectric effect is explained by a particle model of electromagnetic radiation.

23 Dual Wave-Particle Nature of Light: 2) Both a wave model and a particle model are necessary to explain all observations about electromagnetic radiation. Both are used!

24 Atomic Line-Emission Spectrum: 1) The wave model of light cannot explain how is light produced in glowing tubes of neon signs 2) Ground state = lowest state of energy for an atom 3) Excited state = state in which an atom has a 3) Excited state = state in which an atom has a higher potential energy than it has in its ground state

25 Atomic Line-Emission Spectrum: 4) Light in a neon sign is produced when electricity is passed through a tube filled with neon gas and excites the neon atoms a) The excited atoms emit light to release energy.

26 Atomic Line-Emission Spectrum: 5) When an excited atom returns to its ground state or a lower energy state, it gives off energy in the form of ER. 6) Atomic line-emission spectrum = the set of frequencies of the electromagnetic waves emitted by the atoms of the element.

27 Atomic Line-Emission Spectrum: 7) If the light emitted by neon is passed through a glassprism, neon s atomic line-emission spectrum is produced.

28 Atomic Line-Emission Spectrum: 8) Each element s atomic line-emission spectrum is unique and can be used to identify an element.

29 Atomic Line-Emission Spectrum: 9) HYDROGEN S line-emission spectrum is separated into four specific colors of the visible spectrum

30 Atomic Line-Emission Spectrum: 10) Why can you see just the 4 specific frequencies of light (lines) & not a continuous spectrum (rainbow) after being shone through a prism? 11) Niels Bohr attempted to answer this?

31 The Bohr Model: 1) The Bohr Model ONLY explained hydrogen s observed line-emission spectrum and no other element. a) The single electron in a hydrogen atom can circle the nucleus only in allowed paths (orbits). b) A hydrogen atom has a definite fixed energy for each possible electron orbit.

32 The Bohr Model: When a hydrogen atom is in its ground state, its electron is in itslowest energy orbit, closest to the nucleus. When a hydrogen atom is in an excited state, its electron is in a higher energy orbit, farther away from the nucleus.

33 The Bohr Model: 2) How does the Bohr Model explain the lineemission spectrum of hydrogen? a) It is only possible for the electron in a hydrogen atom to gain or lose certain quantities of energy, corresponding to the electron s allowed orbits.

34 The Bohr Model: b) When the electron falls from a higher energy orbit to a lower energy orbit it loses energy. (Energy does not just disappear. Where did it go?)

35 The Bohr Model: The atom emits aphoton whose energy corresponds to the energy difference between the two energy levels.(energy lost by electron = energy of released photon) Each frequency corresponds to a certain color or line on a line-emission spectrum

36 The Bohr Model: 3) Bohr Model = The electrons circle the nucleus only in allowed paths (orbits). 4) Illustration of the Bohr Model of calcium:

37 The Quantum Mechanical Model of the Atom: 1) In 1924 Louis de Broglie proposed that ELECTRONS have a duel wave-particle nature. Other experiments soon demonstrated wave properties of electrons.

38 The Quantum Mechanical Model of the Atom: 2) In 1926 Erwin Schrödinger treated electrons as waves in a model called the quantum mechanical model of the atom. a) Schrödinger s equation applied equally well to elementsother than hydrogen. b) Schrodinger s equation: helps determine probable electron location in an atom

39 The Quantum Mechanical Model of the Atom: 3) orbital = a three dimensional region around the nucleus that indicates the probable location of an electron. (fuzzy electron clouds) a) The cloud has no definite boundary, it is possible that the electron might be found at a considerable distance from the nucleus

40 Summary of the Models of the Atom: Quantum Mechanical Model (1926)

41

42 Quantum Numbers: Electrons are not locked into fixed orbits We can only predict the areas where electrons are most likely to be found Numbers are given to electrons to help with predictions

43 Quantum Numbers: 1) quantum numbers = specify the properties of the atomic orbitals and the properties of electrons in those orbitals

44 Quantum Numbers: 2) The first three quantum numbers result from solutions to Schrodinger s equation and describe the orbital in which an electron is located. Principal quantum number Angular momentum quantum number Magnetic quantum number

45 Quantum Numbers: 3) The fourth quantum number describes an electron in an orbital. Spin quantum number 4) Each electron is given four quantum numbers.

