1. Chapter 2: Atoms, Molecules and Ions Joseph DePasquale

2. Building Blocks of Chemistry 2.1: Early Ideas in Atomic Theory 2.2: Evolution of Atomic Theory 2.3: Atomic Structure and Symbolism 2.4: Chemical Formulas 2.5: The Periodic Table 2.6: Molecular and Ionic Compounds 2.7: Chemical Nomenclature

3. 2.1 Atomic Theory (1807) John Dalton 1) Matter is composed of small particles called atoms. An atom is the smallest unit of an element that can participate in a chemical change. 2) An element consists of only one type of atom. 3) Atoms of one element differ in properties from atoms of all other elements.

4. 2.1 Atomic Theory (1807) John Dalton 4) A Compound consists of atoms of two or more elements combined in a small, whole-number ratio. In a given compound, the number of atoms of each of its elements are always present in the same ratio.

5. 2.1 Atomic Theory (1807) John Dalton 5) Atoms are neither created nor destroyed during a chemical change, but instead rearrange to yield a different type(s) of matter.

6. Fundamental Laws of Matter Dalton’s Atomic Theory is the basis for the three fundamental laws of matter. 1) Law of Conservation of Mass There is no change in mass in an ordinary chemical reaction. If atoms are conserved in a reaction, then mass must also be conserved. Matter can be neither created nor destroyed

7. Fundamental Laws of Matter 2) Law of Definite Proportions (or Constant Composition) All samples of a pure compound always contains the same elements in the same proportions by mass. Example: Pure water has the same composition everywhere.

8. Fundamental Laws of Matter 3) Law of Multiple Proportions Applies to situations where 2 elements form more than one compound. The law states that when two elements react to form more than one compound, a fixed mass of one element will react with masses of the other element in a ratio of small, whole number.

9. The Law of Multiple Proportions: Two different chromium-oxygen compounds Green: 2.167 g Cr/1 g O Red: 1.083 g Cr/1 g O

11. 2.2 Evolution of Atomic Theory Atomic theory raised more questions than it answered. What are atoms composed of? Is there something smaller than an atom? Almost 100 years after atomic theory was proposed, subatomic particles were discovered.

12. Discovery of the Electron J.J. Thomson (1897) Cathode Rays Experiment Cathode rays are emitted by all materials. These cathode ray particles turned out to be electrons. Electrons have a charge of -1 Electrons have a very small mass. Nobel Laureates: J. J. Thompson & Ernest Rutherford

13. Discovery of the Electron 13 Evidence that cathode rays are composed of negatively charged particles, which we now refer to as electrons: 1) Repelled by the negative end of a magnet. 2) Attracted to the positive end of a magnet.

14. Thomson’s Model of the Atom “Plum Pudding Model” Atoms consisted of a positively charged mass with negatively charged electrons embedded in that mass. Raisins = electrons Bread = Positively charge

15. Discovery of the Atomic Nucleus Ernest Rutherford (1911) Gold Foil Scattering experiment Discovered the Atomic Nucleus and determined that it is was positively charged. Nobel Laureates: J. J. Thompson & Ernest Rutherford

16. Rutherford’s Gold Foil Experiment Alpha particles (positively charged) were fired at gold foil. Most passed through the foil Some were reflected at sharp angles

17. Conclusions: 1) The center of an atom contains a small positively charged center that contains most of the atom’s mass. 2) Most of the atom is “empty space”.

18. Structure/Components of an Atom Nucleus: Positively charged center. Composed of two main parts 1) Protons: Charge = +1 2) Neutrons (Discovered by James Chadwick, 1932): Charge = 0 Protons and neutrons are much heavier than electrons. Over 99.9% of an atom’s mass is concentrated in the nucleus. Electron Cloud: Negatively charged outer region. Accounts for most of the atom’s volume, composed of only electrons. Electrons have a charge = -1, very small mass, and spread far apart.

19. 2.3: Atomic Structure and Symbolism

20. Atomic Structure Diameter of an atom ~ 10 -10 m Diameter of a nucleus is 100,000 times smaller ~ 10 -15 m The nucleus accounts for most of the atom’s mass, but very little of it’s volume.

21. 21 Subatomic Particles Small units are needed to describe the properties of subatomic particles. Atomic mass unit (amu) Fundamental unit of charge (e) A proton has a mass about 1800 times greater than that of an electron. Neutrons are just slightly larger than protons.

22. Atomic Number (Z) All atoms of a particular element have the same number of protons in the nucleus. The number of protons in the nucleus of an atom is its atomic number (Z). Therefore, all atoms of the same element have the same atomic number. In a neutral atom: Number of protons = Number of electrons

23. Ions When the number of protons and electrons are not equal, the atom is electrically charged and is called an ion. Atomic charge = Atoms acquire charge by losing or gaining electrons. There is no change in the number of protons in the nucleus when an ion forms. Only the number of electrons increases or decreases.

25. Cations and Anions Positively charged ions are called cations. Cations form by loss of electrons. Negatively charged ions are called anions. Anions form by gain of electrons.

