1. Chapter 11 Modern Atomic Theory

2. 11.1 Rutherford’s Atom  The nuclear atom (atom with a nucleus) resulted from Ernest Rutherford’s Gold Foil Experiment.  Rutherford and his colleagues were able to show that an atom had a nucleus made of protons and neutrons, and that the atom was mostly empty space.

3.  Rutherford was not able to determine what was going on with the electrons.  Rutherford suggested that the electron were rotating around the nucleus like planets revolving around the sun.

4.  The problem with Rutherford’s thoughts on electrons was that he couldn’t explain why the positively charged nucleus did not attract the negatively charged electrons.

5. 11.2 Electromagnetic Radiation  Electromagnetic radiation is energy that is transmitted from one location to another by light.  There are many different types of electromagnetic radiation, all of which are contained on the electromagnetic spectrum.

7.  The different types of electromagnetic radiation are different in two ways: wavelength, frequency.  The wavelength (λ) is the distance between two consecutive peaks.  The frequency (υ) indicates how many peaks pass a certain point per given period of time.  Waves also have speed; the speed of a wave indicates how fast it travels.

9.  Electromagnetic radiation can be thought of as a wave that carries energy through space.  Electromagnetic radiation doesn’t always behave as a wave, sometimes it behaves as a stream of particles.  A beam of light can be thought of as a stream of tiny packets of energy called photons.  Different wavelengths of electromagnetic radiation carry different amounts of energy.

10. 11.3 Emission of Energy by Atoms  When atoms receive energy from some source, they become excited and release this energy by emitting light.  The energy emitted is carried away by a photon.  A photon corresponds to the exact energy change experienced by the atom.  High energy photons correspond to short wavelength light and low energy photons correspond to long wavelength light.

11.   Teacher's Domain Fireworks Teacher's Domain Fireworks

12. 11.4 The Energy Levels of Hydrogen  An excited atom can release some or all of its excess energy by emitting a photon and thus move to a lower energy state.  A photon is a “particle” of electromagnetic radiation.  The lowest possible energy state of an atom is called the ground state.

14.  We can learn a great deal about the energy states of any atom by observing the photons it emits.  Scientists are able to learn because the different wavelengths of light carry different amounts of energy per photon.  The energy contained in a photon corresponds to the change in energy that the atom experiences going from an excited state to a ground state.

15.  The simplest atom to study is hydrogen, because it contains only one proton and one electron.  When scientists studied hydrogen they found that hydrogen only emits certain colors of visible light.  Because only certain photons are emitted, we know that only certain energy changes are occurring.  Meaning hydrogen atoms have discrete energy levels.

16.  Hydrogen atoms always emit photons of the same discrete colors.  This means that hydrogen atoms have the same set of discrete energy levels.  The energy levels are quantized, that only certain values are allowed.  Scientists were surprised by the quantization of the energy levels.

17. Hydrogen Line Emission

18. 11.5 The Bohr Model of the Atom  Niels Bohr constructed a model of the hydrogen atom with quantized energy levels the agreed with the hydrogen emission results.  Bohr’s model had an electron moving in circular orbits corresponding to the allowed energy levels.  He suggested the electron could jump to a different level by absorbing or emitting a photon of light with the correct energy content.

19.  Bohr’s model seemed like it explained everything.  The electrons don’t fall into the nucleus because they have a specific amount of energy.  It explained the line emission spectrum of hydrogen.

20.  Bohr’s model was unfortunately incorrect.  Bohr’s calculations only worked for the hydrogen atom.  His calculations could not explain the line emission spectrum of helium.  The current model of the atom is not the Bohr model.  In the current model, electrons do not orbit the nucleus like planets orbiting the sun.

21.  Teacher’s Domain Video Teacher’s Domain Video Teacher’s Domain Video

22. 11.6 The Wave Mechanical Model of the Atom  Louis de Broglie and Erwin Schrodinger were the next scientist to explore the structure of the atom. These scientists suggested that since light behaves like both a wave and a particle, electrons must do the same.  When Schrodinger performed a mathematical analysis and considered the electron to be both wave-like and particle-like, he found that his data worked. His data worked equally well for all atoms, not just hydrogen.  This model of the atom is called the wave mechanical model (or the quantum mechanical model).

