1. Prentice Hall © 2003Chapter 2 Chapter 2 Atoms, Molecules, and Ions CHEMISTRY The Central Science 9th Edition David P. White

2. Prentice Hall © 2003Chapter 2 John Dalton: –Each element is composed of atoms –All atoms of an element are identical. –In chemical reactions, the atoms are not changed. Compounds are formed when atoms of more than one element combine. Dalton’s law of multiple proportions: When two elements form different compounds, the mass ratio of the elements in one compound is related to the mass ratio in the other by a small whole number. The Atomic Theory of Matter

3. Prentice Hall © 2003Chapter 2 The Atomic Theory of Matter

4. Prentice Hall © 2003Chapter 2 The Discovery of Atomic Structure The ancient Greeks were the first to postulate that matter consists of indivisible constituents. Later scientists realized that the atom consisted of charged entities. Cathode Rays and Electrons A cathode ray tube (CRT) is a hollow vessel with an electrode at either end. A high voltage is applied across the electrodes.

5. Prentice Hall © 2003Chapter 2 The Discovery of Atomic Structure Cathode Rays and Electrons The voltage causes negative particles to move from the negative electrode to the positive electrode. The path of the electrons can be altered by the presence of a magnetic field. Consider cathode rays leaving the positive electrode through a small hole. –If they interact with a magnetic field perpendicular to an applied electric field, the cathode rays can be deflected by different amounts.

6. Prentice Hall © 2003Chapter 2 The Discovery of Atomic Structure Cathode Rays and Electrons –The amount of deflection of the cathode rays depends on the applied magnetic and electric fields. –In turn, the amount of deflection also depends on the charge to mass ratio of the electron. In 1897, Thomson determined the charge to mass ratio of an electron to be 1.76  10 8 C/g. Goal: find the charge on the electron to determine its mass.

7. Prentice Hall © 2003Chapter 2 The Discovery of Atomic Structure Cathode Rays and Electrons

8. Prentice Hall © 2003Chapter 2 The Discovery of Atomic Structure Cathode Rays and Electrons Consider the following experiment: Oil drops are sprayed above a positively charged plate containing a small hole. As the oil drops fall through the hole, they are given a negative charge. Gravity forces the drops downward. The applied electric field forces the drops upward. When a drop is perfectly balanced, the weight of the drop is equal to the electrostatic force of attraction between the drop and the positive plate.

9. The Discovery of Atomic Structure Cathode Rays and Electrons

10. Prentice Hall © 2003Chapter 2 The Discovery of Atomic Structure Cathode Rays and Electrons Using this experiment, Millikan determined the charge on the electron to be 1.60  10 -19 C. Knowing the charge to mass ratio, 1.76  10 8 C/g, Millikan calculated the mass of the electron: 9.10  10 -28 g. With more accurate numbers, we get the mass of the electron to be 9.10939  10 -28 g.

11. Prentice Hall © 2003Chapter 2 The Discovery of Atomic Structure Radioactivity Consider the following experiment: A radioactive substance is placed in a shield containing a small hole so that a beam of radiation is emitted from the hole. The radiation is passed between two electrically charged plates and detected. Three spots are noted on the detector: –a spot in the direction of the positive plate, –a spot which is not affected by the electric field, –a spot in the direction of the negative plate.

12. Prentice Hall © 2003Chapter 2 The Discovery of Atomic Structure Radioactivity

13. Prentice Hall © 2003Chapter 2 The Discovery of Atomic Structure Radioactivity A high deflection towards the positive plate corresponds to radiation which is negatively charged and of low mass. This is called  -radiation (consists of electrons). No deflection corresponds to neutral radiation. This is called  -radiation. Small deflection towards the negatively charged plate corresponds to high mass, positively charged radiation. This is called  -radiation.

14. The Discovery of Atomic Structure The Nuclear Atom From the separation of radiation we conclude that the atom consists of neutral, positively, and negatively charged entities. Thomson assumed all these charged species were found in a sphere.

