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Ch 6 Chemical Names & Formulas

Ch 6 Chemical Names & Formulas

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Ch 6 Chemical Names & Formulas

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  1. Ch 6 Chemical Names & Formulas P132 Addison Wesley Chemistry

  2. Contents • 6.1 Introduction to Chemical Bonding • 6.2 Representing Chemical Compounds • 6.3 Ionic Charges • 6.4 Ionic Compound • 6.5 Molecular Compounds and Acids • 6.6 Summary of Naming + Formula Writing

  3. 6.1 Introduction to Chemical Bonding • Molecules & Molecular Compounds Contents

  4. Monoatomic • Only noble gases exist as isolated atoms • Consist of single atoms • Ne, Ar, Ra

  5. Diatomic • Molecules that exist in pairs naturally • O2, H2, N2

  6. Triatomic • 3 atoms of the same element • O3 - ozone

  7. Molecule • Smallest electrically neutral unit of a substance that still has the properties of the substance and consisting of one or more atoms bonded together • O + O  O2 • H + H + O  H2O • Most atoms exist - in form of “molecule”

  8. Compound • A substance formed from the combination of two or more elements in a definite proportion by weight. • A compound can be one or more molecules.

  9. Molecular compounds • Low melting pt. • Low boiling pt. • Many gas or liquid @ room temp. • Molecules composed of atoms of 2+ nonmetals • Ex carbon monoxide – CO- diatomic • Ex water – H2O – triatomic

  10. Ions & Ionic Compounds • Ions- atoms or groups of atoms w/ ‘+’ or ‘-‘ charge • Form when loss or gain of electron from atom • Na = 11 p+ + 11 e- • If loose 1e- • Na + = 11p+ + 10e-

  11. Cation • Any atoms w/ ‘+’ charge • Fewer e- than atom that formed it

  12. Cation Naming • Metallic elements names • Name cation same as name of element • Sodium atom forms sodium cation • Do Not have same properties as atoms • Na atom explosive w/ H2O • Na very stable in H2O

  13. Anion • Atoms of nonmetallic elements game electrons • Have ‘-’ charge

  14. Anion Naming • Nonmetallic elements name • Not same as name • End in “-ide” • Sulfide-S2- • Bromide-Br-

  15. Ionic Compound • Compounds composed of cation and anions • Metal cation + nonmetal anion • Ex NaCl • Na+ + Cl- • Sodium Chloride – Table Salt

  16. Ionic Compound Properties • Electrically neutral • ‘+’ = ‘-‘ charge • usually solid crystals @ room temp. • melt @ high temp.

  17. Table 6.1 Characteristics of Molecular + Ionic Cmpds.

  18. 6.2 Representing Chemical Compounds Contents

  19. Chemical formulas • Shows kinds + numbers of atoms in the smallest representative unit of the substance

  20. If the molecules of an element have more than 1 atom a number is used as a subscript • Ex O2- Oxygen • H2, F2, O2, N2, Cl2, Br2, I2 • O3- Ozone

  21. Molecular Formulas • Shows kinds + numbers of atoms present in a molecule of a compound • Ex H2O- Water • Ex CO2- Carbon dioxide • Ex C2H6- Ethane • Does not show arrangement

  22. Structural Formula • ammonic gas (NH3) • H-N-H • H

  23. Ball-and-stick molecular model Methane – CH4 Penicillin

  24. Formula Units • Representation of ionic compounds • (Na+) and (Cl-) form the solid NaCl (sodium chloride) • Lowest whole number ratio of the ions • Charges not shown

  25. Magnesium chloride – Mg2+ and Cl – for MgCl2 *Remember there is no such thing as a molecule of NaCl or MgCl2 – they exist as a repeating positively and negatively charged ions in 3D patterns.

