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Atomic Models Handout 5

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  Development of the Atomic Theory Democritus   460 BC - Greek philosopher proposes the existence of the atom   He pounded materials until he made them into smaller and smaller parts   He called them atoma which is Greek for “indivisible”.   His Theory:   All atoms:   Are small hard particles   Are made of a single material formed into different shapes and sizes   Are always moving, and they form different materials by joining together John Dalton   1803 - British chemist; elements combine in specific proportions to form compounds    Solid Sphere Model or Bowling Ball Model   Proposed by John Dalton His Theory:   All substances are made of atoms that cannot be created, divided, or destroyed.   Atoms join with other atoms to make new substances.   Atoms of the same element are exactly alike   Atoms of different elements are different in mass and size.    His ideas account for the law of conservation of mass  (atoms are neither created nor destroyed)    The law of constant composition  (elements combine in fixed ratios).   J.J. Thomson   1897 - English chemist and physicist; discovered 1 st  subatomic particles  Plum Pudding Model or Raisin Bun Model   Proposed by J.J. Thomson   His Theory:   Atoms contain negatively charged particles called electrons and positively charged matter.   Created a model to describe the atom as a sphere filled with positive matter with negative particles mixed in   Referred to it as the plum pudding model Ernest Rutherford   1912 - New Zealand physicist discovered the nucleus Nuclear Model   Proposed by Ernest Rutherford His Theory:   Small, dense, positively charged particle present in nucleus called a proton   Electrons travel around the nucleus, but their exact places cannot be described.   Niels Bohr   1913 - Danish physicist; discovered energy levels Bohr Model or Planetary Model   Proposed by Niels Bohr   His Theory:   Electrons travel around the nucleus in definite paths and fixed distances.   Electrons can jump from one level to a path in another level. rwin Shrodinger   1924 - Austrian physicist; developed the electron cloud model Electron Cloud Model Proposed by Erwin Schrodinger   His Theory:   The exact path of electrons cannot be predicted.   The region referred to as the electron cloud,   Is an area where electrons can likely be found. James Chadwick 1932 - English physicist; discovered neutrons His Theory:   Neutrons have no electrical charge.   Neutrons have a mass nearly equal to the mass of a proton.   Unit of measurement for subatomic particles is the atomic mass unit (amu).   Modern Theory of the Atom     Atoms are composed of three main subatomic particles: the electron, proton, and neutron.   Most of the mass of the atom is concentrated in the nucleus of the atom   The protons and neutrons are located within the nucleus, while the electrons exist outside of the nucleus.   In stable atoms, the number of protons is equal to the number of electrons.   The type of atom is determined by the number of protons it has.   The number of protons in an atom is equal to the atomic number.   The sum of the number of protons and neutrons in a particular atom is called the atomic mass.   Valence electrons are the outermost electrons. Isotopes and Radioisotopes     Atoms of the same element that have different numbers of neutrons are called isotopes.     Due to isotopes, mass #s are not round #s.     Li (6.9) is made up of both 6 Li and 7 Li.     Often, at least one isotope is unstable.     It breaks down, releasing radioactivity.     These types of isotopes are called radioisotopes   Q- Sometimes an isotope is written without its atomic number - e.g. 35 S (or S-35). Why? Q- Draw B-R diagrams for the two Li isotopes.  A- The atomic # of an element doesn’t change although the number of neutrons can vary, atoms have definite numbers of protons.   Law of Conservation of Mass In a chemical reaction, no change in mass takes place. The total mass of the products is equal to the total mass of the reactant. Antoine Lavoisier  , a brilliant French chemist, formulated this law by describing one of his experiments involving mercuric oxide. He placed a small amount of mercuric oxide, a red solid, inside a retort and sealed the vessel tightly. He weighed the system, and then subjected it to high temperature. During the heating, the red solid turned into a silvery liquid. This observation indicated that a chemical reaction took place. After which, the setup was cooled and then weighed. The weight of the system was found to be the same as before heating. Law of Definite Proportion:  A compound always contains the same constituent elements in a fixed or definite proportion by mass. If water samples coming from different sources are analyzed, all the samples will contain the same ratio by mass of hydrogen to oxygen. Law of Multiple Proportions: If two elements can combine to form more than one compound, the masses of one element that will combine with a fixed mass of the other element are in a ratio of small whole numbers. Dalton’s Atomic Theory   This theory was proposed by John Dalton, can be used to explain the laws of chemical change. This theory is based on the following set of postulates: 1. Elements are made up of very small particles known as atoms. 