15. Prentice Hall © 2003Chapter 2 The Discovery of Atomic Structure The Nuclear Atom Rutherford carried out the following experiment: A source of  -particles was placed at the mouth of a circular detector. The  -particles were shot through a piece of gold foil. Most of the  -particles went straight through the foil without deflection. Some  -particles were deflected at high angles. If the Thomson model of the atom was correct, then Rutherford’s result was impossible.

17. Prentice Hall © 2003Chapter 2 The Discovery of Atomic Structure The Nuclear Atom In order to get the majority of  -particles through a piece of foil to be undeflected, the majority of the atom must consist of a low mass, diffuse negative charge  the electron. To account for the small number of high deflections of the  -particles, the center or nucleus of the atom must consist of a dense positive charge.

18. The Nuclear Atom Rutherford modified Thomson’s model as follows: –assume the atom is spherical but the positive charge must be located at the center, with a diffuse negative charge surrounding it. The Discovery of Atomic Structure

19. Prentice Hall © 2003Chapter 2 The Modern View of Atomic Structure The atom consists of positive, negative, and neutral entities (protons, electrons, and neutrons). Protons and neutrons are located in the nucleus of the atom, which is small. Most of the mass of the atom is due to the nucleus. –There can be a variable number of neutrons for the same number of protons. Isotopes have the same number of protons but different numbers of neutrons. Electrons are located outside of the nucleus. Most of the volume of the atom is due to electrons.

20. Prentice Hall © 2003Chapter 2 The Modern View of Atomic Structure

21. Prentice Hall © 2003Chapter 2 The Modern View of Atomic Structure Isotopes, Atomic Numbers, and Mass Numbers Atomic number (Z) = number of protons in the nucleus. Mass number (A) = total number of nucleons in the nucleus (i.e., protons and neutrons). By convention, for element X, we write Z A X. Isotopes have the same Z but different A. We find Z on the periodic table.

22. Prentice Hall © 2003Chapter 2 Atomic Weights The Atomic Mass Scale 1 H weighs 1.6735 x 10 -24 g and 16 O 2.6560 x 10 -23 g. We define: mass of 12 C = exactly 12 amu. Using atomic mass units: 1 amu = 1.66054 x 10 -24 g 1 g = 6.02214 x 10 23 amu

23. Prentice Hall © 2003Chapter 2 Atomic Weights Average Atomic Masses Relative atomic mass: average masses of isotopes: –Naturally occurring C: 98.892 % 12 C + 1.108 % 13 C. Average mass of C: (0.98892)(12 amu) + (0.0108)(13.00335) = 12.011 amu. Atomic weight (AW) is also known as average atomic mass (atomic weight). Atomic weights are listed on the periodic table.

24. Prentice Hall © 2003Chapter 2 The Periodic Table The Periodic Table is used to organize the 114 elements in a meaningful way. As a consequence of this organization, there are periodic properties associated with the periodic table.

25. The Periodic Table

26. Prentice Hall © 2003Chapter 2 The Periodic Table Columns in the periodic table are called groups (numbered from 1A to 8A or 1 to 18). Rows in the periodic table are called periods. Metals are located on the left hand side of the periodic table (most of the elements are metals). Non-metals are located in the top right hand side of the periodic table. Elements with properties similar to both metals and non- metals are called metalloids and are located at the interface between the metals and non-metals.

27. Prentice Hall © 2003Chapter 2 The Periodic Table Some of the groups in the periodic table are given special names. These names indicate the similarities between group members: Group 1A: Alkali metals. Group 2A: Alkaline earth metals. Group 6A: Chalcogens. Group 7A: Halogens. Group 8A: Noble gases.

28. Prentice Hall © 2003Chapter 2 Molecules and Molecular Compounds Molecules and Chemical Formulas Molecules are assemblies of two or more atoms bonded together. Each molecule has a chemical formula. The chemical formula indicates –w hich atoms are found in the molecule, and –in what proportion they are found. Compounds formed from molecules are molecular compounds. Molecules that contain two atoms bonded together are called diatomic molecules.

