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The Components of Matter

The Components of Matter. Chapter 2. Overview of the Atomic Theory. The early civilizations had an idea of the composition of matter. Egyptian: “ kēme ” (earth) Al- kimia : “The art of transformation” History? --> No single, straight history. ORIGINS: Possibly the advent of fire and burning

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The Components of Matter

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  1. The Components of Matter Chapter 2

  2. Overview of the Atomic Theory

  3. The early civilizations had an idea of the composition of matter. • Egyptian: “kēme” (earth) • Al-kimia: “The art of transformation” • History? --> No single, straight history. • ORIGINS: Possibly the advent of fire and burning • Leading to pottery, metal-works (metallurgy), ancient structure-making

  4. The early civilizations had an idea of the composition of matter. Matter was seen as continuous, as the four Greek elements were. No distinct divisions between fire, water, earth, air Mixtures of the four gave the properties of being hot or cold, moist, or dry

  5. The early civilizations had an idea of the composition of matter. • LEUCIPPUS: “There must be tiny particles of water that could not be subdivided.” Observe the SAND. • DEMOCRITUS: Referred to these particles as atomos; Each atom was distinct in size and shape (eg. Water as round balls, Fire as sharp)

  6. The early civilizations had an idea of the composition of matter. • ARISTOTLE: Matter was continuous, not atomistic

  7. Metallurgy and Alchemy was also common practice in ancient civilizations. Metallurgy • Fire led to purification and material creation (eg. Glass and metals) • ALLOYS -> Bronze Age (ca.3000-1200BCE) • Alloys and lasting metals (eg. gold) were precious

  8. Metallurgy and Alchemy was also common practice in ancient civilizations. • Need: more stable materials like gold • Hypothesis: Perhaps there is a way of converting other materials to gold (philosopher’s stone) • MORE THAN GOLD: An early philosphical and spiritual discipline linking all things • “Solve et coagula”- “Separate and join together” • Al-kimia: the art of transformation • Aim to improve life: Elixir of life • Problems: Non-systematic, vague language and concepts, pseudoscience angle, fraudulence • 14th Century weakening Alchemy

  9. Finally in 1661, Robert Boyle re-worked the hypothesis of atoms. • “The ScepticalChymist or Chymico-Physical Doubts & Paradoxes” by Robert Boyle in 1661 • Thesis: “Matter consisted of atoms and clusters of atoms in motion and that every phenomenon was the result of collisions of particles in motion” • Appeal to EXPERIMENT! • Chemistry should stop being subservient to medicine and alchemy and establish itself as a separate science

  10. According to Boyle, substances are made of elements which are in turn composed of “simple bodies” = atom. Element - the simplest type of substance with unique physical and chemical properties. An element consists of only one type of atom. It cannot be broken down into any simpler substances by physical or chemical means. Molecule - a structure that consists of two or more atoms that are chemically bound together and thus behaves as an independent unit.

  11. According to Boyle, substances are made of elements which are in turn composed of “simple bodies” = atom. Compound - a substance composed of two or more elements which are chemically combined. Mixture - a group of two or more elements and/or compounds that are physically intermingled.

  12. According to Boyle, substances are made of elements which are in turn composed of “simple bodies” = atom.

  13. In the 18th century, several experiments led to the different LAWS we know of today. • Antoine Laurent Lavoisier • 1789- “When a chemical reaction is carried out in a closed system, the total mass of the system is not changed.” • Red mercuric oxide  Mercury + OXYGEN • 1st to use systematic names; 1stchem bk. ; “father” • Experiments with burning coal (combustion), and breathing guinea pigs (respiration). • LAW: Matter is neither created nor destroyed in a chemical change. The total mass of the reaction products is always equal to the total mass of the reactants We cannot create from nothing. Chemistry is about transformation.

  14. total mass total mass calcium oxide +carbon dioxide calcium carbonate CaO + CO2 CaCO3 In the 18th century, several experiments led to the different LAWS we know of today. Law of Mass Conservation: The total mass of substances does not change during a chemical reaction. reactant 1 + reactant 2 product = 56.08g + 44.00g 100.08g

  15. In the 18th century, several experiments led to the different LAWS we know of today. Law of Mass Conservation:

  16. In the 18th century, several experiments led to the different LAWS we know of today. • Joseph Louis Proust: Copper carbonate always had the same composition • LAW OF DEFINITE PROPORTIONS or CONSTANT COMPOSITION: A compound always contains the same elements in certain definite proportions and in no other combinations. • J.J.Berzelius: Prepared an extensive list of atomic weights; Lead sulfide experiments • Henry Cavendish: 1783; • Hydrogen gas + Oxygen gas  Water • 1800: Volta designed a powerful battery W.Nicholson and A.Carlisle would use to separate water into its elements.

