2. Early History of Chemistry
Greeks were the first to attempt to explain why chemical changes occur.
Alchemy dominated for 2000 years.
Several elements discovered.
Mineral acids prepared.
Robert Boyle was the first “chemist”.
Performed quantitative experiments.
Developed first experimental definition of an element.
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3. Three Important Laws Law of conservation of mass (Lavoisier):
Mass is neither created nor destroyed in a chemical reaction.
Law of definite proportion (Proust):
A given compound always contains exactly the same proportion of elements by mass.
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4. Three Important Laws (continued)
Law of multiple proportions (Dalton):
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.
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5. Dalton’s Atomic Theory (1808)
Each element is made up of tiny particles called atoms.
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6. Dalton’s Atomic Theory (continued)
The atoms of a given element are identical; the atoms of different elements are different in some fundamental way or ways.
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7. Dalton’s Atomic Theory (continued)
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.
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8. Dalton’s Atomic Theory (continued)
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.
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9. J. J. Thomson (1898—1903)
Postulated the existence of negatively charged particles, that we now call electrons, using cathode-ray tubes.
Determined the charge-to-mass ratio of an electron.
The atom must also contain positive particles that balance exactly the negative charge carried by electrons.
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10. Cathode-Ray Tube
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11. Robert Millikan (1909)
Performed experiments involving charged oil drops.
Determined the magnitude of the charge on a single electron.
Calculated the mass of the electron
(9.11 × kg).
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12. Three types of radioactive emission exist:
Henri Becquerel (1896)
Discovered radioactivity by observing the spontaneous emission of radiation by uranium.
Three types of radioactive emission exist:
Gamma rays (ϒ) – high energy light
Beta particles (β) – a high speed electron
Alpha particles (α) – a particle with a 2+ charge
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13. Ernest Rutherford (1911) Explained the nuclear atom.
The atom has a dense center of positive charge called the nucleus.
Electrons travel around the nucleus at a large distance relative to the nucleus.
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14. Electrons – found outside the nucleus; negatively charged.
The atom contains:
Electrons – found outside the nucleus; negatively charged.
Protons – found in the nucleus; positive charge equal in magnitude to the electron’s negative charge.
Neutrons – found in the nucleus; no charge; virtually same mass as a proton.
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15. The nucleus is:
Small compared with the overall size of the atom. The volume of nucleus (nuclear volume) is very small compared to the volume of the atom (atomic volume); Volume = space occupied.
Extremely dense; accounts for almost all of the atom’s mass. Density of the nucleus is much larger than that of the atom. This is because the masses of the nucleus and the atom are nearly the same but their volumes are very different. Since the atomic volume is much larger than the nuclear volume, the atomic density is much smaller.
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16. Nuclear Atom Viewed in Cross Section
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17. In nature most elements contain mixtures of isotopes.
Atoms with the same number of protons but different numbers of neutrons. So isotopes of the same element will also have different atomic masses.
Show almost identical chemical properties; chemistry of atom is due to its electrons (we will discuss this in chapter 7).
In nature most elements contain mixtures of isotopes.
For example hydrogen occurs in 3 isotopic forms; calcium occurs in 5 isotopic forms.
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18. Two Isotopes of Sodium
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19. Isotopes are identified by:
Atomic Number (Z) – number of protons. This is unique to each element irrespective of the isotopic form
Mass Number (A) – number of protons plus number of neutrons. This number is not unique to an element. Atoms of the same element can have different mass numbers (isotopes).
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20. A certain isotope X contains 23 protons and 28 neutrons.
EXERCISE!
A certain isotope X contains 23 protons and 28 neutrons.
What is the mass number of this isotope?
Identify the element.
Mass Number = 51 ( )
Vanadium; We get this answer from the proton count of 23.
Symbol: 5123V
The mass number is 51. Mass Number = # protons + # neutrons. Mass Number = = 51. The element is vanadium.
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21. Compounds are formed by chemical combinations of atoms
Compounds are formed by chemical combinations of atoms. They are either molecular compounds (combinations of nonmetals or metalloids & nonmetals) or they are ionic compounds (combinations of metals & nonmetals).
What holds atoms together in a compound are forces called chemical bonds.
In a molecular compounds these chemical bonds are called covalent bonds and in an ionic compound they are ionic bonds. (bond means attraction)
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22. Covalent Bonds
Covalent Bonds= attractive forces holding atoms together in molecular (covalent) compounds
Bonds form between atoms by sharing electrons. This bond is a sharing of electron pairs between atoms. (Discussed in detail in chapters 8, 9)
Resulting collection of atoms is called a molecule.