26. Mass Number (A) Mass number (A) – The total number of protons and neutrons in an atom. A = number of protons + number of neutrons Atoms of the same element do not always have the same mass number. Atoms must have the same number of protons, but the number of neutrons may vary.

27. Isotopes Atoms that contain the same number of protons but a different number of neutrons are called isotopes. Isotopes are atoms of the same element that differ in mass.

28. Chemical Symbols Each element is assigned a symbol. Consists of one or two letters. Some newer ones contain three letters. First letter is always capitalized. Second and third letters are always lower case. The symbols are typically derived from the English name. Sometimes the symbol is derived from other languages.

29. Chemical Symbols The composition of an atom can be represented by a chemical symbol. Elemental symbol (from periodic table) Mass number (number of protons + neutrons) Atomic number (number of protons)

30. Three Isotopes of Hydrogen

31. Isotopes of Hydrogen Notice the ice on top of the water. Notice some of the ice on the bottom of the glass.

32. Atomic Mass Each proton and each neutron has a mass of ~ 1 amu Each electron weighs far less. Therefore the atomic mass of a single atom in amu is approx. equal to its mass number.

33. Atomic Mass However, most elements exist naturally as a mixture of two or more isotopes. The periodic table shows the average atomic mass of each element, which represents the average atomic mass of the naturally occurring mixture of isotopes of that element. Therefore, the atomic masses in the periodic table are not whole numbers. Exception:

34. Average Atomic Mass of Carbon Carbon has two main isotopes, C-12 and C-13, which account for basically 100% of naturally occurring carbon. Average atomic masses are calculated as weighted averages.

35. Determination of Atomic Mass and Isotopic Abundance An atom’s atomic mass can be determined to a highly precise value by using a high tech instrument known as a mass spectrometer. The mass spectrometer separates matter based on its mass and charge. This data can be used to determine the abundance and mass of each isotope in a naturally occurring sample of that element.

36. Mass Spectrometer (instrument) Mass spectrum (data)

37. 2.4: Chemical Formulas Molecular Formula – A representation of a molecule or compound which consists of the following: 1) Chemical symbols to indicate the types of atoms. 2) Subscripts after the symbol to indicate the number of each type of atom.

38. Molecular Formulas Examples: Water, H 2 O Ammonia, NH 3 Methane, CH 4

39. Structural Formulas A Structural Formula shows the same information as a molecular formula but also how the atoms are connected.

40. Structural Formulas Showing the Molecule’s Geometry Space-Filling Model Ball-and-Stick Model More on geometries in Ch. 7…

41. Structural Formula: Caffeine

42. Molecular Elements Some elements exist as molecules: Diatomic molecules: H 2, N 2, O 2, F 2, Cl 2, Br 2, I 2, Others: P 4, S 8

43. Empirical Formulas Empirical Formula – Indicates the simplest whole number ratio of the number of atoms in a compound. Molecular formula – Indicates the actual number of atoms in a molecule or compound.

44. Empirical Formula vs. Molecular Formula Dividing the subscripts in the molecular formula by the lowest common denominator gives the empirical formula.

45. Isomers It is possible for the same atoms to be arranged in different ways. Isomers – Compounds with the same molecular formula, but different structural formula.

46. Isomers Isomers – Compounds with the same molecular formula, but different structural formula. Consider the molecular formula, C 2 H 4 O 2

47. 2.5: The Periodic Table The first periodic table: Mendeleev (1869) and Lothar Meyer (1870) First to propose primitive versions of the periodic table.

48. The First Periodic Table Elements listed in order of increasing atomic mass and grouped in columns by properties.

49. Modern Day Periodic Table

50. Elements listed in order of increasing atomic number and grouped in columns by properties. Periodic Law – The properties of the elements are periodic functions of their atomic numbers. Periods (rows) First period: Second period: Third period: Groups (columns) Modern Day Periodic Table

51. Elements are categorized as metals, nonmetals, or metalloids. Metals are good conductors of heat and electricity. Nonmetals are poor conductors of heat and electricity. Metalloids have properties intermediate between that of metals and nonmetals.

52. Metals, Nonmetals, and Metalloids

53. 53 Nonmetals

54. Blocks in the Periodic Table Main group elements (or representative elements) Groups: 1, 2, 13-18 Transition metals Groups: 3-12 Inner transition metals Lanthanides Actinides

56. Common Names for Main Group Elements Alkali Metals, Group 1 (except hydrogen) All soft, reactive metals. Alkaline Earth Metals, Group 2 Harder and less reactive than the alkali metals. Halogens, Group 17 All reactive non-metals. Noble Gases, Group 18 All relatively unreactive gases.

58. 2.6 Molecular and Ionic Compounds Isolated atoms rarely appear in nature. Only the noble gasses consist of individual, non-reactive atoms. Atoms tend to combine with each other in various ways to either form: 1) Molecules 2) Ionic Compounds During the formation of molecules and compounds, only electrons interact. The nucleus remains unchanged

59. Formation of Monoatomic Ions Monoatomic Ions – The periodic table can help us predict the charge of the ion formed by many main group elements. Many main group elements lose or gain electrons so that they have the same number of electrons as a near by noble gas.

