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CHEMICAL FORMULAS AND BONDING

CHEMICAL FORMULAS AND BONDING. Ions and Molecules. Learning Target. Understand octet rule and how it applies to oxidation state of an ion. What do all atoms need in order to be stable?. OCTET RULE. Elements will gain or lose electrons in order to obtain a noble gas valence (full shell).

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CHEMICAL FORMULAS AND BONDING

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  1. CHEMICAL FORMULAS AND BONDING Ions and Molecules

  2. Learning Target • Understand octet rule and how it applies to oxidation state of an ion. • What do all atoms need in order to be stable?

  3. OCTET RULE • Elements will gain or lose electrons in order to obtain a noble gas valence (full shell)

  4. Remember…. • It takes energy to gain or lose electrons • Nature wants to move towards the path of least resistance….. ……because it requires less energy

  5. Gain or Lose Electrons? • Depends on the nearest noble gas. Noble Gases

  6. OXYGEN • Oxygen has 6 valence electrons • Is it easier to gain 2e- or lose 6e-? • Easier to gain 2e-

  7. SODIUM • Sodium has 1 valence e- • Is it easier to gain 7e- or lose 1e-? • Easier to lose 1e-.

  8. Try some…… Gain Gain Lose Neither Both Lose Lose Lose • Sulfur • Fluorine • Potassium • Argon • Carbon • Hydrogen • Calcium

  9. OXIDATION NUMBER • The possible charge an atom could obtain by gaining or losing electrons or the number of electrons an element will donate/accept in a bond. • Remember electrons are negative

  10. OXYGEN • Oxygen gains 2e- • So its charge is 2- O2- Oxidation Number

  11. SODIUM • Sodium loses 1e- • So its charge is 1+ Na+

  12. TRY SOME….. • Magnesium • Boron • Bromine Loses 2e-, Mg2+ Loses 3e-, B3+ Gains 1e-, Br-

  13. 2 TYPES OF CHEMICAL BONDS WE WILL TALK ABOUT • Ionic compounds • Covalent compounds • 2 terms that go with the ionic and covalent: • Empirical • Molecular

  14. Review • What is the oxidation number for Rubidium? Selenium? 2. Do metals have (+) or (–) oxidation numbers? 3. Do non metals have (+) or (–) oxidation numbers?

  15. Review • 1. What rule states that elements will gain or lose electrons to obtain a noble gas valence? • Octet Rule • 2. Do non metals have (+) or (–) oxidation numbers? • (-) • 3. Do metals have (+) or (–) oxidation numbers? • (+) • 5. What is the oxidation number for barium? Iodine? • Ba2+ and I-

  16. Goal of Today • Know how to write formulas for Ionic Compounds.

  17. Ionic Bond • Electrons are everywhere – static is a good example • Positive ion is attracted to a negative ion in an ionic bond • What kind of elements?

  18. Ionic Compound • Made up of ions • Electrically neutral • Charges must equal each other • Bonds between metals (+ charge) and nonmetals (- charge)

  19. Ionic Bonding • Ionic bonds occur between positive metal ions (cations) and negative nonmetal ions (anions) • Made up of ions • Electrically neutral • Charges must equal each other

  20. Properties of Ionic Compounds • Strong bonds • High melting points • Brittle • Soluble in water • Their solutions are good conductors of electricity • Solids at 20°C

  21. Formation of Ionic Compounds • Use • The Octet rule • Lewis Dot Diagrams • Crisscross Method

  22. Monatomic Cations • Positive 1, 2, or 3 • Transition metals can vary, and even have a charge of 4+ • Use element name • Use a Roman numeral with any metal that varies in charge • Example Copper(I) is Cu+1 and Copper (II) is Cu2+

  23. Monatomic Anions • Can be negative 1,2, or 3 • Change element name to “-ide” ending • Example: chloride

  24. Polyatomic Ions • Two or more atoms that are bonded together and carry a single charge • Names are on the handout • Most are negative with one positive • Usually end in “-ate” or “-ite” • Example: NO3- is nitrate

  25. Formulas for Binary Compounds • Contain a monatomic cation (metal) and a monatomic anion (nonmetal) • Metal is first • Charges must add to “0” • Use subscripts to get the value to “0” • Why is sodium chloride NaCl? • Try some

