1. 1 Topic 2 Atoms, Elements, Molecules, Ions, and Compounds Early in the 19 th century John Dalton developed atomic theory. His theory explained the best available experimental data at that time. His theory has been modified since then with the discovery of other data, but his work was the initial ground work that we will examine first.

2. Atomic Theory of Matter Postulates of Dalton’s Atomic Theory 1.)All matter is composed of indivisible atoms. An atom is an extremely small particle of matter that retains its identity during chemical reactions. 2.)An element is a type of matter composed of only one kind of atom, each atom of a given element having the same properties. Mass is one such property. Thus the atoms of a given element have a characteristic mass. 2 Protons, neutrons, electrons Atoms have different isotopes that have the same # protons but different # neutrons and hence different mass. Note: #protons gives identity of atom. Later found indivisible to be untrue. Later found all atoms of the same element does not have to have the same mass.

3. Atomic Theory of Matter Postulates of Dalton’s Atomic Theory 3 3.) A compound is a type of matter composed of atoms of two or more elements chemically combined in fixed proportions. –The relative numbers of any two kinds of atoms in a compound occur in simple ratios. –Water, for example, consists of hydrogen and oxygen in a 2 to 1 ratio (2H: 1O) for all molecules of water.

4. Atomic Theory of Matter Postulates of Dalton’s Atomic Theory 4 4.) A chemical reaction consists of the rearrangements of the atoms present in the reacting substances to give new chemical combinations present in the substances formed by the reaction (new chemical with different properties). 5.) Atoms are not created, destroyed, or broken into smaller particles by any chemical reaction. Na (s) + Cl 2 (g)  2 2 NaCl (s) Protons, neutrons, electrons i.e., solid sodium mixed with chlorine gas forms a new substance, salt, with totally different properties from the starting materials. Once again, later found indivisible to be untrue.

5. Atomic Theory of Matter The Structure of the Atom 5 –Although Dalton postulated that atoms were indivisible, experiments at the beginning of the 1900’s showed that atoms themselves consist of particles. –Experiments by Ernest Rutherford in 1910 showed that the atom was mostly “empty space.”

6. Atomic Theory of Matter 6 –These experiments showed that the atom consists of two kinds of particles: a nucleus, the atom’s central core, which is positively charged and contains most of the atom’s mass, and one or more electrons. –Electrons are very light, negatively charged particles that exist in the region around the atom’s positively charged nucleus. Nucleus ( + ) e-e-

7. Atomic Theory of Matter 7 –In 1897, the British physicist J. J. Thompson conducted a series of experiments that showed that atoms were not indivisible particles but instead made of smaller particles. –From his experiments, Thompson calculated the ratio of the electron’s mass, m e, to its electric charge, e.

8. Atomic Theory of Matter 8 –In 1909, U.S. physicist, Robert Millikan obtained the charge on the electron (1.602 x 10 -19 C). –These two discoveries (Millikan and Thompson) combined provided us with the electron’s mass of 9.109 x 10 -31 kg, which is more than 1800 times smaller than the mass of the lightest atom (hydrogen) thereby proving that the atom is made up of smaller particles. –These experiments showed that the electron was indeed a subatomic particle.

9. Atomic Theory of Matter The nuclear model of the atom. 9 –Ernest Rutherford, a British physicist, put forth the idea of the nuclear model of the atom in 1911, based on experiments done in his laboratory by Hans Geiger and Ernest Morrison. –Rutherford’s famous gold leaf experiment gave credibility to the theory that the majority of the mass of the atom was concentrated in a very small nucleus. –Positively charged alpha particles were directed at a metal foil. Only 1/8000 were deflected indicating that the nucleus was extremely small and positively charged. Only those alpha particles that directly hit the nucleus were deflected; the rest passed through.

10. Atomic Theory of Matter 10 –Most of the mass of an atom is in the nucleus; however, the nucleus occupies only a very small portion of the space in the atom. –The diameter of an atom is approximately 100 pm while the diameter of the nucleus is approximately 0.001 pm. For comparison, if an atom was 3 miles in diameter, the nucleus would be the size of a golf ball. –The nucleus of an atom is composed of two different kinds of particles: protons (+) and neutrons (neutral).. –An important property of the nucleus is its positive electric charge.

