Download Chapter 2: Atoms, Molecules, and Ions - GW

Chapter 2: Atoms, Molecules,
and Ions
Macroscopic
• Objects that are large enough to be seen with
the unaided eye
Microscopic
• Objects that can be seen only under
magnification
Democritus (460 B.C.-370 B.C)
• First to suggest
existence of atoms
• Said atoms are
indivisible and
indestructible
Dalton’s Atomic Theory 1766-1844
• Elements are made up of tiny particles called
atoms
• All atoms of a given element are identical
• All atoms of a given element are different
from those of a different element
• Atoms combine to form compounds. These
compounds contain specific ratios.
• Atoms are indivisible by chemical processes- a
solid indivisible mass
Errors in Dalton’s Theory
• Atoms are now known to be divisible
• Protons
• Neutrons
• Electrons
Daltons Model
1803: atoms are
tiny,
indestructible
particles with no
internal structure
Atoms
• The smallest particle of
an element that retains
its identity in a chemical
reaction
• Are the basic building
blocks of matter.
Iron Atoms on Copper
Element
• Is composed of only one type of atom.
Compound
• The atoms of two or more elements combine
in definite arrangements
Law of Constant Composition
• In a given compound, the relative number and
kind of atoms are constant
Law of Conservation of Mass
• The total mass of materials present after a
chemical reaction is the same as the total
mass before the reaction
Cathode Rays
• High voltage radiation
originating from the
negative electrode
(cathode)
• Cathode rays travel in a
straight line, however they
will bend in the presence of
a magnetic field and will
make a metal plate
negatively charged  the
radiation was composed of
negatively charged particles
(electrons)
Robert Millikan (1868-1953)
• In 1909 Millikan
measured the charge of
an electron using the
“Millikan oil drop
experiment)
Millikan Oil-Drop Experiment
• Small drops of oil which
had picked up electrons
fell between 2 electrically
charged plates
• Millikan monitored the
drops, measuring how the
voltage on the plates
affected the rate of their
fall
• Charges were always
integral multiples of
1.6 x 10 -19 C the charge
of an electron
J.J Thomson 1856-1940
• 1897: Plum Pudding
Model  electrons
embedded in a sphere
of positive electrical
charge
Radioactivity
• Spontaneous emission
of radiation
Ernest Rutherford 1871-1937
• Proposed the nuclear atomelectrons surround a
nucleus
• 1911 directed a particles at
gold foil- some were
reflected or deflected,
indicating a concentrated
positive charge
• Discovered the proton-later
experiments showed the
presence of neutrons
Ernest Rutherford
1911: an atom has a small,
dense, positively charged
nucleus. Electrons move
around the nucleus
Protons
• Positively charged subatomic particles (+1)
• Discovered in 1886 by Eugene Goldstein 
Canal Rays travel in opposite directions as
cathode rays
Neutrons
• Subatomic particles with no charge but mass
nearly equal to that of a proton
Nucleus
• Protons and Neutrons are located in the
nucleus
Niels Bohr
• The electron moves in a
circular orbit at a fixed
distance from the
nucleus
Bohr Model Continued
• Energy levels are quantized
or fixed- you cannot exist
half way between
• To move from one energy
level to another, an electron
must gain or lose energy
• Quantum: the amount of
energy required to move an
electron up 1 level
Bohr Model Continued
• Not all steps are equalenergy levels are not
equally spaced
• Energy levels become closer
together the further out
that you move.
