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Chapters 3 & 4

Chapters 3 & 4. Chemical Bonding. Valence Electrons. Outermost electrons s and p electrons for main group elements Responsible for chemical properties of atoms Participate in chemical reactions. Valence Electron. Core Electrons. Problems.

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Chapters 3 & 4

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  1. Chapters 3 & 4 Chemical Bonding

  2. Valence Electrons • Outermost electrons • s and p electrons for main group elements • Responsible for chemical properties of atoms • Participate in chemical reactions Valence Electron Core Electrons

  3. Problems • Write out the electron configurations for the following elements and identify how many core and valence electrons each has. • Mg • S • Br • Kr

  4. Lewis Dot Structures • LDS: a representation of an atom using its chemical symbol surrounded by dots that signify valence electrons

  5. Li: [He]1s1 Na: [Ne]2s1 K: [Ar]3s1

  6. Octet Rule • Octet Rule: the tendency for atoms to seek 8 electrons in their outer shells • Natural electron configuration of the Noble Gases • Done by gaining, losing, or sharing electrons • Increases stability • H and He seek a “Duet”

  7. 3.2 IonsA. Cations and Anions • Metals, like sodium (Na) and magnesium (Mg), form cations. • By losing one, two, or three electrons, an atom forms a cation with a completely filled outer shell of electrons. • Nonmetals, like chlorine (Cl), form anions. • By gaining one, two, or three electrons, an atom forms an anion with a completely filled outer shell of e−. • The octet rule: a main group element is especially stable when it possesses an octet of e− in its outer shell. octet = 8 valence e−

  8. 3.2 IonsB. Relating Group Number to Ionic Charge forMain Group Elements • Elements in the same group form ions of similar charge. • For metals in groups 1A, 2A, and 3A, the group number = the charge on the cation. • For nonmetals in Groups 6A and 7A, the anion charge = 8 – the group number.

  9. 3.2 Ions

  10. 3.2 IonsC. Metals with Variable Charge

  11. Introduction to Bonding • Bonding is the joining of two atoms in a stable arrangement. • Elements will gain, lose, or share electrons to reach the electron configuration of the noble gas closest to them in the periodic table. • There are two different kinds of bonding: • Ionic bonds result from the transfer of electrons • from one element to another. • Covalent bonds result from the sharing of • electrons between two atoms.

  12. 3.1 Introduction to Bonding Ionic bonds form between: • A metal on the left side of the periodic table. • A nonmetal on the right side of the periodic table.

  13. Ionic Bonding • Ionic Bond: a bond formed through the transfer of one or more electrons from one atom or group of atoms to another atom or group of atoms

  14. Formula Unit

  15. 3.5 Physical Properties of Ionic Compounds • Ionic compounds are crystalline solids with very high meltingand boiling points. • When ionic compounds dissolve in water, they separate into cations and anions, increasing the conductivity of the solution. +  solution NaCl water

  16. 3.3 Ionic Compounds • When the cation and anion have different charges, use the ion charges to determine the number of ions of each needed. Ca2+ Cl− A +2 charge means 2 Cl−anions are needed. A -1 charge means 1 Ca2+ cation is needed. Cl− + CaCl2 Ca2+ 2 Cl− for each Ca2+

  17. Ionic Compounds: compounds composed of oppositely charged ions that are held together by their attraction to each other • Metal + Non-metal • NaCl • Metal + Polyatomic Ion • NaNO3 • Polyatomic Ion + Non-metal • NH4Cl • Polyatomic Ion + Polyatomic Ion • NH4NO3 • Net charge on compound equal to zero

  18. Oxyanions

  19. Rules For Naming Ionic Compounds • Name the cation by its elemental/polyatomic name • If the metal is a transition metal with a variable charge, indicate its charge with a Roman Numeral in parentheses • Next, name the anion and change its ending to “-ide” • If the anion is polyatomic, do not change the ending to “-ide” • Do NOT use prefixes (mono, di, tri etc.) to indicate how many of each atom are present

  20. 3.4 Naming Ionic CompoundsB. Naming Anions Anions are named by replacing the ending of the element name by the suffix “-ide.”