46 The 4 Quantum Numbers: PRINCIPAL 1) The PRINCIPAL QUANTUM NUMBER indicates the main energy level (shell) of the orbital in which a particular electron is located. a) Theprincipal quantum number is symbolized by n.

47 The 4 Quantum Numbers: PRINCIPAL b) Values of n = 1, 2, 3 (7 is the highest principal quantum number for any known element in its ground state). Corresponds to the 7rows on the Periodic Table

48 c) Ex: Energy Levels & the Periodic Table n=1 n=2 n=3 n=4 n=5 n=6 n=7 n=6 n=7

49 The 4 Quantum Numbers: PRINCIPAL d) An atom s LOWEST energy level is assigned a principal quantum number of 1. e) The higher the principal quantum number the further away from the nucleus the electron is f) An electron located in an orbital in a higher main energy level will have a higher energy.

50 The 4 Quantum Numbers: ANGULAR MOMENTUM 2) The ANGULAR MOMENTUM QUANTUM number indicates the shape of the orbital in which a particular electron is located. a)the angular momentum quantum number divides the main energy levels into smaller groups of orbitals called sublevels.

51 The 4 Quantum Numbers: ANGULAR MOMENTUM b) The angular momentum quantum number is symbolized by l. c) Angular momentum quantum numbers are usually designated with letters s, p, d, f d) The order of the sublevels can be remembered as follows: some people don t forget

52 The 4 Quantum Numbers: ANGULAR MOMENTUM e) Each energy sublevel relates to orbitals of different shape. s orbitals are sphere-shaped p orbitals are dumbbell-shaped d orbitals are double dumbbell-shaped f orbitals are flower-shaped

53 Orbital Shapes: f orbitals

54 The 4 Quantum Numbers: ANGULAR MOMENTUM f) Orbitals within each main energy level occupy sublevels (subshells). Ex: An electron located in a p orbital in the 2 nd main energy level would be said to be located in the 2p sublevel.

55 The 4 Quantum Numbers: ANGULAR MOMENTUM Energy levels can be thought of as rows of seats in a theater. The rows that are higher up and farther from the stage contain more seats. Similarly, energy levels related to orbitals farther from the nucleus contain more sublevels.

56 The 4 Quantum Numbers: ANGULAR MOMENTUM g) The two sublevels in main energy level 2 are designated 2s and 2p. The 2s sublevel corresponds to the 2s orbital, which is spherical like. The 2p sublevel corresponds to three dumbbell-shaped p orbital designated 2p x, 2p y, 2p z. Each of the p orbitals related to an energy sublevel has the same energy.

57 The 4 Quantum Numbers: Angular Momentum h) An electron located in different sublevels of the same energy level would differ in energy (for multielectron atoms). Ex: An electron located in the 4s sublevel has a lower energy than an electron located in the 4p sublevel which has a lower energy than an electron located in the 4d sublevel which has a lower energy than an electron located in the 4f sublevel.

58 The 4 Quantum Numbers: MAGNETIC 3) The MAGNETIC QUANTUM NUMBER indicates the spatial orientation of the orbital in which a particular electron is located. a) The magnetic quantum number is symbolized by m. b) The orientation of an orbital is designated using a three-dimensional coordinate system with the nucleus at the center.

59 The 4 Quantum Numbers: MAGNETIC c) Orientation of orbitals: An s orbital has 1 possible orientation (a sphere centered on the nucleus). A p orbital has 3 possible orientations. (p x, p y, p z ). A d orbital has 5 possible orientations. (d xz, d yz, d xy, d x 2 - y2, d z2 ) An f orbital has 7 possible orientations. (orientations not shown)

60 The 4 Quantum Numbers: MAGNETIC d) Ex: Orbitals around the Nucleus of a Neon Atom

61 The 4 Quantum Numbers: SPIN 4) The SPIN QUANTUM NUMBER indicates the spin of electron on its own axis a) The spin quantum number is symbolized by s. b) There are two possible fundamental states (spins) for an electron in an orbital. +½ and ½ are used to indicate the two possible states (spins) of an electron in an orbital.

