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Chapter 2 Atoms, Molecules, and Ions History of Atomic Theory Started with the Greeks and four elements (earth, air, water and fire) Democritus termed.

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Presentation on theme: "Chapter 2 Atoms, Molecules, and Ions History of Atomic Theory Started with the Greeks and four elements (earth, air, water and fire) Democritus termed."— Presentation transcript:

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2 Chapter 2 Atoms, Molecules, and Ions

3 History of Atomic Theory Started with the Greeks and four elements (earth, air, water and fire) Democritus termed “atomos” Aristotle defined elements Robert Boyle provided an experimental definition of elements Lavoisier is the Father of Modern Chemistry

4 Laws of Chemistry Law of Conservation of Mass – Mass cannot be created nor destroyed Law of Definite Proportions – Compounds have a definite composition – Compounds react in ratios by mass Law of Multiple Proportions – Elements combine in small whole number ratios – Water vs. hydrogen peroxide 2 H with 1 O vs. 2 H with 2 O

5 Dalton’s Atomic Theory 1. Each element is made up of atoms 2. Atoms of a given element are identical; atoms of different elements are different 3. Chemical compounds form when atoms combine. Given compounds always have the same relative numbers 4. Chemical reactions involve the rearrangement of atoms (atoms not changed)

6 Helpful Observations Gay-Lussac observed under the same conditions of temperature and pressure, compounds react in whole number ratios by volume Avogadro developed Avogadro’s hypothesis that at the same temperature and pressure, equal volumes of different gases contain the same number of particles

7 Experiments of the Atom J.J. Thomson used cathode ray tubes Thomson’s experiment Voltage source +-

8 Thomson’s experiment Voltage source Passing an electric current makes a beam appear to move from the negative to the positive end. +-

9 Thomson’s experiment Voltage source + - By adding an electric field, he found that the moving pieces were negative

10 Thomson’s Model Found the electron. Couldn’t find positive (for a while). Said the atom was like plum pudding. A bunch of positive stuff, with the electrons able to be removed.

11 Radioactivity Discovered by accident Bequerel Three types alpha- helium nucleus (+2 charge, large mass) Was used in early experiments about the atom beta- high speed electron gamma- high energy light

12 Rutherford’s Experiment Used alpha particles to shoot at gold foil Hypothesized alpha particles should go straight through the foil Found particles were directed in many directions, including back in the original direction Used gold foil because it could be made atoms thin

13 Lead block Uranium Gold Foil Florescent Screen Rutherford’s Experiment

14 What he expected

15 Rutherford’s Experiment Because, he thought the mass was evenly distributed in the atom.

16 Rutherford’s Experiment What he got

17 Rutherford’s Experiment How Rutherford explained it – Atom is mostly empty. – Small dense, positive piece at center. – Alpha particles are deflected by it if they get close enough. – Stated his results could only be explained through a nuclear atom (dense ceter)

18 Modern Atomic View Two main components to an atom – Nucleus Protons are positively charged Neutrons have no charge Contains almost all the mass of an atom – Electron cloud Electrons are negatively charged Contains most of the space of an atom

19 Symbols for elements  Z - atomic number = number of protons determines type of atom.  A - mass number = number of protons + neutrons.  Number of protons = number of electrons if neutral. X A Z Na 23 11

20 Chemical bonds  The forces that hold atoms together.  Covalent bonds share electrons making molecules.  Chemical formula- the number and type of atoms in a molecule.  C 2 H carbon atoms, 6 hydrogen atoms,  Structural formula shows the connections, but not necessarily the shape.

21 Ions  Atoms or groups of atoms with a charge.  Cations- positive ions - get by losing electrons(s).  Anions- negative ions - get by gaining electron(s).  Ionic bonding- held together by the opposite charges.  Ionic solids are called salts.

22 Metals Conductors Lose electrons Malleable and ductile

23 Nonmetals Brittle Gain electrons Covalent bonds

24 Semi-metals or Metalloids

25 Naming Ionic Compounds 1.Cation is named first, anion named second 2.Monatomic cation takes name of the element 3.Monatomic anion has root of element name and -ide suffix 4.Polyatomic ions must be memorized (p. 62) 5.For example, sodium chloride and potassium iodide

26 Naming Ionic Compounds 1.Cations with more than one charge must have the charge designated 2.Use a Roman numeral to indicate the charge 3.Examples are: CuCl - copper (I) chloride PbS - lead (II) sulfide

27 Naming covalent compounds 1.Name the first element first 2.Name the second element as if it were an anion 3.Use prefixes to denote the number of each element (except mono on the first element) 4. Prefixes are on page 63

28 Naming acids 1.Binary acids use the prefix hydro and name the second as if an anion 2.With oxyacids, depends on the number of oxygen If anion ends in -ate, use the suffix -ic If anion ends in -ite, use the suffix -ous


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