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Unit 2-Atomic Structure [Ch.3/11/12 (14)] Ch. 3 PowerPoint Presentation
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Unit 2-Atomic Structure [Ch.3/11/12 (14)] Ch. 3

Unit 2-Atomic Structure [Ch.3/11/12 (14)] Ch. 3

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Unit 2-Atomic Structure [Ch.3/11/12 (14)] Ch. 3

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  1. Unit 2-Atomic Structure [Ch.3/11/12 (14)]Ch. 3 What is an atom? An element? A compound? A molecule?

  2. Structure of the atomDalton’s Atomic Theory (p.56) • Elements are made of tiny particles called _____________. • Atoms of a given element are ______________. • Atoms of a given element are different from those of any other ____________. • Atoms of different elements combine to form chemical ________________. • Atoms are ______________ in chemical processes. Atoms are not created or destroyed in chemical reactions.

  3. Table 3.4 • Page 63. Table 3.4. Mass and charge of subatomic particles

  4. Thomson • All atoms must contain electrons. Those electrons must be balanced by protons

  5. Plum pudding model • Theory that electrons were scattered in a positive “pudding” (p.61). • Draw model

  6. Rutherford • Found that plum pudding model was wrong. Found that an atom is mostly open space. A center that is (+) and electrons (–) circled around it. He also is credited with finding the neutron.

  7. Modern concept of the atom. • Extremely tiny nucleus. Electrons circle around it very fast, but from a very far distance. It is these electrons that make up the chemical properties of an element. • Think about a needle in the middle of a football stadium and the seats are where electrons could be.

  8. IONS • We can make ions by taking a neutral atom and removing or adding more valence electrons. • Positive ion=cation. Electrons are lost. • Negative ion=anion. Electrons are gained. (p 76-78) • Copy charges onto periodic table if not done already.

  9. Unit 2 Ch.11Photoelectric Effect • What is light? • Visible light is a kind of electromagnetic radiation. • Other types of electromagnetic radiation are x-rays, ultraviolet light, infrared light, microwaves, and radio waves. • All light forms travel at a speed of 3.0x108 m/s. • Frequency (v) is the # of waves that pass a given point in a specific time, usually 1 sec. V is expressed in waves/sec which is called a Hertz (Hz). • The speed of light is wavelength multiplied by frequency.

  10. Photoelectric effect • This is the emission of electrons from a metal when light shines on the metal.

  11. Photoelectric effect • Light as a particle. Max Planck thought that when metal gets hot, there is not a continuous emission of electromagnetic energy. Instead, energy is emitted in small, specific amounts called quanta. A quantum is the minimum quantity of energy that can be lost or gained by an atom. • E=hv

  12. Photoelectric effect • E is energy (in joules), v is the frequency of radiation, and h is Planck’s constant = 6.626x10-34 J seconds.

  13. Photoelectric effect • Therefore, light = waves and particles. This particle is called a photon. These photons have no mass, just a quantum of energy.

  14. Hydrogen atom line-emission spectrum • When an excited atom returns to its ground state, it gives off the energy it gained in the form of electromagnetic radiation. (Ex. is light in neon signs) • When electric current was passed through H gas, a pinkish glow occurs. That light can be separated through a prism into the line-emission spectrum. The different wavelengths of light are separated. • Whenever atoms fall back from an excited state to a ground state (or lower energy state), it emits a photon of radiation. This energy is equal to: • Ephoton = E2 – E1 = hv

  15. Bohr model • This lead to the Bohr model of an atom. The electron can circle the nucleus only in allowed paths, or orbits. When the electron is in one of these orbits, the atom has a definite, fixed energy. The electron is in its lowest energy state when it is in the orbit closest to the nucleus. The energy of the electron is higher when it is in orbits that are further from the nucleus.

  16. Energy in H atom • Problem is that this only applied for the H atom, 1 e- • (Page 366) • absorption emission

  17. The Quantum model • Electrons have wavelike properties. They can, like waves, bend or diffract. They can also interfere or overlap. • Heisenberg uncertainty principle states that it is impossible to determine simultaneously both the position and velocity of an electron or any other particle. • Schrodinger took this idea and came up with equations known as quantum theory or showing mathematically the wave properties of electrons and other very small particles. This shows the probability of where an electron traveled. They do not travel in neat orbitals.

  18. Atomic orbitals and quantum #’s Rules for electrons in orbitals. • n = principal quantum number • values of n are 1,2,3,…. • As n increases, Energy increases. • n = 1 is first energy level closest to nucleus • more than 1 electron can have the same n value • total # of orbitals in a shell (main energy level) is n2 • l indicates the shape of the orbital • The # of orbital shapes possible is = to n • Values of l are 0 and all positive integers less than or equal to n-1. (Ex. n = 2, l = 0, l = 1)

  19. Principle energy levels Page 375

  20. Atomic Orbitals • (page 372 shows the shape of the orbitals) • n = 1, it is a s orbital • n = 2, it is an s and p orbital • n = 3, it is an s,p, and d orbital • n = 4, it is an s,p,d, and f orbital

  21. Electron ConfigurationsThe arrangement of electrons in an atom. • RULES • The electron occupies the lowest-energy orbital that can receive it. (Aufbau principle) • Two electrons in each orbital must have opposite spins (Pauli exclusion principle) • Each orbital must be filled by an electron in each division (s,p,d,f) before the whole division is full (Hund’s rule).This allows electron-electron (- , -) repulsion.

