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1 Chapter 7 Atomic Structure. 2 7.1 Types of EM waves They have different and They have different and n Radio waves, microwaves, infra red, ultraviolet,

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Presentation on theme: "1 Chapter 7 Atomic Structure. 2 7.1 Types of EM waves They have different and They have different and n Radio waves, microwaves, infra red, ultraviolet,"— Presentation transcript:

1 1 Chapter 7 Atomic Structure

2 2 7.1 Types of EM waves They have different and They have different and n Radio waves, microwaves, infra red, ultraviolet, x-rays and gamma rays are all examples. n Light is only the part our eyes can detect. Gamma Rays Radio waves Frequency decreases Wavelength increases

3 3 Parts of a wave Wavelength Frequency = number of cycles in one second Measured in hertz 1 hertz = 1 cycle/second

4 4 Frequency = Frequency =

5 5 The speed of light c = in a vacuum is x 10 8 m/s c = c = in a vacuum is x 10 8 m/s c = n What is the wavelength of light with a frequency 5.89 x 10 5 Hz? n What is the frequency of blue light with a wavelength of 484 nm? n (YDVD) EMwave

6 6 7.2 Nature of Matter n Matter and energy were seen as different from each other in fundamental ways. n Matter was particles. n Energy could come in waves, with any frequency. n Max Planck found that as the cooling of hot objects couldnt be explained by viewing energy as a wave.

7 7 Energy is Quantized Planck found E came in chunks with size h Planck found E came in chunks with size h E = n h E = n h n where n is an integer. n and h is Plancks constant n h = x Js These packets of h are called quantum. These packets of h are called quantum.

8 8 Einstein is next n Said electromagnetic radiation is quantized in particles called photons. Each photon has energy = h = h c/ Each photon has energy = h = h c/ n Blue color in fireworks has a wavelength of 450 nm. How much energy does a photon of blue light have? A mole of these photons? n Combine this with E = mc 2 n You get the apparent mass of a photon. m = h /( c) m = h /( c)

9 9 Which is it? n Is energy a wave like light, or a particle? n Yes n Concept is called the Wave - Particle duality. (Flatman) n What about the other way, is matter a wave? Yes n Photoelectric Effect (YDVD) n (BDVD) Blue & Yellow

10 10 Matter as a wave n All matter exhibits both particle and wave properties. De Broglies equation = h/mv De Broglies equation = h/mv n We can calculate the wavelength of an object. n The laser light of a CD is 7.80 x 10 2 m. What is the frequency of this light? n What is the energy of a photon of this light? n What is the apparent mass of a photon of this light? n What is the energy of a mole of these photons?

11 11 What is the wavelength? m = h /( c) m = h /( c) n Of an electron with a mass of 9.11 x kg traveling at 1.0 x 10 7 m/s? n Of a softball with a mass of 0.10 kg moving at 35 mi/hr?

12 Hydrogen spectrum n Emission spectrum because these are the colors it gives off or emits. n Called a line spectrum. n There are just a few discrete lines showing only certain energies are possible. 410 nm 434 nm 486 nm 656 nm

13 13 What this means n Only certain energies are allowed for the hydrogen atom. n Can only give off certain energies. Use E = h = h c/ Use E = h = h c/ n Energy in the in the atom is quantized. n (YDVD) Refraction

14 Niels Bohr n Developed the quantum model of the hydrogen atom. n He said the atom was like a solar system. n The electrons were attracted to the nucleus because of opposite charges. n Didnt fall in to the nucleus because it was moving around.

15 15 The Bohr Ring Atom n He didnt know why but only certain energies were allowed. n He called these allowed energies energy levels. n Putting Energy into the atom moved the electron away from the nucleus. n From ground state to excited state. n When it returns to ground state it gives off light of a certain energy.

16 16 The Bohr Ring Atom n = 3 n = 4 n = 2 n = 1

17 17 The Bohr Model n n is the energy level (an integer). n Z is the nuclear charge, which is +1 for hydrogen. n E = x J (Z 2 /n 2 ) n n = 1 is called the ground state. n When the electron moves from one energy level to another. E = E final - E initial E = E final - E initial E = x J Z 2 (1/ n f 2 - 1/ n i 2 ) E = x J Z 2 (1/ n f 2 - 1/ n i 2 )

18 18 Examples E = x J Z 2 (1/ n f 2 - 1/ n i 2 ) E = x J Z 2 (1/ n f 2 - 1/ n i 2 ) n Calculate the energy need to move an electron from level n=1 to level n=2. n What is the wavelength of light that must be absorbed by a hydrogen atom in its ground state to reach this excited state. n Calculate the energy released when an electron moves from n= 5 to n=3 in a He +1 ion

19 19 The Bohr Model n Doesnt work. n Only works for hydrogen atoms. n Electrons dont move in circles. n The quantization of energy is right, but not because they are circling like planets.