  26. Write the formulas for the following • potassium iodide • barium chloride • lithium bromide • calcium hypochorite • chromium (III) sulfide • gold (III) bisulfate

  27. Review • What are negative ions called? 2. What are positive ions called? 3. Write the formula for niobium(V) phosphate

  28. Goal of Today • Understanding naming of Ionic Compounds and begin our “Ionic Bonding Puzzle Lab”

  29. Naming Ionic Compounds • Use NaCl as a good example • Metal first, then nonmetal • Ends in “-ide” if binary (2 elements only) • Use polyatomic name • Use Roman numeral if necessary • Suspect every transition metal to possibly have a Roman numeral

  30. Review Questions 1. Write the formulas for: beryllium carbonate silver nitrate 2. Name the following ionic compound: Cu(HSO4)2

  31. GOAL OF TODAY Understand how covalent bonds form and know how to draw Lewis structures to represent covalent compounds.

  32. COVALENT BONDING • A covalent bond is formed by a shared pair of electrons between two atoms. • A group of atoms that are united by covalent bonds is called a MOLECULE • Most of what you see is covalently bonded

  33. Describing a molecular bond • Molecular formula: tells you how many atoms and which kind are in each molecule. Glucose C6H12O6 • Empirical formula: gives you the ratio of the atoms in a molecule. Glucose: C1H2O1

  34. Covalent Bonding • Single bonds share 2 electrons. Example: Ammonia (H-N-H) H • Double bonds share 4 electrons: Formaldehyde H-C=O H • Triple bonds share 6 electrons Ethyne H-C=C-H

  35. Exit Questions What is the difference between ionic bonds and covalent bonds? Draw the Lewis structure for N2H2

  36. GOAL OF TODAY • Know that there are exceptions to the octet rule. • Know how to determine and notate polarity of chemical bonds.

  37. Exceptions to the Octet Rule • Some elements are satisfied with fewer than 8 electrons (6 or 4) • Some structures can only be drawn with 7 electrons • Compounds with Beryllium (Be) and Boron (B) may have less than an octet BCl3 • 3rd Row elements or lower may exceed octet rule. SF4. • Odd number of electrons cannot follow the octet rule. I.e., NO or CO • These substances can be short-lived and reactive – called free radicals

  38. Chloroform CHCl3 Draw the Lewis structure for chloroform

  39. Draw the Lewis structure for PF3

  40. Properties of Covalent Bonds • Electrons are not necessarily shared equally. This depends on the electronegativity of an atom (atom’s attraction for electrons)

  41. Polar and Nonpolar Covalent Bonds • Polar bonds: 1 atom is significantly more electronegative than the other one. One side of the bond is slightly positive the other negative. Example: Water • Nonpolar bonds: Both atoms have similar electronegativities. Example:

  42. Bond Type by Electronegativity • First, Find the difference of electronegativities of the atoms.Then apply the information of the table below

  43. Goal for Today • Review bond polarity • Know the difference in naming ionic and covalent compounds.

  44. Look at the table on page 241. • What is the electronegativity difference between hydrogen and oxygen (H20)? • 1.4 (polar covalent) • What is the electronegativity difference between sodium and chlorine (NaCl)? • 2.1 approximately (ionic) • What is the electronegativity difference between 2 nitrogen atoms (N2)? • 0 (non polar covalent)

  45. Properties of Molecular Substances • Weak bonds • Can be solids, liquids, or gases at 20°C • Lower melting points • Poor conductors of electricity • Soft, not hard and brittle • Some are soluble in water.

  46. Naming Molecular Compounds • Similar to naming ions • Numerical prefixes are used. Example: CO2=Carbon Dioxide • Do not use a prefix for one • - ide is added to the more electronegative element • Some elements have common names: Diatomics like O2= oxygen; not dioxide NH3= ammonia; not Nitrogen tetrahydride H2O is water not dihydrogenmonoxide

  47. HYDRATES Ionic compounds that absorb water into their solid structures. Example: CuSO4. 5 H2O (s) Copper (II) Sulfate pentahydrate

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