11. Atomic Theory of Matter 11 –A proton is the nuclear particle having a positive charge equal to that of the electron’s (a “unit” charge) and a mass more than 1800 times that of the electron. It is for this reason that we refer to H as a pure proton. –The number of protons in the nucleus of an atom is referred to as its atomic number (Z) and gives the identity of an element. All species that have same #p have the same properties. neutral species: #p = #e - #p remains constant in species; #e - can vary and dictates the charge of species H Z=1 1p, 1e - Na Z=11 11p, 11e - Cl Z=17 17p, 17e - Cl - Z=17 17p, 18e - Na + Z=11 11p, 10e - + charge, more p than e - - charge, more e - than p

12. Atomic Theory of Matter 12 –An element is a substance whose atoms all have the same atomic number (Z). The #protons defines the identity of an atom and can be found on the periodic table (large number in top of element box). –The neutron is a nuclear particle having a mass almost identical to that of a proton, but no electric charge. The charge of the nucleus comes from the #protons. The atoms may have different masses because of different #neutrons (isotopes). –Summary of masses and charges of the three fundamental particles: particlemass, kgcharge, Crelative charge location electron, e - 9.109 x 10 -31 -1.602 x 10 -19 outside nucleus proton, p1.6726 x 10 -27 1.602 x 10 -19 +1nucleus neutron, n1.6749 x 10 -27 0 0nucleus

13. 13 The mass number (A) is the total number of protons and neutrons in a nucleus. A nuclide is an atom characterized by a definite atomic number and mass number. The shorthand notation for a nuclide consists of its symbol with the atomic number, Z, as a subscript on the left and its mass number, A, as a superscript on the left. A = #p + #n = Z + #n How many neutrons does sodium 23 have? A = 23, Z = 11 (number on periodic table) A = Z + #n23 = 11 + #n #n = 23 - 11 = 12 11p, 12 n, 11e - + 11p, 12 n, 10e -

14. What is the nuclide symbol for a nucleus that consists of 17 protons, 18 neutrons, and 17 electrons? How many protons, neutrons, and electrons are in the following nucleus 14 HW 9 HW 10 What’s the element? 17 p  atomic number on periodic table for chorine A = #p + #n = 17 + 18 = 35 35 p45 n A = #p + #n 80 = 35 + #n#n = 80 – 35 = 45 n 36 e - Note: #e - = #p; therefore, neutral species Note: #e - > #p; therefore, negatively charged species 35e - + one additional e - based on -1 charge = 36 e - code for both: proton

15. Atomic Theory of Matter 15 –The fractional abundance is the fraction of a sample of atoms that is composed of a particular isotope. –Isotopes are atoms whose nuclei have the same atomic number (Z) but different mass numbers (A); that is, the nuclei have the same number of protons but different numbers of neutrons thereby causing them to have different masses. –Chlorine, for example, exists as two isotopes: chlorine-35 and chlorine-37. 17p, 18 n, 17e - 17p, 20 n, 17e - (0.75771)(34.97 amu) + (0.24229) (36.97 amu) = 35.45 amu Frac abund = 75.771% Mass = 34.97 amu Frac abund = 24.229% Mass = 36.97 amu Note: The mixture of isotope masses make up the actual mass of the element given on periodic table. Cl has a mass of 35.45 amu which is based on the two isotopes of Cl-35 and Cl-37.

