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The Components of Matter

The Components of Matter. Chapter 2. Chapter 2: The Components of Matter. 2.1 Elements, Compounds, and Mixtures: An Atomic Overview 2.2 The Observations That Led to an Atomic View of Matter 2.3 Dalton’s Atomic Theory 2.4 The Observations That Led to the Nuclear Atom Model

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The Components of Matter

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  1. The Components of Matter Chapter 2

  2. Chapter 2: The Components of Matter 2.1Elements, Compounds, and Mixtures: An Atomic Overview 2.2 The Observations That Led to an Atomic View of Matter 2.3 Dalton’s Atomic Theory 2.4 The Observations That Led to the Nuclear Atom Model 2.5 The Atomic Theory Today

  3. Chapter 2: The Components of Matter 2.6 Elements: A First Look at the Periodic Table

  4. Atoms, Molecules and Ions • Dalton’s Atomic Theory • Atomic Structure • Atomic Mass • Chemical Formulas • Chapters 2.1-2.5, 3.1 (Silberberg)

  5. Goals & Objectives • The student will understand the characteristics of the three types of matter: elements, compounds, and mixtures. (2.1) • The student will understand the significance of the three mass laws: conservation of mass, definite composition and multiple proportions. (2.2)

  6. Goals & Objectives • The student will understand the postulates of Dalton’s atomic theory and how it explains the mass laws. (2.3) • The student will understand the major contributions of experiments by Thomson, Millikan, and Rutherford concerning the atomic structure. (2.4)

  7. Goals & Objectives • The student will understand the structure of the atom along with the characteristics of the proton, neutron, and electron and the importance of isotopes. (2.5) • The student will understand the meaning and usefulness of the mole. (3.1)

  8. Goals & Objectives • The student will understand the relationship between the formula mass and molar mass. (3.1) • The student will understand the relationship among the amount of substance (in moles), mass (in grams), and the number of chemical atoms. (3.1)

  9. Goals & Objectives • The student will understand the information in a chemical formula. (3.1)

  10. Master these Skills • Distinguish between elements, compounds and mixtures at the atomic scale. (SP 2.1) • Using the mass ratio of an element to a compound to find the mass of an element in a compound. (SP 2.2) • Visualizing the mass laws. (SP 2.3)

  11. Master these Skills • Using atomic notation to express the subatomic makeup of an isotope. (SP 2.4) • Calculating an atomic mass from isotopic composition. (SP 2.5) • Calculating the molar mass of any substance. (3.1; SP 3.3 and 3.4)

  12. Master these Skills • Converting between amount of substance (in moles), mass (in grams), and the number of atoms. (SP 3.1-3.3)

  13. Matter • A scheme for the classification of matter

  14. Definitions for Components of Matter 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. Figure 2.1

  15. Definitions for Components of Matter Compound - a substance composed of two or more elements which are chemically combined. Figure 2.1 Mixture - a group of two or more elements and/or compounds that are physically intermingled.

  16. Table 2.1 Some Properties of Sodium, Chlorine, and Sodium Chloride.

  17. Figure 2.2 The law of mass conservation. The total mass of substances does not change during a chemical reaction.

  18. reactant 1 + reactant 2 product total mass total mass calcium oxide + carbon dioxide calcium carbonate CaO + CO2 CaCO3 Law of Mass Conservation The total mass of substances present does not change during a chemical reaction. = 56.08 g + 44.00 g 100.08 g

  19. Law of Definite (or Constant) Composition No matter the source, a particular compound is composed of the same elements in the same parts (fractions) by mass. Figure 2.3

  20. 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 20.0 g 1.00 part by mass 100% by mass Calcium carbonate Analysis by Mass (grams/20.0 g) Mass Fraction (parts/1.00 part) Percent by Mass (parts/100 parts)

  21. Law of Multiple Proportions If elements A and B react to form two compounds, the different masses of B that combine with a fixed mass of A can be expressed as a ratio of small whole numbers. 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

  22. 57.1 g O = = 1.33 42.9 g C g O 72.7 = = 2.66 g C 27.3 2.66 g O/g C in II 2 = 1.33 g O/g C in I 1 Assume that you have 100 g of each compound. In 100 g of each compound: g O = 57.1 g for oxide I & 72.7 g for oxide g C = 42.9 g for oxide I & 27.3 gfor oxide II For oxide I: For oxide II:

  23. Dalton postulated that: All matter consists of atoms; tiny indivisible particles of an element that cannot be created or destroyed. Atoms of one element cannot be converted into atoms of another element. Atoms of an element are identical in mass and other properties and are different from the atoms of any other element. Compounds result from the chemical combination of a specific ratio of atoms of different elements. Dalton’s AtomicTheory

  24. Dalton’sAtomicTheory explains the mass laws 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 the same atoms are present but in different combinations; therefore there is no mass change overall.

