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Unit 4: Electronic Structure and Periodic Trends

Unit 4: Electronic Structure and Periodic Trends. Mapping out electron configuration. Electrons are located in energy levels or shells around the nucleus of the atom Principle Energy Level (n) – Energy level where the electron(s) is(are) found (“ n” is the row of the PT)

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Unit 4: Electronic Structure and Periodic Trends

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  1. Unit 4: Electronic Structure and Periodic Trends

  2. Mapping out electron configuration • Electrons are located in energy levels or shells around the nucleus of the atom • Principle Energy Level (n)– Energy level where the electron(s) is(are) found (“n” is the row of the PT) Examples: n= 1, 2, 3, 4, 5, 6, 7 • As “n” increases, electrons: • Are farther from the nucleus • Have higheramounts of energy • Electrons want to be as close to the nucleus as they can so they can expend the LEAST amount of energy.

  3. Mapping out electron configuration Each energy level has a maximum capacity for electrons. Maximum in each level = 2n2; where n= 1, 2, 3, 4, 5, 6, 7 • Example: for n =1: max = 2(1)2 electrons • You solve for: n = 2 n = 3 n = 4 = 2 electrons

  4. Mapping out electron configuration • Each energy level contains sublevels or subshells • Sublevels– subdivisions of “n” • The first sublevel in any level is always s; the second is p; third is d; fourth is f. • The number of sublevels in each energy level is equal to n. Therefore: n = 1 has one sublevel (s) n = 2 has two sublevels (s, p) n = 3 has three sublevels (s, p, d) n = 4 has four sublevels (s, p, d, f) The maximum number of sublevels is 4.

  5. Mapping out electron configuration • Orbitals– make up each sublevel and are regions where electrons are located • Each sublevel has specific orbitals for electrons. • Orbitals are like seats, in a row (sublevel) in a section (energy level), in an arena (atom) • Each orbital can hold 2 electrons. • S has 1 orbital; maximum electrons = • P has 3 orbitals; maximum electrons = • D has 5 orbitals; maximum electrons = • F has 7 orbitals; maximum electrons =

  6. Orbitals P orbitals Sorbital Orbitals represent different orientations of each sublevel (or shape). D orbitals F orbitals

  7. Energy Level = Section Sublevels = row in a section Orbitals = seats in a row in a section Mapping out electron configuration Nucleus

  8. Energy Levels, Sublevels, and Orbitals

  9. Energy Levels, Sublevels, and Orbitals

  10. Definition of Electron Configuration • Is a form of notation showing how electrons are distributed among atomic orbitals and energy levels. • The format for writing electron configurations includes a series of [Number][letter] [superscript number] 1s2

  11. Standard Notation of Fluorine Number of electrons in the sub level 2,2,5 1s22s22p5 Principle Energy Level Numbers 1, 2, 2 Subshell

  12. Decoding a Complicated Message • The energy sub levels are filled out in a specific order • Start at the beginning of each arrow and follow it to the end. • Fill the sub levels as the arrow passes through it. • The order is as follows: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d,7p.

  13. Blocks in the Periodic Table s starts at 1, p starts at 2, d starts at 3, and f starts at 4

  14. Writing an Electron Configuration- Selenium • Find the element you are looking for on the periodic table • Always start your configuration at hydrogen

  15. Selenium Continued 3. Write the energy number and letter of the subshell, then as a superscript write the number of electrons held in that subshell 1s22s22p63s23p64s23d104p4

  16. Practice Time • Write the electron configuration of the following elements • Beryllium • Cadmium • Bromine • Uranium • Iron

  17. Electron Configuration with Ions • When doing electron configuration with ions, write the configuration and then add (if it is an anion) or subtract (if it is a cation) the charge from the number of electrons (superscript) • The electron configuration of an ion should match a noble gas configuration • Example: O-2 • Oxygen: 1s22s22p4 • Oxygen ion: 1s22s22p6 Add two electrons

  18. Practice Time • Write the electron configuration of the following ions • Ca+2 • P-3 • Mg+2 • F-1

  19. Noble Gas Notation(AKA shorthand notation) • Use the last noble gas that is located in the periodic table right before the element. • Write the symbol of the noble gas in brackets. • Write the remaining configuration after the brackets. • Ex: • Fluorine: • [He] 2s2 2p5

  20. Practice Noble Gas Notation(AKA shorthand notation) • Write the noble gas configuration for the following elements: • Chlorine • Mercury • Lanthanum (La) • Argon • K+ • F-

  21. Electrons and Energy Levels • Electrons in their energy levels are considered to be in their ground state • If electrons are given enough energy, they absorb the energy and they can jump up in energy levels • We call this the excited state • Electrons usually absorb energy from sources such as light and or fire • When they fall back to their ground state energy level, they release a photon of light • A photon of light is a tiny particle of light • The color of the light corresponds to the amount of energy the electron released when it goes from its excited state to its ground state

  22. Energy absorbed Energy released as a photon

  23. Energy Levels • Electrons can only move to certain energy levels • Every element releases a photon or photons with unique amounts of energies • Each element releases specific wavelength(s) of light because each element has a unique electron configuration • The color(s) of light released can be used to identify the element • Ex: Was used to identify elements in the sun

  24. Low energy High energy Low Frequency High Frequency Electromagnetic Spectrum- the range of wavelengths or frequencies over which electromagnetic radiation extends. X-Rays Radio waves Microwaves Ultra-violet Gamma Rays Infrared . Visible Light (ROY G BIV) Long Wavelength Short Wavelength

