Transcript Chapter 2 Atoms, Molecules, and Ions

Chapter 2
Atoms, Molecules, and Ions
Krishna Trehan
AP Chemistry 2008-2009
Mrs.Molchany
2.1 - The Atomic Theory of Matter
• The idea of atoms was formulated when
scientists tried to explain properties of gasses.
– Gas composed of invisible objects in constant
motion.
• Isaac Newtown proposed the idea that atoms
were the chemical building blocks of nature.
Daltons Atomic Theory
• 1. Each element is composed of small particles
called atoms
• 2. All atoms of a given element are identical
• 3. Chemical reactions do not alter atoms, they
can be neither created nor destroyed.
• 4. Compounds are formed when atoms of
more than one element combine.
Daltons Theory
• Atoms are the most basic building blocks of
matter that retain chemical identities of
elements.
• Law of constant composition – in a compound
the relative numbers and kinds of atoms are
constant.
• Law of conservation of mass – the total mass
of materials present is the same before and
after a reaction occurs.
Daltons Theory
• Law of multiple proportions – When elements
combine, they do so in the ratio of small
whole numbers.
i.e. When Carbon and Oxygen react, the produce CO
or CO₂ and not something like CO1.6
2.2 The Discovery of Atomic Structure
• The atom is composed of subatomic particles.
• There are only two kinds of charges : positive,
and negative
– Law of Electrostatic Attraction
like charges repel one another, unlike charges
attract.
Cathode Rays and Electrons
• An tube is pumped almost empty of air. High
voltage produces radiation, which is called a
cathode ray, because it comes from the
negative electrode, or cathode.
• The rays could not be seen, but detected.
Depending on which gas is in the tube, the ray
will give off a certain colored light.
J.J.Thompson on Cathodes and Rays
• J.J. Thompson concluded that cathode rays are
not waves, but particle masses.
• Discovered the ratio of the electrons charge to
mass as 1.76 x 10⁸ coulombs per gram. He did
this by bending the path of a cathode using
magnets. By observing the magnitude of the path
change, he was able to derive this number.
Robert Millikan
• Robert Millikan tried to find the mass of an
electron by using the Oil Drop Experiment
• Millikan allowed small drops of oil (which had
obtained extra electrons) to fall on top of
electrically charged plates.
• From these observations he found that the
charge on an electron was 1.6 *1019C
Millikan
1.6 x1019C
 28
Mass 

9
.
1
x
10
g
8
1.76 x 10 C / g
• By using J.J.
Thomson’s value for
the electrical charge
to mass ratio,
Millikan found the
presently accepted
mass of the electron.
Ernest Rutherford
• Rutherford discovered three types of radiation: alpha,
beta, and gamma.
• Alpha and beta are bent by electric fields, while gamma
rays aren't.
• Alpha: are larger than beta rays, and have a positive
charge (attracted to negatively charged plates).
• Beta: have a negative charge (attracted to positively
charged plates.
– Considered the radioactive equivalent of cathode rays.
The Nuclear Atom
• Thompson proposed the
idea that since electrons
were relatively small, they
held a small fraction in the
total area of the atom.
– He then made the “Plum
Pudding” model where
individual negatively
charged electrons were
spread throughout a
positive sphere.
Gold Foil Experiment
• Thompson sent a beam of alpha particles
through a thing piece of gold foil. He
discovered that all of the alpha particles
passes through. But after reviewing the
experiment, he saw that some alpha particles
were deflected in other directions, some even
bounced back. This contradicted the Plum
Pudding model.
Gold Foil Experiment
• He postulated from this
experiment that all of the
positive charge in the
atom was concentrated
in a small dense region,
called the nucleus.
• This meant that most
alpha particles simply
passed the nucleus, but
the ones that bounced
back, hit the nucleus and
were repelled by its
positive charge.
2.3 The Modern View of Atomic
Structure
•
•
•
•
•
Charge of an electron is -1.602 x 10 -19 C
Charge of a proton is 1.602 x 10 -19 C
1.602 x 10 -19 is called the electronic charge.
C = coulomb
Atoms have the same number of electrons
and protons, so they have no net electrical
charge
Atomic Structure
• Protons and Neutrons reside in the nucleus
• A large majority of an atoms area is outside the
nucleus, where the electrons are.
• Atoms have very little mass, so we use the atomic
mass unit (amu) to measure mass.
24
1
.
66054
x
10
g
• One amu is equivalent to:
• A proton is 1.0073 amu and a neutron is 1.0087
amu
• Angstroms (Å), which are 10-10 m, are used to
measure diameter
Isotopes, Atomic Numbers, and Mass
Numbers
• Protons are what make elements unique
– All atoms of an element have the same number of
protons in the nucleus.
• Isotopes are atoms of a given element that
differ in the number of neutrons.
• Atomic number – the number of protons.
• Mass number – total number of protons and
neutrons in an element
2.4 The Periodic Table
• The arrangement of elements in order of
increasing atomic number, with elements
having similar properties placed in vertical
columns is known as the periodic table.
• Each column is a group.
– Elements in the same group and have similar
physical and chemical properties.
Metallic Elements
• All the elements on the left side and in the
middle of the periodic table are metals.
• Metals are generally lustrous and are good
conductors of heat and electricity.
• All metals (besides mercury Hg) are solid at
room temperature.
