1. Chapter 2 Atoms, Molecules and Ions

2. Atomic Theory of Matter

3. Law of Conservation of Mass
The total mass of substances present at the end of a chemical process is the same as the mass of substances present before the process took place.

4. Law of Definite ( or constant ) composition
No matter what its source, a particular chemical compound is composed of the same elements in the same proportions by mass.

5. LAW OF MULTIPLE PROPORTIONS
If two elements can combine to form more than one compound, the masses of one element that combine with a fixed mass of the other element are in the ratio of small whole numbers.
Carbon monoxide g C to g O
Carbon dioxide g C to g O

6. The Electron
Streams of negatively charged particles were found to emanate from cathode tubes, causing fluorescence.
J. J. Thomson is credited with their discovery (1897).

7. The Electron
Thomson measured the charge/mass ratio of the electron to be 1.76  108 coulombs/gram (C/g).

8. Millikan Oil-Drop Experiment
Once the charge/mass ratio of the electron was known, determination of either the charge or the mass of an electron would yield the other.
Robert Millikan (University of Chicago) determined the charge on the electron in 1909.

9. Radioactivity
Radioactivity is the spontaneous emission of radiation by an atom.
It was first observed by Henri Becquerel.
Marie and Pierre Curie also studied it.

10. Radioactivity
Three types of radiation were discovered by Ernest Rutherford:
 particles
 particles
 rays

11. The Atom, circa 1900
The prevailing theory was that of the “plum pudding” model, put forward by Thomson.
It featured a positive sphere of matter with negative electrons imbedded in it.

12. Discovery of the Nucleus
Ernest Rutherford shot  particles at a thin sheet of gold foil and observed the pattern of scatter of the particles.

13. The Nuclear Atom
Since some particles were deflected at large angles, Thomson’s model could not be correct.

14. The Nuclear Atom
Rutherford postulated a very small, dense nucleus with the electrons around the outside of the atom.
Most of the volume of the atom is empty space.

15. Other Subatomic Particles
Protons were discovered by Rutherford in 1919.
Neutrons were discovered by James Chadwick in 1932.

16. Subatomic Particles
Protons and electrons are the only particles that have a charge.
Protons and neutrons have essentially the same mass.
The mass of an electron is so small we ignore it.

17. Symbols of Elements
Elements are symbolized by one or two letters.

18. Symbols of Elements
All atoms of the same element have the same number of protons, which is called the atomic number, Z.

19. Symbols of Elements
The mass of an atom in atomic mass units (amu) is the total number of protons and neutrons in the atom.

20. X H H (D) H (T) U Atomic number (Z) = number of protons in nucleus
Mass number (A) = number of protons + number of neutrons
= atomic number (Z) + number of neutrons
Isotopes are atoms of the same element (X) with different numbers of neutrons in their nuclei
Mass Number X A Z
Element Symbol
Atomic Number H 1
H (D) 2
H (T) 3 U
235 92
238

21. Isotopes Isotopes are atoms of the same element with different masses.
Isotopes have different numbers of neutrons.

22. Do You Understand Isotopes?
How many protons, neutrons, and electrons are in C 14 6 ?
6 protons, 8 (14 - 6) neutrons, 6 electrons
How many protons, neutrons, and electrons are in C 11 6 ?
6 protons, 5 (11 - 6) neutrons, 6 electrons

23. Atomic Mass
Atomic and molecular masses can be measured with great accuracy using a mass spectrometer.

24. Average Mass
Because in the real world we use large amounts of atoms and molecules, we use average masses in calculations.
Average mass is calculated from the isotopes of an element weighted by their relative abundances.

25. Atomic weight = (0.7578)(34.969 amu) + (0.2422)(36.966 amu)
Sample Exercise 2.4 Calculating the Atomic Weight of an Element from Isotopic Abundances
Naturally occurring chlorine is 75.78% 35Cl (atomic mass amu) and 24.22% 37Cl (atomic mass amu).Calculate the atomic weight of chlorine.
Solution
We can calculate the atomic weight by multiplying the abundance of each isotope by its atomic mass and summing
these products. Because 75.78% = and 24.22% = , we have
Atomic weight = (0.7578)( amu) + (0.2422)( amu)
= amu amu
= amu
This answer makes sense: The atomic weight, which is actually the average atomic mass, is between the masses of the two isotopes and is closer to the value of 35Cl, the more abundant isotope.
Practice Exercise
Three isotopes of silicon occur in nature: 28Si (92.23%), atomic mass amu; 29Si (4.68% ), atomic mass amu; and 30Si (3.09%), atomic mass amu. Calculate the atomic weight of silicon.
Answer: amu

26. Periodic Table
The periodic table is a systematic catalog of the elements.
Elements are arranged in order of atomic number.

27. Periodic Table The rows on the periodic chart are periods.
Columns are groups.
Elements in the same group have similar chemical properties.

28. Periodicity a repeating pattern of reactivities.
When one looks at the chemical properties of elements, one notices
a repeating pattern of reactivities.