23.  In the wave mechanical model of the atom electrons are found in orbitals, which are nothing like orbits.  An orbital shows the areas where the electron is most likely to be found 90% of the time. Frequently the electron is close to the nucleus, but occasionally it drifts further away.  The exact path of the electron is not predictable. Schrodinger work allowed him to predict the areas where it most probable to find the electron in an atom.

25. 11.7 The Hydrogen Orbitals  The probability map for an electron is called an orbital. The probability of finding an electron decreases at greater distances from the nucleus, but it never reaches zero.  The edge of an orbital is fuzzy, because an orbital does not have an exact defined size. Chemists decided to define the area that contains 90% of the electron probability.

26.  For a hydrogen electron the ground state is the 1s orbital. When the electron absorbs more energy it moves to a higher state. The higher energy orbitals have different shapes.  The hydrogen atom has discrete energy levels, called principal energy levels. Principal energy levels are labeled with numbers (like the ground state is 1). Each principal energy level is divided into sublevels. Sublevels are notes with letters (like s, p, d, and f).

27.  Within each sublevel are orbitals which hold electrons.  The principal energy level (1) contains one sublevel, or one type of orbital. The only orbital, the 1s is spherically shaped.

28.  The second principal energy level has two sublevels. These two sublevels are indicated by the letters s and p.  The 2s level contains one circular orbital. The 2p contains three “peanut” shaped orbitals. Each orbital holds at most two electrons.

29. Orbital Labels  The number tells the principal energy level  The letter tells the shape.  The letter s means spherical orbital; the letter p means a peanut shaped orbital. The subscript x, y, or z denotes the coordinate axes where the peanut lies.

31.  As the principal energy level number increases the distance from the nucleus increases.  An orbital should be thought of as a potential space for an electron.  Hydrogen may be in its ground state in the 1s orbital, but many other “potential spaces” exist for the one electron to go when it is excited.

32. 11.8 The Wave mechanical Model: Further Development  The wave mechanical model of the atom can be applied to all atoms. The wave mechanical model of the atom explains the arrangement of the periodic table.  The wave mechanical model of the atom explains the similarities that occur between elements in the same group, based on electron arrangements.

33.  Each electron has a spin. The electron can spin in two directions. This spin is represented with an arrow. Two electrons in the same orbital must have opposite spins.  The Pauli exclusion principal states that an atomic orbital can hold a maximum of two electrons, and those electrons must have opposite spins.

34. 11.9 Electron Arrangements in the First 18 atoms in the Periodic Table  Electron fill the orbital with the lowest energy first; this is called the Aufbau principle.  The arrangement of electron in an atom can be represented two ways: electron configuration or an orbital diagram.  An electron configuration lists the orbitals that contain electrons and how many electrons those orbitals contain.  An orbital diagram uses a box to represent each orbital and an arrow to represent each electron.

35.  Example Electron Configuration Nitrogen: 1s 2 2s 2 2p 3Nitrogen: 1s 2 2s 2 2p 3  Example Orbital Diagram:

36.  When making orbital diagrams there is another rule to follow: Hund’s Rule.  Hund’s Rule states that when more than one orbital of equal energy is available electrons go into each orbital before pairing up.

37. 11.11 Atomic Properties and the Periodic Table  Dynamic Periodic Table Dynamic Periodic Table Dynamic Periodic Table

38. Trends in Atomic Size  The atom does not have a sharply defined boundary that limits its size; therefore the radius of an atom cannot be determined directly.  The atomic radius is one half the distance between the nuclei of two like atoms in a diatomic molecule.  The atomic radius of an atom of an element indicates the relative size.

39. Atomic Size in Groups  Atomic size generally increases as you move down a group.  As you move down a group electrons are added to a higher principle energy level and the nuclear charge increases.  The enlarging effect of the greater distance from the nucleus overcomes the shrinking effect of the increasing charge of the nucleus.  The outermost orbital is larger as you move downward, therefore the size of the atom increases.