  17. 8.0 g calcium 2.4 g carbon 9.6 g oxygen 0.40 calcium 0.12 carbon 0.48 oxygen 40% calcium 12% carbon 48% oxygen 1.00 part by mass 100% by mass 20.0 g In the 18th century, several experiments led to the different LAWS we know of today. Law of Definite Composition: No matter the source, a particular compound is composed of the same elements in the same parts (fractions) by mass. Analysis by Mass (grams/20.0g) Mass Fraction (parts/1.00 part) Percent by Mass (parts/100 parts)

  18. In the 18th century, several experiments led to the different LAWS we know of today. • Elements could combine in in more than one set of proportions. • If elements A and B react to form two different compounds, the masses of B combined with a fixed mass of A, can be expressed as a ratio of small whole numbers

  19. g O g O 72.7 57.1 g C g C = 27.3 42.9 = 1.33 2.66 g O/g C in II 2 = 1.33 g O/g C in I 1 = 2.66 In the 18th century, several experiments led to the different LAWS we know of today. Law of Multiple Proportions: Example: Carbon Oxides A & B Carbon Oxide I : 57.1% oxygen and 42.9% carbon Carbon Oxide II : 72.7% oxygen and 27.3% carbon Assume that you have 100g of each compound. In 100 g of each compound: gO = 57.1 g for oxide I & 72.7 g for oxide II gC = 42.9 g for oxide I & 27.3 gfor oxide II 2 =

  20. Finally in 1808, John Dalton presented the atomic theory of matter. Dalton’s Atomic THEORY: 1. All matter consists of atoms. 2. Atoms of one element cannot be converted into atoms of another element. 3. Atoms of an element are identical in mass and other properties and are different from atoms of any other element. 4. Compounds result from the chemical combination of a specific ratio of atoms of different elements.

  21. Dalton’s Atomic Theory explained the different mass laws that were established before. Mass conservation Atoms cannot be created or destroyed postulate 1 or converted into other types of atoms. postulate 2 Since every atom has a fixed mass, postulate 3 during a chemical reaction atoms are combined differently and therefore there is no mass change overall.

  22. Dalton’s Atomic Theory explained the different mass laws that were established before. Definite composition Atoms are combined in compounds in specific ratios postulate 3 postulate 4 and each atom has a specific mass. So each element has a fixed fraction of the total mass in a compound.

  23. Dalton’s Atomic Theory explained the different mass laws that were established before. Multiple proportions Atoms of an element have the same mass postulate 3 and atoms are indivisible. postulate 1 So when different numbers of atoms of elements combine, they must do so in ratios of small, whole numbers.

  24. However, Dalton’s Atomic theory did not fully explain why atoms bond as they do or how charged particles arise in experiments.

  25. In the 19th century Cathode rays were studied to learn more about electricity. Phosphor coated plate detects position of the CATHODE RAY Anode Cathode + - Evacuated tube CATHODE RAY! Power Supply

  26. In the 19th century Cathode rays were studied to learn more about electricity. N Anode S Cathode + - 1. Magnetic Field bends the cathode ray!

  27. In the 19th century Cathode rays were studied to learn more about electricity. Anode + Cathode + - - 2. In Electric field, ray bends toward the (+) plate

  28. In the 19th century Cathode rays were studied to learn more about electricity. Anode Cathode + - 3. Changing the metal in the cathode doesn’t matter.

  29. In the 19th century Cathode rays were studied to learn more about electricity. 1. Ray bends in magnetic field. 2. Ray bends towards positive plate in electric field. consists of charged particles consists of negative particles 3. Ray is identical for any cathode. particles found in all matter

  30. In 1897, J. J. Thompson measured the mass-to-charge ratio of this cathode ray particle… MASS OF PARTICLE <<<<<< MASS OF HYDROGEN THERE IS SOMETHING SMALLER THAN THE SMALLEST ELEMENT?!?! O.o >_< x_x Atoms are divisible! ^___^

  31. In 1909, Robert Millikan measured the charge and mass of this small particle (later called electrons)

  32. determined by J.J. Thomson and others mass charge mass of electron = X charge In 1909, Robert Millikan measured the charge and mass of this small particle (later called electrons) = (-5.686x10-12 kg/C) X (-1.602x10-19C) = 9.109x10-31kg = 9.109x10-28g

  33. If matter is electrically neutral, atoms must have a positively charged part to counter act the electrons…

  34. Before the turn of the century, Becquerel discovered radioactivity, (a term which was coined by Marie Sklodowska) the spontaneous emission of radiation from unstable elements. • MARIE SKLODOWSKA married PIERRE CURIE, a French Physicist and discover radioactive polonium and radium • 1903 Nobel in Physics (Becquerel, Curie, Curie) • Marie Curie: 2nd Nobel prize in 1911

  35. Ernest Rutherford named the radioactive emissions as alpha (+), beta (-) and gamma

  36. In 1910, Rutherford (and his colleagues Hans Geiger and Ernest Marsden) discovered that the positive part of the atom is not dispersed but rather concentrated at the center. • atoms positive charge is concentrated in the nucleus • proton (p) has opposite (+) charge of electron (-) • mass of p is 1840 x mass of e- (1.67 x 10-24g)