2. All the atoms of an element are identical in mass and size, and are different from the atoms of another element. Dalton used the different shapes or figures to represent different elements, as follows: 3. Compounds are composed of atoms of more than one element, combined in definite ratios with whole number values. 4. During a chemical reaction, atoms combine, separate, or rearrange. No atoms are created and no atoms disappear. Atomic number   = number of protons = number of electrons in a neutral atom Mass number   = number of protons + number of neutrons Isotopes    –  atoms of an element having the same atomic number but different mass number. The existence of isotopes was shown by mass spectroscopy experiments, wherein elements were found to be composed of several types of atoms, each with different masses. a. The atomic number identifies an element. The atoms of isotopes of an element have the same number of protons and electrons. b. The atoms of isotopes of an element differ in the number of neutrons. Ions  can be made up of only one atom ( monoatomic  ) or more than one type of atom (  polyatomic  ).  Monoatomic ions are named based on the element. a. For cations , the name of the element is unchanged. If an element can form two ions of different charges, the name, which is usually derived from its Latin name, is modified by the suffix  – ic    for the ion with the higher charge, and  – ous   for that with the lower charge. b. For anions , the name of the element is modified by the suffix  – ide .  Several anions are polyatomic and are named based on the atomic constituents and the suffix - ide . The most common examples are: a. OH -  –  hydroxide ion b. CN -  –  cyanide ion  A number of polyatomic anions containing oxygen atoms are named based on the root word of the central (or non-oxygen) atom and the suffix  – ate   for the one with more oxygen atoms and  – ite   for the one with less oxygen atom. a. NO 3 -    –  nitrate ion   b. NO 2 -    –  nitrite ion   c. SO 3   2-    –  sulfite ion   d. SO 4 2-    –  sulfate ion   e. PO 4 3-  –  phosphate ion Some anions have common names ending with the suffix  – ate . a. C 2 H 3 O 2 -    –  acetate ion   b. C 2 O 4 2-    –  oxalate ion Chemical Formula.  The formula consists of the symbols of the atoms making up the molecule. If there is more than one atom present, a numerical subscript is used. Examples are the following: a. O 2  –  oxygen gas b. H 2 O  –  water c. NaOH  –  sodium hydroxide (liquid Sosa) d. HCl  –  hydrochloric acid (muriatic acid) Molecular formula    –  gives the composition of the molecule, in terms of the actual number of atoms present. Examples are the following: i. C 6 H 12 O 6 ii. K 3 PO 4 iii. Na 2 C 2 O 4 Empirical formula    –  gives the composition of the molecule, in terms of the smallest ratio of the number of atoms present. Examples are the following: i. CH 2 O ii. NaCO 2 Binary compounds    –  made up of two elements. Ionic compounds    –  made up of a cation and an anion. They are named by giving the name of the cation first, followed by the name of the anion. Ex: NaI  –  sodium iodide MgCl2  –  magnesium chloride  Molecular compounds    –  made up of two non-metals. They are named by giving the name of the first nonmetal and then that of the second nonmetal modified by the ending - ide . Molecular compounds are usually gases. Ex: CO2  –  carbon dioxide SO3  –  sulfur trioxide Ternary Compounds    –  made up of three elements. The naming of ternary compounds follows the same rule as that of the binary ionic compound: the name of the cation is given first, followed by the name of the anion. Ex: BaCrO4  –  barium chromate K2SO4  –  potassium sulfate Acids    –  yield hydrogen ions in aqueous solutions. Binary acids    –  composed of hydrogen and another element, usually a nonmetal. The first part of the name starts with the prefix hydro- followed by the name of the element, modified by the ending  – ic  . The second part consists of the word ‘acid’. Name = hydro -  (root name of element) - ic    + acid Ex: HCl  –  hydrochloric acid H2S  –  hydrosulfuric acid Ternary acids    –  made up of hydrogen and an anion, usually containing oxygen. The first part of the name consists of the root word of the name of the element, modified by the ending  – ic  . The second part consists of the word ‘acid’. If there is another acid with the same atoms, the suffix  – ous   is used to denote the one with less number of atoms. Name = (root name of element) - ic    (or  – ous ) + acid Ex: H2SO3  –  sulfurous acid H3PO4  –  phosphoric acid
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