29. Groups
These five groups are known by their names.

30. Nonmetals are on the right side of the periodic table (with the exception of H).
Metalloids border the stair-step line (with the exception of Al, Po, and At
Metals are on the left side of the chart.

31. Chemical Formulas
The subscript to the right of the symbol of an element tells the number of atoms of that element in one molecule of the compound.
Molecular compounds are composed of molecules and almost always contain only nonmetals.

32. Diatomic Molecules
These seven elements occur naturally as molecules containing two atoms:
Hydrogen
Nitrogen
Oxygen
Fluorine
Chlorine
Bromine
Iodine

33. Types of Formulas
Empirical formulas give the lowest whole-number ratio of atoms of each element in a compound.
Molecular formulas give the exact number of atoms of each element in a compound.
Structural formulas show the order in which atoms are bonded.
Perspective drawings also show the three-dimensional array of atoms in a compound.

34. H2O
molecular
empirical
H2O
C6H12O6
CH2O O3 O
N2H4
NH2

36. Ions When atoms lose or gain electrons, they become ions.
Cations are positive and are formed by elements on the left side of the periodic chart.
Anions are negative and are formed by elements on the right side of the periodic chart.

37. cation – ion with a positive charge
An ion is an atom, or group of atoms, that has a net positive or negative charge.
cation – ion with a positive charge
If a neutral atom loses one or more electrons
it becomes a cation. Na
11 protons
11 electrons
Na+
11 protons
10 electrons
anion – ion with a negative charge
If a neutral atom gains one or more electrons
it becomes an anion.
Cl-
17 protons
18 electrons Cl
17 protons
17 electrons

38. How many protons and electrons are in
Do You Understand Ions?
How many protons and electrons are in Al 27 13 ? 3+
13 protons, 10 (13 – 3) electrons
How many protons and electrons are in Se 78 34 2- ?
34 protons, 36 (34 + 2) electrons

39. Solution Practice Exercise
Sample Exercise 2.7 Writing Chemical Symbols for Ions
Give the chemical symbol, including superscript indicating mass number, for (a) the ion with 22 protons, 26 neutrons, and 19 electrons; (b) the ion of sulfur that has 16 neutrons and 18 electrons.
Solution
(a) The number of protons is the atomic number of the element. A periodic table or list of elements tells us that the element with atomic number 22 is titanium (Ti). The mass number (protons plus neutrons) of this isotope of titanium is = 48. Because the ion has three more protons than electrons, it has a net charge of 3+: 48Ti3+.
(b) The periodic table tells us that sulfur (S) has an atomic number of 16. Thus, each atom or ion of sulfur contains 16 protons. We are told that the ion also has 16 neutrons, meaning the mass number is = 32. Because the ion has 16 protons and 18 electrons, its net charge is 2– and the ion symbol is 32S2–.
In general, we will focus on the net charges of ions and ignore their mass numbers unless
the circumstances dictate that we specify a certain isotope.
Practice Exercise
How many protons, neutrons, and electrons does the 79Se2– ion possess?
Answer: 34 protons, 45 neutrons, and 36 electrons

40. Ionic Bonds
Ionic compounds (such as NaCl) are generally formed between metals and nonmetals.

41. Sample Exercise 2.9 Identifying Ionic and Molecular Compounds
Which of these compounds would you expect to be ionic: N2O, Na2O, CaCl2, SF4?
Solution
We predict that Na2O and CaCl2 are ionic compounds because they are composed of a metal combined with a nonmetal. We predict (correctly) that N2O and SF4 are molecular compounds because they are composed entirely of nonmetals.
Practice Exercise
Which of these compounds are molecular: CBr4, FeS, P4O6, PbF2?
Answer: CBr4 and P4O6 41

42. Writing Formulas
Because compounds are electrically neutral, one can determine the formula of a compound this way:
The charge on the cation becomes the subscript on the anion.
The charge on the anion becomes the subscript on the cation.
If these subscripts are not in the lowest whole-number ratio, divide them by the greatest common factor.

43. The ionic compound NaCl
ionic compounds consist of a combination of cations and an anions
they only have empirical formulas, no molecular formula
the sum of the charges on the cation(s) and anion(s) in each formula unit must equal zero
The ionic compound NaCl

44. Formula of Ionic Compounds
2 x +3 = +6
3 x -2 = -6
Al2O3
Al3+
O2-
1 x +2 = +2
2 x -1 = -2
CaBr2
Ca2+
Br-
1 x +2 = +2
1 x -2 = -2
Na2CO3
Na+
CO32-

45. Common Cations

46. Common Anions

47. Inorganic Nomenclature
Write the name of the cation.
If the anion is an element, change its ending to -ide; if the anion is a polyatomic ion, simply write the name of the polyatomic ion.
If the cation can have more than one possible charge, write the charge as a Roman numeral in parentheses.

48. Patterns in Oxyanion Nomenclature
When there are two oxyanions involving the same element:
The one with fewer oxygens ends in -ite.
The one with more oxygens ends in -ate.
NO2− : nitrite; SO32− : sulfite
NO3− : nitrate; SO42− : sulfate

49. Central atoms on the second row have bond to at most three oxygens; those on the third row take up to four.
Charges increase as you go from right to left.

