Molecules, Ions and Their Compounds

1. Molecules, Ions and Their Compounds Chemistry 101 Chapter 3 Virginia State University Dr. Victor Vilchiz Summer 2008

2. Chemical Substances; Formulas and Names • Naming simple compounds • Chemical compounds are classified as organic or inorganic. • Organic compounds are compounds that contain carbon combined with other elements, such as hydrogen, oxygen, and nitrogen; they do not contain metals. • Inorganic compounds are compounds composed of elements other than carbon and usually contain at least one metal atom.

3. Chemical Formulas; Molecular Substances • Organic compounds • An important class of molecular substances that contain carbon is the organic compounds. • Organic compounds make up the majority of all known compounds. • The simplest organic compounds are hydrocarbons, or compounds containing only hydrogen and carbon. • Common examples include methane, CH4, ethane, C2H6, and propane, C3H8.

4. 0 Naming Covalent Compounds • A covalent compound as we said before is formed by sharing electrons between 2 nonmetals or metalloids. • These compounds are usually molecular and are named using a prefix system. • When naming these compounds name the element further to the left (in the periodic table) first, then the one on the right.

5. Naming Covalent Compounds • You name the first element using the exact element name. • Name the second element by writing the root of the element’s name and add the suffix “–ide.” • If there is more than one atom of any given element, you add the Greek prefix denoting how many atoms of that element are present. Table lists the Greek prefixes used. • If only one atom of the second element is present it gets the prefix “mono”

6. Naming Covalent Compounds • Here are some examples of prefix names for binary molecular compounds. • PF5 phosphorus pentafluoride • SO2 sulfur dioxide • SF6 sulfur hexafluoride • N2O4 dinitrogen tetroxide • CO carbon monoxide

7. Naming Acids • Acidsare traditionally defined as compounds that could donate an H+; however, they are acids only in the presence of water. In other words before they enter the liquid they are covalent compounds and they are NOT acids. • There are two main types of acids: • Binary acids consist of a hydrogen ion and any single anion. For example, HCl is hydrochloric acid. • An oxoacidis an acid containing hydrogen, oxygen, and another element. An example is a HNO3, nitric acid. (see Figure 2.23)

8. Naming Acids • Binary Acids • Start with the prefix “Hydro” which represents the Hydrogen, followed it with the root of the name of the second element and append the ending –oic acid. • Oxoacids • Use the root of the “E” element if the ion taking part in the acid had an ending in –ate to the root append the ending –ic acid, if it ends on –ite then append the ending –ous acid. If the ion had a prefix use the same prefix.

9. Naming Acids • Examples: • HCl(g) Hydrogen Chloride • HCl(aq) HydroChloric Acid • H2S(g) Dihydrogen Sulfide • H2S(aq)HydroSulfic Acid • H3PO4(aq)Phosphoric Acid • HClO4(aq)Perchloric Acid • HClO(aq)Hypochlorous Acid

10. 0 Ionic Compounds Formulas • How do we know how many atoms of each ion we need? • A simple crossing of the charges can answer that question about 90% of the time. • Example: Mg2+ and PO43- Mg3(PO4)2 Check the charges… 3 x (+2) = +6 2 x (-3) = -6 • When they combined they cancel to yield a neutral compound.

11. 0 Ionic Compounds Formulas • The crossing technique does not work if the magnitude of the charges is the same • Example: Mg2+ and CO32- Mg2(CO3)2 This is incorrect since we want the lowest ratio possible which is 1:1 to yield MgCO3

12. Ionic Compounds Properties • Ionic compounds have properties completely different from their component elements. • Example: Table Salt (NaCl) • Sodium (Na) in the presence of water reacts violently heating up the water and producing hydrogen if the temperature of the water is high enough the hydrogen can ignite explosively. • Chloride (Cl) Green poisonous and corrosive gas. If inhaled will destroy the nasal passages then dissolve in the stomach producing high concentration of hydrochloric acid which will destroy the stomach lining producing ulcers. • Salt (NaCl) posses none of the properties mentioned above.