  26. Laws of Definite Proportions • In any sample of any chemical compound, the masses of the elements are always in the same proportion. • Means a sample of a chemical compound in our lab is the same as a sample in someone else’s lab by mass of elements in the compound. • Magnesium sulfide 100g =43.13g Mg + 56.87g S

  27. Law of Multiple Proportions • Whenever 2 elements form more than one compound, the different masses of 1 element that combine w/ the same mass of the other element are in the ratio of small whole numbers. • Example • A+B=C and 2A+B=D • Then A is a proportion of 1:2

  28. Multiple Proportions • Example • H2O 8gO • H2O2 16gO • O is a proportion 8:16 or 1:2 by mass • Helps if have mass of 1 proportion allow easy way to obtain mass of other

  29. 6.3 Ionic Charges Contents

  30. Monatomic Ions • ions consisting of only 1 atom

  31. Metallic Elements • tend to lose electrons (became positive) • Group 1A => 1+ charge cation • Group 2A => 2+ charge cation • Group 3A => 3+ charge cation • Aluminum only common group 3A metal

  32. Group A Nonmetals • tend to form anions • They tend to gain electrons (become negative) • Subtract the group # from 8 and make negative • Group 7A => 1- charge anion • Group 6A => 2- charge anion • Group 5A => 3- charge anion • 3 nonmetals in this group • N3-, P3-, As3-

  33. Group 4A • Group 4A and O generally don’t form ions • Transition metals + the Z metals in 4A • (Sn-tin + Pb-Lead) • Tend to form more than 1 cation

  34. Transition metals • Transition metals + the Z metals in 4A • (Sn-tin + Pb-Lead) • Tend to form more than 1 cation

  35. Table 6.3 p144 • Formulas + Names of Common Metal Ions w/ more than 1 Ionic Charge

  36. Polyatomic Ions • Tightly bound groups of atoms that behave as a unit and carry a charge. • Most end in ‘-ite’ or ‘-ate’

  37. 3 exceptions • ammonium – (NH4+) • cyanide – (CN-) • hydroxide- (OH-)

  38. ‘ite’ vs. ‘ate’ ending • ‘ite’ signifies 1 less oxygen than ‘ate’ ion • does not tell # of O atoms

  39. Hydrogen at Beginning • It is H+ cation combined w/ polyatomic anion • It is algebraic sum of ionic charges • H++ CO32- →HCO3- hydrogen carbonate • H++ PO43- → HPO42- hydrogen phosphate • H++ HPO42-→H2PO- dihydrogen phosphate

  40. 6.4 Ionic Compound Contents

  41. Writing Formulas for Binary Ionic Compounds • Binary compounds • Compounds composed of 2 elements • KCL – potassium chloride (K+) + (Cl-) • Anion charge = cation charge

  42. In writing formulas • The positive charge cation must balance the negative charge anion • The net ionic charge is zero • Cation (+) is written first • Anion (-) is written second

  43. Binary Compound Examples • Calcium bromide • Ca2+ +Br- => CaBr2 • 2-2=0 • took 2 Br to equal zero • 1:2 • rust-Iron (III) oxide • Fe3+ + O2- =>Fe2O3 • Make sure it is lowest whole # ratio • 2:3 ratio

  44. Naming Binary Ionic Compounds • CuO • Metallic cation- copper • Nonmetallic anion- oxide • 1:1 ratio • Oxide anion always 2- so copper cation must balance it w/2- • Copper(П) oxide= CuO

  45. Transition metals may have 2 common ions • Cu- 1++2+ - copper • Sn- 2++4+ - tin

  46. Prefixes & Suffixes p 152

  47. Binary Compounds • All end in ‘ide’ • Contain only 2 elements • One type-“molecular compound” • Prefixes represent # of atoms • Carbon monoxide • Carbon dioxide

  48. Polyatomic Ions • End in ‘ite’ or ‘ate’ • Most consist of 3 or more atoms • ‘ite’ or ‘ate’ relate to the # of O atoms • ‘ite’ is 1 less O than ‘ate’ • Nitrite ion-NO2- or nitrate ion NO3- • Neither ending represents a specific #

  49. Ternary Ionic Compounds • A compound containing atoms of 3 different elements • Usually contain a polyatomic ion

  50. Ternary Ionic Compound Formula • 1st write formula (symbol+ charge) for each ion • then balance charges