Transcript Chapter 2

Chapter 2
Atoms, Molecules, and Ions
LAW OF CONSERVATION OF
MASS
• Antoine Lavoisier (1743-1794)
• Carefully measured and provided a
quantitative interpretation of the chemical
reaction associated with combustion.
• Law: During a chemical change, the total
mass remains constant.
LAW OF DEFINITE
PROPORTIONS
• Joseph Proust (1754-1826)
• Law: Different samples of a pure
compound always contain the same
proportion of elements by mass.
• Also called Law of Constant Composition
LAW OF MULTIPLE
PROPORTIONS
• John Dalton (1766-1844)
• Law: When two or more elements form
more than one compound, the ratio of the
masses of one element in these compounds
for a fixed mass (i.e. 1 gram) of the other
element is a small whole number.
ATOMIC THEORY OF
MATTER (1808)
• John Dalton (1766-1844)
• Elements (matter) are composed of atoms
• The atoms of a given element are identical.
Each element is characterized by the mass
of its atoms.
• Compounds are formed when atoms of
different elements combine with each other
(Multiple Proportions).
ATOMIC THEORY OF
MATTER (2)
• A given compound is a chemical
combination of the same atoms in the same
relative numbers. (Definite Proportions).
• A chemical reaction is the rearrangement of
atoms leading to new compounds. Atoms
are neither destroyed nor created in a
chemical reaction (Conservation of Mass).
AVOGADRO’S HYPOTHESIS
• Amadeo Avogadro (1776-1856)
• Equal volumes of two gases at the same
pressure and temperature contain the same
number of particles (atoms or molecules).
ATOMIC STRUCTURE
Subatomic Particles (Table 2.1)
• ELECTRON (cathode ray)
– 1899 J.J. Thomson: electron charge = -1 and e/m = 1.76E+8 C/g
– 1909 R. Millikan: electron mass = 9.11E-31 kg or
about 1/1836 of proton or neutron mass
• PROTON
1911 Rutherford
– Positive charge, +1
• NEUTRON
– No charge, 0
Relative mass = 1 amu
1932 Chadwick
Relative mass = 1 amu
Table 2.1 The Mass and Charge of
the Electron, Proton, and Neutron
MODELS OF THE ATOM
• Thomson (Plum Pudding) Model: positive
mass with electrons embedded in it
• Rutherford Model (1911): positive charge
in small volume with (diameter = 1E-15 m)
electrons occupying mostly empty space (d
= 1E-10 m) around the nucleus
• Bohr Atom - Chapter 7
• Quantum Mechanical Atom - Chapter 7
Figure 2.13 a & b
(a) Expected Results of the Metal Foil
Experiment if Thomson's Model Were Correct
(b) Actual Results
ATOMIC STRUCTURE
• Atomic Symbol
– Shorthand notation for element
– One or two letters on Periodic Table
• Atomic Structure
– Atomic Number (Z) = #protons, uniquely
defines an atom
– Mass Number (A) = #protons + #neutrons
– If atom is neutral, Z = #electrons
NUCLIDE SYMBOL
• Atomic symbol, E; symbol in the middle of each
element box on the Periodic Table.
• Z (left subscript); number on the top of each
element box on the Periodic Table.
• A (left superscript)
• If species is an ion (has a charge), add + or charge (right superscript)
•
A
ch
E
Z
23
11
Na has 11 p+, 11 e-, 12 no
IONS
• A charged species with unequal numbers of
protons and electrons.
• If # protons > # electrons, the ion has a net
positive charge and is called a cation
• If # protons < # electrons, the ion has a net
negative charge and is called a anion
• An ion may consist of an atom or a group of atoms
• 2311Na+ has 11 p+, 10 e-, 12 no
MOLECULES
• Molecules or compounds form when atoms are connected
by chemical bonds in which electrons act as the “glue”
between atoms.
– If electrons are shared between two atoms, the bond is a covalent
bond. I.e., the bond between two non-metal atoms.
– If electrons are transferred to produce ions, the bond is ionic.
• Ions are charged particles which form via the gain (anion, commonly
formed from nonmetal elements) or loss (cation, commonly formed
from metal elements) of electrons.
• Oppositely charged ions attract and form an ionic bond.
• Type of bond between a metal and a non-metal atom.
• Polyatomic ions are charged groups of atoms; they can also form
ionic bonds.
ISOTOPE
• Atoms which have the same Z (same #
p+)but a different A (different # n0)
• Most elements have isotopes that occur in
nature in precise proportions (fractional
abundances, %).
• A few elements have no naturally occurring
isotopes.
Figure 2.15 Two Isotopes of
Sodium
PERIODIC TABLE
• An arrangement of elements according to
increasing atomic number which shows the
periodic or regularly repeating nature of
elemental properties.
– Rows = periods
– Columns = groups or families; note similarity
of properties
– Metals
Nonmetals
Semimetals
– Main group (A)
Transition Metals
Figure 2.21 The Periodic Table
NOMENCLATURE or
NAMING COMPOUNDS
• Binary Ionic Compounds
– Metal atoms tend to lose electrons and form cations.
– Nonmetal atoms tend to gain electrons and form anions.
– Use Periodic Table to determine charges and number of
each ion in the compound. Note that the ionic
compound must be neutral overall.
– Name cation first as element and anion second with
“ide” ending.
• Binary Ionic Compound Type I (T2.3)
– Some atoms form only a single type of ion (Binary
Ionic Cmp Type I)
Figure 2.22 Common Cations and
Anions
NOMENCLATURE (con’t)
• Binary Ionic Compound Type II (T2.4, Fig
2.22)
– Some transition metal elements form more than
one common ion. Designate charge with
Roman numeral (II)
• Polyatomic Ions (Table 2.5)
– Memorize
– Oxoanion = nonmetal + oxygen
Table 2.5 Common Polyatomic
Ions
NOMENCLATURE (con’t)
• Binary Covalent Compound Type III (Table 2.2)
– Compounds formed from two nonmetals in which
electrons are shared in chemical bond.
– Name more “cation-like” first, then the more “anionlike) second with “ide” ending. Hydrogen is almost
always named first.
– Indicate number of each using prefix as needed. (T2.6)
– Note historic names (water, ammonia)
Table 2.6
Prefixes
Used to
Indicate
Number in
Chemical
Names