Download CHAPTER 3, ATOMS: THE BUILDING BLOCKS OF MATTER

CHAPTER 3, ATOMS: THE BUILDING BLOCKS OF MATTER
Section 1, The Atom: From Philosophical Idea to Scientific Theory
The law of conservation of mass states that mass is neither created nor destroyed
during ordinary chemical reactions or physical changes.
The law of definite proportions states that a chemical compound contains the same
elements in exactly the same proportions by mass regardless of the size of the
sample or source of the compound.
The law of multiple proportions states that 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 a certain mass of the first element is always a ratio of small
whole numbers.
Dalton’s atomic theory:
1. All matter is composed of extremely small particles called atoms.
2. Atoms of a given element are identical in size, mass, and other properties;
atoms of different elements differ in size, mass, and other properties.
3. Atoms cannot be subdivided, created, or destroyed.
4. Atoms of different elements combine in simple whole-number ratios to
form chemical compounds.
5. In chemical reactions, atoms are combined, separated, or rearranged.
Not all aspects of Dalton’s atomic theory have proven to be correct. Today we
know that atoms are divisible into even smaller particles. Also, we know that a
given element can have atoms with different masses.
Section 2, The Structure of the Atom
An atom is the smallest particle of an element that retains the chemical properties
of that element.
All atoms consist of two regions.
1. The nucleus is a very small region located at the center of the atom. The
nucleus is made up of at least on proton (positively charged particle) and
usually one or more neutrons (neutral particle).
2. The region surrounding the nucleus contains electrons. This region is
very large compared with the size of the nucleus
Cathode-ray experiment conclusions
! All cathode rays are composed of identical negatively charged particles,
which were name electrons.
! Electrons are present in atoms of all elements.
! Atoms are divisible.
! The electron has a very large charge-to-mass ratio
Discovery of the Charge of an Electron: 1909 by Robert A. Millikan via oil drop
experiment
Further conclusions
! Because atoms are electrically neutral, they must contain a positive charge
to balance the negative electrons.
! Because electrons have so much less mass than atoms, atoms must contain
other particles that account for most of their mass.
The plum pudding model, proposed by Thomson, showed that negative electrons
were spread evenly throughout the positive charge of the rest of the atom (i.e.,
seeds in a watermelon).
Discovery of the Nucleus: 1911 by Ernest Rutherford w/ Hans Geiger and Ernest
Marsden via gold foil experiment
Gold Foil Conclusions
! The deflected alpha particles experienced some powerful force within the
atom.
! The source of the force must occupy a very small amount of space
because so few alpha particles had bee affected by it.
! The force must be caused by a very densely packed bundle of matter with
a positive electric charge, which he called a nucleus.
The Bohr model, proposed by Niels Bohr, showed electrons surrounding the
positively charged nucleus, as the planets surround the sun. (See Chapter 4.)
Table 1. Properties of Subatomic Particles
Protons, neutrons, and electrons are often referred to as subatomic particles.
Discovery of the Electron: 1897 by Joseph John Thomson via cathode-ray
experiments
The nuclear force is short-range proton-neutron, proton-proton, neutron-neutron
force that holds nuclear particles together.
The radius of an atom is the distance from the center of the nucleus to the outer
portion of the electron cloud. Expressed in picometers.
Section 3, Counting Atoms
Isotopes are atoms of the same element that have different masses. Although isotopes
have different masses, they do not differ significantly in their chemical behavior.
A nuclide is a general term for a specific isotope of an element.
The mass number (A) is the total number of protons and neutrons that make up the
nucleus of an isotope.
The atomic number (Z) of an element is the number of protons of each atom of that
element.
Represent the nucleus of an atom of an element through the following symbol:
A
Z
!
X
A = mass number
Z = atomic number
X = element symbol
One atomic mass unit (amu) is exactly 1/12 the mass of a carbon-12 atom.
Average atomic mass is the weighted average of the atomic masses of the naturally
occurring isotopes of an element.
A mole (mol) is the amount of a substance that contains as many particles as there
are atoms in exactly 12 g of carbon-12, OR, the amount of a substance that
contains Avogadro’s number of particles.
Avogadro’s number (6.022 x 1023) is the number of particles in exactly one mole of
a pure substance.
The mass of one mole of a pure substance is called the molar mass of that
substance.
Figure 11. Solving Mole Problems