  25. Orbital Diagrams Orbital Diagrams • Orbital diagrams are very similar to electron configurations • They show the placement of electrons in an atom • Boxes or lines are used to represent orbitals S has 1 orbital P has 3 orbitals D has 5 orbitals F has 7 orbitals • The energy level and sublevel are written under the boxes (example 1s or 2p) • Electrons are represented by half arrows • ONLY 2 arrows per box pointing opposite directions

  26. Three rules are used to build the Orbital Diagrams: • Aufbau principle: Electrons occupy orbitals of lower energy levels first • Electrons are attracted to the nucleus and will occupy orbitals closest to the nucleus first • Hund’s Rule: In a set of orbitals, the electrons will fill the orbitals in a way that would give the maximum number of parallel spins (maximum number of unpaired electrons). • Analogy: Students could fill each seat of a school bus, one person at a time, before doubling up • Pauli Exclusion Principle: No two electrons can have the same set of quantum numbers • An orbital can hold only two electrons and they must have opposite spin • One arrow points up, the other points down

  27. Orbital Diagrams Orbital Diagram Example • Draw the orbital diagram for Nitrogen: • Write the configuration for nitrogen • Draw the boxes you need and label them 3. Fill in arrows following Aufbau’s, Hund’s, and Pauli’s rules Why are the arrows in 2p in separate boxes? 1s 2s 2p

  28. Practice • Draw the orbital diagram for Boron • Draw the orbital diagram for Bromine (you may use short hand just don’t for get to include the noble gas before the element in brackets [ ]) • Draw the orbital diagram for Titanium

  29. Quantum Numbers • 4 Quantum Numbers • The 4 quantum numbers describe the exact location of an electron • It is like the “address” of an electron • Quantum numbers can also be written as ranges to describe a set of electrons

  30. 1st Quantum Number • Principal Quantum Number (n) – Describes the size and energy level of an orbital • The values of n are integers ≥ 1 • The greater the n… the greater the Energy… the greater the distance from nucleus

  31. 2ndQuantum Number • Angular Momentum Quantum Number (l) - Defines the subshell or the shape of the orbital • Allowed values of l are integers ranging from 0 to n-1 • s subshell l = 0 • p subshell l = 1 • dsubshell l = 2 • fsubshell l = 3

  32. 3rd Quantum number • Magnetic Quantum Number (ml) – describes the 3D orientation of an orbital • Values of ml range from - l to + l • Each box in the orbital diagram indicates a different mlvalue • The center box is always zero • Example: l value for p is 1 so ml range from -1 to 1

  33. 4th Quantum number • Spin Quantum Number (ms) – describes the electron’s magnetic spin direction • Values for ms are + ½ (spin up/ clockwise) or – ½ (spin down/ counterclockwise) • The direction of the arrow in the orbital diagram determines the ms value

  34. Quantum Numbers - Practice • What are the 4 quantum numbers of: • The last electron in P • Sodium’s valance electron • The 9th electron in Cl • The 3rd electron in O • The last electron in Cu • The last electron in Fe

  35. Summary for electrons… • Draw all of the following for H • Lewis Dot Diagram—simplified notation showing only valence electrons • Electron Configuration—the shorthand (in numbers and letters) that represent the locations of electrons in the atom • Orbital Diagram— a visual representation of electron configuration • Quantum Numbers- The exact location of an electron or of valance electrons

  36. Periodic Trends • There are trends that occur periodically… that’s why it is called the Periodic Table • These trends are general patters, they do have some exceptions

  37. Atomic Radius • Atomic Radius is defined ½ the distance between nuclei of atoms bonded together

  38. Atomic Radius Periodic Trends • Atomic radius increases moving down a group • Atomic radius decreases moving across a period • Where is the element with the largest atomic radius located? What element is it? “9:30 trend” (arrows point in the increasing direction)

  39. Explaining Atomic Radius Trend • Why are atoms larger going down a group? • Moving down a group the amount of energy levels and the number of electrons increases. • Why do atoms get smaller going to from left to right across a period? • Moving across a period there is an increases in nuclear charge (more protons). • This pulls the electrons closer to the nucleus.

  40. Practice- Atomic Radius Which of the following has the largest atomic radius? • Cobalt or nickel • Phosphorous or nitrogen • Potassium or oxygen List the following in order of increasing atomic radius. • Fluorine, gallium, and carbon • Barium, iodine, and gold List the following in order of decreasing atomic radius. • Aluminum, sulfur, and sodium

  41. Ionic Radius • Ionic radius is the radius of the atom once it becomes an ion • A cation has a smaller radius than its atom • There are less electrons in the ion than the atom, which take up less place • An anion has a larger radius than its atom • There are more electrons in the ion than the atom, which repel each other and take up more space

  42. Practice • Which of the following has a larger radius? • Calcium atom or calcium ion? • Manganese atom or manganese ion? • Selenium atom or selenium ion? • Chlorine atom or chlorine ion?

  43. Ionization Energy “3 o’clock trend” • Ionization Energy is the energy required to completely remove an electron from a atom or an ion. • A large ionization energy means a lot of energy is required to remove the electron (hard to remove). • Decreases moving down a group • Increases moving across a period *Ignore noble gases, they have the highest ionization energies (they do not become ions)

  44. Explaining Ionization Energy Trend • Why does ionization energy decrease moving down a group? • Moving down a group the atoms are larger and electrons are further from the nucleus • The further away electrons are, the less they feel the attraction from the nucleus and are easier to remove • Why does ionization energy increase moving across a period? • Moving across a period there is an increases in nuclear charge • The greater the nuclear charge the stronger the attraction is to the electrons, making them harder to remove

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