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

2. Chapter 2: Topics Early history of chemistry Fundamental chemical laws
Dalton’s atomic theory
Early experiments to characterize the atom
The modern view of atomic structure
Molecules and ions
An introduction to the Periodic Table
Naming simple compounds

3. 2.1 The early history of chemistry
Greeks
Democritus and others - atomos
Alchemy
Robert Boyle- experimental definition of element.
Lavoisier- Father of modern chemistry.

4. Greeks Matter is composed of fire, earth, water and air
The Greek philosopher Democritus (460 B.C. – 370 B.C.) was among the first to suggest the existence of atoms (from the Greek word “atomos”)
He believed that atoms were indivisible and indestructible
His ideas did agree with later scientific theory, but did not explain chemical behavior, and was not based on the scientific method – but just philosophy

5. Alchemy Turning Cheep metals into gold
Alchemists discovered several elements and prepared mineral acids

6. 17th Century
Robert Boyle: First “chemist” to perform quantitative experiments
He published his first book: “The Skeptical Chemist” in 1661.
He talked about elements

7. 18th Century George Stahl: Phlogiston flows out of a burning material.
Joseph Priestley: Discovers oxygen gas, “dephlogisticated air, i.e., low in phlogistone”

8. 2.2 Fundamentals chemical Laws
Law of Conservation of Mass
Law of Definite Proportion
Law of Multiple

9. Law of Conservation of Mass
It was discovered by Antoine Lavoisier
It was the basis for development of chemistry in the 19th century
Mass is neither created nor destroyed
Combustion involves oxygen, not phlogiston

10. Law of Definite Proportion (Proust’s Law)
A given compound always contains exactly the same proportion of elements by mass.
Water is composed of 11.1% H and 88.9% O (w/w)

11. Law of Multiple Proportions
When two elements form a series of compounds, the ratios of the masses of the second element that combine with 1 gram of the first element can always be reduced to small whole numbers.
The ratio of the masses of oxygen that combine with 1g of H in H2O and H2O2 will be a small whole number (“2”).

12. Example Water, H2O has 8 g of oxygen per 1g of hydrogen.
Hydrogen peroxide, H2O2, has 16 g of oxygen per 1g of hydrogen.
16/8 = 2/1
Small whole number ratios.
This fact could be explained in terms of atoms

13. 2.3 Dalton’s Atomic Theory (1808)
Elements are made up of small particles called atoms
Atoms of each element are identical. Atoms of different elements are different.
Compounds are formed when atoms combine. Each compound has a always same type and relative number of atoms
Chemical reactions are rearrangement of atoms but atoms are never changed into atoms of other element. , or created or destroyed.

14. Gay-Lussac hypothesis (1809)
Provided basics to determining absolute formulas of compounds
Gay-Lussac- under the same conditions of temperature and pressure, compounds always react in whole number ratios by volume.
2volumes of H react with one volume of O to form 2volumes of gaseous water and

15. Avogadro’s Hypothesis (1811)
At the same temperature and pressure, equal volumes of different gases contain the same number of particles.
5 liters of oxygen
5 liters of nitrogen
Same number of particles!

16. If Avogadro's hypothesis is correct, Gay-Lussac’s can be interpreted as follows:
2 molecules of H react with 1 molecule of O molecules of H2O

17. 2.4 Early experiments to characterize the atom
Based on Dalton, Gay-Lussac, Avogadro and others, work started to identify the nature of the atom
What is an atom made of? How do atoms of various elements differ?

18. The electron
J. J. Thomson - postulated the existence of electrons using cathode ray tubes.
Ernest Rutherford - explained the nuclear atom, containing a dense nucleus with electrons traveling around the nucleus at a large distance.

19. - + Thomson’s Experiment Voltage source
When high voltage is applied to the tube a ray emanates from the cathode is called cathode ray.

20. Thomson’s Experiment
Voltage source - +

21. - + Thomson’s Experiment Voltage source
Passing an electric current makes a beam appear to move from the negative to the positive end.