16. Atomic Weights Calculate the atomic weight of boron, B, from the following data: ISOTOPE ISOTOPIC MASS (amu) FRACTIONAL ABUNDANCE B-10 10.013 0.1978 (19.78%) B-11 11.009 0.8022 (80.22%) 16 B-10: 10.013 amu x 0.1978 = 1.980 B-11: 11.009 amu x 0.8022 = 8.831 10.811 = 10.811 amu ( = atomic wt.) HW 11 Note: fractional abundances must add to 1 (100%) Note: mass on periodic table matches 10.811 amu (weighted average of isotopes) code: amu

17. Atomic Weights Dalton’s Relative Atomic Masses 17 –Since Dalton could not weigh individual atoms, he devised experiments to measure their masses relative to the hydrogen atom. –Hydrogen was chosen as it was believed to be the lightest element. Daltons assigned hydrogen a mass of 1 (1 Dalton = mass of H). –For example, he found that carbon weighed 12 times more than hydrogen. He therefore assigned carbon a mass of 12 ( mass of carbon = 12 Daltons).

18. Atomic Weights Dalton’s Relative Atomic Masses 18 –Dalton’s atomic weight scale was eventually replaced in 1961, by the present carbon–12 mass scale. –One atomic mass unit (amu) is, therefore, a mass unit equal to exactly 1/12 the mass of a carbon–12 atom. –On this modern scale, the atomic weight of an element is the average atomic mass for the naturally occurring element, expressed in atomic mass units. Periodic table is based on atom mass with units of amu. Na - 23.1 amu  mass of 1 atom of sodium

19. The Periodic Table In 1869, Dmitri Mendeleev discovered that if the known elements were arranged in order of atomic mass (A), they could be placed in horizontal rows such that the elements in the vertical columns had similar properties. 19 –periodic table - tabular arrangement of elements in rows and columns, highlighting the regular repetition of properties of the elements. –periodic law – states that certain sets of physical and chemical properties recur at regular intervals (periodically) when the elements are arranged according to increasing atomic number (Z). –Note: eventually changed from atomic mass to atomic number because of a couple of anomalies.

20. Figure: A modern form of the periodic table. 20 anomalies

21. The Periodic Table Periods and Groups 21 –A period consists of the elements in any one horizontal row of the periodic table. –A group consists of the elements in any one column of the periodic table (similar properties/structure). –The groups are usually numbered (North American uses roman numbers and A/B; IUPAC 1-18). –The eight “A” groups are called main group (or representative) elements.

22. The Periodic Table Periods and Groups 22 –The “B” groups are called transition elements. –The two rows of elements at the bottom of the table are called inner transition elements. –Elements in any one group have similar properties because their outer shells have the same number of valence electron (discuss in later sections).

23. The Periodic Table Periods and Groups 23 –The elements in group IA (except H) - alkali metals –The group VIIA elements - halogens –The elements in group IIA - alkaline earth metals, –The group VIIIA elements – noble gases (monoatomic) –Diatomic elements – H 2, N 2, O 2, F 2, Cl 2, Br 2, I 2 –Most species are solids at room temperature; H 2, N 2, O 2, F 2, Cl 2, and noble gases are gases; Br 2 and Hg are liquids.

24. 24 HW 12 metals nonmetals Metallic character Metals, Nonmetals, and Metalloids – generally, left of staircase are metals, touching staircase are metalloids, right of staircase are nonmetals. This is important for determining bond type, using proper terminology, and making decisions. code: table

25. 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. 25 –aluminum oxide, Al 2 O 3 2Al:3O ratio –sodium chloride, NaCl 1Na:1Cl ratio –calcium nitrate, Ca(NO 3 ) 2 1Ca:2NO 3 - ratio or 1Ca:2N:6O ratio

26. 26 –A molecule is a definite group of atoms that are chemically bonded together through sharing of electrons (covalent bonding, generally nonmetal- nonmetal including H). Chemical Formulas; Molecular and Ionic Substances Molecular substances –A molecular substance is a substance that is composed of molecules, all of which are alike. –A molecular formula gives the exact number of atoms of elements in a molecule (i.e. C 2 H 6 O ). –Structural formulas show how the atoms are bonded to one another in a molecule. i.e. ethanol ( C 2 H 6 O ) has a structural formula of CH 3 CH 2 OH involves covalent bond – share electrons between atoms – typically nonmetal/nonmetal involves ionic bond – transfer electrons between atoms – attraction between charged particles – typically metal/nonmetal or polyatomic ions C : C Na + Cl -

