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Early History of Chemistry

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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. Copyright © Cengage Learning. All rights reserved

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. Copyright © Cengage Learning. All rights reserved

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. Copyright © Cengage Learning. All rights reserved

5 Dalton’s Atomic Theory (1808)
Each element is made up of tiny particles called atoms. Copyright © Cengage Learning. All rights reserved

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. Copyright © Cengage Learning. All rights reserved

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. Copyright © Cengage Learning. All rights reserved

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. Copyright © Cengage Learning. All rights reserved

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. Copyright © Cengage Learning. All rights reserved

10 Cathode-Ray Tube Copyright © Cengage Learning. All rights reserved

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). Copyright © Cengage Learning. All rights reserved

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 Copyright © Cengage Learning. All rights reserved

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. Copyright © Cengage Learning. All rights reserved

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. Copyright © Cengage Learning. All rights reserved

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. Copyright © Cengage Learning. All rights reserved

16 Nuclear Atom Viewed in Cross Section
Copyright © Cengage Learning. All rights reserved

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. Copyright © Cengage Learning. All rights reserved

18 Two Isotopes of Sodium Copyright © Cengage Learning. All rights reserved

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). Copyright © Cengage Learning. All rights reserved

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. Copyright © Cengage Learning. All rights reserved

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) Copyright © Cengage Learning. All rights reserved

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. Copyright © Cengage Learning. All rights reserved

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. Copyright © Cengage Learning. All rights reserved

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). Copyright © Cengage Learning. All rights reserved

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 Copyright © Cengage Learning. All rights reserved

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 Copyright © Cengage Learning. All rights reserved

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) Copyright © Cengage Learning. All rights reserved

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. Copyright © Cengage Learning. All rights reserved

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 Copyright © Cengage Learning. All rights reserved

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. Copyright © Cengage Learning. All rights reserved

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 Copyright © Cengage Learning. All rights reserved

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 Copyright © Cengage Learning. All rights reserved

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 Copyright © Cengage Learning. All rights reserved

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 Copyright © Cengage Learning. All rights reserved

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 Copyright © Cengage Learning. All rights reserved

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 Copyright © Cengage Learning. All rights reserved

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 Copyright © Cengage Learning. All rights reserved

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 Copyright © Cengage Learning. All rights reserved

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). Copyright © Cengage Learning. All rights reserved


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