22. Thomson’s Experiment
Voltage source
By adding an electric field

23. Thomson’s Experiment Voltage source + -
By adding an electric field, he found that the moving particles were negatively charged

24. Results of Thomson Experiment
Electrons are produced from electrodes made from various types of metals, all atoms must contain electrons.
Since atoms are electrically neutral, they must contain positively charged particles.
Thomson determined charge-to-mass ratio of an electron:
e/m = -1.76X108C/g

25. Thomson’s Model
Atom consisted of a diffuse cloud of positive charge with negative electrons embedded randomly
Atom was like plum pudding.
Thomson believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model.

26. Millikan’s Experiment
Atomizer
Oil droplets - +
Oil
Telescope

27. Millikan’s Experiment
X-rays
X-rays give some electrons a charge.

28. Millikan’s Experiment
From the mass of the drop and the charge on
the plates, the mass of an electron is calculated

29. Radioactivity Certain elements produce high energy radiation
Discovered by accident and was a result of spontaneous emission by uranium
Bequerel (1896) found that a piece of mineral containing uranium could produce an image on a photographic plate in the absence of light.
Three types of radiation were known:
alpha- helium nucleus (+2 charge, 7300 times that of the electron)
beta- high speed electron
gamma- high energy light

30. The nuclear atom Rutherford’s Experiment
Aimed at testing Thomson’s plum pudding model
Used uranium to produce alpha particles.
Alpha particles are directed at gold foil through hole in lead block.
Since the mass is evenly distributed in gold atoms alpha particles should go straight through.
Used gold foil because it could be made atoms thin.

31. Florescent
Screen
Lead block
Uranium
Gold Foil

32. What he expected

33. Because

34. Because, he thought the mass was evenly distributed in the atom.

35. What he got

36. How he explained it
Atom is mostly empty
Small dense, positive particle at
center.
Alpha particles are deflected by
it if they get close enough. +

37. Proof for nuclear atom +

38. Nuclear atom model According to Rutherford:
The atom consists of a dense center of positive charge (Nucleus) with electrons moving around it at distance that is large relative to the nuclear radius

39. 2.5 The modern view of an atomic structure: An introduction
The atom is mostly empty space.
Two regions
Nucleus- protons and neutrons.
It is characterized by small size and high density
Electron cloud- region where you might find an electron.
The chemistry of atom
Results mainly from electrons

40. A cross section of nuclear atom

41. Mass and charge of nuclear particles
Particle Mass (Kg) Charge
Electron 9.11X
Proton 1.67X
Neutron 1.67X None

42. Why atoms of different elements have different properties?
Atoms of different elements have different number of protons and electrons.
Number and arrangement of electrons around nucleus differ from one element to another.

43. Sub-atomic Particles
Z - atomic number = number of protons determines type of atom.
A - mass number = number of protons + neutrons.
Number of protons = number of electrons if atom is neutral.

44. Symbols A
Mass number X
Atomic number Z 23 Na 11
Na-23

45. Isotopes
Atoms of the same element (same atomic number) with different mass numbers
Atoms with the same number of protons, but different numbers of neutrons.
Isotopes of chlorine
35Cl 37Cl
chlorine chlorine – 37
Cl Cl-37

46. Two isotopes of sodium
Isotopes show almost identical chemical properties. Why?
They possess same number of electrons

47. 2.6 Molecules and ions Introduction to chemical bonding
The forces that hold atoms together are called chemical bonding
Covalent bonding - sharing electrons.
Collection of atoms by covalent bonding lead to molecules
Molecules can be represented by formulas
Chemical formula- Symbol relates number and type of atoms in a molecule.
Diatomic molecule: two atoms of same element are connected by a covalent bond.

