Chapter 2 Atoms, Molecules, Ions

1. Chapter 2Atoms, Molecules, Ions HW: 4 11 23 25 31 35 39 45 50 53 55 59 61 63 65 67 69 71 102

2. 2.1 – Atomic Theory Dalton’s Atomic Theory • Elements are made of extremely small particles called atoms • Atoms of an element are identical • Atoms are not created or destroyed • Compounds are combinations of atoms (1766-1844)

3. 2.1 – Atomic Theory Law of Definite Proportions (ConstantComposition) -Joseph Proust -A compound always has the same proportion of its elements Law of Conservation of Mass (Matter) = Matter cannot be created or destroyed Law of Multiple Proportions -Two or more compounds with the same elements must have different proportions of the elements -Example: Water vs. Hydrogen peroxide

4. 2.2 – Discovery of Atomic Structure Atom = Basic unit of an element that can enter into a chemical reaction Subatomic Particle = Particles that make up an atom (protons, neutrons, electrons are as small as we will study)

5. 2.2 – Discovery of Atomic Structure -Cathode (Ray) Tube / Crooke’s Tube / = Glass tube w/2 metal plates. Connected to high voltage source. Emits ray. Crooke – Determined that the ray was made of negative particles

6. 2.2 – Discovery of Atomic Structure JJ Thomson = 1897 – Credited w/ finding electrons Determined charge to mass ratio to be -1.76  108 coulombs/g. Millikan = 1909 - Performed the Oil Drop Experiment. Found the charge of an electron (-1.60 x 10-19C). Could then calculate mass of the electron

7. 2.2 – Discovery of Atomic Structure Radioactivity = Spontaneous emission of radiation a = + charge = Helium nucleus = High mass b = - charge (high speed electrons) = Low mass g = No charge. High energy = No mass

8. 2.2 – Discovery of Atomic Structure • 1900 - “Plum pudding” model, put forward by JJ Thompson. • Positive sphere of matter with negative electrons imbedded in it.

9. 2.2 – Discovery of Atomic Structure Protons = + charge particles in the nucleus Rutherford = 1919- Gold Foil Experiment -Most alpha particles when through, some deflected

10. 2.2 – Discovery of Atomic Structure This proved that Thomson’s model was incorrect Results of Gold Foil Experiment: Rutherford postulated a very small, dense nucleus with the electrons around the outside of the atom. -Most of the volume of the atom is empty space.

11. 2.2 – Discovery of the Structure of the Atom -We will only discuss protons, neutrons and electrons because they are the only ones that affect chemical behavior -Charge on proton = +1.602 x 10-19 C. Assigned a +1 charge. -Charge on an electron = -1.602 x 10-19 C. Assigned a -1 charge. -Neutrons – Electrically neutral. Similar mass to protons. Discovered by Chadwick. Mass of electron so small it is ignored

12. 2.3 – Modern Atomic Theory amu = atomic mass unit. 1 amu = 1.66054 x10-24 g. Based on 12 amu = mass of one atom of Carbon-12. 1 amu is approximately equal to the mass of a proton/neutron. Angstrom (Å) = Unit used to measure atom size. 1 Angstrom = 1 x 10-10 m

13. 2.3 – Modern Atomic TheoryAtomic Numbers, Mass Numbers and Isotopes Symbols: 1. 2. Carbon - 12 Mass Number

14. 2.3 – Atomic Number, Etc. Atomic Number: All atoms of the same element have the same number of protons: The atomic number (Z)

15. 2.3 – Atomic Number, Etc. Mass Number: The total number of protons and neutrons in the atom.

16. 11 6 12 6 13 6 14 6 C C C C 2.3 – Atomic Number, Etc. Isotopes = Atoms of the same element with different masses. -Isotopes have different numbers of neutrons. -Isotopes have different mass numbers. Carbon-11 Carbon-13 Carbon-12 Carbon-14

17. 2.4 – Atomic Weights Average Atomic Masses – Based on abundance of isotopes and mass of each isotope. =Atomic Weight Example 2.4 – Page 47

18. 2.5 - Periodic Table Periodic Table – Arranged by atomic number with elements with similar properties in same column. Mendeleev – Organized first accepted Periodic Table • The rows on the periodic table are periods. • Columns are groups. • Elements in the same group have similar chemical properties.