29. Prentice Hall © 2003Chapter 2 Molecules and Molecular Compounds Molecules and Chemical Formulas

30. Prentice Hall © 2003Chapter 2 Molecules and Molecular Compounds Molecular and Empirical Formulas Molecular formulas –give the actual numbers and types of atoms in a molecule. –Examples: H 2 O, CO 2, CO, CH 4, H 2 O 2, O 2, O 3, and C 2 H 4.

31. Prentice Hall © 2003Chapter 2 Molecules and Molecular Compounds Molecular and Empirical Formulas Molecular formulas

32. Prentice Hall © 2003Chapter 2 Molecules and Molecular Compounds Molecular and Empirical Formulas Empirical formulas –give the relative numbers and types of atoms in a molecule. –That is, they give the lowest whole number ratio of atoms in a molecule. –Examples: H 2 O, CO 2, CO, CH 4, HO, CH 2.

33. Prentice Hall © 2003Chapter 2 Molecules and Molecular Compounds Picturing Molecules Molecules occupy three dimensional space. However, we often represent them in two dimensions. The structural formula gives the connectivity between individual atoms in the molecule. The structural formula may or may not be used to show the three dimensional shape of the molecule. If the structural formula does show the shape of the molecule, then either a perspective drawing, ball-and- stick model, or space-filling model is used.

34. Prentice Hall © 2003Chapter 2 Molecules and Molecular Compounds Picturing Molecules

35. Prentice Hall © 2003Chapter 2 Ions and Ionic Compounds When an atom or molecule loses electrons, it becomes positively charged. –For example, when Na loses an electron it becomes Na +. Positively charged ions are called cations.

36. Prentice Hall © 2003Chapter 2 Ions and Ionic Compounds When an atom or molecule gains electrons, it becomes negatively charged. For example when Cl gains an electron it becomes Cl . Negatively charged ions are called anions. An atom or molecule can lose more than one electron.

37. Prentice Hall © 2003Chapter 2 Ions and Ionic Compounds In general: metal atoms tend to lose electrons to become cations; nonmetal ions tend to gain electrons to form anions. Predicting Ionic Charge The number of electrons an atom loses is related to its position on the periodic table.

38. Prentice Hall © 2003Chapter 2 Ions and Ionic Compounds Predicting Ionic Charge

39. Prentice Hall © 2003Chapter 2 Ionic Compounds The majority of chemistry involves the transfer of electrons between species. Example: –To form NaCl, the neutral sodium atom, Na, must lose an electron to become a cation: Na +. –The electron cannot be lost entirely, so it is transferred to a chlorine atom, Cl, which then becomes an anion: Cl -. –The Na + and Cl - ions are attracted to form an ionic NaCl lattice which crystallizes. Ions and Ionic Compounds

40. Prentice Hall © 2003Chapter 2 Ions and Ionic Compounds Ionic Compounds

41. Prentice Hall © 2003Chapter 2 Ions and Ionic Compounds Ionic Compounds Important: note that there are no easily identified NaCl molecules in the ionic lattice. Therefore, we cannot use molecular formulas to describe ionic substances. Consider the formation of Mg 3 N 2 : Mg loses two electrons to become Mg 2+ ; Nitrogen gains three electrons to become N 3-. For a neutral species, the number of electrons lost and gained must be equal.

42. Prentice Hall © 2003Chapter 2 Ions and Ionic Compounds Ionic Compounds However, Mg can only lose electrons in twos and N can only accept electrons in threes. Therefore, Mg needs to lose 6 electrons (2  3) and N gain those 6 electrons (3  2). I.e., 3Mg atoms need to form 3Mg 2+ ions (total 3  2+ charges) and 2 N atoms need to form 2N 3- ions (total 2  3- charges). Therefore, the formula is Mg 3 N 2.

