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Periodic Trends Bellringer

Periodic Trends Bellringer. 1 ) Which has a larger size, a Mg atom or a Mg ion? 2) smaller ionization energy, K or Br? 3) smaller size, F atom or an I atom? 4) smaller electronegativity, O or Se? 5) less shielding, Ca or Sr ?. Answers. 1) Mg atom 2) K 3) F 4) Se 5) Ca.

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Periodic Trends Bellringer

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  1. Periodic Trends Bellringer • 1) Which has a larger size, a Mg atom or a Mg ion? • 2) smaller ionization energy, K or Br? • 3) smaller size, F atom or an I atom? • 4) smaller electronegativity, O or Se? • 5) less shielding, Ca or Sr?

  2. Answers • 1) Mg atom • 2) K • 3) F • 4) Se • 5) Ca

  3. Ch. 7 Ionic and Metallic Bonding

  4. Valence Electrons • Valence electrons = electrons in the highest occupied energy level of an element’s atoms • This # largely determines the chem. properties of the element • To find the # of valence e-’s in an atom of a representative element, look at its group # • Group 1 has 1, group 2 has 2, group 13 or 3A has 3, … • Electron dot structures (Lewis dot diagrams) = diagrams that show valence electrons as dots • *Draw Lewis dot structures*

  5. Octet Rule • Octet rule = atoms gain or lose electrons to achieve a stable level of usually 8 • atoms of metals tend to lose their valence electrons leaving a complete octet in the next-lowest energy level • Atoms of some nonmetals tend to gain electrons or to share electrons w/ another nonmetal to achieve a complete octet

  6. Cations • Cations (+ charged atom) are formed when they lose one or more valence electrons (ionization) in order to become stable • Usually lose from 1 to 3 valence electrons tops • Can use electron configurations to illustrate the point but let’s use electron dot diagrams for ease • Exceptions are due to having to lose or gain too many valence electrons to achieve a noble gas state, so: • Some atoms attain a pseudo noble-gas electron configurations – Cu - copper, Ag - silver, Au - gold, Cd - cadmium Two atoms are walking down the street. Says one atom to the other, "Hey! I think I lost an electron!" The other says, "Are you sure??" "Yes, I'm positive!"

  7. Anions • Anions (- charged atoms) are formed when they gain 1 or more valence electrons • Typically suffix is –ide • Halide ions (F-, Cl-, Br-, I-) are halogens that gain 1 e- • CP 7.1, PP 1-2 pg. 193, sect. assessment 7.1 pg. 193 3-11

  8. Common ions Cation Name Anion Name* H 1+ hydrogen H 1-hydride Li 1+ lithium F 1-fluoride Na 1+ sodiumCl1-chloride K 1+ potassium Br 1-bromide Cs 1+ cesium I 1-iodide Be 2+ beryllium O 2-oxide Mg 2+ magnesium S 2-sulfide Al 3+ aluminum Ag 1+ silver *The root is given in color.

  9. Ionic Compounds • Anions and cations are held together by opposite charges. • Ionic compounds are called salts • Electrically neutral • Simplest ratio is called the formula unit • The bond is formed through the transfer of electrons (called an ionic bond) • Electrons are transferred to achieve noble gas configuration • Most are crystalline solids at room temp. • High melting points – large attractive forces result in a very stable structure • Good conductors of electricity when melted or dissolved in water Overheard at the mall Teen 1: Did you hear oxygen and magnesium got together?? Teen 2: OMg! Your mama's so ugly Your mama's so ugly...even Fluorine won't bind to her

  10. Formulas • Chemical formula = shows the kinds and #s of atoms in the smallest representative unit of a substance • NaCl is the chem. formula for sodium chloride • Formula unit = the lowest whole-number ratio of ions in an ionic compound • MgCl2, NaCl, AlBr3 • CP 7.2, PP 12-13 pg. 196, • 7.2 sect. assessment pg. 199 14-22

  11. Ch. 7 Bellringer • Write the Lewis electron-dot symbol for each of the following • A) sodium • B) fluorine • C) magnesium ion (Mg2+) • Write the chemical formula that results when the following pairs of ions combine to form an ionic bond • D) Mn4+ and O2- • E) Li1+ and Cl1-

