Presentation is loading. Please wait.

Presentation is loading. Please wait.

Chapter 2: Chemical Foundations The Early History of Chemistry Fundamental Chemical Laws Dalton’s Atomic Theory Early Characterizations of Atoms The Modern.

Similar presentations


Presentation on theme: "Chapter 2: Chemical Foundations The Early History of Chemistry Fundamental Chemical Laws Dalton’s Atomic Theory Early Characterizations of Atoms The Modern."— Presentation transcript:

1 Chapter 2: Chemical Foundations The Early History of Chemistry Fundamental Chemical Laws Dalton’s Atomic Theory Early Characterizations of Atoms The Modern View of Atomic Structure Molecules and Ions Introduction to the Periodic Table Naming Simple Compounds

2 History Greeks (400 BC): first to explain why chemical changes occur.  Democritus and Leucippus - atomos  Aristotle - elements. Alchemy Robert Boyle (1660): experimental definition of element & perform truly quantitative experiments. A. Lavoisier (1789): Father of modern chemistry.  He wrote the first modern chemistry book.

3 Laws Conservation of MassConservation of Mass  Antoine Lavoisier finally explained the true nature of combustion.

4 Laws Law of Definite ProportionLaw of Definite Proportion  Joseph Proust: compounds have a constant composition.  They react in specific ratios by mass.

5 Laws Multiple Proportions - When two elements form more than one compound, the ratios of the masses of the second element that combine with one gram of the first can be reduced to small whole numbers.Multiple Proportions - When two elements form more than one compound, the ratios of the masses of the second element that combine with one gram of the first can be reduced to small whole numbers. John Dalton John Dalton

6 What?! Water has 8 grams of oxygen per gram of hydrogen.Water has 8 grams of oxygen per gram of hydrogen. Hydrogen peroxide has 16 grams of oxygen per gram of hydrogen.Hydrogen peroxide has 16 grams of oxygen per gram of hydrogen. 16/8 = 2/116/8 = 2/1 Small whole number ratios.Small whole number ratios.

7 Dalton’s Atomic Theory 1)Elements are made up of atoms. 2)Atoms of each element are identical. Atoms of different elements are different.

8 Dalton’s Atomic Theory 3)Compounds are formed when atoms combine. Each compound has a specific number and kinds of atom. 4)Chemical reactions are rearrangement of atoms. Atoms are not created or destroyed.

9 Gay-Lussac - under the same conditions of temperature and pressure, compounds always react in whole number ratios by volume.Gay-Lussac - under the same conditions of temperature and pressure, compounds always react in whole number ratios by volume. Avagadro - interpreted that to mean at the same temperature and pressure, equal volumes of gas contain the same number of particles (called Avagadro’s Hypothesis).Avagadro - interpreted that to mean at the same temperature and pressure, equal volumes of gas contain the same number of particles (called Avagadro’s Hypothesis). Additional Observations

10 The Electron Streams of negatively charged particles were found to emanate from cathode tubes.Streams of negatively charged particles were found to emanate from cathode tubes. J. J. Thompson is credited with their discovery (1897).J. J. Thompson is credited with their discovery (1897).

11 The Electron Thompson measured the charge/mass ratio of the electron to be 1.76  10 8 coulombs/g.Thompson measured the charge/mass ratio of the electron to be 1.76  10 8 coulombs/g.

12 Millikan Oil Drop Experiment Once the charge/mass ratio of the electron was known, determination of either the charge or the mass of an electron would yield the other.

13 Millikan Oil Drop Experiment Robert Millikan (University of Chicago) determined the charge on the electron in 1909.

14 The Atom, circa 1900 “Plum pudding” model, put forward by Thompson.“Plum pudding” model, put forward by Thompson. Positive sphere of matter with negative electrons imbedded in it.Positive sphere of matter with negative electrons imbedded in it.

15 Discovery of the Nucleus Ernest Rutherford shot  particles at a thin sheet of gold foil and observed the pattern of scatter of the particles.