23. Ionic Bonds
Bonds formed due to force of attraction between oppositely charged ions = ionic bonds
Ion – atom or group of atoms that has a net positive or negative charge.
Ions may be monoatomic (consisting of a single atom) or polyatomic (containing more than 1 atom)
A Cation is a positively charged ion; A monoatomic caltion is formed when electron(s) is (are) lost.
An Anion is a negatively ion; A monoatomic anion is formed when electron (s) is (are) gained.
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24. Metal atoms tend to lose electrons to form monoatomic cations.
The charge of cation = the number of electrons lost
In reactions with metals, nonmetal atoms tend to gain electrons to form monoatomic anions.
The charge of the anion = the number of electrons gained.
All ionic compounds are charge neutral. They contain cations & anions. The charge due to the cations neutralizes the charge due to the anions

25. A certain isotope X+ contains 54 electrons and 78 neutrons.
EXERCISE!
A certain isotope X+ contains 54 electrons and 78 neutrons.
What is the mass number of this isotope?
133
The mass number is 133. The plus charge in X+ means that the ion has lost an electron, therefore the number of protons is 55 (54+1). The ion is Cs+ with a mass number of 133 (55+78).
Note: Use the red box animation to assist in explaining how to solve the problem.
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26. CONCEPT CHECK!
Which of the following statements regarding Dalton’s atomic theory are still believed to be true?
Elements are made of tiny particles called atoms.
All atoms of a given element are identical.
A given compound always has the same relative
numbers and types of atoms.
IV. Atoms are indestructible.
Statements I and III are true. Statement II is not true (due to isotopes and ions). Statement IV is not true (due to nuclear chemistry).
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27. The Periodic Table Metals vs. Nonmetals
Groups or Families – elements in the same vertical columns; have similar chemical properties
Periods – horizontal rows of elements
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28. The Periodic Table

29. Groups or Families
Table of common charges formed when creating ionic compounds.
Group or Family
Charge
Alkali Metals 1+
Alkaline Earth Metals 2+
Halogens 1–
Noble Gases
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30. Labelling of Groups in Periodic Table
IUPAC (International Union of Pure & Applied Chemistry) Scheme: Groups labelled from left to right from 1 to 18; Alkali metals = group 1; noble gases = group 18
North American Scheme: Number from I to VIII followed by letter A (tall groups), or letter B (short groups); Alkali metal = group IA; Noble Gases = Group VIII A.

31. Naming Compounds: The naming rules depend on the type of compound
Binary Compounds
Composed of two elements
Ionic and covalent (molecular) compounds included
Binary Ionic Compounds
Metal—nonmetal (FeO, Fe2O3)
Binary Covalent (Molecular) Compounds
Nonmetal—nonmetal (HCl, Br2O3) or metalloid-nonmetal (SiCl4)
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32. Binary Ionic Compounds (Type I)
1. The cation is always named first and the anion second. 2. A monatomic cation takes its name from the name of the parent element. 3. A monatomic anion is named by taking the root of the element name and adding –ide.
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33. Names & Formulas of Monoatomic Anions (Memorize)
H-: Hydride C-4: Carbide
F-: Fluoride
Cl-: Chloride
Br-: Bromide
I-: Iodide
O-2: Oxide
S-2: Sulfide
Se-2: Selenide
N-3: Nitride
P-3: Phosphide

34. Naming Cations
Type I Metal Cations: Cations with fixed charges: Name of cation = name of element. Example: Na+ = sodium ion; Ca+2 = calcium ion; Al+3 = Aluminum ion. Alkali metals, Alkaline earth metals, Al, Zn, Cd, Ag all form only one cation

35. Names & Formulas of Type I Metal Cations (memorize)
Alkali metal cations: Li+ = Lithium ion; Na+ = sodium ion; K+ = potassium ion; Rb+ = Rubidium ion; Cs+ = Cesium ion
Alkaline earth metal cations: Mg+2 = Magnesium ion; Ca+2 = Calcium ion; Sr+2 = strontium ion; Ba+2 = Barium ion
Ag+ = silver ion
Al+3 = Aluminum ion
Zn+2 = Zinc ion
Cd+2 = Cadmium ion

36. Naming Cations
Type II Metal Cations: Cations with variable charges: Name of cation = name of element(charge in roman numerals). An alternate naming is when the ending is changed to ous for the lower charged ion and the ending is changed to ic for the higher charged ion.
Most of the transition metals, post-transition metals form more than one cation that vary in their charge.
Examples: Cu+ = Copper(I) ion or cuprous ion;
Cu+2 = copper(II) ion or cupric ion; Fe+3 = Iron(III) ion or Ferric ion. Fe+2 = Iron(II) ion or Ferrous ion.

