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

2. 2.1 Early Chemistry 3. Early Chemistry

The Greeks the first to try and explain chemical changes

• Proposed that all matter was composed of four fundamental substances.

• • Fire

• • Earth

• • Water

• • Air

4. Early Chemistry

Demokritos (also Greek) believed that matter was composed of small, indivisible particles.

• Atomos term used to describe these particles.

• Greeks were without the ability to test these hypothesis and no definite conclusion was reached.

Alchemy predominate pseudoscience that dominated the next 2000 years.

5. Modern Chemistry

The foundations were laid in the sixteenth century

The first chemist to perform quantitative experiments was Robert Boyle (1627-1691)

• Published relationship between pressure and volume

• • Boyles Law

6. 2.2 Fundamental Chemistry Laws 7. Fundamental Chemical Laws

Law of Conservation of Mass Mass is neither created nor destroyed in a chemical reaction.

• Verified by Lavoisier (1743 1794) through his quantitative studies of combustion.

• Lavoisier presented the most unified and complete knowledge of chemistry to date.

8. Fundamental Chemical Laws

Law of Definite Proportion a given compound always contains exactly the same proportion of elements by mass.

• Prousts (1754-1826) discovery was made through careful experiments regarding composition.

9. Fundamental Chemical Laws

Law of Multiple Proportions when two elements form a series of compounds, the ratios of the masses of the second element that combine with mass of the first element can always be reduced to small whole numbers.

• Dalton (1766-1844)discovered that carbon and oxygen form two different compounds that contain different relative amounts of carbon and oxygen.

10. Daltons Atomic Theory 11. Daltons Atomic Theory

Each element is made up of tiny particles called atoms.

The atoms of a given element are identical; the atoms of different elements are different in some fundamental way or ways.

12. Daltons Atomic Theory

Chemical compounds are formed when atoms of different elements combine with each other.A given compound always has the same relative numbers and types of atoms.

Chemical reactions involve reorganization of the atoms- changes in the way they are bound together.The atoms themselves are not changed in a chemical reaction.

13. Daltons Table of Atomic Masses

Dalton prepared the first table of atomic masses

• Many of the masses were later found to be wrong

• Dalton still provided an important step with table construction.

14. Gay-Lussac

Joseph Gay-Lussac (1778-1850) provided the keys to determining the absolute formulas for compounds by his experimental work.

• Performed experiments under the same conditions of temperature and pressure and determined the amount of gases that would react.

15. Amadeo Avogadro

Avogadros hypothesis at the same temperature and pressure, equal volumes of different gases contain the same number of particles.

• The volume of gas is determined by the number of molecules present and not the size of the individual particles.

• Avogadros number 6.02 x 10 23

16. 2.4 Early Experiments to Characterize the Atom 17. The Electron

Cathode-Ray Tube a high voltage applied to a tube produced a ray (cathode ray).The ray was produced at the negative electrode and was repelled by the negative pole of an applied electric current.Thomson (1898-1903) concluded that the ray was a stream of negatively charged particles.

18. Deflection of Cathode Rays by an Applied Electric Field 19. Video

http://www.youtube.com/watch?v=7YHwMWcxeX8

20. The Electron

Millikan discovered the Charge-to-mass ratio of an electron with his oil drop experiment.

• e/m = -1.76 x 10 8C/g

21. A Schematic Representation of the Apparatus Millikan Used to Determine the Charge on the Electron 22. The Electron

Plum Pudding Model Thomson postulated that atoms must have a cloud of positive charge in order to counter the random negatively charged electrons.

23. The Plum Pudding Model of the Atom 24. Radioactivity

Three types of radioactivity:

• Gamma ( ) rays is high-energy light

• Beta () particles high speed electron

• Alpha () particles Helium nucleus (+2 charge)

25. The Nuclear Atom

Rutherford tested Thomsons plum pudding model by sending a stream of particles through a thin sheet of metal foil.

• The alpha particles were scattered and reflected concluding that a dense center of the atom existed.

26. Rutherford's Experiment Ona -Particle Bombardment of Metal Foil 27. (a) Expected Results of the Metal Foil Experiment if Thomson's Model Were Correct(b) Actual Results 28. 2.5 The Modern View of Atomic Structure: Intro 29. Atomic Structure

Protons 1 amu (1.673 x 10 -27kg);+1 charge

Neutrons 1 amu ; no charge

Electrons (9.109 x 10 -31kg) ; -1 charge

30. Two Isotopes of Sodium 31. Writing Symbols for Atoms

Atomic number(number of protons) is written as a subscript.

Mass number ( total number of protons and neutrons) written as a superscript.

32. 2.6 Molecules and Ions 33. Molecules

Chemical Bonds - the forces that hold atoms together in compounds

Chemical formula symbols for the elements are used to indicate the types of atoms present and subscripts are used to indicate the relative numbers of atoms.

34. Molecules

Structural formula individual bonds are shown (indicated by lines) between element symbols.

Space- filling model shows the relative sized of the atoms as well as their relative orientation in the molecule.

Ball-and-stick model also used to represent molecules.

35. Space-Filling Model of Methane 36. Ball-and-Stick Model of Methane 37. Ball-and-Stick Models of the Ammonium Ion and the Nitrate Ion 38. The Structural Formula for Methane 39. Molecules

Covalent bonds sharing electrons to form amolecule .

