College Prep' Chemistry Ch' 3 p' 1

1
College Prep. Chemistry Ch. 3 p. 1

• Atom-smallest particle of an element that retains
  its identity in a chemical reaction.
• Democritus- A Greek Philosopher who first
  described the notion of the atom, 4th Century
  B.C. Called Atomists He believed that atoms
  were indivisible and indestructible. No
  Experimental Support.
• Law of Conservation of Mass- Mass is neither
  created nor destroyed during any ordinary
  chemical or physical change.
• Law of Definite Proportions -a chemical compound
  has the same proportions by mass regardless of
  the size of the sample. Ex. NaCl. 11 ratio

2
College Prep. Chemistry Ch. 3 p. 2

• Law of Multiple Proportions- If two or more
  different compounds are composed of the same two
  elements, then the ratio of the masses of the
  second element combined with the first element is
  in the ratio of small, whole numbers. Ex. CO and
  CO2.
• John Dalton- British scientist who did
  experiments to explain his theories in the
  1840s.
• Daltons Atomic Theory
• 1. All elements are composed of indivisible
  particles called atoms.
• 2. Atoms of the same element are identical.
• 3. Atoms of different elements are different.

3
College Prep. Chemistry Ch. 3 p. 3

• Atoms of different elements can physically mix
  together or can chemically combine with one
  another in simple, whole number ratios to form
  compounds.
• Chemical reactions occur when atoms are
  separated, joined or rearranged. However atoms
  of one element are never changed into atoms of
  another element as a result of a chemical
  reaction.
• Section Review p. 71 1-3

4
College Prep. Chemistry Ch. 3 p. 4

• Subatomic Particles- includes electrons, protons,
  neutrons.
• Protons are positively charged particles that are
  found in the nucleus. Neutrons are neutral
  particles found in the nucleus. Electrons are
  negatively charged particles found on the outside
  of the nucleus.
• The nucleus is the small, dense, positively
  charged region in the center of the atom.
• Electrons were discovered in 1897 by J.J.
  Thomson.
• He performed experiments that involved passing
  electric current through gases at low pressure.

5
College Prep. Chemistry Ch. 3 p. 5

• There were sealed tubes with metal disks called
  electrodes at each end. P. 73
• Anode- positively charged electrode.
• Cathode-negatively charged electrode.
• Cathode Ray- Electron beam that travels from the
  cathode to the anode. Electron beam is attracted
  to the positive plate of the magnet.
• Electrons, e- have 1 unit of negative charge.
• Mass 1/1840 amu.
• Millikans Oil Drop Experiment in 1916 reported
  the charge to mass ratio of the electron.

6
College Prep. Chemistry Ch. 3 p. 6

• Rutherfords Experiment-Gold Foil Experiment,
  he directed a narrow beam of alpha particles at a
  very thin sheet of gold foil. P. 74. Alpha
  particles are He nuclei.
• A small amount of alpha particles were deflected,
  most went through the gold foil. He proposed
  that all of the mass and all of the positive
  charge are in the small, dense part of the atom
  called the nucleus.
• Nucleus- center of the atom, contains n and p,
  small dense part of the atom.
• Electrons occupy most of the volume of the atom,
  scattered outside the nucleus. Electrons have
  very little mass.

7
College Prep. Chemistry Ch. 3 p. 7

• Nuclear Forces- short range proton-neutron,
  proton-proton, neutron-neutron forces hold the
  nuclear particles together.
• Sizes of atoms are measured in picometers.
  1x1012 pm 1 m.
• Section Review p. 76 1-5
• Lab Constructing a Model (handout)
• Atomic Number- number of protons in the nucleus
  of the atom. (Z)
• A. It identifies the element., Elements are
  arranged by increasing atomic number on the
  periodic table.
• Isotopes- Atoms that have the same number of

8
College Prep. Chemistry Ch. 3 p. 8

• Protons, but different numbers of neutrons.
• Isotopes have different mass numbers.
• 11H Hydrogen-1 (0 neutrons), 21H Deuterium-2 (1
  neutrons), 31H Tritium-3 (2 neutrons).
• Another Example Carbon-12 126C, Carbon-13
  136C This is called a nuclear symbol.
• Mass Number- is the total number of protons and
  neutrons in the nucleus of an isotope. Isotopes
  are identified by their mass numbers. Carbon-13.
• Mass Number Atomic Number Number of Neutrons
• Nuclide- General term for an isotope of an
  element.

9
College Prep. Chemistry Ch. 3 p. 9

• Sample Problem A p. 79
• Practice p. 80 1-3
• Atomic Mass- Masses are difficult to work with of
  individual atoms. We use a comparative scale.
  All atoms are compared relative to carbon.
• Carbon-12 12 a.m.u. Hydrogen-1 is 1.007
  a.m.u. 1.66 x 10-27 kg.
• A.m.u. atomic mass unit
• Isotopes have similar chemical behavior, even if
  they have different masses of the same element.

10
College Prep. Chemistry Ch. 3 p. 10

• Average atomic mass is the weighted average of
  the atomic masses of the naturally occurring
  isotopes of an element.
• Ex. For Boron (10amu x .1991) (11amu x
  .8009) 10.810 a.m.u.
• .1991 19.91 .8009 80.09
• Mole- SI unit for the amount of substance. It is
  the amount of a substance that contains as many
  particles as there are atoms in exactly 12 g of
  carbon-12.

11
College Prep. Chemistry Ch. 3 p. 11

• The number of particles in a mole is known as
  Avogadros Number. 6.022 x 1023. This is
  representative particles which includes
  molecules, atoms, ions, and formula units.
• 1 mole 6.022 x 1023 rep. particles
• Molar Mass the mass of one mole of a pure
  substance is called the molar mass of that
  substance. Ex. Li is 6.94 g/mole, Hg is 200.59
  g/mole
• Sample Problem B p. 84
• ? g 3.5 moles Cu x 63.5 g Cu/mole Cu 222 g Cu

12
College Prep. Chemistry Ch. 3 p. 12

• Practice p. 85 top 1-4, Bottom 1-3
• Sample Problem C p. 85
• ?moles Al 11.9 g Al x 1 mole Al/27 g Al .441
  moles Al
• Sample Problem D p. 86
• ? Moles Ag 3.01 x 1023 atoms Ag x 1 mole/ 6.02
  x 1023 atoms Ag .5 moles Ag
• Practice p. 86 1-3

13
College Prep. Chemistry Ch. 3 p. 13

• Sample Problem E p. 86-87
• ? g Cu 1.2 x 108 atoms Cu x 1 mole Cu/ 6.02 x
  1023 atoms Cu x 63.5 g Cu/mole Cu 1.27 x 10-14
  g Cu
• Practice p. 87 1-3, Section Review p. 87 1-7
• Review p. 89-90 1, 3-9, 11, 12, 14, 15, 17, 19,
  21-23, 28