Nonmetallic Elements
• The metals are separated for nonmetals by a
diagonal staircase line that runs from boron
(B) to astatine (At).
• Hydrogen, though a nonmetal, is on the left
side of the periodic table.
• The state of matter at which nonmetals are at
room temperature, vary from element to
element.
Metalloids
• Elements that have the characteristics of both
metals and nonmetals are called metalloids.
• They can also be called semi-metals.
2.5 Molecules and Molecular
Compounds
• Molecule – an assembly of two or more atoms
tightly bound together
• Diatomic – molecule composed of two atoms,
both of which are the same element.
• Molecular compounds are compounds
composed of molecules.
– For example, in water there are two hydrogens,
and one oxygen. The two hydrogens are a
diatomic molecule.
Molecular and Empirical Formulas
• Molecular Formulas – chemical formulas that
indicate the actual numbers and types of atoms
in a molecule.
• Empirical Formulas – chemical formulas that give
only the relative number of atoms of each type in
a molecule.
– Always show the smallest whole number ratio
• i.e. C₂H₄ = CH₂
• Molecular formulas provide more/accurate
information about a molecule
Picturing Molecules
• Structural formula - shows which elements
are attached to which in a chemical formula.
• Generally, structural formulas do not
represent actual images, but simply provide a
sketch of the structure.
H₂O
H₂O₂
2.6 Ions and Ionic Compounds
• If an electron is added or removed from an
atom, then the atom now has a charge, or has
become an ion.
• Cation = positively charged = loss of electron
• Anion = negatively charged = gain of electron
• Polyatomic ion – ions that consist of atoms
joined as in a molecule, but have a net
positive/negative charge : NO₃⁻1 or SO₄⁻2
Predicting Ionic Changes
• Atoms gain or lose electrons so as to end up
with the same number of electrons as the
noble gas closest to them.
• The noble gases are non-reactive and form
few compounds.
Ionic Compounds
• Ionic compound - compounds that contains
positively charged ions and negatively charged
ions
– When sodium and chlorine react, an electron from
sodium goes to chlorine which leaves us with Na⁺+ Cl⁻
After that, the opposite charges attract one another
and Na and Cl are held together.
• Ionic compounds are combinations of metals and
nonmetals.
• Molecular compounds are generally composed of
nonmetals only.
Chemical Compounds
• Chemical compounds are always electrically
neutral, so we need to make sure the ions in an
ionic compounds occur in a ratio where the total
positive charge is EQUAL to the total negative
charge.
Mg2+ + N3-  Mg₃N₂
3 (number of Mg) x 2 (charge of Mg) = 6
2 (number of N ) x -3 (charge of N) = -6
-6 + 6 = 0 which means we have an
overall neutral charge.
2.7 Naming Inorganic Compounds
Chemical Nomenclature
On page 56 of the old textbook you can find a
list of many different cations.
On page 57 of the old textbook you can find a
list of many different anions.
Most of these are found on the pink sheets we
got at the beginning of the year.
Positive Ions (Cations)
• Cations formed from metal atoms have the
same name as the metal.
Example : Na⁺ = sodium ion
• IF a metal can form cations of differing
charges, the positive charge is given by a
Roman numeral in parentheses following the
name of the metal:
Example : Fe2+ = iron(II) ion
Fe3+ = iron(III) ion
• Most elements that have multiple charges are
transition metals
– An older form of naming elements was using the
latin name (these are found on the pink sheets
Mrs.Molchany gave us)
• Fe2+ = ferrous ion
Fe3+ = ferric ion
• Cations formed from nonmetal atoms have
names that end in –ium.
NH₄⁺ = ammonium ion
H₃0⁺ = hydronium ion
Negative Ions (Anions)
• Monatomic anions have names formed by
dropping the ending of the name of the
element and adding the ending –ide.
H⁻ = hydride ion
CN⁻ = cyanide ion.
• Polyatomic anions containing oxygen have
names ending in –ate or –ite.
NO₃⁻1 = nitrate
NO₂⁻1 = nitrite
Negative Ions (Anions) (cont.)
• Anions derived by adding H⁺ to an oxyanion
are named by adding as a prefix the word
hydrogen or dihydrogen.
CO₃2- (carbonate ion)  HCO₃⁻ (hydrogen carbonate ion)
Ionic Compounds
• Naming ionic compounds follow the format:
• cation name followed by the anion name
BaBr₂ = barium bromide
Al(NO₃)₃ = aluminum nitrate
Cu(ClO₄)₂ = copper (II) perchlorate
Names and Formulas of Acids
• Acids based on anions whose names end in –ide.
Anions whose names end in –ide have associated
acids that have the –hydro prefix and an –ic.
– Cl⁻(chloride)  HCL (hydrochloric acid)
• Acids based on anions whose names end in ite or
ate. Anions whose names end in ate have acids
ending in ic, while those ending in ite, have acids
ending in ous.
– ClO⁻ (hypochlorite) = HClO (hypochlorous acid)
Names and Formulas of Binary
Molecular Compounds
1. The name of the element farthest to the left
in the periodic table is written first
2. If both elements are in the same group, the
lower one is named first
3. The name of the second element is given an
–ide ending.
4. Greek prefixes are used to indicate the
number of each element (ie. Mono = 1)
Cl₂O = dichlorine monoxide