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Ch. 5: Molecules and Compounds. Dr. Namphol Sinkaset Chem 152: Introduction to General Chemistry. I. Chapter Outline. Introduction Chemical Formulas Views of Elements/Compounds Naming “Type I” Compounds Naming “Type II” Compounds Polyatomic Ions Naming Acids Naming “Type III” Compounds

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Ch. 5: Molecules and Compounds

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    1. Ch. 5: Molecules and Compounds Dr. Namphol Sinkaset Chem 152: Introduction to General Chemistry

    2. I. Chapter Outline • Introduction • Chemical Formulas • Views of Elements/Compounds • Naming “Type I” Compounds • Naming “Type II” Compounds • Polyatomic Ions • Naming Acids • Naming “Type III” Compounds • Molecular Masses

    3. I. Sugar • Sugar is composed of carbon, oxygen, and hydrogen atoms. • Properties of sugar completely different than elements from which it’s made.

    4. I. Sodium and Chlorine…

    5. I. Versus Sodium Chloride

    6. I. Elements in Compounds • When an element forms a compound, its properties change completely. • Generally, properties of the compound have no correlation to the original elements. • In this chapter, we see how elements become compounds and cover chemical nomenclature.

    7. I. Molecules of a Compound Are the Same • Law of Constant Composition: all samples of a given compound have the same proportions of their constituent elements. • Generally, this is expressed as a mass ratio.

    8. I. Water’s Mass Ratio • If 18.0 g of water is decomposed into it’s elements O and H, there would be 16.0 g of O and 2.0 g of H. • The O:H mass ratio is thus 8.0:1.0. • Any sample of water would have this exact same ratio.

    9. II. Representing Compounds • Chemical formulas are used to refer to compounds. • chemical formula: a way to show the elements present in a compound and the relative numbers of each elemental atom. • The most common is the molecular formula.

    10. II. Chemical Formulas • There are three types of formulas. • molecular: gives the actual number of atoms of each element in a molecule of a compound (e.g. H2O2) • empirical: gives the relative number of atoms of each element in a compound (e.g. HO) • structural: uses lines to represent covalent bonds and shows interconnectivity

    11. II. Writing Molecular Formulas • The more metallic element is generally listed first. • Metallic character increases to the left and down on the periodic table. • Subscripts indicate the number of that type of atom in the compound. • If groups of atoms behave as an independent entity, parentheses are used.

    12. II. Molecular Formulas

    13. II. Chemical Models • Formulas lead to models which give an idea of the 3-D shape of a molecule.

    14. II. From Macroscopic to Symbolic

    15. III. Pure Substances

    16. III. Atomic Elements • If element exists as individual atoms, it is named as “atomic.” • e.g. atomic mercury

    17. III. Molecular Elements • Some elements occur naturally as groups of two or more atoms. • These are named “molecular” or “diatomic” (for two).

    18. III. Molecular Compounds • Compounds formed from two or more nonmetals.

    19. III. Ionic Compounds • Comprised of cations and anions. • A formula unit is the smallest electrically-neutral collection of ions.

    20. IV. Chemical Nomenclature • Like any specialized field, chemistry has its own language. • The ability to name and recognize names of chemical entities is very important. • The naming system is LOGICAL!! • The periodic table is indispensable when you are first learning nomenclature.

    21. IV. Type I Compounds • Type I compounds are ionics that have a metal from Groups 1 or 2 and a nonmetal from Groups 14-17. • Examples: • NaCl = sodium chloride • MgBr2 = magnesium bromide • K2S = potassium sulfide

    22. IV. Type I Compounds • To get a formula from a name, remember that a compound must be neutral. • Ion charges can be found by locating the element on the periodic table. • “The charge on one becomes the subscript of the other.”

    23. IV. Type I Compounds

    24. IV. Sample Problem • Give the correct name or formula for the compounds below. • sodium nitride • CaCl2 • potassium sulfide • MgO

    25. V. Transition Metals • Transition metals are found in the “Valley,” Groups 3-12, of the periodic table. • Transition metal cations often can carry different charges, e.g. Fe2+ and Fe3+. • Thus, a name like “iron chloride” is ambiguous.