  21. 3.6 Polyatomic IonsA. Writing Formulas for Ionic Compounds withPolyatomic Ions • When a cation and anion of unequal charge combine, use the ionic charges to determine the relative number of each ion that is needed. Mg2+ + Mg(OH)2 OH− zero overallcharge −1 charge means 1 Mg2+anion is needed. +2 charge means 2 OH− anions are needed.

  22. Problems Write the name for the following compounds: • KI • MgBr2 • Al2O3 • FeCl2 • CaSO4 • Ba(NO2)2 • Cu(NO3)2

  23. Write the Formula for the following ionic compounds: • Sodium Fluoride • Calcium Sulfite • Calcium Chloride • Iron (III) Oxide • Cobalt (II) Hydroxide • Ammonium Bromide • Ammonium Carbonate • Aluminum Carbonate

  24. Iron (II) Chloride Iron (III) Chloride

  25. Covalent Compounds • Covalent Compounds: compounds composed of atoms bonded to each other through the sharing of electrons • Electrons NOT transferred • No + or – charges on atoms • Non-metal + Non-metal • Also called “molecules” • Examples: • H2O • CO2 • Cl2 • CH4

  26. orH-H Duet or

  27. Name the first non-metal by its elemental name Add a prefix to indicate how many Name the 2nd non-metal and change its ending to “-ide” Add a prefix to indicate how many Naming Covalent Compounds

  28. Problems Write the name of the following compounds: • CO • NI3 • N2O • SF6 • B2O3

  29. Write the formula for the following compounds: • Phosphorous Pentachloride • Nitrogen Monoxide • Dinitrogen Tetroxide • TetraphosphorousDecoxide

  30. Problems • KCl • Na2S • H2O • SO2 • K3PO4 • FeCl3 • (NH4)2SO4 • SCl2 • Cu(OH)2 • P2O5

  31. Sodium Iodide • Aluminum Sulfate • Phosphorous Pentabromide • Magnesium Nitride

  32. Naming Acids • Acids that do not contain oxygen • Begin the name with “hydro” • Name the anion, but change the ending to “-ic” • Add “acid” on the end • HCl • HF

  33. Acids that contain oxygen • Do not put “hydro” at the beginning • Begin the name with the anion • If the anion has the ending “-ate,” change this to “-ic acid” • If the anion has the ending “-ite,” change this to “-ous acid” • HClO4 • HClO3 • HClO2 • HClO

  34. Problems • Name the following • HBr(g) • HBr(aq) • HNO2(aq) • HNO3(aq) • HI (aq) • HI (g) • H2CO3 (aq) • H3PO4 (aq) • H3PO3 (aq) • HCN (aq)

  35. Molecular Structures

  36. Ball & Stick Models Space-Filling Models Water Methane

  37. Ethanol

  38. C6H6

  39. Lewis Dot Structures • Count the total number of valence electrons in the molecule. Ex: PCl3 • Use atomic symbols to draw a proposed structure with shared pairs of electrons. • Atoms don’t tend to bond to other atoms of the same element when they can avoid it • Exception: Carbon

  40. Place lone pair electrons around each (except H) to satisfy the octet rule, beginning with the terminal atoms • Place any leftover electrons on the central atom • If the number of electrons around the central atom is less than 8, change single bonds to the central atom to multiple bonds (double or triple). • Ex: CH2O

  41. Problems Draw the LDS’s for the following molecules: • Cl2O • C2H4 • C2H6O

  42. What Things Like To Do • Halogens • Like to be terminal • Like to have one single bond and 3 lone pairs (non-bonding electrons) • Carbon • Likes to have 4 single bonds and no lone pairs • A double bond counts as two singles • A triple bond counts as three singles • Likes to be central • Likes to bond to other carbons

  43. Silicon • Likes to do what carbon does • Oxygen • Likes to have two single bonds and 2 lone pairs • Sulfur • Likes to do what oxygen does • May expand its octet • Nitrogen • Likes to have 3 single bonds and one lone pair

  44. Phosphorous • Likes to do what nitrogen does • May expand its octet • Hydrogen • Likes to be terminal with only one single bond • No lone pairs! • Boron • Likes 3 bonds and no lone pairs (sextet) https://phet.colorado.edu/en/simulation/molecule-shapes

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