19. Periodic Table Metals are on the left side of the chart. -Low electronegativity -Good conductors of heat and electricity -Malleable, Ductile, Luster -Readily lose electrons

20. Periodic Table Nonmetals are on the right side of the periodic table (with the exception of H). -High electronegativity -Readily gain electrons

21. Periodic Table Metalloids border the stair-step line (with the exception of Al and Po). -Properties are intermediate between metals and nonmetals.

22. Groups 1 2 16 17 18 These five groups are known by their names.

23. 2.6 – Molecules and Molecular Compounds Chemical Formula = Shows the atoms present in a substance Molecule – Two or more atoms in a definite arrangement held by covalent bonds. Can be the same or different atoms. ex – H2, H2O, C6H12O6 Diatomic Molecule = Molecule made of only 2 atoms total Molecular Compound – Must have DIFFERENT elements ex- H2O is a compound, but H2 is NOT

24. 2.6 – Molecules REVIEW: Bonds – -Nonpolar Covalent – Equal sharing of electrons -No partial charges -Electronegativity of two atoms is close -Difference of electronegativities is <0.3 -Polar Covalent – Unequal sharing of electrons -Partial charges exist on atoms -Electronegativity difference is 0.3-1.7

25. 2.6 - Molecules Molecular Formula = Gives actual number and type of atoms in a molecule Empirical Formula = Simplest whole number ration of atoms Example = Glucose

26. 2.6 - Molecules Structural Formula – Similar to a Lewis Structure. Shows bonds, but not shape.

27. 2.6 - Molecules

28. 2.7 – Ions and Ionic Compounds • When atoms lose or gain electrons, they become ions. • Cations are positive and are formed by elements on the left side of the periodic chart. • Anions are negative and are formed by elements on the right side of the periodic chart.

29. 2.7 - Ions • Examples:

30. 2.7 - Ions Ionic compounds (such as NaCl) are formed by ionic bonds. Contain both + and – ions. Ionic bonds are generally formed between metals (cations) and nonmetals (anions). Monatomic – one atom only with a charge Polyatomic – more than one atom with an overall charge

31. Metallic Bond • Metal ions with an electron sea.

32. 2.8 – Naming Inorganic CompoundsIonic Names and Formulas of Ionic Compounds: • Cations • One charge only = name only Zn+2 = Zinc ion Ca+2 = Calcium ion • Multiple charges possible = use Roman Numeral Cu+1 = Copper (I) Cu+2 = Copper (II) -Old system uses –ous / -ic • Polyatomics – hydronium, ammonium

33. Common Cations

34. 2.8 – Inorganic CompoundsIonic Names and Formulas of Ionic Compounds: 2. Anions: • Monotomic – Change ending to –ide • Polyatomic – MEMORIZE 8-ates and rules! • H+ added ions = use bi- or hydro-prefix Example – bicarbonate = hydrogen carbonate • Carbonate CO3-2 • Nitrate NO3-1 • Sulfate SO4-2 • Chlorate ClO3-1 • Chromate CrO4-2 • Bromate BrO3-1 • Phosphate PO4-3 • Iodate IO3-1

35. Common Anions

36. 2.8 – Inorganic CompoundsIonic Ionic Compounds – Cation + Anion -Binary – Compound of 2 “atoms” -ex – NaCl, Al2O3 -Ternary – Compound of 3 “atoms” -ex – Mg(NO3)2

37. 2.8 – InorganicAcids

38. 2.8 – Inorganic CompoundsBinary Molecular Molecular Compounds: Discrete molecular units. Contain covalent bonds. Use prefixes.

39. 2.9 - Organic