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Chapter 2 Atoms, Molecules, and Ions

Chapter 2 Atoms, Molecules, and Ions. History Lesson. Democritus (460-370 BC): indivisible particles called atoms Plato and Aristotle challenged this view believing that matter was continuous Newton (1642-1727 AD) proposed the idea of invisible particles in the air called atoms

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Chapter 2 Atoms, Molecules, and Ions

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  1. Chapter 2Atoms, Molecules, and Ions

  2. History Lesson • Democritus (460-370 BC): indivisible particles called atoms • Plato and Aristotle challenged this view believing that matter was continuous • Newton (1642-1727 AD) proposed the idea of invisible particles in the air called atoms • Antoine Lavoisier (1743-1794 AD) conducted experiments demonstrating mass of products = mass of reactants • John Dalton (1766-1844 AD) proposed a model of matter • Dalton’s Atomic Theory

  3. Dalton’s Atomic Theory Assumption: Matter is discontinuous! 1: Matter is composed of atoms. An atom is the smallest unit of an element that has all the properties of that element. 2: An element is composed entirely of one type of atom. Postulates

  4. Dalton’s Atomic Theory (cont’d) 3: A compound contains atoms of two or more different elements. The relative number of atoms of each element in a compound is always the same. 4: Atoms do not change identity in chemical reactions; only the way in which they are joined together changes.

  5. Law of Constant Composition • Law of constant composition: All samples of a pure substance contain the same elements in the same proportions by mass. • This observation follows from Dalton’s third postulate (the relative numbers of atoms are the same in the same compound).

  6. Law if Multiple Proportions • Law of multiple proportions: When the same elements form more than one compound, the masses of one element that combines with a fixed mass of a second element are in a ratio of small whole numbers. • This follows from the postulate that individual atoms enter into chemical combination.

  7. Law of Conservation of Mass • Law of Conservation of Mass: There is no detectable change in mass when a chemical reaction occurs. • Dalton’s fourth postulate accounts for this law. The atoms do not change mass or identity when a chemical reaction takes place.

  8. Atomic Composition and Structure • Experiments over many years showed that atoms are not simple particles, but are composed of the subatomic particles listed below: • Electrons • Protons • Neutrons

  9. Cathode Rays • The application of a high voltage across a partially evacuated tube produces cathode rays.

  10. Electrons • J. J. Thomson demonstrated that cathode rays were negatively charged by applying magnetic and electric fields to cathode rays. • Cathode rays are electrons, negatively charged particles that are one of the components of an atom.

  11. Millikan Oil Drop Experiment • Robert A. Millikan performed experiments that determined the charge of the electron as 1.60 x 10-19 coulombs.

  12. Mass of electron • Thompson, using Milikan’s data for the charge of an electron, determined the mass to charge ratio of an electron. • This allowed him to calculate the mass of an electron. • 9.11 x 10-31 kg (mass of electron)

  13. Scattering of Alpha Particles by Gold

  14. The Nuclear Model of the Atom • Rutherford concluded that the results of the scattering experiment required that atoms consist of: • a nucleus that is very small compared to the atom, has a high positive charge and contains most of the mass of the atom. • the remainder of the space in an atom contains enough electrons to give a neutral atom.

  15. Atomic View of Rutherford Experiment

  16. The Proton • Rutherford proposed that the hydrogen nucleus was a fundamental particle called the proton, which has apositive charge equal in magnitude to the negative charge of the electron. • Protons account for the charge on the nucleus of all atoms. • The mass of the proton (1.673 x 10-27 kg) is 1836 times that of the electron.

  17. The Neutron • The number of protons in a nucleus, as determined by its positive charge, accounts for half or less of the nuclear mass. • Scientists inferred there must be a massive, neutral particle also present in the nucleus. • This neutral particle is called the neutron; its mass is almost the same as that of the proton.

  18. Particles in the Atom

  19. Definitions • Atomic number (Z) is the number of protons in the nucleus of an atom. • Mass number (A) is the sum of the numbers of protons and neutrons in the nucleus. • The number of protons (the atomic number) determines the identity of the element; all H atoms contain 1 proton, all He atoms contain 2 protons, etc.

  20. Isotopes • Isotopes are atoms of one element whose nuclei contain different numbers of neutrons (same Z, different A). Isotopes of Hydrogen What about Dalton’s Postulate #2?

  21. Symbols of Isotopes • A symbol to identify a specific isotope is where A = mass number, Z = atomic number, and X is the one or two letter symbol of the element. • The three isotopes of hydrogen are:

  22. Symbols of Isotopes • Oxygen also has three isotopes, containing 8, 9, and 10 neutrons respectively. The symbols are: • Since the value of Z, and the symbol, both identify the element, Z is often omitted from the symbol:

  23. Example: Symbols of Atoms • Write the symbol for the isotopes with:(a) 15 protons and 16 neutrons.(b) 21 protons and 24 neutrons.

  24. Ions • In many chemical reactions, atoms gain or lose electrons, producing charged particles called ions. • A cation has a positive charge and forms when an atom loses one or more electrons. • An anion has a negative charge and forms when an atom gains one or more electrons.