  25. Dalton’sAtomicTheory explains the mass laws Definite composition Atoms are combined in compounds in specific ratios postulate 3 and each atom has a specific mass. postulate 4 Each element constitutes a fixed fraction of the total mass in a compound.

  26. Dalton’sAtomicTheory explains the mass laws Multiple proportions Atoms of an element have the same mass postulate 3 and atoms are indivisible. postulate 1 When different numbers of atoms of elements combine, they must do so in ratios of small, whole numbers.

  27. The Structure of the Atom • Fundamental Particles in Atoms • ParticleMass(amu)Charge • electron 0.00054858 -1 • proton 1.0073 +1 • neutron 1.0087 0

  28. The Structure of the Atom • Electrons

  29. Figure 2.4 Observations that established the properties of cathode rays.

  30. Figure 2.5 Millikan’s oil-drop experiment for measuring an electron’s charge.

  31. determined by J.J. Thomson and others mass x charge charge = (-5.686x10-12 kg/C) x (-1.602x10-19 C) Millikan’s findings were used to calculate the mass on an electron. mass of electron = = 9.109x10-31 kg = 9.109x10-28 g

  32. Figure 2.6 Rutherford’s a-scattering experiment and discovery of the atomic nucleus.

  33. Figure 2.7 General features of the atom. The atom is an electrically neutral, spherical entity composed of a positively charged central nucleus surrounded by one or more negatively charged electrons. The atomic nucleus consists of protons and neutrons.

  34. Table 2.2 Properties of the Three Key Subatomic Particles * The coulomb (C) is the SI unit of charge. † The atomic mass unit (amu) equals 1.66054x10-24 g.

  35. 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.

  36. Atomic Symbol, Number and Mass Figure 2.8 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

  37. Isotopes Isotopes are atoms of an element with the same number of protons, but a different number of neutrons. Isotopes have the same atomic number, but a different mass number. Figure 2.8

  38. PROBLEM: Silicon (Si) has three naturally occurring isotopes: 28Si, 29Si, and 30Si. Determine the number of protons, neutrons, and electrons in each silicon isotope. PLAN: The mass number (A) is given for each isotope and is equal to the number of protons + neutrons. The atomic number Z, found on the periodic table, equals the number of protons. The number of neutrons = A – Z, and the number of electrons equals the number of protons for a neutral atom. Sample Problem 2.4 Determining the Number of Subatomic Particles in the Isotopes of an Element 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)

  39. Tools of the Laboratory Figure B2.1 Formation of a positively charged neon particle (Ne+).

  40. Tools of the Laboratory Figure B2.2 The mass spectrometer and its data.

  41. Isotopes and Neutrons • atoms of the same element having different atomic mass numbers

  42. The Complete Symbol • Competency III-1 • MZX where • M = atomic mass number(p + n) • Z = atomic number (p) • = electrons in a neutral atom • Number of neutrons = M - Z • 73Li p =___ e =____ n =____ • 7434Se p = ____ e = ____ n = ___

  43. Ions • Charged atoms obtained by adding or removing electrons to an atom. • Cations-positively charged ions • Anions-negatively charged ions • 3919K+ p = ___ e = ___ n = ___ • 3216S-2 p = ___ e = ___ n = ___

  44. Atomic Weights(Masses) • a relative scale based upon 126C having a mass of exactly 12 amu (atomic mass units) • The atomic weight of an element is the weighted average of all the isotopes of that element.

  45. PROBLEM: Silver (Ag, Z = 47) has two naturally occurring isotopes, 107Ag and 109Ag. From the mass spectrometric data provided, calculate the atomic mass of Ag. Isotope Mass (amu) Abundance (%) 107Ag 106.90509 51.84 109Ag 108.90476 48.16 PLAN: Find the weighted average of the isotopic masses. mass (g) of each isotope portion of atomic mass from each isotope atomic mass Sample Problem 2.5 Calculating the Atomic Mass of an Element multiply by fractional abundance of each isotope add isotopic portions

  46. Sample Problem 2.5 SOLUTION: mass portion from 107Ag = 106.90509 amu x 0.5184 = 55.42 amu mass portion from 109Ag = 108.90476amu x 0.4816 = 52.45amu atomic mass of Ag = 55.42amu + 52.45amu = 107.87amu

  47. Atomic Weights(Masses) • Naturally occurring copper consists of two isotopes. It is 69.1% 6329Cu which has a mass of 62.9 amu and 30.9% 6529Cu which has a mass of 64.9 amu. Determine the atomic weight of copper.

  48. Atomic Weights(Masses) • Naturally occurring chromium consists of four isotopes. It is 4.31% chromium-50 ,mass = 49.946amu, 83.76% chromium-52, mass = 51.941amu, 9.55% chromium 53, mass = 52.941amu 2.38% chromium 54, mass = 53.939amu. Determine the atomic mass of chromium.

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