13. Ionic Structure • Ions form a 3-D lattice. • The coulombic (electrostatic) attraction is so high that in order to separate one ion from the lattice requires a lot of energy (DHlatt). • The lattice energy depends on charge and size of the ions.

14. Lattice Energy • Since the lattice energy is an electrostatic interaction the more separated the charges are the weaker the interaction is. • Bigger ions have lower lattice energies • The higher the charge of the ions the stronger they will attract ions of the opposite charge. • When size and charge point to opposite trends the charge will outweigh the size. • From smallest atom to biggest atom there is only 1.7x factor. From a +1 to +2 that is already a 2x factor.

15. Properties of Ionic Substances • Dues to the charged interaction a blow to a crystal leads to the possibility of splitting the crystal since we will force like charged particles to interact. • Ionic compounds have high melting/boiling points since in order to move the ions from their respective spots it will require breaking the lattice.

16. Ionic Solutions • However, if we do melt an ionic compound it will be able to conduct current. • When ionic compounds are placed in a solvent the produced solution conducts electricity. The higher the number of ions the higher the conductivity. • More when we cover chapter 4.

17. 0 Naming Hydrates • A hydrateis a compound that contains water molecules weakly bound in its crystals. • Hydrates are named from the anhydrous (dry) compound, followed by the word “hydrate” with a Greek prefix to indicate the number of water molecules per formula unit of the compound. • For example, CuSO4. 5H2O is known as copper(II)sulfate pentahydrate. (see Figure 2.24)

18. Determining Chemical Formulas • Determining both empirical and molecular formulas of a compound from the percent composition. • The percent composition of a compound leads directly to its empirical formula. • An empirical formula(or simplest formula) for a compound is the formula of the substance written with the smallest integer (whole number) subscripts.

19. Determining Chemical Formulas • The percent composition of a compound is the mass percentage of each element in the compound. • We define the mass percentage of “A” as the parts of “A” per hundred parts of the total, by mass. That is,

20. Mass Percentages from Formulas • Let’s calculate the percent composition of butane, C4H10. • First, we need the molecular mass of C4H10. • Now, we can calculate the percents.

21. Determining Chemical Formulas • Determining the empirical formula from the percent composition. • Benzoic acid is a white, crystalline powder used as a food preservative. The compound contains 68.8% C, 5.0% H, and 26.2% O by mass. What is its empirical formula? • In other words, give the smallest whole-number ratio of the subscripts in the formula Cx HyOz

22. Determining Chemical Formulas • Determining the empirical formula from the percent composition. • For the purposes of this calculation and making calculations simpler, we will assume we have 100.0 grams of sample benzoic acid. • Then the percentage of each element equals the mass of each element in the sample. • Since x, y, and z in our formula represent mole-mole ratios, we must first convert these masses to moles.

23. Determining Chemical Formulas • Determining the empirical formula from the percent composition. • Our 100.0 grams of benzoic acid would contain: This isn’t quite a whole number ratio, but if we divide each number by the smallest of the three, a better ratio might emerge.

24. Determining Chemical Formulas • Determining the empirical formula from the percent composition. • Our 100.0 grams of benzoic acid would contain: now it’s not too difficult to see that the smallest whole number ratio is 7:6:2. The empirical formula is C7H6O2.

25. Determining Chemical Formulas • Determining the “true” molecular formula from the empirical formula. • An empirical formula gives only the smallest whole-number ratio of atoms in a formula. • The “true” molecular formula could be a multiple of the empirical formula (since both would have the same percent composition). • To determine the “true” molecular formula, we must know the “true” molecular weight of the compound.

