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Ch. 13 & 14 Notes

Ch. 13 & 14 Notes. 13.1 Evolution of Atomic Models. Dalton Atom was a solid indivisible mass. Cu. Ag. Thompson Plumb Pudding model of the atom Electrons stuck into a lump of positively charged material, similar to raisins stuck in dough. Rutherford

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Ch. 13 & 14 Notes

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  1. Ch. 13 & 14 Notes

  2. 13.1 Evolution of Atomic Models • Dalton • Atom was a solid indivisible mass

  3. Cu Ag

  4. Thompson • Plumb Pudding model of the atom • Electrons stuck into a lump of positively charged material, similar to raisins stuck in dough.

  5. Rutherford • Proposed that atoms contained a dense positive nucleus surrounded by electrons • Thought the atom was mostly empty space

  6. Bohr • Electrons are arranged in concentric circular paths (orbits) around the nucleus • Called the planetary model • Electrons have particular paths with fixed energy, which he called its energy level

  7. Quantum Mechanical Model – Modern description of the electrons in atoms • It comes from the mathematical solutions to the Schodinger equation • The probability of finding an electron within a certain volume of space surrounding the nucleus can be represented by a cloud (electron cloud) • The cloud represents where the electron can be found approximately 90% of the time.

  8. Summary…

  9. Energy level • region around the nucleus where the electron is likely to be moving • electrons can “jump” from one energy level to another • Quantum • the amount of energy required to move an electron from its present energy level to the next higher energy level

  10. Atomic orbital • the region in space where the electron is likely to be found • Locations of Electrons in an Atom • Electrons can be described by a series of four quantum numbers

  11. Quantum Numbers • Principal quantum number (n) • Describes the principal energy level an electron occupies • It has values of 1,2,3,4…

  12. The maximum number of orbitals in a particular energy level is equal to n2. • The maximum number of electronsin a particular energy level is equal to 2n2.

  13. Color Your Periodic Table Make a key that includes all 4 orbital letters. “Stupid Dogs Pee on Fences”

  14. 2 1 1 2 4 8 3 9 18

  15. 13.2 Electron Configurations • Electron configuration – ways in which electrons are arranged around the nuclei of atoms

  16. Three rules for finding electron configurations: • Aufbau Principle – electrons enter orbitals of lowest energy first. • Pauli exclusion principle – an atomic orbital can contain at most two electrons. They must have opposite spins. • Hund’s Rule – When electrons occupy orbitals of equal energy, one electron enters each orbital until all orbitals contain one electron with parallel spins. Second electrons are added and are paired so that each orbital contains two electrons with opposite spins.

  17. s 1 2 p 3 6 d 5 10 f 7 14

  18. Complete the orbital diagrams below: 1s 2s 2p 3s • H __ __ __ __ __ • He __ __ __ __ __ • O __ __ __ __ __ • Na __ __ __ __ __ __

  19. EXAMPLES: Now, write the electron configurations for the following: He Na Ti 1s2 1s2 2s2 2p6 3s1 1s2 2s2 2p6 3s2 3p6 4s2 3d2

  20. Valence Electrons • Valence electrons are electrons in the outermost energy level. • The group number tells you the number of valence electrons that each atom has

  21. Electron Dot Diagrams • Consists of; an elemental symbol, and a group of 1-8 dots which shows the configuration of the valence electrons. • Valence Electrons – the outer-most electron shell of the atom • Below is an example of the proper Lewis dot diagram for the element oxygen. 

  22. Octet Rule • The octet rule is a chemical rule of thumb that states that atoms tend to combine in such a way that they each have eight electrons in their outer energy level • So it is the magic number 8!

  23. Oxidization numbers/Charges • Because all atoms want 8 valence electrons then they have an oxidization number that tells how many they are missing or how many they want to get rid of.

  24. X 1s 2s 2p 3s 3p • Na+ __ __ __ __ __ __ __ __ __ • Al3+ __ __ __ __ __ __ __ __ __ • O2- __ __ __ __ __ __ __ __ __ • Fe3+ will be done on board. X X X Now, they are STABLE b/c their orbitals are full!!!

  25. 1s2 2s2 2p6 • Now write their electron configurations. • Na+ = • Al3+ = • Fe3+ = • O2- = 1s2 2s2 2p6 1s2 2s2 2p6 3s2 3p6 4s2 3d3 1s2 2s2 2p6

  26. There are two exceptions to the three rules governing orbital theory. 1s 2s 2p 3s 3p 4s 3d Cu _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ Cr _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ It is more stable this way because the electrons are on a lower energy level!