  37. In 1910, Rutherford (and his colleagues Hans Geiger and Ernest Marsden) discovered that the positive part of the atom is not dispersed but rather concentrated at the center. • RUTHERFORD: The smallest positive-ray particle is the unit of positive charge in the nucleus. This is the PROTON, with a charge equal in magnitude to a an electron, and nearly the same mass as a hydrogen atom

  38. In 1932, James Chadwick measured the mass of He vs. mass of H. JAMES CHADWICK and the NEUTRON H atoms - 1 p; He atoms - 2 p mass He/mass H should = 2 measured mass He/mass H = 4 neutron (n) is neutral (charge = 0) n mass ~p mass = 1.67 x 10-24g

  39. Revisiting the model: atomic radius ~ 100 pm = 1 x 10-10m nuclear radius ~ 5 x 10-3 pm = 5 x 10-15m If the atom were the blue eagle gym, the nucleus would be a microscopic speck of dust in the center, but weighing millions of tons.

  40. Location in the Atom Relative (amu)† Name (Symbol) Relative Absolute (g) Absolute (C)* Proton (p+) +1.60218 x 10-19 1.00727 1.67262 x 10-24 Nucleus 1+ 0 0 1.00866 1.67493 x 10-24 Nucleus Neutron (n0) Outside Nucleus Electron (e-) -1.60218 x 10-19 9.10939 x 10-28 0.00054858 1- Properties of the Three Key Subatomic Particles Table 2.2 Charge Mass * The coulomb (C) is the SI unit of charge. † The atomic mass unit (amu) equals 1.66054 x 10-24 g.

  41. Atomic Symbols, Isotopes, Numbers A X The Symbol of the Atom or Isotope Z X = atomic symbol of the element A = mass number; A = Z + N Z = atomic number (the number of protons in the nucleus) N = number of neutrons in the nucleus Isotopes = atoms of an element with the same number of protons, but a different number of neutrons Figure 2.8

  42. PROBLEM: Silicon (Si) is essential to the computer industry as a major component of semiconductor chips. It has three naturally occurring isotopes: 28Si, 29Si, and 30Si. Determine the number of protons, neutrons, and electrons in each silicon isotope. Sample Problem 2.4 Determining the Number of Subatomic Particles in the Isotopes of an Element PLAN: We have to use the atomic number and atomic masses. SOLUTION: The atomic number of silicon is 14. Therefore, 28Si has 14p+, 14e- and 14n0 (28 - 14) 29Si has 14p+, 14e- and 15n0 (29 - 14) 30Si has 14p+, 14e- and 16n0 (30 - 14)

  43. PROBLEM: Silver (Ag: Z = 47) has 46 known isotopes, but only two occur naturally, 107Ag and 109Ag. Given the following mass spectrometric data, calculate the atomic mass of Ag: Isotope Mass (amu) Abundance (%) 107Ag 106.90509 51.84 109Ag 108.90476 48.16 Calculating the Atomic Mass of an Element Sample Problem 2.5 PLAN: We have to find the weighted average of the isotopic masses, so we multiply each isotopic mass by its fractional abundance and then sum those isotopic portions.

  44. Calculating the Atomic Mass of an Element Sample Problem 2.5 continued SOLUTION: mass portion from 107Ag = 106.90509 amu x 0.5184 = 55.42 amu mass portion from 109Ag = 108.90476 amu x 0.4816 = 52.45 amu atomic mass of Ag = 55.42 amu + 52.45 amu = 107.87 amu

  45. The Modern Reassessment of the Atomic Theory 1. All matter is composed of atoms.The atom is the smallest body thatretains the unique identityof the element. 2. Atoms of one element cannot be converted into atoms of another element in a chemical reaction. Elements can only be converted into other elements in nuclear reactions. 3. All atoms of an element have the same number of protons and electrons, which determines the chemical behavior of the element.Isotopes of an element differ in the number of neutrons, and thus in mass number. A sample of the element is treated as though its atoms have anaverage mass. 4. Compounds are formed by the chemical combination of two or more elements in specific ratios.

  46. As instruments became sophisticated, more and more elements were discovered. There came a need to organize elements.

  47. Compounds: Introduction to Bonding

  48. Electrons of atoms interact to form chemical bonds  COMPOUNDS. Electrons of atoms can interact in two ways: An atom can transfer an electron to another atom forming ions. Oppositely charged ions are attracted to each other forming Ionic Compounds Cl- Na+ • Cations • positive charges • Usually form from metals that lose electrons • Anions • Negative charges • Usually form from nonmetals that gain electrons

  49. Electrons of atoms interact to form chemical bonds  COMPOUNDS. Electrons of atoms can interact in two ways: The resulting ions aggregate (due to electrostatic forces of attraction) and form a regular pattern of stable solids.

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