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Chapter 5: The Periodic Law

Chapter 5: The Periodic Law. CP Chemistry Mrs. Klingaman. #13) Valence Electrons. Packet pg. 4. Definition : The electrons available to be lost, gained or shared in the formation of chemical compounds The OUTERMOST (highest energy level) electrons in an atom

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Chapter 5: The Periodic Law

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  1. Chapter 5: The Periodic Law CP Chemistry Mrs. Klingaman

  2. #13) Valence Electrons Packet pg. 4 • Definition: • The electrons available to be lost, gained or shared in the formation of chemical compounds • The OUTERMOST (highest energy level) electrons in an atom • For main-group elements, the valence electrons are the electrons in the highest energy level s & p atomic orbitals

  3. Do you see a trend among EC???

  4. #14) Valence Electrons Packet pg. 4

  5. #3) Valence Electrons Packet pg. 10 • Definition: • The OUTERMOST (highest energy level) electrons in an atom • Electrons involved in the chemical bonding of atoms

  6. # 4) Periodic Table & Valence Electrons 1 2 7 8 3 4 5 6

  7. #5) Bohr Diagrams Packet pg. 10 • Show total number of electrons in an atom and to which energy levels those electrons belong • Drawn like the Bohr Model, where each ring is an energy level (“n”) • Electrons fill in the lowest energy levels first (closest to nucleus)

  8. #5) Bohr Diagrams Packet pg. 10 • Each energy level has a different maximum number of electrons which it can contain

  9. #5) Bohr Diagrams Packet pg. 10 • Nucleus: • Determine # of protons & neutrons • Energy level rings: • The total number of energy levels in an atom is the same as the period # the atom is in • Fill in all electrons on the appropriate energy level (keep in mind the maximum # of electrons allowed on each level

  10. Oxygen - Bohr Diagram • p+ _______; n0 _______; • e– _______ • Group #: ____________ • Period #: ____________ • How many valence electrons does oxygen have? _____________ Hint: Look at the outer most ring, this is the HIGHEST energy level (the 2nd level)

  11. Silicon- Bohr Diagram Silicon (Si) : 1s22s22p63s23p2 • p+ _______; n0 _______; • e– _______ • Group #: ____________ • Period #: ____________ • How many valence electrons does oxygen have? _____________ Hint: Look at the outer most ring, this is the HIGHEST energy level (the 2nd level)

  12. #6) Lewis Dot Structures Packet pg. 10 • Show only the valence electrons in an atom (represented by dots) • No atom can have more than 8 valence electrons X 3 7 Dots (# of valence electrons) are added around the element symbol so that one electron is on each side first 6 5 1 2 8 4

  13. Lewis Diagrams: • Examples • Na – Cl • C – Ne • S

  14. Trends in Valence e–

  15. Bohr & Lewis - Practice Chlorine • p+ _______; n0_______; e– _______ • Group Name: __________________________ • Group #: ________ Period #: ________ • Draw Bohr Diagram & Lewis Dot Structure

  16. Chapter 5: Periodic Trends CP Chemistry Mrs. Klingaman

  17. Periodic Trends • Group Trends • Describes the general trend for elements looking down a group/family • from top to bottom • Period Trends • Describes the general trend for elements looking across a period • from left to right

  18. Periodic Trends • There are 4 periodic trends we will discuss: • Atomic Radius • Ionization Energy • Electron Affinity • Electronegativity

  19. #7) Atomic Radius Packet pg. 2 • Definition: • ½ the distance between the nuclei of identical atoms bonded together • The relative size of an atom

  20. #7) Atomic Radius (AR) Packet pg. 2 • Group Trend: • generally increases down a group. • looking top-to-bottom down a group, atoms get larger in size (measured by atomic radius) • Why?- Increasing Energy Level • As you move down a column, each successive element has one additional energy level surrounding the nucleus of the atom • This means electrons in the atom are farther away from the nucleus, creating a larger atomic radius

  21. Atomic RadiusGroup Trend:Increasesdown a column due to increasing energy level (electrons are farther from nucleus)

  22. #7) Atomic Radius (AR) Packet pg. 2 • Period Trend: • generally decreasesacross a period. • looking left-to-right across a period, atoms get smaller in size (measured by atomic radius) • Why?- Increasing Nuclear Charge • As you move across a period, each successive element has one more proton in its nucleus than the previous atom. • This increasing positive nuclear charge will cause a greater force of attraction and pull the electrons in closer to the nucleus.

  23. Atomic RadiusPeriod Trend:Decreases across a period due to increasing nuclear charge(electrons are pulled closer to the nucleus)

  24. Atomic Radius: Hint Element Francium (Fr) = “Frankenstein”  the biggest Fluorine (F)  smallest

  25. Skip to #10) Ions Packet pg. 3 • What are Ions? • An atom or group of bonded atoms that have a positive or negative CHARGE • Unlike atoms which are electrically NEUTRAL, ions have a net CHARGE (either positive or negative) • How are Ions Formed? • Formed by either adding or removing electrons to/from a neutral atom • Adding electronswill form a NEGATIVE ion • Removing electronswill form a POSITIVE ion

  26. #11) Cation Packet pg. 3 Cation: • Positively charged ion • Formed by the loss on 1 or more electrons from a neutral atom • On the periodic table, metals on the left side form cations

  27. #12) Anion Packet pg. 3 Anion: • Negatively charged ion • Formed by the addition of 1 or more electrons from a neutral atom • On the periodic table, nonmetals on the right side form anions

