Chemistry and Society

1. Chemistry and Society Bonding Dr. Victor H. Vilchiz Spring 2011

2. Bonding • There are three types of bonds • Ionic: involves the exchange of electrons and usually occurs between a metal and a non-metal • Covalent: involves the sharing of electrons and usually occurs between two non-metals or a non-metal and a metalloid • Metallic: involves the pooling of electron and involves two metals.

3. Chemical Formulas; Molecular and Ionic Substances • The chemical formula of a substance is a notation using atomic symbols with subscripts to convey the relative proportions of atoms of the different elements in a substance. • Consider the formula of aluminum oxide, Al2O3. This formula implies that the compound is composed of aluminum atoms and oxygen atoms in the ratio 2:3.

4. Chemical Formulas; Molecular and Ionic Substances • Molecular substances • A molecule is a definite group of atoms that are chemically bonded together – that is, tightly connected by attractive forces. • A molecular substance is a substance that is composed of molecules, all of which are alike. • Chemical formula provides information regarding to the composition of the compound.

5. Chemical Formulas • There are at least four types of chemical formulas • Empirical formula shows the smallest ratio of elements possible in a compound. • A molecular formula gives the exact number of atoms of elements in a molecule. • Structural formulas show how the atoms are bonded to one another in a molecule. • Condensed formula shows every atom present in a formula showing the bonding atoms but not the exact structural positioning. (molecules)

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

7. Chemical Formulas; Ionic Substances • Ionic substances • Although many substances are molecular, others are composed of ions. • An ion is an electrically charged particle obtained from an atom or chemically bonded group of atoms by adding or removing electrons. • Sodium chloride is a substance made up of ions. (NaCl)

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

9. Chemical Formulas; Ionic Substances • Ionic substances • When an atom picks up extra electrons, it becomes a negatively charged ion, called an anion. • An atom that loses electrons becomes a positively charged ion, called a cation. • An ionic compound is a compound composed of cations and anions. • An ionic compound is held together by the coulombic attractionsof the opposite sign charges.

10. Chemical Formulas; Ionic Substances • Ionic substances • The formula of an ionic compound is written by giving the smallest possible whole-number ratio of different ions in the substance that will yield a neutral species. • Ionic substances do not have molecular formulas since a cation is surrounded by many anions to be able to distinguish where a molecule ends and the next starts. • Ionic compounds instead have formula unitsof the substance is the group of atoms or ions explicitly symbolized by its formula.

11. 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.

12. 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.

13. Chemical Substances; Formulas and Names • Ionic compounds • Most ionic compounds contain metal and nonmetal atoms; for example, NaCl. • You name an ionic compound by giving the name of the cation followed by the name of the anion. • A monatomic ion is an ion formed from a single atom. • Table lists some common monatomic ions of the main group elements.

14. Chemical Substances; Formulas and Names • Rules for predicting charges on main group monatomic ions • Most of the main group metals form cations with the charge equal to their group number. • Example: Aluminum is in group IIIA; therefore an aluminum ion will be Al3+ • The charge on a monatomic anion for a nonmetal equals the group number minus 8. • Example: Nitrogen is in group VA; therefore a nitrogen ion will be N(5-8),N3-

15. Common Ions of the Main Group Elements Return to Lecture

16. Transition Metal Ions • A lot of the transition metal ions are confused as to how many electrons they wish to donate; therefore, they form more than one ion, each with a different charge. (see Table). • You should know the ones for Cu, Fe, Pb, and Hg.

17. Common Ions of the transition metals Return to Lecture

18. 0 Naming Cations • Rules for naming cations • There are two types of cations: elemental (Monatomic) and Ammonium Ion • Monatomic cations are named after the element and the word Ion is appended. For example, Al3+ is called the aluminum ion. • If there is more than one cation of an element, a Roman numeral in parentheses denoting the charge on the ion is used. This often occurs with transition elements. Example: Fe2+ and Fe3+; Iron(II) ion and Iron(III) ion

19. Naming Anions • Rules for naming anions • Monatomic Anions: Use the root of the element’s name and change ending to – ide and append the word ion. • example, Br Br – Bromine becomes Bromide ion. • Polyatomic Anions: Use the table provided.