  22. PRACTICE • 1. He 2. Li • 3. C 4. O • 5. Ne 6. Mg • 7. Si 8. Cl • 9. Ar

  23. Now list each in its electron configuration notation. • 1. He 2. Li • 3. C 4. O • 5. Ne 6. Mg • 7. Si 8. Cl • 9. Ar

  24. Noble-Gas Notation • Meant to shorten the long configuration notion. • Uses the previous noble gas and then adds only that period. (ex-Sodium is [Ne]3s1 )

  25. Now list each in its noble-gas notation. • 1. C 2. O • 3. Mg 4. Cl • 5. Ca 6. Br • 7. Fe 8. Ba • 9. Au

  26. More Practice Write both the long and Noble Gas electron configuration for: 1 W 2 Y 3 Hg 4 La 5 Fr

  27. Chapter 11 TRENDS • Atomic radius is defined as one-half the distance between the nuclei of identical atoms that are bonded together. Periodic handout I gave you shows the trend of the size of the atomic radii of the elements. • TREND is ? (Answer on left side of Cornell notes)

  28. Ionization Energy • Ions are an atom or group of bonded atoms that has a positive or negative charge. For example, Na wants to lose an e-, meaning it will be +1. It forms a Na+ ion. • Ionization energy (IE) is the energy required to remove one electron from a neutral atom. This is the first ionization IE1. The next electron it loses would be _______. Notice on Handout I gave you what the trend is. • TREND is ? (Answer on left side of Cornell notes)

  29. Ions • In general, nonmetals have higher ionization energies than metals do. • A positive ion is known as a cation. • A negative ion is known as an anion.

  30. Electronegativity • Is a measure of the ability of an atom in a chemical compound to attract electrons. • The most electronegative element is ______________. • TREND is ? (Answer on left side of Cornell notes)

  31. Unit 2 Chapter 12Chemical Bonding Questions? 1. When representing a water molecule as H-O-H, what do the lines between the letters symbolize? 2. What are the three types of bonding and what causes bonding in the first place?

  32. Chemical bonding Covalent bonding • Molecule is a neutral group of atoms that are held together by covalent bonds. • Chemical bonding happens naturally because the atoms are at a lower potential energy when bonded then alone. Remember, atoms want to be stable.

  33. Octet Rule (p 414): List how many val.e’s for each and then draw the Lewis dot structure for each pair. • Ca O • C H

  34. Electron dot notation = Lewis Structures • Only valence electrons of an atom of an element are shown. • 1 dot is 1 e-. • In a molecule, 2 dots would be 1 line, representing a bond.

  35. Practice 1) H H 5) C H H H H 2) O O 3) F F 4) Br Br

  36. Practice • 2 lines represent a double bond, 3 represent a __________ bond. • Now try (2) C and (4) H bonded together.

  37. Ionic bonding • An Ionic compound is composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal. • Crystalline structures -Solids -NaCl is a good example (table salt) -Caused from ionic bonding

  38. Ionic compounds • Most ionic compounds bond in a crystal lattice configuration (see page 411). • Na + Cl • Na+ + Cl-

  39. Polyatomic ions • A charged group of covalently bonded atoms (page 412)

  40. Metallic Bonding • Different from covalent and ionic bonding. Metal atoms are attracted to each other and their electrons make up a “sea of electrons” that surrounds the metal atoms. • This accounts for metals conducting electricity and heat as well as absorbing light and then emitting it which creates the shiny appearance.

  41. Metallic bonding • These properties also account for metals being malleable (hammered) and ductile (pulled into wire).

  42. Molecular Shapes/Geometry • VSEPR theory (Valence-Shell, Electron Pair Repulsion) says that valence electrons want to be as far apart as possible. • Main possible shapes: 5 (p.431) which depends on how many atoms are bonding together.

  43. Intermolecular Forces • 3 main types. • 1) Dipole-dipole. Occurs in polar molecules. Strongest intermolecular force. • (Example: top page 490) Draw

  44. Intermolecular forces • 2) Hydrogen bonding-occurs where H atom is bonded to a highly electronegative atom (ex. water). Because of this, these molecules have a high boiling point. • (Example: bottom page 490) Draw

  45. Intermolecular forces • 3) London dispersion-happens because of constant motion of electrons which cause instantaneous dipoles. • (Example: bottom of page 491) Draw

  46. Ch 14 Changes of State • Solids, liquids, and gases • Due to equilibrium = which is a dynamic condition in which 2 opposing changes occur at equal rates in a closed system. • Happens because of E entering or leaving a system. • Page 492, page 497 #1-6