27. 27 –Although many substances are molecular, others are composed of ions (charged particles) that have transferred electrons and have ionic bonding; occurs generally with metal-nonmetal interactions. Ionic substances –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. Na Cl +  1e -  - Na + Cl -

28. 28 –The formula of an ionic compound is written by giving the smallest possible whole-number ratio of different ions in the substance. Chemical Formulas; Molecular and Ionic Substances Ionic substances –The formula unit of the substance is the group of atoms or ions explicitly symbolized by its formula. Covalent bond (share e - )Ionic bond (transfer e - / attraction charged particles nm –nmm – nm and charged ions MoleculesFormula unit Molecular substanceIonic substance Molecular formulaformula C : O Na + Cl -

29. 29 –When an atom gains extra electrons, it becomes a negatively charged ion, called an anion (more electrons than protons). i.e, Cl - Ionic substances –An atom that loses electrons becomes a positively charged ion, called a cation (more protons than electrons). i.e., Na +. –An ionic compound is a compound composed of cations and anions.NaCl CaBr 2 Na 2 SO 4 CO 2 Answer the following questions for species below: ionic or molecular substance; formula unit or molecule; ionic or covalent bonds involved? ionic substance; formula unit; ionic bond ionic substance; formula unit; ionic bonds ionic substance; formula unit; ionic and covalent bonds in SO 4 2- molecular substance; molecule; covalent bonds

30. 30 Ions in Aqueous Solution Many (not all) ionic compounds (ionic bond/m-nm) dissociate into independent ions when dissolved in water NaCl (s)  Na + (aq) + Cl - (aq) Soluble ionic compounds dissociate 100% - referred to as strong electrolytes – breaks into charged particles until reaches saturation point. Soluble salt charges particles

31. 31 Ions in Aqueous Solution Most molecular (covalent bond/nm-nm) compounds dissolve but do not dissociate into ions, exception acids. C 6 H 12 O 6 (s)  C 6 H 12 O 6 (aq) These compounds are referred to as nonelectrolytes; no charged particles; soluble to saturation point but no ions formed. How would sodium sulfate dissolve based on bonding? Na 2 SO 4 (s)  2Na + (aq) + SO 4 2- (aq) no charges particles; remains whole ionic bond dissociates while covalent bonds in sulfate remain intact

32. 32 –Most ionic compounds contain metal and nonmetal atoms (as well as polyatomic ions); for example, NaCl. Chemical Substances; Formulas and Names Ionic compounds –You name an ionic compound by giving the name of the cation followed by the name of the anion with -ide. Sodium chloride, NaCl Calcium Iodide, CaI 2 Potassium Bromide, KBr –We give the monatomic ion name for the cations and anions when naming compounds. A monatomic ion is an ion formed from a single atom.

33. 33 –Most of the main group metals form cations with the charge equal to their roman group number. How do we get the charge for ions? Rules for predicting charges on monatomic ions –The charge on a monatomic anion for a nonmetal equals the roman group number minus 8. –Most transition elements form more than one ion, each with a different charge (exceptions Cd 2+, Zn 2+, Ag + ). –Other important elements with variable charge Pb 4+, Pb 2+ Sn 4+, Sn 2+ As 5+, As 3+ Sb 5+, Sb 3+ 1+ 2+3+4+ 0 1-2-3- 4- varies

34. 34 –Monatomic cations are named after the element. For example, Al 3+ is called the aluminum ion. Rules for naming monatomic ions –If there is more than one cation of an element (charge), a Roman numeral in parentheses denoting the charge on the ion is used. This often occurs with transition elements. Na + sodium ionCa 2+ calcium ion Fe 2+ iron (II) ionFe 3+ iron (III) ion Older name: higher ox state (charge) – ic, / lower, -ous Fe 3+ ferric ion Fe 2+ ferrous ion Cu 2+ cupric ion Cu + cuprous ion Hg 2+ mercuric ion Hg 2 2+ mercurous ion also done with Pb 4+, Pb 2+ ; Sn 4+, Sn 2+ ; As 5+, As 3+ ; Sb 5+, Sb 3+. For the names of the monatomic anions, use the stem name of the element followed by the suffix – ide. For example bromine, the anion is called bromide ion, Br -.