• The higher the energy level,
the easier it is for an
electron to escape from the
atom
Erwin Schrodinger 1887-1961
• Electron Cloud Model
(Quantum Mechanical
Model)
• Mathematical model
• Estimates the probability of
finding an electron in a
certain position
• Probability is represented as
a fuzzy cloud- an electron
cloud
Atomic Number
• The number of protons in the nucleus of an
atom
• Elements are different because they contain
different numbers of protons
Mass Number
• The total number of protons and neutrons in
an atom
Mass
Number
Calculating Neutrons
6 Neutrons
Isotopes
• Atoms that have the
same number of
protons but different
numbers of neutrons
• Because they have
different numbers of
neutrons they also have
different mass numbers
Nuclide
• An atom of a specific isotope
Atomic Mass
• Is the weighted average mass of the atoms in
a naturally occurring sample of the element
• Reflects both the mass and the relative
abundance of the isotopes as they occur in
nature
Atomic Mass (AMU)
• To calculate the atomic mass of an element,
multiply the mass of each isotope by its
natural abundance, expressed as a decimal,
and then add the products
Example
98.89%
Atomic Mass = 12.00 amu X 0.9889
+ 13.003 amu X .0111
12.011 amu
1.11%
Valence Electrons
• The electrons in the highest occupied energy
level of an element’s atoms
• The number of valence electrons largely
determines the chemical properties of an
element
• To find the number of valence electrons look
at the group number (except noble gases)
Electron Dot Structure
• Are diagrams that show valence electrons as
dots
Octet Rule
• In forming compounds, atoms tend to achieve
the electron configuration of a noble gas
(except helium)
• General electron configuration- ns2np6
Octet Rule
• Metals tend to lose electrons
• Nonmetals tend to gain electrons
Formation of Cations
• An atom is electronically neutral because it
has equal numbers of protons and electrons
• An atom’s loss of valence electrons produces a
cation- a positively charged ion
• For metals, the name of the cation is the same
as the name of the element (Na+ = Sodium)
Cations
Cations of Transition Metals
• Charges of cations may vary
• Fe  Fe2+ + 2e• Fe  Fe3+ + 2e-
Formation of Anions
• An atom’s gain of valence electrons produces
a anion- a negatively charged ion
• The name of an anion ends in –ide (Cl- =
chloride)
Polyatomic Ions
• An ion composed of two or more elementsbehaves as a unit with a set positive or
negative charge
Formation of Ionic Compounds
• Compounds composed of cations and anions
• Although they are composed of ions they are
electrically neutral
Ionic Bond
• Anions and cations have opposite charges and
attract one another by means of electrostatic
forces
Formula Unit
• The smallest whole-number ratio of ions in an
ionic compound
• Na+ + Cl-  NaCl (Formula Unit)
• Mg2+ + 2Cl-  MgCl2 (Formula Unit)
• Al3+ + 3Cl-  AlCl3 (Formula Unit)
Properties of Ionic Compounds
• Most ionic compounds are crystalline solids at
room temperature
• Component ions are arranged in repeating 3dimensional patterns
• Large attractive forces result in a very stable
structure
• High melting points
Ionic Compounds
NaCl
FeS2
CaF2
CaCO3
HgS
BaSO4
Naming Ionic Compounds
• Always name the cation first element name
• Name the anion second element name
ending in –ide
– NaCl
– MgO
– K2O
Sodium Chloride
Magnesium Oxide
Potassium Oxide
Naming Ionic Compounds Containing
Transition Metals
• Always name the cation first element name
• Use Roman Numerals to represent the charge
of the cation
• Name the anion second element name
ending in –ide
– FeCl2
– FeCl3
Iron (II) Chloride
Iron (III) Chloride
Group or Family
Periodic Table
Period
Periodic Law
• In the modern periodic table, elements are
arranged in order by increasing Atomic
Number
• 7 horizontal rows called periods
• Each period corresponds to a principle energy
level
• Elements within a column (group) have similar
properties
Periodic Law
• States: when elements are arranged by
increasing atomic number, there is a periodic
repetition of their physical and chemical
properties
Metals, Nonmetals, Metalloids
Metals
• 80% of the periodic table is
metals
• Good conductors of heat
and electricity
• Malleable- can be shaped
• Ductile- can be drawn into
wires
• Lustrous
• Solids at room temperature
(except Hg)
Nonmetals
•
•
•
•
Poor conductors of heat and electricity
Not malleable
Not ductile
dull
Metalloids
• Generally has some properties of a metal,
some of a nonmetal