20 The Quantum Mechanical Model n A totally new approach. n De Broglie said matter could be like a wave. n De Broglie said they were like standing waves. n The vibrations of a stringed instrument.

21 21 Whats possible? n You can only have a standing wave if you have complete waves. n There are only certain allowed waves. n In the atom there are certain allowed waves called electrons. n 1925 Erwin Schroedinger described the wave function of the electron. n A lot of math, but what is important are the solutions.

22 22 Schroedingers Equation The wave function ( is a (x, y, z) function. The wave function ( is a (x, y, z) function. n Solutions to the equation are called orbitals. n These are not Bohr orbits. n Each solution is tied to a certain energy. n These are the energy levels.

23 23 The Heisenberg Uncertainty Principle. n There is a fundamental limitation on how precisely we can know both the position and momentum of a particle at a given time. n MATHEMATICALLY: x · (mv) > h /4 x · (mv) > h /4 x is the uncertainty in the position. x is the uncertainty in the position. (mv) is the uncertainty in the momentum. (mv) is the uncertainty in the momentum. the minimum uncertainty is h /4 the minimum uncertainty is h /4

24 24 What does the wave Function mean? n Nothing. It is not possible to visually map it. n The square of the function is the probability of finding an electron near a particular spot. n Best way to visualize it is by mapping the places where the electron is likely to be found.

25 25 Defining the size n The nodal surface. n The size that encloses 90% to the total electron probability. n NOT at a certain distance, but a most likely distance. n For the first solution it is a a sphere. This is of course and s orbital.

26 Quantum Numbers n There are many solutions to Schroedingers equation n Each solution can be described with quantum numbers that describe some aspect of the solution. n Principal quantum number (n) - size and energy of of an orbital. n Has integer values > 0

27 27 Quantum numbers n Angular momentum quantum number l. n shape of the orbital. n integer values from 0 to n - 1 n l = 0 is called s n l = 1 is called p n l =2 is called d n l =3 is called f n l =4 is called g

28 28 Quantum numbers n Magnetic quantum number (m l ) n integer values between - l and + l n tells direction in each shape. n Electron spin quantum number (m s ) n Can have 2 values. n either +1/2 or -1/2 n Q: What are the possible quantum numbers for the last electron of a sulfur atom? Iron?

29 29 Example: Describe the electrons defined by the following quantum numbers: n l m s electron p electron d electron not allowed (l must be < n) not allowed (ml must be between -l and l)

30 30 1. Wha t are the shapes of s, p, and d orbitals respectively? 2. How many 1s orbitals are there in an atom? 4p orbitals? 4d orbitals? 3. What is the maximum number of orbitals with: n = 4 l = 1 n = 2 l = 2 n = 3 l = 2 n = 5 l = 1 ml = Which orbitals cannot exist? 2p 3p 4d 3f 6s 2d 5. Write a set of quantum numbers for a 4f orbital.

31 31 1. Wha t are the shapes of s, p, and d orbitals respectively? s= spherical p = dumbbell d = cloverleaf 2. How many 1s orbitals are there in an atom? 4p orbitals? 4d orbitals? 1s: 1 4p: 3 4d: 5 3. What is the maximum number of orbitals with: n = 4 l = 1 3 (the 4p orbitals) n = 2 l = 2 none (l must be < n) n = 3 l = 2 5 (the 3d orbitals) n = 5 l = 1 ml = -1 1 (3 q.n. define a unique orbital) 4. Which orbitals cannot exist? 2p 3p 4d 3f 6s 2d 3f and 2d 5. Write a set of quantum numbers for a 4f orbital. n = 4 l = 3 ml = 3, 2, 1, 0, -1, -2, -3

32 32 S orbitals (BDVD) orbitals

33 33 P orbitals

34 34 D orbitals

35 35 F orbitals

36 36 Pauli Exclusion Principle n In a given atom, no two electrons can have the same set of four quantum numbers ( n, l, m l, m s ). n Therefore, an orbital can hold only two electrons, and they must have opposite spins.

37 Polyelectronic Atoms n More than one electron. n Three energy contributions. n The kinetic energy of moving electrons. n The potential energy of the attraction between the nucleus and the electrons. n The potential energy from repulsion of electrons.

38 38 Polyelectronic atoms n Electron shielding occurs. This affects the properties of these atoms. n Electrons are attracted to the nucleus. n Electrons are repulsed by other electrons. n Electrons would be bound more tightly if other electrons were not present.

39 The Periodic Table n Developed independently by German Julius Lothar Meyer and Russian Dmitri Mendeleev (1870s). n Didnt know much about atom. n Put in columns by similar properties. n Predicted properties of missing elements.