40. Atomic Size in Periods  Atomic size generally decreases as you move from left to right across a period.  Each element has one more proton and one more electron than the preceding element.  The electrons are added to the same principle energy level.  The effect of the increasing nuclear charge on the outermost electrons is to pull them closer to the nucleus, therefore the size of the atom decreases.

41. Electron Shielding  The effect of the increasing nuclear charge is less pronounced in periods where there are electrons in energy levels between the nucleus and the outermost electrons.  This is because the inner electron shields the charge of the nucleus. The effect of shielding is constant within a period

42. http://images.google.com/imgres?imgurl=http://www.camsoft.co.kr/CrystalMaker/support/tutorials/crystalmaker/resources/VFI_Atomic_Radii_sm.jpg&imgrefurl=http://www.camsoft.co.kr/Cryst alMaker/support/tutorials/crystalmaker/AtomicRadii.htm&h=450&w=680&sz=80&hl=en&start=3&um=1&usg=__dzl5lBD9l8fzu_rjxzUFr7QtUYo=&tbnid=kjwcl42ApwWlLM:&tbnh=92&tbnw=139 &prev=/images%3Fq%3Dperiodic%2Btable%2Bshowing%2Batomic%2Bradius%26um%3D1%26hl%3Den%26client%3Dfirefox-a%26channel%3Ds%26rls%3Dorg.mozilla:en- US:official%26sa%3DN

43. Trends in Ionization Energy  When an atom gains of loses an electron it becomes an ion.  The energy required to overcome the attraction of the nucleus and remove an electron from a gaseous atom is called ionization energy.

44. Groups and Ionization Energy  The energy required to remove the first electron is called first ionization energy.  The first ionization energy generally decreases as you move down a group.  This happens because the size of the atom increases and the outermost electron is further from the nucleus, making it more easily removed.

45. Periods and Ionization Energy  The first ionization energy increases as you move from left to right across a period.  The nuclear charge increases and the shielding effect is constant, therefore there is a greater attraction of the nucleus for the electron and a higher ionization energy.

46. Ionization Energy http://images.google.com/imgres?imgurl=http://www.800mainstreet.com/4/0004-000-IE.GIF&imgrefurl=http://www.800mainstreet.com/4/0004-002- Periodic.html&h=340&w=510&sz=39&hl=en&start=19&um=1&usg=__2Ai1eH6TjvM51NHjO4Hyr8FJsIM=&tbnid=Rm_BRv8Igob7pM:&tbnh=87&tbnw=131&prev=/images%3Fq%3Dperiodic%2Btable %2Bof%2Bionization%2Benergy%26start%3D18%26ndsp%3D18%26um%3D1%26hl%3Den%26client%3Dfirefox-a%26channel%3Ds%26rls%3Dorg.mozilla:en-US:official%26sa%3DN

47. Ionic Size  The atoms of metallic elements have low ionization energies and thus form positive ions easily.  Remember a positive ion is formed when electrons are lost.  Atoms of nonmetallic elements however form negative ions.  A negative ion is formed when an atom gains electrons.

48. Cations (Positive Ions)  Positive ions are always smaller than the neutral atoms from which they form. This is because the loss of outer-shell electrons results in an increased attraction by the nucleus for the remaining electrons.

49. Anions (Negative Ions)  Negative ions are always larger than the neutral atoms from which they form. This is a result of the smaller effective nuclear charge.

50. Trends in Ionic Size  Going from left to right across a row, there is a gradual decrease in the size of positive ions.  Then beginning with group 5, the negative ions (which are much bigger) gradually decrease in size as you continue to move right.  Going down a group there is a general increase in ionic size.

51. Trends in Ionic Size http://www.chem.umass.edu/~botch/Chem111F04/Chapters/Ch8/IonicRadii.jpg

52. Electronegativity  The electronegativity of an element is the tendency for the atoms of the element to attract electrons when they are chemically combined with atoms of another element.  Electronegativity is used to predict the type of bond that occurs between atoms.

53. Trends in Electronegativity  Electronegativity generally decreases as you move down a group.  As you move across a period from left to right the electronegativity increases.  Fluorine is the most electronegative element, while cesium is the least electronegative element.

54. Trends in Electronegativity

55. All Periodic Trends http://www.meta-synthesis.com/webbook/35_pt/general_pt.gif