37. Names & Formulas of Type II Metal Cations (memorize)
Cu+ = Copper(I) ion or cuprous ion; Cu+2 = copper(II) ion or cupric ion; Fe+2 = Iron(II) ion or Ferrous ion.; Fe+3 = Iron(III) ion or Ferric ion. Sn+2 = Tin(II) ion or Stannous ion; Sn+4 = Tin(IV) ion or Stannic ion. Pb+2 = Lead(II) ion or Plumbous ion; Pb+4 = lead(IV) ion or Plumbic ion. Hg2+2 = Mercury(I) ion or Mercurous ion; Hg+2 = Mercury(II) ion or Mercuric ion. Au+ = Gold(I) ion or Aurous ion; Au+3 = Gold(III) ion or Auric ion.

38. Formulas of Ionic Compounds
The formula of an ionic compound is obtained by combining cation & anion in that ratio that will make it charge neutral.
If the magnitudes of the charges of the cation and anion are the same, the ratio of cation:anion in the formula = 1:1
If the magnitudes of charges of the cation and anion are different, swap the charges to obtain the ratio of cation:anion in the formula.
Formula of cation is written first followed by that of the anion;
Do not write the charges of the ions when you write the formula of the ionic compound

39. Binary Ionic Compounds (Type I) ratio = 1:1
Examples:
KCl Potassium chloride
AlN Aluminum Nitride
CaO Calcium oxide
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40. Binary Ionic Compound (Type I) ratio is not 1:1
Combining Ag+ & S-2 : Compound formed = Ag2S = silver sulfide
Combining Ag+ & P-3 : Compound formed = Ag3P = silver Phosphide
Name the compound: Ba3N2: Barium Nitride
Name the compound: Al2Se3: Aluminum Selenide

41. Binary Ionic Compounds (Type II)
Metals in these compounds form more than one type of positive ion.
Charge on the metal ion must be specified.
Roman numeral indicates the charge of the metal cation.
Transition metal cations usually require a Roman numeral.
Elements that form only one cation do not need to be identified by a roman numeral.
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42. Binary Ionic Compounds (Type II)
Examples:
CuBr Copper(I) bromide or cuprous bromide
CuBr2 Copper(II) bromide or cupric bromide
FeS Iron(II) sulfide or Ferrous sulfide
Fe2S3 Iron(III) sulfide or Ferric sulfide
PbO2 Lead(IV) oxide or Plumbic oxide
? Lead(II) oxide or Plumbous oxide
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43. Polyatomic Ions (memorize)
Only 1 polyatomic cation to memorize: NH4+: Ammonium ion
Memorize list of polyatomic anions given in the following slides
Examples of compounds containing polyatomic ions:
NaOH Sodium hydroxide
Mg(NO3)2 Magnesium nitrate
(NH4)2SO4 Ammonium sulfate
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44. Polyatomic Anions that are oxyanions: XOm-n; X= central atom
Names of oxyanions are obtained by changing the ending of the name of the central atom (X) to either ite or ate depending on the number of oxygen atoms (m) in the formula
If m= 1 or 2, the ending is always ite.
If m= 4, the ending is always ate
If m= 3, the ending may be ite or ate.

45. Partial list of Oxyanions (memorize)
CO3-2: Carbonate
NO3-: Nitrate; NO2-: Nitrite
SO4-2: Sulfate; SO3-2: Sulfite
PO4-3: Phosphate; PO3-3: Phosphite
MnO4-: Permanganate

46. A prefix hypo is included for the oxyanion with 1 O atom
Naming Oxyanions
If oxyanions with 1 , 2 , 3 & 4 O atoms exist for a central atom, then, the ones with 1 & 2 O atoms have names ending in ite and the ones with 3 & 4 O atoms have names ending in ate
A prefix hypo is included for the oxyanion with 1 O atom
A prefix per is included for the oxyanion with 4 O atoms.