20. 0 List of Polyatomic Ions Return to Lecture

21. Naming Ionic Compounds • NaF - • LiCl - • MgO -

22. 0 Naming Ionic Compounds • NaF- Sodium Ion and Fluoride ion • LiCl- Lithium Ion and Chloride Ion • MgO- Magnesium Ion and Oxide Ion

23. 0 Naming Ionic Compounds • NaF - Sodium Fluoride • LiCl - Lithium Chloride • MgO - Magnesium Oxide

24. Cd(OH)2 Cadmium Hydroxide Ca3(PO4)2 Calcium Phosphate AgCN Silver Cyanide Na2SO4 Sodium Sulfate Na2SO3 Sodium Sulfite KNO3 Potassium Nitrate 0 More Practice

25. 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.

26. 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

27. 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.

28. 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”

29. Greek Prefixes for Covalent Compounds Nomenclature Return to Lecture

30. 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

31. Bonding • I mentioned earlier that the periodic table was arranged according to how many electrons were in the outer shell. • It is the same electrons in the outer shell that dictate how a compound interacts with the rest of the substances. • It is the electrons that are involved in bonding and determine how a compound is held together.

32. Electrons in Bonding • Electrons in an atom feel an attraction towards the nucleus, where the positive charged molecules are found. • However, there are other electrons in the outer shell as well as inner shells that will repel each other due to the similar charge. • The same is true of protons feeling the proximity of other protons in the nucleus.

33. Bonding • At the end what we have is a tug-o-war between the electrostatic attraction between the electrons and protons on one side and the proton-proton and electron-electron repulsion on the other. • Once those three forces are balanced we end up with a bond between atoms.

34. How does bonding take place? • Sometimes in order to learn how something happens it is easier to look at the case of when it does not happen. • Lets look at the case of the noble gases which rarely form compounds covalent or otherwise. • What is it that noble gases have that the rest of the elements do not. • First notice their position in the periodic table. • They are at the very far right of the table.

35. Noble gases and bonding • The far right means that the “outermost” subshell is full and no more electrons can be added. • To a certain extent gases are happy with the number of electrons they have. They do not want more they do not want less. • The rest of the elements strive to reach this ideal “electron configuration.”

36. Distributing electrons • Lets us see what takes place when atoms meet. • Imagine there are 15 of us and that we get together to play basketball… do we make 1 team of 5 one of 4 and one of 6? • That does not seem to make sense… the logical thing to do is to make 3 teams of 5 • The same thing happens in bonding… when an atom with too many electrons meets up with one that is short an atom they send an electron to the other team and even up the game. IONIC BONDING

37. Covalent Bonding • There are times when there is not enough electrons for everyone to have their fill… • Imagine playing baseball and having two teams of 8 players and one extra person… • Have the extra person catch for both teams. • The last person is shared… • COVALENT BONDING • There are times that sharing electrons is a better pathway than exchanging electrons.

38. Lewis Symbols • Now that we have discussed how bonding takes place lets start thinking about writing some chemical formulas. • First let us start with Lewis Symbols. • Lewis symbols are ONLY used with main group elements. • A Lewis symbol tells you how many electrons are available for bonding.

39. Lewis Symbols • There are 4 steps to follow when writing Lewis Symbols • 1. Write Down the Element Symbol • 2. Determine the # of e-’s in the valance shell (group #). • 3. Place 1 dot per electron around the symbol for each electron in the valance shell; go around the block before starting to pair electrons. • 4. Make sure the number of electrons on the symbol equals the group number. (never have more than 8 electrons around a symbol).

40. Lewis Symbols • Carbon Correct C Incorrect C C • Nitrogen • Correct N N N N . . .. . . . .. .. . .. . . . . . . . . . .. .. . . . ..

41. Lewis Formulas • When we take Lewis Symbols and bring them together to represent a reaction we make Lewis Formulas • H + H  H H = H-H = H2 . . ..

42. Electronegativity • Electronegativity: The ability of an atom to pull SHARED electrons towards itself. • Electronegativity increases from left to right and decreases from top to bottom. • Noble gases do not form bonds therefore are not included in EN tables. • The most EN element is F the least EN element is Fr

43. Polarity • A bond between two ions gives an ionic compound • All ionic bonds are polar. • Polarity: A charge separation between one side of the molecule (bond) and the other side. • In an ionic compound one ion (side) is positive the other ion is negative, clearly we have charge separation.

44. Polarity • All ionic bonds are polar… but are all covalent bonds non-polar? • NO… • There are very few bonds that yield NO-POLAR bonds. • In order for us to obtain a fully non-polar bond we must form a bond between atoms of the same element. • Bond polarity in covalent molecules depends on the DEN between the two participating atoms.

45. Electronegativities

46. Making and Acid Return to Lecture

47. Molecular model of nitric acid. Return to Lecture

48. Naming Flow Chart II Return to Lecture