  28. Skip to #15) Ions & Valence Electrons Packet pg. 4 • The maximum number of valence electrons an atom may have is ______. When this occurs the atom is said to have a full valence shell. • The elements within group _______ (_____________) do have a full valence shell; This full valence shell is the reason why elements in this group are stable and chemically unreactive. 8 18/8 Noble Gases

  29. Skip to #15) Ions & Valence Electrons Packet pg. 4 • In fact, all chemically bonding occurs because every element wants to be stable and have a full valence shell. In order to achieve this, some atoms will lose electrons (to form ________) and some atoms will gain electrons ( to form _______) cations anions

  30. #8) Ionization Energy (IE) Packet pg. 2 • Definition: • the energy required to REMOVE an electron from a neutral atom of an element • Equation to represent this: A + energy  A+ + e– Atom + energy  ion w/positive charge + electron that was aka: “cation” removed

  31. #8) Ionization Energy (IE) Packet pg. 2 • Group Trend: • generally DECREASESdown a group. • it takes LESS energy to remove electrons from the atoms farther down in a group • Why?- Increasing Energy Level • As you move down a column, each successive element is larger and has electrons in higher energy levels so the electrons are farther away from the nucleus • Electrons farther away from the nucleus, are not as greatly attracted/pulled in by the nucleus, thus LESS energy is required to rip an electron away from the bigger atoms

  32. #8) Ionization Energy (IE) Packet pg. 2 • Period Trend: • generally INCREASESacross a period. • it takes MORE energy to remove electrons from the atoms across a period • Why?- Increasing Nuclear Charge • Again we see that moving across a period, each successive element has one more proton in its nucleus than the previous atom. • This increasing positive nuclear charge will cause a greater force of attraction and pull the electrons in closer to the nucleus, making it HARDER to remove the electrons and thus requiring MORE energy

  33. Ionization Energy: Hint Element Helium (He) the HIGHEST ionization energy b/c it takes the most energy to remove electrons In general, the Noble Gas Family has the HIGHEST ionization energies of any family. Why?

  34. #9) Electron Affinity (EA) Packet pg. 3 • Definition: • the energy associated with the addition of an electron to a neutral atom • Equation to represent this: A + e– A– + energy Atom + add an electron ion w/negative charge + energy aka: “anion” released

  35. #9) Electron Affinity (EA) Packet pg. 3 • Difference between EA and IE: • EA is basically the opposite of IE • Electron Affinity = energy is released when an electron is ADDED to an atom • Ionization Energy = energy is requiredwhen an electron is REMOVED from an atom

  36. #9) Electron Affinity (EA) Packet pg. 3 • Group Trend: • generally DECREASESdown a group. • Down a column, electrons added to atoms will have LESS energy released than the atoms above them • Why?- Increasing Energy Level • Down a group, atoms get larger with more energy levels; larger atoms contain more electrons inside them • The more electrons there are in the atom, the LESS likely the nucleus will be to pull/attractadditional electrons into the atom; when finally an electron is gained by the atom, LESS energy is released because it was harder to gain that electron

  37. #9) Electron Affinity (EA) Packet pg. 3 • Period Trend: • generally INCREASESacross a period. • Across a period, MOREenergy is released when electrons are added • Why?- Increasing Nuclear Charge • Again we see that moving across a period, each successive element has one more proton in its nucleus than the previous atom. • This increasing positive nuclear charge will cause a greater force of attraction and pull the electrons in closer to the nucleus, making itEASIER to gain an electron, thus releasing MORE energy when the electron is added to the atom

  38. Electron Affinity: Hint Element Fluorine (F) the MOST energy released when an electron is added to an atom The Halogens have the highest EA because it is really beneficial if they gain an electron  = lots of energy released

  39. #16) Electronegativity (EN) Packet pg. 5 • Definition: • the ability of an atom in a chemical compound to attract or pull electrons to itselffrom another atom

  40. #16) Electronegativity (EN) Packet pg. 5 • Group Trend: • generally DECREASESdown a group. • Going down a column, atoms are LESS able to attract electrons towards itself and away from other atoms in a compound • Why?- Increasing Energy Level • Down a column, atoms are larger with more & more electrons surrounding the nucleus • The more electrons surrounding the nucleus, the LESS of an attractive force the nucleus will have in order to pull electrons away from other atoms.

  41. #16) Electronegativity (EN) Packet pg. 5 • Period Trend: • generally INCREASESacross a period. • Across a period, atoms have a GREATER ability to attract electrons to itself • Why?- Increasing Nuclear Charge • Again we see that moving across a period, each successive element has one more proton in its nucleus than the previous atom. • This increasing positive nuclear charge will cause a stronger force of attraction between the nucleus and electrons (even electrons in other atoms) • The higher the nuclear charge, the GREATER the pull the atom will have in order to attract electrons towards itself

  42. #16) Electrongeativity (EN) Packet pg. 5 • To rank atoms based on this ability, each element is given an EN value between 0.0 and 4.0 (the Linus Pauling EN Scale) • The higher the value (closer to 4.0) the more EN the atom • The Noble Gases are not even listed or given EN values because they have ZERO ability to attract electrons toward themselves!!!

  43. Electronegativity: Hint Element Fluorine (F) the GREATEST ability to attract electrons to itself (very tiny w/a high nuclear charge) Linus Pauling EN Scale

  44. Electron Affinity and Electronegativity • EA & EN are similar to one another: • Both have the same group & period trends • Both deal with gaining or attracting electrons • Difference between EA and IE: • Electron Affinity = the ENERGY CHANGE associated with gaining and electron • Electronegativity = the ABILITY, or how well, and atom can attract an electron

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