47. Partial List of Oxyanions (memorize)
ClO-: hypochlorite; ClO2-: chlorite; ClO3-: chlorate; ClO4-: perchlorate
BrO-: hypobromite; BrO2-: bromite; BrO3-: bromate; BrO4-: perbromate
IO-: hypoiodite; IO2-: iodite; IO3-: iodate; IO4-: periodate

48. Remaining Polyatomic Anions (memorize)
CN-: Cyanide;
N3-: Azide; (do not confuse with N-3: nitride)
OH-: hydroxide;
HS-: hydrogen sulfide or bisulfide
O2-2: peroxide; (do not confuse with O-2: oxide)
C2H3O2- or CH3COO-: Acetate
S2O3-2: thiosulfate
HCO3-: bicarbonate or hydrogencarbonate
HSO4-: bisulfate or hydrogensulfate
HPO4-2: biphosphate or hydrogenphosphate
H2PO4-: dihydrogen phosphate

49. Write the formula for ferric cyanide: Fe(CN)3; If you have more than of a polyatomic ion in a formula, you need enclose the formula of the polyatomic ion within brackets.
Write the formula for Ammonium carbonate: (NH4)2CO3
Write the formula for Ammonium nitrate: NH4NO3
Formula = FeSO4; Name =? ; Answer: Ferrous sulfate or Iron(II) sulfate
Formula = Pb3(PO4)2; Name =?; Answer: Lead(II) phosphate or plumbous phosphate
Formula = AgHCO3; Name =?; Answer: Silver bicarbonate or silver hydrogencarbonate

50. Acids
Acids can be recognized by the hydrogen that appears first in the formula—HCl.
An acid is formed when H+ is combined with any anion to form a neutral molecule
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51. Acids
If the anion does not contain oxygen, the acid is named with the prefix hydro– and the ending –ic acid.
Examples:
HCl Hydrochloric acid
HCN Hydrocyanic acid
H2S Hydrosulfuric acid
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52. Acids If the anion does contain oxygen:
The ending –ic acid is added to the root name if the anion name ends in –ate.
Examples:
HNO3 Nitric acid
H2SO4 Sulfuric acid
HC2H3O2 Acetic acid
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53. Acids If the anion does contain oxygen:
The ending –ous acid is added to the root name if the anion name ends in –ite.
Examples:
HNO2 Nitrous acid
H2SO3 Sulfurous acid
HClO2 Chlorous acid
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54. I. Name the following acids:
HClO: hypochlorous acid
HClO2: chlorous acid
HClO3: chloric acid
HClO4: perchloric acid
II. Write the formulas for the following acids
Hydrobromic acid: HBr
Hypobromous acid:HBrO
Bromous acid: HBrO2
Bromic acid: HBrO3
Perbromic acid: HBrO4

55. Flowchart for Naming Acids
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56. Binary Covalent Compounds (Type III)
Formed between two nonmetals or metalloid & nonmetal.
1. The first element in the formula is named first, using the full element name.
2. The second element is named with the ending changed to ide.
3. Prefixes are used to denote the numbers of atoms present.
The prefix mono- is never used for naming the first element.
If the prefix ends in an a or o and the name of the element begins with a vowel, then the a or o in the prefix is omitted
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57. Prefixes Used to Indicate Number in Chemical Names
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58. Binary Covalent Compounds (Type III)
Examples:
CO2 Carbon dioxide
CO Carbon monoxide
SF6 Sulfur hexafluoride
N2O4 Dinitrogen tetroxide
P2O5 Diphosphorus pentoxide
P2S5 Diphosphorus pentasulfide
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59. Flowchart for Naming Binary Compounds
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60. Overall Strategy for Naming Chemical Compounds
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61. Which of the following compounds is named incorrectly?
EXERCISE!
Which of the following compounds is named incorrectly?
KNO3 potassium nitrate
TiO2 titanium(II) oxide
Sn(OH)4 tin(IV) hydroxide
PBr5 phosphorus pentabromide
CaCrO4 calcium chromate
The correct answer is “b”. The charge on oxygen is 2-. Since there are two oxygen atoms, the overall charge is 4-. Therefore, the charge on titanium must be 4+ (not 2+ as the Roman numeral indicates).
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