62 Principal Quantum Number: Main Energy Level (n) Summary of the First 4 Energy Levels Type(s) of Sublevel (orbital shapes) # of Orbitals per main energy level Maximum # of Electrons per sublevel Number of Electrons per Main Energy Level (2n 2 ) 1 s s p s p d s p d f

63 WHERE ARE THE ELECTRONS (e-) LOCATED? 1) Electron Configuration = Arrangement of electrons in an atom a) Use atomic number of an element to indicate the number of electrons. b) AufbauPrinciple = an electron occupies the lowest-energy orbital that can receive it

64 WHERE ARE THE ELECTRONS (e-) LOCATED? c) Hund s Rule: Orbitals of equal energy are each occupied byone electron before any orbital is occupied by a second electron All electrons in singly occupied orbitals must have the same spin

65 WHERE ARE THE ELECTRONS (e-) LOCATED? d) Pauli s Exclusion Principle: No two electrons in the same atom can have the same four quantum numbers As a result, two electrons of opposite spin can occupy the same orbital

66 WHERE ARE THE ELECTRONS (e-) LOCATED? 2) Ground State Electron Configurations: a) Ground state electron configuration = lowest energy arrangement of electrons in an atom

67 WHERE ARE THE ELECTRONS (e-) LOCATED? b)according to the Quantum Mechanical Model, electrons are said to be located in orbitals. How should this be interpreted? How is this represented? c) Electrons in atoms tend to assume arrangements that have the LOWEST possible energies. d)the Aufbau principle, Hund s rule, and the Pauli s exclusion principle can be used to determine the lowest energy, ground state electron configuration for atoms. These are your rules/guidelines!

68 WHERE ARE THE ELECTRONS (e-) LOCATED? 3) Example Electron Configuration notation of Hydrogen: Hydrogen has 1 electron Main Energy Level 1s 1 # Electrons in a sublevel Sublevel

69 WHERE ARE THE ELECTRONS (e-) LOCATED? 4) Diagonal Rule = a tool used to help write electron configurations. oalways start with the LOWEST energy level (1s 2 ) and work your way through to the HIGHEST energy level (7p 6 ) by following the tail to head of each arrow. 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 4f 14 5s 2 5p 6 5d 10 5f 14 6s 2 6p 6 6d 10 7s 2 7p 6

70 WHERE ARE THE ELECTRONS (e-) LOCATED? 5) Example Electron Configuration of Na (11): 1s 2 2s 2 2p 6 3s 1

71 WHERE ARE THE ELECTRONS (e-) LOCATED? 6) Diagonal rule is based on trends of the Periodic Table.

72 WHERE ARE THE ELECTRONS (e-) LOCATED? Use the Diagonal Rule to find the electron configuration notation for the following: (Remember Aufbau sprinciple= an electron occupies the lowest-energy orbital that can receive it)

73 Ex: Mg (12): 1s 2 2s 2 2p 6 Start here and move along the arrows one by one. 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 4f 14 5s 2 5p 6 5d 10 5f 14 6s 2 6p 6 6d 10 7s 2 7p 6

74 Ex: P (15): 1s 2 2s 2 2p 6 Start here and move along the arrows one by one. 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 4f 14 5s 2 5p 6 5d 10 5f 14 6s 2 6p 6 6d 10 7s 2 7p 6

75 Ex: Ni (28): 1s 2 2s 2 2p 6 Start here and move along the arrows one by one. 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 4f 14 5s 2 5p 6 5d 10 5f 14 6s 2 6p 6 6d 10 7s 2 7p 6

76 Orbital Notation: 1) Shows the number of orbital(s) and the spin of electron(s). 2) Complete electron notation first 3) Orbitals are represented by circle(s): 4) Each orbital can only hold a maximum of 2 electrons 5) Electrons are represented by slash (/) marks

77

78 6) Remember: Orbital Notation: a) Paul s Exclusion Principle: Two electrons of opposite spin can occupy the same orbital. b) Hund s Rules: Orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron. All electrons in singly occupied orbitals must have the same spin.

79 Orbital Notation: Ex: Electron notation of sodium (11) = 1s 2 2s 2 2p 6 3s 1 Orbital notation of sodium =

80 Orbital Notation: 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 4f 14 5s 2 5p 6 5d 10 5f 14 Ex: O (8): 6s 2 6p 6 6d 10 7s 2 7p 6

81 Orbital Notation: 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 4f 14 5s 2 5p 6 5d 10 5f 14 Ex: Fe (26): 6s 2 6p 6 6d 10 7s 2 7p 6