16 The Nuclear Atom Since some particles were deflected at large angles, Thompson’s model could not be correct.

17 The Nuclear Atom Rutherford postulated a very small, dense nucleus with the electrons around the outside of the atom.Rutherford postulated a very small, dense nucleus with the electrons around the outside of the atom. Most of the volume of the atom is empty space.Most of the volume of the atom is empty space.

18 Radioactivity The spontaneous emission of radiation by an atom.The spontaneous emission of radiation by an atom. First observed by Henri Becquerel.First observed by Henri Becquerel. Also studied by Marie and Pierre Curie.Also studied by Marie and Pierre Curie.

19 Radioactivity Three types of radiation were discovered by Ernest Rutherford:Three types of radiation were discovered by Ernest Rutherford:   particles   particles   rays

20 Other Subatomic Particles Protons were discovered by Rutherford in 1919.Protons were discovered by Rutherford in Neutrons were discovered by James Chadwick in 1932.Neutrons were discovered by James Chadwick in 1932.

21 Subatomic Particles Protons and electrons are the only particles that have a charge.Protons and electrons are the only particles that have a charge. Protons and neutrons have essentially the same mass.Protons and neutrons have essentially the same mass. The mass of an electron is so small we ignore it.The mass of an electron is so small we ignore it.

22 Symbols of Elements Elements are symbolized by one or two letters.

23 Atomic Number All atoms of the same element have the same number of protons: The atomic number (Z)

24 Atomic Mass The mass number (A) of an atom in atomic mass units (amu) is the total number of protons and neutrons in the atom.

25 Isotopes: Atoms of the same element with different masses.Atoms of the same element with different masses. Isotopes have different numbers of neutrons.Isotopes have different numbers of neutrons C 12 6 C 13 6 C 14 6 C

26 Atomic Mass Atomic and molecular masses can be measured with great accuracy with a mass spectrometer.

27 Average Mass Because in the real world we use large amounts of atoms and molecules, we use average masses in calculations.Because in the real world we use large amounts of atoms and molecules, we use average masses in calculations. Average mass is calculated from the isotopes of an element weighted by their relative abundances.Average mass is calculated from the isotopes of an element weighted by their relative abundances.

28 Periodic Table: A systematic catalog of elements.A systematic catalog of elements. Elements are arranged in order of atomic number.Elements are arranged in order of atomic number.

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

30 Periodic Table The rows on the periodic chart are periods.The rows on the periodic chart are periods. Columns are groups.Columns are groups. Elements in the same group have similar chemical properties.Elements in the same group have similar chemical properties.

31 Groups These five groups are known by their names.

32 Periodic Table Nonmetals are on the right side of the periodic table (with the exception of H).

33 Periodic Table Metalloids border the stair-step line (with the exception of Al and Po).

34 Periodic Table Metals are on the left side of the chart.

35 Chemical Formulas The subscript to the right of the symbol of an element tells the number of atoms of that element in one molecule of the compound.

36 Diatomic Molecules These seven elements occur naturally as molecules containing two atoms.

37 Types of Formulas Empirical formulas give the lowest whole-number ratio of atoms of each element in a compound.Empirical formulas give the lowest whole-number ratio of atoms of each element in a compound. Molecular formulas give the exact number of atoms of each element in a compound.Molecular formulas give the exact number of atoms of each element in a compound.

38 Types of Formulas Structural formulas show the order in which atoms are bonded.Structural formulas show the order in which atoms are bonded. Perspective drawings also show the three-dimensional array of atoms in a compound.Perspective drawings also show the three-dimensional array of atoms in a compound. Ball-and-stick modelBall-and-stick model Space-filling modelSpace-filling model

39 Ions When atoms lose or gain electrons, they become ions.When atoms lose or gain electrons, they become ions.  Cations are positive and are formed by elements on the left side of the periodic chart.  Anions are negative and are formed by elements on the right side of the periodic chart.