48. Molecular formula and structural formula
Molecular formulas
give the actual numbers and types of atoms in a molecule.
Examples: H2O, CO2, CO, CH4, H2O2, O2, O3, and C2H4.
Structural formula: bonds are shown as lines H H H C C H H H

49. Representing Structure in Molecules
Accurately represents the angles at which molecules are attached.

50. Ions Atoms or groups of atoms with a charge.
Cations- positive ions - get by losing electrons(s).
Anions- negative ions - get by gaining electron(s).
Ionic bonding- Force of attraction between oppositely charged ions.
Ionic solids are called salts.

51. Formation of Cations

52. Formation of Anions

53. Examples of ions Cation: A positive ion Mg2+, NH4+
Anion: A negative ion
Cl-, SO42-
Polyatomic ions

54. 2.7 Introduction to the Periodic Table
Elements are classified by:
- properties
- atomic number
Groups (vertical)
1A = alkali metals
2A = alkaline earth metals
7A = halogens
8A = noble gases
Periods (horizontal)

55. Periodic Table
Groups /Families
Periods

56. Metals
Conductors
Lose electrons
Malleable and ductile

57. Nonmetals
Brittle
Gain electrons
Covalent bonds

58. Semi-metals or Metalloids

59. Alkali Metals

60. Alkaline Earth Metals

61. Halogens

62. Transition metals

63. Noble Gases

64. Inner Transition Metals
Lanthanides
Lanthanides

65. Select an element ( )
= Internet link

67. 2.8 Naming Simple Compounds
Binary compounds are composed of two electrons
Both ionic and covalent compounds will be considered

68. Binary Ionic Compounds(Type I)
1. Cation first, then anion
2. Monatomic cation takes its name from the name of the element
Ca2+ = calcium ion
3. Monatomic anion takes its name from the root of the element + -ide
Cl- = chloride
CaCl2 = calcium chloride

69. Name the following compounds
CsF
Calcium fluoride
AlCl3
Aluminum chloride
LiH
Lithium hydride
Calcium hydroxide
Ca(OH)2

70. Some Common Cations

71. Some Common Anions

72. Binary Ionic Compounds (Type II):
- metal forms more than one cation
- use Roman numeral in name
PbCl2
Pb2+ is cation
PbCl2 = lead (II) chloride
PbCl4 = lead (IV) chloride

73. Example FeCl3 Iron(III) chloride Ferric chloride FeCl2
Iron(II) chloride
Ferrous chloride

74. Example Write the name of the compound Fe2O3
Oxidation state of Fe = +3
Iron(III) oxide
Group 1A, Group 2A and Al3+ do not take Roman numerals
Silver, Ag forms more than one oxidation state but Roman numerals are not used.

75. Common Cations and Anions

76. Ionic compounds with polyatomic ions
Polyatomic ions are assigned special names that need to be memorized
Oxyanions: anions that contain an atom of a given element and different numbers of oxygen atoms
Polyatomic anions (with many atoms) containing oxygen end in -ate or -ite. (The one with more oxygen is called -ate.)
Examples: NO3- is nitrate, NO2- is nitrite.
(Exceptions: hydroxide (OH-), cyanide (CN-), peroxide (O22-).)

77. Examples:
NaNO3 Sodium nitrate
K2SO4 Potassium sulfate
Al(HCO3)3 Aluminum bicarbonate
or aluminum hydrogen
carbonate

78. Common Polyatomic Ions

79. Binary Covalent compounds (Type III):
Compounds between two nonmetals
Although they do not contain ions, they are named similarly to binary ionic compounds
- First element in the formula is named first
using full element name
- Second element is named as if it were an anion.
- Use prefixes to denote number of atoms
- Never use mono- for naming first element
P2O5 = diphosphorus pentoxide

81. Acids Substances that produce H+ ions when dissolved in water.
All acids begin with H.
Two types of acids:
Oxyacids
Non-oxyacids

82. Naming acids If the formula has oxygen in it
write the name of the anion, but change
ate to -ic acid
ite to -ous acid
Watch out for sulfuric and sulfurous
HNO3
HNO2
HC2H3O2

83. Naming acids If the acid doesn’t have oxygen add the prefix hydro-
change the suffix -ide to -ic acid
HCl
H2S
HCN