40 Aufbau Principle n Aufbau is German for building up. n As the protons are added one by one, the electrons fill up hydrogen-like orbitals. n Fill up in order of energy levels.

41 41 Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f He with 2 electrons

42 42 Details n Valence electrons - the electrons in the outermost energy levels (not d). n Core electrons - the inner electrons. n Hunds Rule - The lowest energy configuration for an atom is the one have the maximum number of of unpaired electrons in the orbital. n C 1s 2 2s 2 2p 2

43 43 Details n Elements in the same column have the same electron configuration. n Put in columns because of similar properties. n Similar properties because of electron configuration. n Noble gases have filled energy levels. n Transition metals are filling the d orbitals

44 44 Exceptions n Ti = [Ar] 4s 2 3d 2 n V = [Ar] 4s 2 3d 3 n Cr = [Ar] 4s 1 3d 5 n Mn = [Ar] 4s 2 3d 5 n Cu = [Ar] 4s 1 3d 10 n Half filled orbitals. n Scientists arent sure of why it happens

45 Periodic Trends in Properties n Ionization Energy - The energy necessary to remove an electron from a gaseous atom. n The ionization energy for a 1s electron from sodium is 1.39 x 10 5 kJ/mol. n The ionization energy for a 3s electron from sodium is 4.95 x 10 2 kJ/mol. n Demonstrates shielding.

46 46 Shielding n Electrons on the higher energy levels tend to be farther out. n Have to look through the other electrons to see the nucleus. n They are less effected by the nucleus.

47 47 How orbitals differ n The more positive the nucleus, the smaller the orbital. n A sodium 1s orbital is the same shape as a hydrogen 1s orbital, but it is smaller because the electron is more strongly attracted to the nucleus. n The helium 1s is smaller as well. n This provides for better shielding.

48 48 Periodic Trends n Highest energy electron removed first. n First ionization energy (I 1 ) is that required to remove the first electron. n Second ionization energy (I 2 ) - the second electron

49 49 Trends in ionization energy n for Mg I 1 = 735 kJ/moleI 1 = 735 kJ/mole I 2 = 1445 kJ/moleI 2 = 1445 kJ/mole I 3 = 7730 kJ/moleI 3 = 7730 kJ/mole n The effective nuclear charge increases as you remove electrons. n It takes much more energy to remove a core electron than a valence electron because there is less shielding.

50 50 Explain this trend n For Al I 1 = 580 kJ/moleI 1 = 580 kJ/mole I 2 = 1815 kJ/moleI 2 = 1815 kJ/mole I 3 = 2740 kJ/moleI 3 = 2740 kJ/mole I 4 = 11,600 kJ/moleI 4 = 11,600 kJ/mole

51 51 Across a Period n Generally from left to right, I 1 increases because there is a greater nuclear charge with the same shielding. n As you go down a group I 1 decreases because electrons are farther away.

52 52 It is not that simple n Z eff changes as you go across a period, so will I 1 n Half filled and filled orbitals are harder to remove electrons from. n Heres what it looks like.

53 53 First Ionization energy Atomic number

54 54 First Ionization energy Atomic number

55 55 First Ionization energy Atomic number

56 56 The information it hides n Know the special groups n It is the number and type of valence electrons that determine an atoms chemistry. n You can get the electron configuration from it. n Metals lose electrons have the lowest IE n Non metals- gain electrons most negative electron affinities.

57 57 Electron Affinity & Atomic Radii n Electron affinity - The energy change associated with the addition of an electron to a gaseous atom. X(g) + e - X - (g) X(g) + e - X - (g) n Affinity tends to increase across a period and decrease as you go down a group. n Atomic Radii: n Decrease going from left to right across a period and increase going down a group.

58 The Alkali Metals n Doesnt include hydrogen - it behaves as a non-metal. n Decrease in IE. n Increase in radius. n Decrease in density. n Decrease in melting point. n Behave as reducing agents.

59 59 Reducing ability n Lower IE = better reducing agents. n Cs>Rb>K>Na>Li n Works for solids, but not in aqueous solutions. n In solution Li>K>Na n Why? n Its the water - there is an energy change associated with dissolving.

60 60 Hydration Energy n It is exothermic. n for Li kJ/mol n for Na kJ/mol n for K kJ/mol n Li is so big because of it has a high charge density, a lot of charge on a small atom. n Li loses its electron more easily because of this in aqueous solutions.

61 61 The reaction with water n Na and K react explosively with water. n Li doesnt. Even though the reaction of Li has a more negative H than that of Na and K. Even though the reaction of Li has a more negative H than that of Na and K. n Na and K melt. H does not tell you speed of reaction. H does not tell you speed of reaction. n More in Chapter 12.


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