40 Ionic Bonds Ionic compounds (such as NaCl) are generally formed between metals and nonmetals.

41 Writing Formulas Because compounds are electrically neutral, one can determine the formula of a compound this way:Because compounds are electrically neutral, one can determine the formula of a compound this way:  The charge on the cation becomes the subscript on the anion.  The charge on the anion becomes the subscript on the cation.  If these subscripts are not in the lowest whole- number ratio, divide them by the greatest common factor.

42

43

44 Inorganic Nomenclature Write the name of the cation.Write the name of the cation. If the anion is an element, change its ending to -ide; if the anion is a polyatomic ion, simply write the name of the polyatomic ion.If the anion is an element, change its ending to -ide; if the anion is a polyatomic ion, simply write the name of the polyatomic ion. If the cation can have more than one possible charge, write the charge as a Roman numeral in parentheses.If the cation can have more than one possible charge, write the charge as a Roman numeral in parentheses.

45 The less electronegative atom is usually listed first.The less electronegative atom is usually listed first. A prefix is used to denote the number of atoms of each element in the compound (mono- is not used on the first element listed, however.)A prefix is used to denote the number of atoms of each element in the compound (mono- is not used on the first element listed, however.) Nomenclature of Binary Compounds

46 The ending on the more electronegative element is changed to -ide.The ending on the more electronegative element is changed to -ide.  CO 2 : carbon dioxide  CCl 4 : carbon tetrachloride Nomenclature of Binary Compounds

47 If the prefix ends with a or o and the name of the element begins with a vowel, the two successive vowels are often turned into one: N 2 O 5 : dinitrogenpentoxide

48 When there are two oxyanions involving the same element:When there are two oxyanions involving the same element:  The one with fewer oxygens ends in -ite NO 2 − : nitrite ; SO 3 2− : sulfiteNO 2 − : nitrite ; SO 3 2− : sulfite  The one with more oxygens ends in -ate NO 3 − : nitrate; SO 4 2− : sulfateNO 3 − : nitrate; SO 4 2− : sulfate Patterns in Oxyanion Nomenclature

49 The one with the second fewest oxygens ends in -ite  ClO 2 − : chlorite The one with the second most oxygens ends in -ate  ClO 3 − : chlorate

50 The one with the fewest oxygens has the prefix hypo- and ends in -iteThe one with the fewest oxygens has the prefix hypo- and ends in -ite  ClO − : hypochlorite The one with the most oxygens has the prefix per- and ends in -ateThe one with the most oxygens has the prefix per- and ends in -ate  ClO 4 − : perchlorate Patterns in Oxyanion Nomenclature

51 Acid Nomenclature If the anion in the acid ends in -ide, change the ending to -ic acid and add the prefix hydro- :If the anion in the acid ends in -ide, change the ending to -ic acid and add the prefix hydro- :  HCl: hydrochloric acid  HBr: hydrobromic acid  HI: hydroiodic acid

52 Acid Nomenclature If the anion in the acid ends in -ite, change the ending to -ous acid:If the anion in the acid ends in -ite, change the ending to -ous acid:  HClO: hypochlorous acid  HClO 2 : chlorous acid

53 Acid Nomenclature If the anion in the acid ends in -ate, change the ending to -ic acid:If the anion in the acid ends in -ate, change the ending to -ic acid:  HClO 3 : chloric acid  HClO 4 : perchloric acid

54

55

56 Chapter 2: Chemical Foundations The Early History of Chemistry Fundamental Chemical Laws Dalton’s Atomic Theory Early Characterizations of Atoms The Modern View of Atomic Structure Molecules and Ions Introduction to the Periodic Table Naming Simple Compounds


Download ppt "Chapter 2: Chemical Foundations The Early History of Chemistry Fundamental Chemical Laws Dalton’s Atomic Theory Early Characterizations of Atoms The Modern."

Similar presentations


Ads by Google