35. 35 The formula of an ionic compound is written by giving the smallest possible whole-number ratio of different ions in the substance. Sodium chloride Na + Cl - Iron (III) sulfate Fe 3+ SO 4 2- Chromium (III) oxide Cr 3+ O 2- Calcium nitrate Ca 2+ NO 3 - Sodium phosphate Na + PO 4 3- Strontium oxide Sr 2+ O 2- NaCl SrO Na 3 PO 4 Ca(NO 3 ) 2 Cr 2 O 3 Fe 2 (SO 4 ) 3 Based on the charge of the ions and balancing the overall charge on the compound by adjusting the number of ions, a formula is written. Note the sum of all the charges must equal zero, and you do not display the charges in the final formula. ions and charges formula 2Fe 3+ = 6+ charge 3SO 4 2- = 6- charge balanced 1Na + = 1+ charge 1Cl - = 1- charge balanced Generally, you can crisscross the charge of one ion as the subscript on the second ion, reducing when possible. Roman number tells charge of transition metal 2 3 HW 13 & 14code for both: formula

36. 36 Naming Ionic Binary Compounds NaF- -lithium chloride MgO- MnBr 2 - -cobalt (III) oxide - copper (II) chloride or cupric chloride sodium fluoride LiCl magnesium oxide manganese (II) bromide Co 2 O 3 CuCl 2 To name a compound, you must know if it is a molecular or ionic compound so that you know which rules to follow. If you have a metal-nonmetal (or polyatomic ion), it is an ionic compound where you name the metal first then the nonmetal with changing the ending to –ide. If it is a transition metal, you must include the charge of the metal (Roman numbers). If you have nonmetal- nonmetal, it is a molecular compound which we haven’t discussed yet. The charge on Mn must be 2+ to balance out the 2Br - charges. The Roman number 3 tells us the charge on Co is 3+ which helps us determine the formula knowing that O is 2-.

37. 37 –A polyatomic ion is an ion consisting of two or more atoms chemically bonded together and carrying a net electric charge. We name the compounds the same way we just discussed except each polyatomic ion has a particular name. –Books typically have a table that lists common polyatomic ions. Most are oxo anions – consists of oxygen with another element (central element). Chemical Substances; Formulas and Names Polyatomic ions NO 3 - nitrate SO 4 2- sulfate NO 2 - nitrite SO 3 2- sulfite Most groups have –ate, -ite endings and differ by #O. Mn, Br, Cl, I have per- -ate, -ate, -ite, hypo- -ite.

38. 38 Ions You Should Know Polyatomic ions NH 4 + - Ammonium OH - - Hydroxide CN - - Cyanide SO 4 2- - Sulfate SO 3 2- - Sulfite ClO 4 - - perchlorate ClO 3 - - chlorate ClO 2 - - chlorite ClO - - hypochlorite Hg 2 2+ - mercury (I) or mecurous S 2 O 3 2- - thiosulfate SCN - - thiocyanate CNO - - cyanate MnO 4 - - permanganate O 2 2- - Peroxide PO 4 3- - Phosphate PO 3 3- - Phosphite CO 3 2- - Carbonate HCO 3 - - Bicarbonate or Hydrogen Carbonate N 3 - - azide NO 3 - - nitrate NO 2 - - nitrite C 2 H 3 O 2 - or CH 3 COO - - acetate Cr 2 O 7 2- - dichromate CrO 4 2- - chromate C 2 O 4 2- - oxalate HSO 4 - - bisulfate or hydrogen sulfate H 2 PO 4 - - dihydrogen phosphate

39. 39 SnSO 4 sodium sulfite Ca(ClO) 2 barium hydroxide potassium perchlorate Cr 2 (SO 4 ) 3 magnesium nitride Fe 3 (PO 4 ) 2 titanium (IV) nitrate tin (II) sulfate or stannous sulfate Na 2 SO 3 calcium hypochlorite Ba(OH) 2 KClO 4 chromium (III) sulfate Mg 3 N 2 iron (II) phosphate or ferrous phosphate Ti(NO 3 ) 4 Note: Not a polyatomic ion; monoatomic anion of N.