60. Formation of Cations Many main group metals lose electrons to form cations. Group 1  +1 Group 2  +2

62. Formation of Anions Many main group non-metals gain electrons to form anions. Group 16  -2 Group 17  -1

64. Transition Metal Cations Several transition metals can form cations of varying charge. These metals typically DO NOT form ions that have the same number of electrons as a noble gas. Examples: Iron commonly forms Fe 2+ and Fe 3+ Chromium commonly forms Cr 3+ and Cr 6+

65. Polyatomic Ions All ions thus far have been monoatomic. Many important ions in chemistry contain more than one atom. These are known as polyatomic ions. OH -, hydroxide ion NH 4 +, ammonium ion You can think of polyatomic ions as “charged molecules” Most of the polyatomic anions contain one or more oxygen atoms and are referred to as oxyanions.

66. Table 2.5: Polyatomic Ions

67. Oxyanions There is a system to help you remember the oxyanions. When a nonmetal forms two oxyanions -ate is the suffix used for the ion with the larger number of oxygen atoms. -ite is the suffix used for the ion with the smaller number of oxygen atoms. When a nonmetal forms more than two oxyanions, prefixes are used in addition to -ate and -ite per- (largest number of oxygens) hypo- (smallest number of oxygens)

68. Oxyanions of Nitrogen, Sulfur and Chlorine

69. When electrons are transferred and ions form, an ionic bond results. Ionic bond – Electrostatic attraction that holds ions together. Ionic bonds form as a result of the transferring of electrons While covalent bonds form as a result of sharing of electrons. Ionic Bonds

70. Compounds that contains ions held together by ionic bonds are called ionic compounds. Ionic compounds typically contain a metal and nonmetal. Ionic Compounds

71. Formulas of Ionic Compounds The compound’s formula shows the simplest ratio of cations and anions needed to produce an overall neutral compound. Charge balance – The total positive charge of the cations must equal the total negative charge of the anions.

72. There are no discrete NaCl molecules, only Na + and Cl - ions in a continuous network. NaCl, An Ionic Compound

73. Ionic compounds with polyatomic ions

74. Properties of Ionic Compounds Ionic compounds have very high melting and boiling points. Melting requires the strong ionic bonds to be broken. Oppositely charged particles need to be separated. The stronger the bond, the more energy required to break that bond. Ionic compounds tend to have high solubility in water. Molten ionic compounds and water solutions of dissolved ionic compounds can conduct electricity.

75. Many compounds consist of discrete neutral molecules. Molecular compounds form when atoms share electrons, forming covalent bonds. Molecular compounds usually consist of all nonmetals. Properties of molecular compounds: Molecular Compounds

76. 2.7 Chemical Nomenclature Compounds are identified by their formulas and names. Nomenclature – A collection of rules for naming things. The rules for naming a compound depends on its type. Ionic compounds – Metal and nonmetal Molecular compounds – All nonmetals Simple binary molecules Acids

77. Naming Ionic Compounds The name of the cation is always listed first and the name of the anion is listed second. Monatomic Cations - Same name as the metal they are derived from Na + : Al +3 : For transition metals, roman numerals in parentheses after the name are used to indicate the ion’s charge : Fe 2+ : Fe 3+ :

78. Naming Ionic Compounds Monatomic anions are named by using the name of the element, but with its ending replaced with the suffix, –ide. O 2- is named S 2- is named Cl - is named Polyatomic ions (both cations and anions): Just use the name of the ion.

79. Binary Molecular Compounds Consist of two non-metals. Unlike ionic compounds, there is no simple way to deduce the formula of a binary molecular compound. Systematic naming 1. The first word is the name of the first element in the formula, with a Greek prefix if necessary 2. The second word consists of The appropriate Greek prefix The name of the second element with its ending replaced with the suffix, -ide

80. These prefixes must be committed to memory. The “mono” prefix, meaning one atom, is used only with the second word in the name but never the first word.

81. These names are NOT derived from the naming rules, but are commonly used. What is water’s actual systematic name? Common Names for Molecular Compounds

83. Acids Some molecular compounds contain H atoms that ionize in water to produce H + ions. These compounds are called acids. The H atom(s) are always listed first in the molecular formula. Example: HCl As a molecule, HCl is named hydrogen chloride. When put in water, HCl is named hydrochloric acid Special acid naming rules when in water.

84. Binary Acids When in water: 1) Change the word hydrogen to hydro- 2) The name of the second element with its ending replaced with the suffix, –ic acid

85. Oxyacids Many acids contain oxygen in addition to hydrogen and are referred to as oxyacids. Oxyacids typically consist of hydrogen combined with an oxyanion. The name of the oxyacid is derived from the name of the oxyanion. Never use the prefix hydro-.

86. Oxyacid Naming Rules If the oxoanion ends with –ate Replace –ate with –ic acid If the oxoanion ends with –ite Replace –ite with –ous acid The prefixes, per- and hypo- are retained in the acid name.