50. The one with the second fewest oxygens ends in -ite.
ClO2− : chlorite
The one with the second most oxygens ends in -ate.
ClO3− : chlorate

51. The one with the fewest oxygens has the prefix hypo- and ends in -ite.
ClO− : hypochlorite
The one with the most oxygens has the prefix per- and ends in -ate.
ClO4− : perchlorate

52. 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:
• CO32– is the carbonate anion.
• HCO3– is the hydrogen carbonate
(or bicarbonate) anion.
• PO43– is the phosphate ion.
• H2PO4– is the dihydrogen phosphate anion.
HPO42– is the hydrogen phosphate anion.

53. Ionic Compounds (salts)
often a metal + nonmetal (or polyatomic ion)
anion (nonmetal), add “ide” to element name
BaCl2
barium chloride
K2O
potassium oxide
Mg(OH)2
magnesium hydroxide
KNO3
potassium nitrate

54. Transition metal ionic compounds
indicate charge on metal with Roman numerals
FeCl2
iron(II) chloride
2 Cl- -2 so Fe is +2
FeCl3
3 Cl- -3 so Fe is +3
iron(III) chloride
Cr2S3
3 S-2 -6 so Cr is +3 (6/2)
chromium(III) sulfide

55. Acid Nomenclature

56. If the anion in the acid ends in -ide, change the ending to -ic acid and add the prefix hydro- .
HCl: hydrochloric acid
HBr: hydrobromic acid
HI: hydroiodic acid
If the anion in the acid ends in -ate, change the ending to -ic acid.
HClO3: chloric acid
HClO4: perchloric acid

57. An acid can be defined as a substance that yields
hydrogen ions (H+) when dissolved in water.
HCl
Pure substance, hydrogen chloride
Dissolved in water (H+ Cl-), hydrochloric acid
An oxoacid is an acid that contains hydrogen, oxygen, and another element.
HNO3
nitric acid
H2CO3
carbonic acid
H2SO4
sulfuric acid

58. Solution Practice Exercise
Sample Exercise 2.14 Relating the Names and Formulas of Acids
Name the acids (a) HCN, (b) HNO3, (c) H2SO4, (d) H2SO3.
Solution
(a)The anion from which this acid is derived is CN–, the cyanide ion. Because this ion has an -ide ending, the acid is given a hydro- prefix and an -ic ending: hydrocyanic acid. Only water solutions of HCN are referred to as hydrocyanic acid. The pure compound, which is a gas under normal conditions, is called hydrogen cyanide. Both hydrocyanic acid and hydrogen cyanide are extremely toxic.
(b) Because NO3– is the nitrate ion, HNO3 is called nitric acid (the -ate ending of the anion is replaced with an -ic ending in naming the acid).
(c)Because SO42– is the sulfate ion, H2SO4 is called sulfuric acid.
(d) Because SO32– is the sulfite ion, H2SO3 is sulfurous acid (the -ite ending of the anion is replaced with an -ous ending).
Practice Exercise
Give the chemical formulas for (a) hydrobromic acid, (b) carbonic acid.
Answer: (a) HBr, (b) H2CO3 58

59. Nomenclature of Binary Compounds
The less electronegative atom is usually listed first.
A prefix is used to denote the number of atoms of each element in the compound (mono- is not used on the first element listed, however) .

60. Nomenclature of Binary Compounds
The ending on the more electronegative element is changed to -ide.
CO2: carbon dioxide
CCl4: carbon tetrachloride

61. Nomenclature of Binary Compounds
If the prefix ends with a or o and the name of the element begins with a vowel, the two successive vowels are often elided into one.
N2O5: dinitrogen pentoxide

62. Molecular Compounds HI hydrogen iodide NF3 nitrogen trifluoride SO2
sulfur dioxide
N2Cl4
dinitrogen tetrachloride
NO2
nitrogen dioxide
N2O
dinitrogen monoxide

63. Solution Practice Exercise
Sample Exercise 2.15 Relating the Names and Formulas of Binary Molecular Compounds
Name the compounds (a) SO2, (b) PCl5, (c) Cl2O3.
Solution
The compounds consist entirely of nonmetals, so they are molecular rather than ionic. Using the prefixes in Table 2.6,
we have (a) sulfur dioxide, (b) phosphorus pentachloride,
and (c) dichlorine trioxide.
Practice Exercise
Give the chemical formulas for (a) silicon tetrabromide,
(b) disulfur dichloride.
Answer: (a) SiBr4, (b) S2Cl2 63

64. Nomenclature of Organic Compounds
Organic chemistry is the study of carbon.
Organic chemistry has its own system of nomenclature.

65. The ending denotes the type of compound. An alcohol ends in -ol.
The simplest hydrocarbons (compounds containing only carbon and hydrogen) are alkanes.
The first part of the names just listed correspond to the number of carbons (meth- = 1, eth- = 2, prop- = 3, etc.).
When a hydrogen in an alkane is replaced with something else (a functional group, like -OH in the compounds above), the name is derived from the name of the alkane.
The ending denotes the type of compound.
An alcohol ends in -ol.