  12. Metallic Bonds and Properties • Metallic bonds = consist of the attraction of the free-floating valence electrons for the positively charged metal ions • How atoms are held together in the solid • Metals hold onto their valence electrons very weakly • Think of them as positive ions floating in a “sea of electrons” • Electrons are free to move through the solid • Metals conduct electricity • Malleable - hammered into shape (bend) • Ductile - drawn into wires • Electrons allow cations to slide by each other under pressure • Metals are crystalline structures and atoms are arranged in very compact and orderly patterns *Cu vs. Cu compounds hammer demo – pg. 202*

  13. Alloys • Alloys = mixtures composed of 2 or more elements, at least one of which is metal • Brass – copper and zinc • Sterling silver – silver and copper • Bronze – copper and tin • Steel – Fe, Cr, and others • Impt. b/c their properties are often superior to those of their component elements (usually cheaper as well) *Making an alloy DEMO* - pg. 205

  14. Ch. 8 Covalent Bonding

  15. Molecules • Covalent bond = formed by sharing electrons b/w 2 or more atoms • Molecule = a neutral group of atoms joined by covalent bonds • Diatomic molecule = a molecule consisting of 2 of the same atoms • H2, N2, O2, F2, Cl2, Br2, I2 are the diatomic molecules in nature • Molecular compounds tend to have relatively lower melting and boiling points than ionic compounds • Most molecular compounds are composed of atoms to 2 or more nonmetals • *Make a table comparing covalent bonding and ionic bonding*

  16. Molecular Formula • Molecular formula = the chemical formula for a molecular compound • Shows how many atoms of each element a molecule contains • H2O, CO2, C2H6, O2 • 8.1 Sect. assessment, Pg. 216 1-6

  17. Covalent Bonding • Electron sharing usually occurs so that atoms attain the electron configuration of noble gases • Combos of atoms of the nonmetals and metalloids in 4A, 5A, 6A, and 7A are likely to form covalent bonds • Single covalent bond = 2 atoms sharing 1 pair of electrons • H2 • 2 dots in an electron dot diagram represents this bond • A dash in a structural formula represents this bond • A molecular formula does NOT show this bond only the # of atoms • Halogens form these bonds in their diatomic molecules • Unshared pair = pair of valence electrons not shared in an electron dot diagram • *CP 8.1, PP 7-8 pg. 220

  18. Double and Triple Covalent Bonds • Double covalent bond = involves 2 shared pairs of electrons • oxygen • Triple covalent bond = involves 3 shared pairs of electrons • nitrogen

  19. Coordinate covalent bonds • The shared electron pair comes from one of the bonding atoms • CO – look at pg. 223 • Polyatomic ion = a tightly bound group of atoms that has a positive or negative charge and behaves as a unit • NH4+ • *c.p. 8.2, p.p. 9-12 pg. 225 A sign outside the chemistry hotel reads "Great Day Rates, Even Better NO3-'s"

  20. Bond Dissociation Energies • Bond dissociation energy = the energy required to break the bond b/w 2 covalently bonded atoms • a large bond dissociation energy corresponds to a strong covalent bond • H2 = 435 kJ/mol, C-C single bond = 347 kJ/mol, C=C double bonds = 657 kJ/mol, and triple bonds = 908 kJ/mol

  21. Resonance Structure • Resonance structure = a structure that occurs when it is possible to draw 2 or more valid electron dot structures that have the same # of electron pairs for a molecule or ion • Double-headed arrows are used to connect • Double bonds are usually shorter than single bonds but they are the same lengths b/c it is an avg. of the 2 structures • resonance

  22. Exceptions to the Octet Rule • Cannot be satisfied in molecules whose total # of valence electrons is an odd #. • NO2, ClO2, NO • Sometimes w/ an even # as well (Fewer or more) • BF3 • PCl5 • SF6

  23. *8.2 sect. assessment 13-22 pg. 229* *Exceptions to the octet rule: A resonance hybrid teacher DEMO*

  24. Molecular Orbitals • Molecular orbitals = orbitals that apply to the entire molecule • just as an atomic orbital belongs to a particular atom, a molecular orbital belongs to a molecule as a whole • Bonding orbital = a molecular orbital that can be occupied by 2 electrons of a covalent bond

  25. Sigma Bonds • Sigma bonds = formed when 2 atomic orbitals combine to form a molecular orbital that is symmetrical around the axis connecting 2 atomic nuclei • Atomic orbitals overlap end to end • Two s orbitals can combine to form a molecular orbital • H2 • Two p orbitals • F2 • The attractions b/w electrons and nuclei of two atoms overpower the repulsions b/w the 2 nuclei or b/w the 2 sets of electrons = covalent bond (stable molecule)