    26. V. Type II Compounds • Type II compounds are ionics that have a transition metal (Groups 3-12) and a nonmetal (Groups 14-17). • Examples: • FeCl2 = iron(II) chloride • FeCl3 = iron(III) chloride

    27. V. Sample Problem • e.g. Give the correct name or formula for the compounds below. • MnO2 • copper(II) chloride • AuCl3 • molybdenum(VI) fluoride • W2O3

    28. V. Some Transition Metal Cations

    29. VI. Additional Complications • To make naming ionic compounds harder, sometimes polyatomic ions are involved. • polyatomic ion: two or more atoms that are bonded covalently and have a net positive or negative charge

    30. VI. Common Polyatomic Ions

    31. VI. Oxyanion Families • oxyanion: anion containing oxygen • There are families of oxyanions, and they have a systematic naming system. • Have either two- or four-member families. • e.g. NO2- and NO3- • e.g. ClO-, ClO2-, ClO3-, and ClO4-

    32. VI. Two-Member Families • For a two-member family, oxoanion with fewer O atoms is given the “–ite” suffix while the one with more O atoms is given the “–ate” suffix. • e.g. NO2- = nitriteand NO3- = nitrate

    33. VI. Four-Member Families • For the four-member families, the prefixes “hypo-” and “per-” are used to indicate fewer or more oxygen atoms. • e.g. the chlorine oxyanions • ClO- = hypochlorite • ClO2- = chlorite • ClO3- = chlorate • ClO4- = perchlorate

    34. VI. Oxoanion Naming Summary

    35. VI. Sample Problem • e.g. Give names or formulas for the following compounds. • Na2CO3 • magnesium hydroxide • potassium nitrate • CoPO4 • nickel(II) sulfate • NaClO2

    36. VII. Acids • Acids are special ionic compounds that have H+ as the cation. • There are two categories of acids that have different naming rules. • Binary acids contain only hydrogen and a nonmetal. • Oxyacids contain hydrogen, a nonmetal, and oxygen.

    37. VII. Naming Binary Acids • Examples: • HCl = hydrochloric acid • HBr = hydrobromic acid • H2Se = hydroselenic acid

    38. Set 1 HNO3 = nitric acid H2SO4 = sulfuric acid HClO3 = chloric acid HClO4 = perchloric acid H2CO3 = carbonic acid H3PO4 = phosphoric acid Set 2 HNO2 = nitrous acid HClO2 = chlorous acid HClO = hypochlorous acid H2SO3 = sulfurous acid VII. Naming Oxyacids Examples of oxyacids:

    39. VII. Naming Oxyacids -ate oxyanions become –ic acids. -ite oxyanions become –ous acids.

    40. VIII. Type III Compounds • Type III compounds are covalent (nonmetal bonded to nonmetal). • Naming rules: • More metallic element is named 1st using the normal element name EXCEPT when halogens are bonded to oxygen. • Second element is named using its root and the “-ide” suffix. • #’s of atoms indicated with Greek prefixes EXCEPT when there is only one atom of the first element.

    41. VIII. Greek Prefixes

    42. VIII. Type III Compounds • Some examples: • ClO2 = chlorine dioxide • N2O5 = dinitrogen pentoxide • S2Cl2 = disulfur dichloride • SeF6 = selenium hexafluoride

    43. VIII. Naming Practice • e.g. Indicate the “Type” and give the correct formula or name of the compounds below. • CoCl3 • dichlorine heptaoxide • SrO • magnesium hydroxide • carbon tetrachloride • HF(aq) • sodium hydride • V2O5 • Ru(ClO4)3 • hydrosulfuric acid • H2SO4 • titanium(IV) oxide • N2F2

    44. IX. Masses of Compounds • Atomic masses are readily accessible via the periodic table, e.g. H = 1.008 amu. • Formula masses (a.k.a. molecular masses or molecular weights) are calculated by adding up the masses of each atom in the compound.

    45. IX. Molecular Mass of Water • The formula for water is H2O, so it is comprised of 2 H atoms and 1 O atom.

    46. IX. Formula Mass • e.g. What is the formula mass of barium nitrate?