40. 40 Molecular compounds –Binary compounds composed of two nonmetals are usually molecular and are named using a prefix system (name same as ionic except must indicate how many atoms are present using mono, di, tri, etc.). No charges (share electrons) involved with molecular compounds, but we typically put more metallic compound first. Which way is the correct way to write the following formula based on putting the more metallic compound first? Chemical Substances; Formulas and Names NF 3 F 3 N

41. 41 –The name of the compound has the elements in the order given in the formula. Binary molecular compounds –You name the first element using the exact element name. –Name the second element by writing the stem name of the element with the suffix “–ide.” –If there is more than one atom of any given element, you add a prefix (di, tri, tetra, penta, hexa, hepta, octa, etc.) Chemical Substances; Formulas and Names

42. 42 Binary molecular compounds –N2O3–N2O3 –SF 4 – chlorine dioxide – sulfur hexafluoride –Cl 2 O 7 –HCl (g) Name this compound but think about bonding: MgCl 2 Older names: water - H 2 O, ammonia – NH 3, hydrogen sulfide – H 2 S, nitric oxide – NO, hydrazine – N 2 H 4 dinitrogen trioxide sulfur tetrafluoride ClO 2 SF 6 dichlorine heptoxide hydrogen chloride magnesium chloride; ionic comp, no prefix To name a compound, you must know if it is a molecular or ionic compound so that you know which rules to follow. If you have a metal-nonmetal (or polyatomic ion), it is an ionic compound where you name the metal first then the nonmetal with changing the ending to –ide. If you have nonmetal-nonmetal, it is a molecular compound which you do similarly as the ionic compound except that you must use prefixes to indicate the number of atoms. Drop the “a” on prefix if you encounter double vowel in name. Since this is a gas, we name using molecular rules; however, if acid we have other rules.

43. 43 –Acids are traditionally defined as compounds with a potential H + as the cation. Acids –Binary acids consist of a hydrogen ion and any single anion in aqueous solution. For example, HCl (aq) is hydrochloric acid. Binary acid: hydrostemic acid –An oxoacid is an acid containing hydrogen, oxygen, and another element. An example is HNO 3, nitric acid. The oxoacids are a derivation of the oxoanions we discussed earlier. Chemical Substances; Formulas and Names

44. 44 oxoacids Anion prefix/suffixacid prefix/suffic per- -ate ionper- -ic acid -ate ion -ic acid -ite ion -ous acid hypo- -ite ionhypo- -ous acid NO 3 - nitrate ionHNO 3 nitric acid NO 2 - nitrite ion HNO 2 nitrous acid ClO 4 - perchlorate ion HClO 4 perchloric acid SO 4 2- sulfate ion H 2 SO 4 sulfuric acid PO 4 3- phosphate ion H 3 PO 4 phosphoric acid If you learn the oxoanions, you can easily adapt to naming the oxoacids: -ate  -ic and –ite  -ous For some species there is a change in spelling in the name.

45. Chemical Substances; Formulas and Names Hydrates 45 –A hydrate is 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 prefix to indicate the number of water molecules per formula unit of the compound. CuSO 4. 5H 2 O Magnesium sulfate heptahydrate HW 15 - 18 copper(II)sulfate pentahydrate MgSO 4. 7H 2 O code for all: names

46. Chemical Substances; Formulas and Names Naming simple compounds 46 –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. –Inorganic compounds are compounds composed of elements other than carbon.

47. Chemical Formulas; Molecular and Ionic Substances Organic compounds 47 –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 - compounds containing only hydrogen and carbon. –Common examples include methane, CH 4, ethane, C 2 H 6, and propane, C 3 H 8.