  26. Pi Bonds • Pi bond = a covalent bond in which the bonding electrons are most likely to be found in sausage-shaped regions above and below the bond axis of the bonded atoms • Orbitals overlap side by side • Atomic orbitals in pi bonding overlap less than in sigma bonding – weaker than sigma bonds • A typical double bond consists of 1 sigma & 1 pi bond, triple bond is 1 sigma & 2 pi bonds • In special cases, they form w/o any sigma bonds

  27. VSEPR Theory • VSEPR theory (valence-shell electron-pair repulsion theory) = the repulsion b/w electron pairs causes molecular shapes to adjust so that the valence-electron pairs stay as far apart as possible • Explains the actual 3-D shapes of molecules • http://gold.chem.wwu.edu/sdgchem121/Docs/WWUDocs/geometry.pdf • http://www.youtube.com/watch?v=i3FCHVlSZc4

  28. Hybrid Orbitals • Hybridization = several atomic orbitals mix to form the same total number of equivalent hybrid orbitals • Single bonds – one 2s and three 2p orbitals mix to form four sp3 hybrid orbitals • Double bonds – one 2s and two 2p orbitals mix to form three sp2 hybrid orbitals • Triple bonds – one 2s and 1 2p orbitals mix to form two sp hybrid orbitals

  29. Electronic Geometry • 2 electron densities – Linear • 3 e- densities – Trigonal Planar • 4 e- densities – Tetrahedral • 5 e- densities – TrigonalBipyramidal • 6 e- densities – octahedral

  30. Linear • Atoms connected in a straight line • All molecules w/ 2 atoms and some w/ 3 • 180⁰ bond angle • HCl, CO2 • Hybridization – sp • 2 bonds/0 lone pairs

  31. Trigonal Planar • Triangular flat • 120⁰ bond angle • BCl3 • Hybridization – sp2 • 3 bonds/0 lone pairs • Bent – 2 bonds/1 lone pair • 118⁰ • SO2

  32. Tetrahedral • 4 surfaces • 109.5⁰ bond angle • CH4 • Hybridization – sp3 • 4 bonds/0 lone pairs • Pyramidal (trigonal pyramidal) – 3 bonds/1 lone pair • 107⁰ • NH3 • Bent – 2 bonds/2 lone pairs • 105⁰ • H2O

  33. TrigonalBipyramidal • 90, 120, and 180⁰ bond angles • PF5 • Hybridization – sp4 • 5 bonds/0 lone pairs • See-saw – 4 bonds/0 lone pairs • 90, 120, and 180⁰ • SF4 • Tee-shaped – 3 bonds/2 pairs • 90 and 180⁰ • ClF3 • Linear – 2 bonds/3 lone pairs • 180⁰ • XeF2

  34. Octahedral • 90 and 180⁰ bond angles • SF6 • Hybridization – sp5 • 6 bonds/0 lone pairs • Square pyramidal – 5 bonds/1 lone pair • 90 and 180⁰ • BrF5 • Square Planar – 4 bonds/2 lone pairs • 90 and 180⁰ • XeF4 • 8.3 sect. assessment pg. 236 23-29

  35. .. .. .. S O O O C O O S N F O O F F F F F F F F F F P S Xe F F F F F Cl F F F F F The VSEPR Model The Shapes of Some Simple ABn Molecules SO2 Linear Bent Trigonal planar Trigonal pyramidal AB6 T-shaped Square planar Trigonal bipyramidal Octahedral

  36. Bond Polarity • Nonpolar covalent bonds = equal sharing of electrons • H2, O2, N2, Cl2 • Polar covalent bond = unequal sharing of electrons • HCl, H2O • The more electronegative atom attracts electrons more strongly and gains a slightly negative charge, the less electronegative atom has a slightly positive charge

  37. Table 8.3 pg. 238 • C.P. 8.3, P.P. 30-31 pg. 239

  38. Polar Molecules • Polar molecule = one end of the molecule is slightly negative, the other is slightly positive • Dipole = a molecule that has 2 poles w/opposite charges • The shape of a molecule and the polarity of its bonds together determine whether the molecule is polar or nonpolar • Equal and opposite directions arrows cancel = nonpolar • Arrows same direction = polar

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