43. Prentice Hall © 2003Chapter 2 Naming Inorganic Compounds Naming of compounds, nomenclature, is divided into organic compounds (those containing C) and inorganic compounds (the rest of the periodic table). Cations formed from a metal have the same name as the metal. Example: Na + = sodium ion. If the metal can form more than one cation, then the charge is indicated in parentheses in the name. Examples: Cu + = copper(I); Cu 2+ = copper(II).

44. Prentice Hall © 2003Chapter 2 Naming Inorganic Compounds Cations formed from non-metals end in -ium. Example: NH 4 + ammonium ion.

45. Prentice Hall © 2003Chapter 2 Naming Inorganic Compounds

46. Prentice Hall © 2003Chapter 2 Naming Inorganic Compounds Monatomic anions (with only one atom) are called  ide. Example: Cl  is chloride ion. Exceptions: hydroxide (OH  ), cyanide (CN  ), peroxide (O 2 2  ). Polyatomic anions (with many atoms) containing oxygen end in -ate or -ite. (The one with more oxygen is called  ate.) Examples: NO 3 - is nitrate, NO 2 - is nitrite.

47. Prentice Hall © 2003Chapter 2 Naming Inorganic Compounds Polyatomic anions containing oxygen with more than two members in the series are named as follows (in order of decreasing oxygen): per-….-ate -ate -ite hypo-….-ite

48. Prentice Hall © 2003Chapter 2 Naming Inorganic Compounds

49. Prentice Hall © 2003Chapter 2 Naming Inorganic Compounds Polyatomic anions containing oxygen with additional hydrogens are named by adding hydrogen or bi- (one H), dihydrogen (two H), etc., to the name as follows: CO 3 2- is the carbonate anion HCO 3 - is the hydrogen carbonate (or bicarbonate) anion. H 2 PO 4 - is the dihydrogen phosphate anion.

50. Prentice Hall © 2003Chapter 2 Naming Inorganic Compounds Name the anion then cation for the ionic compound. Example: BaBr 2 = barium bromide.

51. Prentice Hall © 2003Chapter 2 Naming Inorganic Compounds

52. Prentice Hall © 2003Chapter 2 Naming Inorganic Compounds Names and Formulas for Acids The names of acids are related to the names of anions: -ide becomes hydro-….-ic acid; -ate becomes -ic acid; -ite becomes -ous acid.

54. Prentice Hall © 2003Chapter 2 Naming Inorganic Compounds Names and Formulas of Binary Molecular Compounds Binary molecular compounds have two elements. The most metallic element is usually written first (i.e., the one to the farthest left on the periodic table). Exception: NH 3. If both elements are in the same group, the lower one is written first. Greek prefixes are used to indicate the number of atoms.

56. Prentice Hall © 2003Chapter 2 Some Simple Organic Compounds Alkanes Organic chemistry: the study of the chemistry of carbon compounds. Alkanes contain only C and H and are called hydrocarbons. The names of alkanes all end in the suffix –ane. Alkanes are named according to the number of C atoms in their backbone chain: Methane has one C atom (CH 4 ) Ethane has two C atoms (CH 3 CH 3 ) Propane has three C atoms (CH 3 CH 2 CH 3 ), etc.

57. Prentice Hall © 2003Chapter 2 Some Simple Organic Compounds Some Derivatives of Alkanes When H atoms in alkanes are replaced by heteroatoms (atoms other than C or H), then we have introduced a functional group into the alkane. When an H is replaced by –OH, then we form an alcohol. Alcohols are also named by the number of C atoms. Consider propanol: there are two places for the OH: on an end C or the middle C. When the OH is located on the end C, we call the substance 1-propanol. When the OH is on the middle C, we have 2-propanol.

58. Prentice Hall © 2003Chapter 2 Some Simple Organic Compounds Some Derivatives of Alkanes When the single bonds in an alkane are replaced by one or more double bonds, then we form alkenes. When a double bond is formed between C and O, we form carboxylic acids, ketones, aldehydes, and esters. Any organic molecule with double or triple bonds is called unsaturated.

59. Prentice Hall © 2003Chapter 2 End of Chapter 2 Atoms, Molecules, and Ions