  25. Symbols for Ions • The number of protons in the nucleus determines the symbol used for an ion. • The element’s symbol is followed by a superscript number and a sign that shows the charge on the ion in electron charge units. • If the ionic charge is one unit, the number is omitted, e.g. Na+ is the symbol for a sodium cation.

  26. Example: Symbols of Ions • Write the symbol for the ions that contain:(a) 9 protons, 10 neutrons, 10 electrons.(b) 19 protons, 20 neutrons, 18 electrons.

  27. Practice • Write the symbols for the particles containing:(a) 8 protons, 9 neutrons, 10 electrons(b) 13 protons, 14 neutrons, 13 electrons

  28. Test Your Skill • Write the symbols for the particles containing:(a) 8 protons, 9 neutrons, 10 electrons(b) 13 protons, 14 neutrons, 13 electrons Answer: (a) (b)

  29. Example: Components of Ions • Fill in the blanks.Symbol Atomic number ____Mass number ____Charge ____no. of protons ____no. of neutrons ____no. of electrons ____

  30. The Atomic Mass Unit (u) • A relative mass scale has been established to express the masses of atoms. • The atomic mass unit (u) is 1/12 the mass of one 12C atom. Experimentally to three significant digits: 1 u = 1.66 x 10-27 kg

  31. Masses of Atoms in u • The masses of both the proton and the neutron are approximately 1 u. • A 24Mg atom has a mass approximately twice that of the 12C atom, so its mass is 24 u. • A 4He atom has a mass approximately 1/3 that of the 12C atom, so its mass is 4 u.

  32. Atomic Mass and Mass Number • Factors other than the mass of the protons and neutrons affect the mass of atoms, so the actual mass of atoms are not whole numbers. (24Mg = 23.98504 u; 4He = 4.002603 u) • When the accurate atomic mass of an atom is rounded to a whole number, it equals the mass number.

  33. Natural Distribution of Isotopes • About 75% of the elements occur in nature as mixtures of isotopes. • Usually, the relative abundance of isotopes of an element is the same throughout nature. • In all natural samples of Li, 7.42% of the atoms are 6Li and the remaining 92.58% are 7Li.

  34. Atomic Masses of the Elements • Isotopic mass is the mass in u, of a particular isotope of an element. • Different isotopes of an element all react essentially the same, so a weighted average of isotopic masses can be used in calculations. • The atomic mass is the weighted average mass, of the naturally occurring element.atomic mass = fractionA x isotopic massA + fractionB x isotopic massB + . .

  35. Example: Calculating Atomic Mass • A mass spectrometer was used to determine that gallium is 60.11% 69Ga (isotopic mass = 68.9256 u) and 39.89% 71Ga (isotopic mass = 70.9247 u). Calculate the atomic mass of Ga.

  36. The Periodic Table • Proposed independently by Dimitri Mendeleev and Lothar Meyer. • Periodic table: arranges the elements in rows that place elements with similar properties in the same column. • Period: a horizontal row • Group: a column - contains chemically similar elements

  37. Atomic Number and Atomic Mass • The atomic number and atomic mass for each element is given on the periodic table. 38 Atomic number Sr Atomic mass 87.62

  38. Important Groups of Elements • Metal: a material that is shiny and is a good electrical conductor; metallic elements are on the center and left side of the periodic table. • Nonmetal: an element that is typically a nonconductor; nonmetals are in the top right part of the periodic table. • Metalloid: an element that has properties of both metals and nonmetals.

  39. Important Groups of Elements • Representative Elements: the elements in the A groups (1,2, 13-18). • Transition Metals: the elements in B groups (3-12). • Inner Transition Metals: the two rows of metals (lanthanides and actinides) set at the bottom of the periodic table.

  40. Important Groups of Elements • Alkali Metals: soft, reactive metals in group 1A. • Alkaline Earth Metals: elements in group 2A. • Halogens (salt formers): reactive nonmetals in group 7A. • Noble Gases: the stable, largely inert, gases in group 8A.

  41. Elements and Biology

  42. Molecules • A molecule is a combination of atoms joined so strongly that they behave as a single particle. • The simplest molecules are diatomic - they contain two atoms.

  43. Elements • If all the atoms in a molecule are the same, the substance is an element.

  44. Molecules • If two or more elements form a molecule, it is a molecular compound.

  45. Molecular Formulas • A molecular formula gives the number of every type of atom in the molecule. • The elements present in the molecule are identified by their symbols. • A subscript numberfollows each symbol, giving the number of atoms of that element present in the molecule; the subscript is omitted if only one atom of the element is present. • A structural formulashows how the atoms are connected in the molecule.

  46. Molecular Formulas

  47. Molecular Mass • The molecular mass is the sum of the atomic masses of all atoms present in the molecular formula, expressed in atomic mass units (u). • The diagram shows the strategy for calculating molecular mass.

  48. Example: Calculate Molecular Mass • One substance present in smog is dinitrogen tetroxide (N2O4). Calculate its molecular mass.

  49. Practice • What is the molecular mass of the fuel propane (C3H8 )?

  50. Ionic Compounds • An ionic compound is composed of cations and anions joined to form a neutral species. • Ionic compounds generally form from the combination of metals with nonmetals. • In ionic compounds each cation is surrounded by several anions and vice versa.

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