48. Classifying Compounds Organic vs. Inorganic in the 18 th century, compounds from living things were called organic; compounds from the nonliving environment were called inorganic organic compounds easily decomposed and could not be made in 18 th century lab inorganic compounds very difficult to decompose, but able to be synthesized 48

49. Modern Classifying Compounds Organic vs. Inorganic today we commonly make organic compounds in the lab and find them all around us organic compounds are mainly made of C and H, sometimes with O, N, P, S, and trace amounts of other elements the main element that is the focus of organic chemistry is carbon 49

50. Carbon Bonding carbon atoms bond almost exclusively covalently – compounds with ionic bonding C are generally inorganic when C bonds, it forms 4 covalent bonds – 4 single, 1 double + 2 singles, 2 double, or 1 triple + 1 single – carbon is unique in that it can form limitless chains of C atoms, both straight and branched, and rings of C atoms 50

51. Examples of Carbon Compounds 51

52. Classifying Organic Compounds there are two main categories of organic compounds, hydrocarbons and functionalized hydrocarbons hydrocarbons contain only C and H most fuels are mixtures of hydrocarbons 52

53. Classifying Hydrocarbons hydrocarbons containing only single bonds are called alkanes hydrocarbons containing one or more C=C double bonds are called alkenes hydrocarbons containing one or more C  C triple bonds are called alkynes hydrocarbons containing C 6 “benzene” ring are called aromatic 53

54. 54

55. Naming Straight Chain Hydrocarbons consists of a base name to indicate the number of carbons in the chain, with a suffix to indicate the class and position of multiple bonds – suffix –ane for alkane, –ene for alkene, –yne for alkyne 55 Base NameNo. of CBase NameNo. of C meth-1hex-6 eth-2hept-7 prop-3oct-8 but-4non-9 pent-5dec-10

56. Functionalized Hydrocarbons functional groups are non-carbon groups that are on the molecule substitute one or more functional groups replacing H’s on the hydrocarbon chain generally, the chemical reactions of the compound are determined by the kinds of functional groups on the molecule 56

57. Functional Groups 57

58. Chemical Reactions: Equations Writing chemical equations 58 –The reactants (consumed; left side of reaction) are starting substances in a chemical reaction. The arrow means “yields.” The formulas on the right side of the arrow represent the products (produced). –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 Reactants (consumed)  Products (produced)

59. Chemical Reactions: Equations Writing chemical equations 59 –In many cases, it is useful to indicate the states of the substances in the equation (s, g, l, aq). –When you use these labels, the previous equation becomes We write above the arrow any conditions for the reaction such as pressure, catalyst, heat, etc. A reaction gives a recipe for the amount of reactants needed to produce the amount of products. Species with no coefficient have an understood coefficient of 1.

60. Chemical Reactions: Equations Writing chemical equations 60 –The law of conservation of mass dictates that the total number of atoms of each element on both sides of a chemical equation must match. The equation is then said to be balanced. 22 We must have the same number of atoms on both sides for a reaction to be considered balanced and obeying the law of conservation of mass. To balance a reaction: 1.First, balance the atoms for elements that occur in only one substance on each side of the reaction. In this problem, O is involved with two substances on the product side; therefore, I will wait on balancing O until later. C & H are only in one species on both sides so I will balance them first. C needs no changes because there are one on each side, but H needs a 2 in front of H 2 O to balance the 4H on the reactants side. 2.Now that we have changed the coefficient of one of the O on the product side, it is easier to balance the O. We determine that we need a 2 coefficient on the O 2 to balance the O on both sides at 4. Now the equation is balanced with 1C, 4O, and 4H on both sides.

61. Chemical Reactions: Equations 61 Fe 2 (SO 4 ) 3 : has 2-Fe, 3x1 = 3-S, 3x4 = 12-O Caution: For formulas that have subscripts, you must account for all atoms especially when dealing with parentheses for polyatomic species. For example, Caution: Remember that you can’t change the subscripts in formulas to balance equations; you may only change coefficients. If you change the subscripts, you are changing the substance.

62. Chemical Reactions: Equations 62 HW 19 22 6 6 3 2 [] 22 Technique to handle odd numbers: determine number needed and divide by subscript of species. Next, you multiple the entire equation by the subscript to obtain whole numbers. code: balance