1. Atoms, Molecules, and Ions
Chapter 2
Atoms, Molecules, and Ions
熊同銘

2. Contents
The Atomic Theory of Matter
The Discovery of Atomic Structure
The Modern View of Atomic Structure
Atomic Weights
The Periodic Table
Molecules and Molecular Compounds
Ions and Ionic Compounds
Naming Inorganic Compounds
Some Simple Organic Compounds

3. 1. The Atomic Theory of Matter Early Chemical Discoveries
Law of Conservation of Mass (Lavoisier 1774): The total mass of substances present after a chemical reaction is the same as the total mass of substances before the reaction.
Solution:

4. Law of Constant Composition (Proust 1799): All samples of a compound have the same composition-the same proportions by mass of the constituent elements. Example: Water always contains 88.81% oxygen (O) and hydrogen (H) (w/w)
Solution:

5. Dalton’s Atomic Theory (Dalton 1766-1844)
Law of Conservation of Mass
Law of Constant Composition

6. The law of multiple proportions (deduced by Dalton): If two elements form more than a single compound, the masses of one element combined with a fixed mass of the second are in the ratio of small whole numbers.

7. 2. The Discovery of Atomic Structure
Discovery of Subatomic Particles
In Dalton’s view, the atom was the smallest particle possible. Many discoveries led to the fact that the atom itself was made up of smaller particles.
Cathode rays and Electrons
Radioactivity
Nucleus: Protons and Neutrons

8. Electromagnetic Effect on charged particles
Electrostatics
Effect of a magnetic field
on charged particles

9. The Discovery of Electrons
Cathode rays: The negatively charged particles (electrons) emitted at the negative electrode (cathode) to anode in the passage of electricity through gases at very low pressures.
*The cathode rays do not depend on the composition of cathode, e.g., cathode rays from aluminum are same as from silver.
德國科學家提出的假說認為陰極射線是如光波之類的電磁波
英國科學家提出的假說認為陰極射線是可能是殘餘氣體經撞擊後形成的負價粒子
Characteristics of cathode rays:
Do not depend on the composition of cathode, e.g., cathode rays from aluminum are same as from silver.

10. e/m ratio for cathode rays by J.J. Thomson
Electric field and magnetic field were adjusted to the cathode undeflected, and e/m is calculated, that is x 108 C/g
Fluorescence: the radiation emission by a phosphor when it is struck by an energetic radiation.

11. Millikan’s oil-drop experiment
The magnitude of the charge on a droplet is an integral multiple of the electric charge,
Thus, e = x C
Therefore, mass of an electron= x g
e/m = x 108 C/g
e = x C
mass of an electron= x g

12. Thomson’s plum-pudding model of the atom
Rutherford and Marsden proved this model wrong.

13. The Discovery of Radioactivity
Radioactivity is the spontaneous emission of high-energy radiation by unstable atomic nuclei. Three types of radiation were discovered by Ernest Rutherford:
 particles (2+ charged, 7400 times heavier than electron)
 particles (1 charged, high-speed electrons)
 rays (uncharged)
Radioactivity is not a type of fluorescence

14. Nucleus: Protons and Neutrons
Rutherford’s -Sacttering experiment

15. Modern Nuclear Model of the Atom
Most of the mass and all of the positive charge of an atom are centered in a very small region called the nucleus (proton + neutron).
There are as many electrons outside the nucleus as there are units of positive charge on the nucleus. The atom as a whole is electrically neutral.

16. Proton
The positively charged subatomic particle found in the nucleus of an atom called proton.
Neutron: An electrically neutral subatomic particle found in the nucleus of an atom called neutron.
* Except for hydrogen, an atom does not have enough protons to account for the mass of the atom, e.g., mass of helium/mass of hydrogen is 4:1, not 2:1.
Subatomic particles-
illustrated by the
helium atom

17. Basic Forces Four basic forces are known in nature:
(1) gravitational force
(2) electromagnetic force
(3) strong nuclear force
(4) weak nuclear force
* Gravitational forces between atoms or between subatomic particles are so small that they are of no chemical significance.

18. 3. The Modern View of Atomic Structure The Structure of the Atom
1 Å (angstrom)
= 1 x 1010 m
Atoms have diameters between 1 – 5 Å (100 – 500 pm).
A cloud of rapidly moving electrons occupies most of the volume of the atom.
The nucleus occupies a tiny region at the center of the atom and is composed of the protons and neutrons.
The nucleus contains virtually all the mass of the atom.

19. Subatomic Particles Summary
* The quantity x C is called the electronic charge. For convenience, the charges of atomic and subatomic particles are usually expressed as multiples of this charge rather than in coulombs. Thus, the charge of an electron is 1- and that of a proton is 1+. Neutrons are electrically neutral (0).