Ionic bond a force of attraction between oppositely charged ions that form acompound.

40. Ions

Ion is an atom or group of atoms that has a net positive or negative charge.

• Anion ion with a negative charge.

• Cation ion with a positive charge.

Polyatomic Ion a compound consisting of many ions.

41. 2.7 An Introduction to the Periodic Table 42. Periodic Table

Most elements are metals.

• Efficient conduction so heat and electricity

• Malleability

• Ductility

• Often lustrous

• Tend to lose electrons

43. Periodic Table

Nonmetals the relative few appear on the upper right corner of the table. (right of the heavy line)

• Tend to gain electrons

• Bond together forming covalent bonds.

44. Periodic Table

Groups/Families (vertical columns in table)Have similar chemical properties due to their similar atomic structure.

• Alkali metals group 1A readily form +1 ions

• Alkaline earth metals group 2A readily form +2 ions

• Halogens group 7A all form diatomic molecules and react with metals to form salts readily form-1 ions.

• Noble gases all exist under normal conditions as monatomic and have little chemical reactivity.

45. Periodic Table

Periods horizontal rows that represent the energy levels of the elements.

46. The Periodic Table 47. 2.8 Naming Simple Compounds 48. Binary Ionic Compounds Type I

The cation is always named first and the anion second

A monatomic (one atom) cation takes its name from the name of the element.

• Na +is sodium

A monatomic anion is named by taking the root of the element name and adding ide.

• Cl -is chloride

49. Binary Ionic Examples

NaCl

KI

CaS

Li 3 N

CsBr

MgO

Sodium Chloride

Potassium Iodide

Calcium Sulfide

Lithium Nitride

Cesium Bromide

Magnesium Oxide

50. Binary Examples

H Hyrdrogen

• H -= Hydride

• LiH

• • Lithium Hydride

51. Common Monatomic Cations and Anions 52. Binary Ionic Compounds Type 2

The charge on the metal ion must be specified.

Roman numerals indicate the charge of the cation.

53. Binary (Type 2) Examples

CuCl

HgO

Fe 2 O 3

MnO 2

PbCl 2

Copper (I) chloride

Mercury (II) oxide

Iron (III) oxide

Manganese (IV) oxide

Lead (II) chloride

54. Common Cations and Anions 55. Common Type 2 Cations 56. Ionic Compounds: Polyatomic Ions

Oxyanions anions that contain an atom of a given element and different numbers of oxygen atoms.

57. Oxyanions

-ite the name of the one with the smaller number of oxygen atoms.

• SO 3- Sulfite

-atethe name of the one with the larger number of oxygen atoms

• SO 4- Sulfate

58. Oxyanions

When more than two oxyanions make up a series the following prefixes apply:

• • Hypo- (less than) the least amount of oxygen atoms

• • • Hypochorite(note-ite)

• • Per- (more than) the most amount of oxygen atoms

• • • Perchlorate (note ate)

59. Polyatomic Ions

Na 2 SO 4

KH 2 PO 4

CsClO 4

NaOCl

Fe(NO 3 ) 3

Sodium sulfate

Potassium dihydrogen phosphate

Cesium perchlorate

Sodium hypochlorite

Iron (III) nitrate

60. Common Polyatomic Ions 61. Binary Covalent Compounds

The first element in the formula is named first, using the full element name.

The second element is named as if it were an anion

Prefixes are used to denote the numbers of atoms present

The prefix mono- is never used for naming the first element.

62. Prefixes Used in Covalent Compounds Prefix Number Indicated mono 1 di 2 tri 3 tetra 4 penta 5 hexa 6 hepta 7 octa 8 nona 9 deca 10 63. Binary Covalent Compounds

N 2 O

NO 2

N 2 O 4

N 2 O 5

NO

Dinitrogen monoxide

Nitrogen dioxide

Dinitrogen tetroxide

Dinitrogen pentoxide

Nitrogen monoxide

64. A Flowchart for Naming Binary Compounds 65. Naming Chemical Compounds 66. Acids

Binary acids the acid is named with the prefix hydro- and the anion ends in ic.Add the name acid on the end.

67. Names of Acids* that Do Not Contain Oxygen 68. Acids with Oxygen

If the anion ends in ate, the suffix ic is added to the root.

If the anion has an ite, the suffix ous is added to the root.

Add the name acid to the end.

69. Acids

HClO 4

HClO 3

HClO 2

HClO

HCl

Perchloric acid

Chloric acid

Chlorous acid

Hypochlorous acid

Hydrochloric

70. Names of Some Oxygen-Containing Acids 71. Figure 2.25Naming Acids 72. Figure 2.7A Cathode-Ray Tube 73. Plant is Newly Discovered Source of Gold 74. Figure 2.1The Priestley Medal is the Highest Honor Given by the American Chemical Society 75. A Silicon Chip 76. Atomic Nucleus 77. Crystals of Copper(II) Sulfate 78. Various Chromium Compounds Dissolved in Water 79. Table 2.1The Mass and Charge of the Electron, Proton, and Neutron 80. Table 2.2The Symbols for the Elements That Are Based on the Original Names 81. Table 2.6Prefixes Used to Indicate Number in Chemical Names