  27. For a shorter way to write electron configuration, write the nearest noble gas and then continue. AKA “Shorthand Notation”. Ex: Ti can be written as OR 1s2 2s2 2p6 3s2 3p6 4s2 3d2 [Ar] 4s2 3d2

  28. Diagram of a wave: • Amplitude – wave’s height from origin the crest. • Wavelength() – the distance between crests. • Frequency (f) – the number of wave cycles to pass a given point per unit time. (Hz or s-1) • f = c/ • Canalso be written: c=f amplitude

  29. EXAMPLE: • A wave of yellow light has a frequency of 2.73 x 1016s-1. Calculate its wavelength. v = 2.73 x 1016 s-1 c = 3.00 x 108 m/s  = ? c = v  • = c/v • = (3.00 x 108 m/s)/(2.73 x 1016 s-1) = 1.10 x 10-8 m

  30. Quantum Concept & the Photoelectric Effect • Max Planck believed that energy was emitted or absorbed by small units called quanta. He found that the amount of energy released or absorbed was proportional to the frequency of the radiation. • E = h f • E = Energy • f = frequency • h = Planck’s constant: 6.63 x 10-34 J∙s • Einstein proposed that light could be viewed as a stream of particles called photons.

  31. EXAMPLE: • Calculate the energy of an individual photon of yellow light having a frequency of 2.73 x 1016s-1. E = ? v = 2.73 x 1016 s-1 h = 6.63 x 10-34 J∙s E=hv E = (6.63 x 10-34 J∙s)(2.73 x 1016s-1) = 1.81 x 10-17 J

  32. Light and Atomic Spectra • Electromagnetic Radiation – a series of energy waves that includes radio waves, microwaves, visible light, infrared and ultraviolet light, x-rays, gamma rays.

  33. Spectrum – series of colors produced when white light is separated by a diffraction gradient (prism). • ROY G. BIV • Red light has the longest wavelength and the shortest frequency. LOWEST ENERGY. • Violet light has the shortest wavelength and the highest frequency. HIGHEST ENERGY.

  34. Atomic emission spectrum – series of lines produced by passing the light emitted by an excited atom through a diffraction gradient. These spectra can be use in element identification. Sun Hydrogen Helium Mercury Uranium

  35. Ch. 14Chemical Periodicity

  36. 14.1 – Classification of the Elements • History of the Periodic Table • Mendeleev- arranged elements in order of increasing atomic mass • Moseley- arranged elements in order of increasing atomic number • Periodic Law- When the elements are arranged in order of increasing atomic number, there is a periodic pattern in their physical and chemical properties.

  37. Columns are called groups or families. Rows are called periods or series. • Be able to locate noble gases, representative elements, transition metals, inner transition metals.

  38. The noble gases are the elements in which the outer most s and p sublevels are filled. They belong to group 0. • The representative elements have outermost s or p orbitals only partially filled. They are usually called group A elements. • Group 1 A – alkali metals • Group 2 A – alkaline earth metals • Group 7 A – halogens

  39. The transition metals are metallic elements in which the outermost s sublevel and nearby d sublevel contain electrons. They can be found in the middle of the Periodic Table. • The inner transition metals are metallic elements in which the outermost s sublevel and nearby f sublevel generally contain electrons. They can be found at the bottom of the Periodic Table.

  40. 14.2 – Periodic Trends • Trends in Atomic Size • Atomic radius – one-half the distance between the nuclei of two like atoms in a diatomic molecule (ex: Cl2) • Trend:

  41. Increasing Atomic Radius SMALLEST Increasing Atomic Radius BIGGEST Atomic Radius

  42. Ionization energy - energy required to overcome the attraction of the nuclear charge and remove one electron from a gaseous atom • 1st ionization energy: the energy required to remove the first electron • 2nd ionization energy: the energy required to remove the second electron • 3rd ionization energy: the energy required removing the second electron • Trend:

  43. Increasing Ionization Energy Increasing Ionization Energy Ionization Energy

  44. Electronegativity – the tendency for the atoms of the element to attract electrons when they are chemically combined with atoms of another element • Note: Nobel gases don’t have any electronegativity b/c they are full and don ‘t bond with other elements. • Trend:

  45. Increasing Electronegativity HIGHEST Increasing Electronegativity Electronegativity

  46. Ion Size • Anions are larger than the atoms from which they were formed. • The negative charge means more electrons are present causing the size of the ion to be larger. • Cations are smaller than the atoms from which they were formed. • The positive charge means fewer electrons are surrounding the nucleus, thus pulling the existing electrons closer and causing the ion to be smaller. • Trend:

  47. Increasing Ionic Radius Increasing Ionic Radius Ionic Radius

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