20. Symbols of Elements
Elements are represented by a one or two letter symbol. This example is the symbol for carbon.
All atoms of the same element (symbol) have the same number of protons, which is called the atomic number. It is written as a subscript before the symbol.
The mass number is the total number of protons and neutrons in the nucleus of an atom. It is written as a superscript before the symbol.
For (neutral charge) atom, atomic number = number of protons = number of electrons

21. Isotopes
Isotopes are atoms of the same element with different mass number.
Isotopes have different numbers of neutrons, but the same number of protons.

22. 4. Atomic Weights Atomic mass unit (amu)
It is defined as exactly 1/12 the mass of a carbon-12 atom by assigning a mass of exactly 12 amu to a atom of the carbon-12 isotope.
1 amu = x 10−24 g
1 g = x1023 amu
* g/ x1023 = x 10−24 g
* For C-12, amu/atom as well as g/mole

23. Atomic Weight (average atomic mass) Atomic Weight =
Percent natural abundances: The relative proportions, expressed as percentage by number (of atoms), in which the isotopes of an element are found in natural sources.
Fractional abundance is the percent abundance divided by 100%. Thus, a 98.93% abundance is a fractional abundance.
Atomic Weight (average atomic mass)
Atomic Weight =
Ʃ [(isotope mass) × (fractional natural abundance)].

24. 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
Atomic weight = (0.7578) ( amu) + (0.2422)( amu)
= amu amu
= amu

25. Atomic Weight Measurement
Atomic and molecular weight can be measured with great accuracy using a mass spectrometer.

26. 5. The Periodic Table

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

28. Groups

29. Metals (metallic elements) are on the left side of the periodic table
Metals (metallic elements) are on the left side of the periodic table. Some properties of metals include
shiny luster
conducting heat and electricity.
solidity (except mercury).
Nonmetals (nonmetallic elements) are on the right side of the periodic table (with the exception of H). They can be solid (like carbon), liquid (like bromine), or gas (like neon) at room temperature.
Elements on the steplike line are metalloids. Their properties are sometimes like metals and sometimes like nonmetals.

30. 6. Molecules and Molecular Compounds
Notes
Noble-gas elements are normally found in nature as isolated atoms.
A molecule is a group of two or more atoms held together by covalent bonds.
Some elements exist in molecular forms. Diatomic molecules: H2, N2, O2, F2, Cl2, and I2. Example of polyatomic molecule: S8, P4.
Compound is a substance composed of two or more elements united chemically in definite proportions.
Molecular compound is compound that consists of molecules in which molecules are the smallest characteristic entities of molecular compound.
Most molecular substances contain only nonmetals.

31. Types of Chemical Formula Expression
Acetic acid for example:
a) Empirical formula: CH2O
d) Structural formula:
b) Molecular formula: C2H4O2
c) Condensed structural formula: CH3COOH
e) Ball-and-stick model
f) Space-filling models

32. Some examples of molecular models

33. Different representations of the methane (CH4) molecule

34. 7. Ions and Ionic Compounds
Notes
Ion: Electrically charged atom or group of atoms (polyatomic ion).
Cation: An ion with a positive charge, Example: Na+.
Anion: An ion with negative charge, Example Cl.
Monoatomic ion: The ions formed by losing or gaining electrons are removed from a atom. Example: Na+ and Cl.
Polyatomic ion: The ion consist of atoms joined as in a molecule, but carrying a net positive or negative charge. Example: NH4+ and SO42.

35. Continued
An ionic compound is formed when oppositely charged ions are attracted to each other in such a way as to neutralize the charges.
Mg3N2 for example:

36. Monatomic ions
Charge of metal ions:
- Group 1A (alkali metals): +1
- Group 2A (alkaline earth metals): +2
(Maximum positive charge: Group A #)
Charge of nonmetal ions:
- Group 6A (chalcogens): −2
- Group 7A (halogen): −1
(Maximum negative charge: 8 – Group A #)

37. Common monatomic ions and their charges

38. Formation of an Ionic Compound NaCl for example:
* Because there is no discrete “molecule” of ionic compounds, we are able to write only an empirical formula for ionic compounds.

39. 8. Naming Inorganic Compounds Names and Formulas of Ionic Compounds
Monatomic cation:
Na+ sodium ion
Ca2+ calcium ion
Fe2+ iron (II) ion ferrous ion
Fe3+ iron (III) ion ferric ion
Sn2+ tin (II) ion stannous ion
Sn4+ tin (IV) ion stannic ion
Cu+ copper (I) ion cuprous ion
Cu2+ copper (II) ion cupric ion
Hg22+ mercury (I) ion mercurous ion
Hg2+ mercury (II) ion mercuric ion
Polyatomic cation
NH4+ ammonium ion
H3O+ hydronium ion

41. Monatomic anion
Name the anion with the ending of –ide
Cl– chloride O2– oxide
A few polyatomic anion also with the ending of –ide
OH– hydroxide CN– cyanide
O22– peroxide
Oxoanion (oxyanion)
Polyatomic anions that contain oxygen
Less oxygen: suffixes is –ite
More oxygen: suffixes is –ate,
e.g. NO2−, nitrite
e.g. NO3−, nitrate