26. Determining Chemical Formulas • Determining the “true” molecular formula from the empirical formula. • For example, suppose the empirical formula of a compound is CH2O and its “true” molecular weight is 60.0 g/mol. • The molar weight of the empirical formula (the “empirical weight”) is only 30.0 g/mol. • This would imply that the “true” molecular formula is actually the empirical formula doubled 2(CH2O) or C2H4O2

27. Molecular and structural formulasand molecular models. Return to Lecture

28. A model of a portion of a Sodium Chloride crystal. Return to Lecture

29. Common Ions of the transition metals Return to Lecture

30. 0 List of Polyatomic Ions Return to Lecture

31. Greek Prefixes for Covalent Compounds Nomenclature Return to Lecture

32. Making and Acid Return to Lecture

33. Molecular model of nitric acid. Return to Lecture

34. 0 Figure 2.24: Copper (II) sulfate. Photo courtesy of James Scherer. Return to Slide 44

35. 0 Naming Flow Chart Return to Lecture

36. Naming Flow Chart II Return to Lecture

37. Naming Acids Flow Chart Return to Lecture

38. 0 Quantities of Reactants and Products Chapter 4 Dr. Victor Vilchiz

39. What is a mole? • A mole is a unit of measurement used to specified amounts of chemical substances. • It is not a unit of mass. • It is similar to “a dozen” • A dozen eggs is not the same as a dozen cars but they are still both a dozen.

40. What is a mole? • A mole is defined as the number of atoms of carbon in 12 g of Carbon-12. (examples) • 1mole=6.022x1023 atoms and can be applied to any moiety • 6.022x1023 is also known as Avogadro’s Number (NA) • 1mol of Carbon=12g Carbon = 6.022x1023 C atoms • 1mol of water= 18g H2O =6.022x1023 water molecules

41. Why the mole? • The mole helps determine amounts of substances and allows for conversion between species. • CaCO3(s) + 2HCl(aq)  CaCl2(aq) + H2CO3(aq) • From this we cannot say 1g of CaCO3 will react with 2 grams of HCl; however, we can say 1mol of CaCO3 reacts with 2 moles of HCl. • 1 mole of CaCO3 is not the same as 1 mole of HCl mass wise, but both have 6.022x1023 molecules.

42. Molar Mass Examples • Molar Mass of Ca(C2H3O2)2, Calcium Acetate. 2x(40.1)+4(12.0)+6(1.01)+4(16.0)=198.3g/mol • Molar Mass of Ethylene Glycol, C2H4O2. 2x(12.0)+4(1.01)+2(16.0)=60.0g/mol • Molar Mass of Ammonium Oxalate, (NH4)2C2O4. 2x(14.0)+8x(1.01)+2(12.0)+4(16.0)=124.1g/mol

43. Stoichiometry: Quantitative Relations in Chemical Reactions • Stoichiometry is the calculation of the quantities of reactants and products involved in a chemical reaction. • It is based on the balanced chemical equation and on the relationship between mass and moles. • Such calculations are fundamental to most quantitative work in chemistry.

44. 0 Chemical Reactions: Equations • A chemical equation is the symbolic representation of a chemical reaction in terms of chemical formulas. • For example, the burning of sodium and chlorine to produce sodium chloride is written • The reactants are starting substances in a chemical reaction. The arrow means “yields.” The formulas on the right side of the arrow represent the products. • Writing chemical equations

45. 0 Chemical Reactions: Equations • In many cases, it is useful to indicate the states of the substances in the equation. • When you use these labels, the previous equation becomes • s=solid, l=liquid, g=gas, aq=aqueous • Writing chemical equations

46. Molar Interpretation of a Chemical Equation • A balanced chemical equation: 2H2 +1O2 2H2O can be interpreted to read 2 moles of Hydrogen react with one mole of oxygen to produce 2 moles of water. • In the balanced equation the 2, 1, and 2 are known as the stoichiometric coefficients. • At the molecular level they refer to the number of molecules reacting.

47. Molar Interpretation of a Chemical Equation • Because moles can be converted to mass, you can also give a mass interpretation of a chemical equation. 2H2 +1O2 2H2O 2(2.02g)H2 react with 1(32.0g) O2 to yield 2(18.0g)H2O 4.04g H2 react with 32.0g O2 to yield 36.0g H2O