42. Extension of oxoanion
Prefix hypo–: one fewer oxygen than –ite oxoanion
Prefix per– : more oxygen than the –ate oxoanion
Example:
ClO– : hypoclorite ion (次氯酸根)
ClO2–: chlorite ion (亞氯酸根)
ClO3–: chlorate ion (氯酸根)
ClO4–: perchlorate ion (過氯酸根)
Anions derived by adding H+
Adding as a prefix the word hydrogen or dihydrogen
CO32- carbonate ion HCO3- hydrogen carbonate ion
PO43- phosphate ion H2PO4- dihydrogen phosphate ion

44. Naming Ionic Compounds
Naming the cation first, followed by the name of the anion Example:
BaH2: barium hydride
MgO: magnesium oxide
KBr potassium bromide
Li2S: Lithium sulfide
SnO: Stannous oxide tin (II) oxide
SnO2: stannic oxide tin (IV) oxide
Hg2Cl2: mercurous chloride mercury (I) chloride
HgCl2: mecuric chloride mercury (II) chloride
NH4OCOCH3 ammonium acetate
CuSO4 copper sulfate
Li2SO3: lithium sulfite
K2Cr2O7: potassium dichromate
Mg(ClO4)2: magnesium perchlorate
*Do not use prefix for the element in naming ionic compounds

45. Names and Formulas of Acids
Naming hydroacids
For hydroacid, that is relate to the molecule which containing “H” atom and suffix with –ide. Start with hydro- as prefix and -ic as the suffix, then add a acid.
Examples:
Molecular compound
Correspond acid name
hydrogen fluoride (HF)
hydrofluoric acid
hydrogen chloride (HCl)
hydrochloric acid
hydrogen bromide (HBr)
hydrobromic acid
hydrogen iodide (HI)
hydroiodic acid
hydrogen cyanide (HCN)
hydrocynanic acid
hydrogen sulfide (H2S)
hydrosulfuric acid

46. Naming oxoacids
Name of oxoacid, a ternary (three element) acids, is derived from correspond polyatomic anion portion which containing “O” atom.
Convert –ate to –ic, then adding acid
Convert –ite to –ous, then adding acid
Examples:
Anion name
Correspond acid name
phosphate (PO43–)
phosphoric acid (H3PO4)
carbonate (CO32–)
carbonic acid (H2CO3)
sulfate (SO42–)
sulfuric acid (H2SO4)
sulfite (SO32–)
sulfurous acid (H2SO3)
nitrate (NO3–)
nitric acid (HNO3)
nitrite (NO2–)
nitrous acid (HNO2)

47. Anion name Corresponding acid cyanide (CN–) hydrocyanic (HCN)
Extended Naming oxoacids
Anion name
Corresponding acid
hypochlorite (ClO–)
hypochlorous acid (HClO)次
chlorite (ClO2–)
chlorous acid (HClO2)亞
chlorate (ClO3–)
chloric acid (HClO3)
perchlorate (ClO4–)
perchloric acid (HClO4)過
Naming ternary acid other than oxoacids
Anion name
Corresponding acid
cyanide (CN–)
hydrocyanic (HCN)

48. Naming Bases
Similar to naming the ionic compounds, OH− named hydroxide.
Examples:
NaOH sodium hydroxide
Ba(OH)2 barium hydroxide

49. Names and Formulas of Binary Molecular Compounds
The element closer to the beginning of this path is generally written first in the formula of a binary molecular compound.

50. General Rules for Naming
Consists of two words, one for each element in the compound.
The first word is the name of the element that appears first in the formula.
The second word retains the stem of the second element name and replaces the ending by –ide.
Adding prefixes to denote the numbers of atoms of each element in the molecule.

51. Examples of binary molecular compound
Do not use “mono” for the first-named element, but do use for the second

52. Exceptions of few compounds containing “H”
B2H6 diborane
CH4 methane
SiH4 silane
NH3 ammonia
PH3 phosphine
H2O water

53. 9. Some Simple Organic Compounds
Carbon compounds containing one or more of the elements H, O, N, or S are common. Most organic compounds are molecular compounds.

54. Alkanes
Hydrocarbons: The molecules that contain only hydrogen and carbon atoms.
Alkanes: The saturated hydrocarbons (have the maximum number of hydrogen atoms possible for the number of carbon atoms). The general formula is CnH2n+2

55. Prefixes for Number of Carbon
Used for simple organic molecules. For alkane ending with “ane”
Examples:
methane
ethane
propane

56. Some Derivatives of Alkanes

57. Functional Groups
Functional group is an atom or group of atoms attached or inserted in a hydrocarbon chain or ring.
Functional group confers the characteristic properties to the molecule.
Examples of functional groups:
Alcohols: contain a hydroxyl group (-OH)
Carboxylic Acids: contain a carboxyl group (-COOH)
Haloalkanes: contain a halogen group (-X)
Ether: contain a ether group (R-O-R’)

58. Isomers
Compounds with the same molecular formula but different structural formulas.
Example:

61. End of Chapter 02