Download CHEM A Note Guides Unit 2

Note Guide 2.1 Defining the atom
•Atom = smallest particle of an element that keeps its identity in a
chemical reaction
all matter made from atoms
not visible to the naked eye
•Democritus = 1st to say “Hey, what about atoms?”
1st to suggest the existence
believed they were indivisible and indestructible
Ideas did line up with later ideas on atoms, with exceptions:
A. Did not explain chemical behavior
B. Did not have basis of scientific research…best guess!
•Dalton took Democritus idea, applied the scientific method (through
experimentation) and devised 4 ideas to his theory
1.
2.
All elements made up of tiny indivisible particles (atoms)
Atoms of the same element are identical…..
EX: gold atom looks just like the gold atom next to it
…..but atoms of one element are different that atoms of another
element
3.
Atoms of different elements can physically mix together or can
chemically combine in simple whole number ratios to form
compounds
Chemical reactions take place when atoms get separated, joined, or
rearranged.
BUT…atoms of one element are NEVER changed into atoms of
another element because of chemical reaction.
4.
Dalton’s Theory
•Size: small 
typical radii of atom ( radius = center to edge)
5 x 10-11
2 x 10-10
See em???
Naked eye = HA HA HA , right!
Microscope = Yes (scanning tunneling microscope)
Note Guide 2-2 Structure of the Atom
3 Subatomic Particles
1. Electrons  negatively charged particles
--discovered by Thomson (Plum pudding model) using a cathode ray tube
(see pg 104)
--each electron carries one unit of negative charge and its mass is
1/1840th the mass of a hydrogen atom
makes mass of electron about 0
--symbol = e–
--actual mass = 9.11 x 10-28
--found in the energy levels
2. Protons  positively charged particles
--Eugene Goldstein provided evidence using cathode ray tube
--Relative mass = 1
--Actual mass = 1.67 x 10-24
--symbol = p+
--found in the center of the atom
3. Neutrons  neutral charged particles
--Chadwick confirms the existence
--relative mass = 1
--actual mass = 1.67 x 10-24
--symbol = n°
--found in the center of the atom.
Nucleus  center of the atom which houses protons and neutrons
--idea of nucleus comes from Rutherford who did the gold foil exp.
(pg. 107) This experiment involved sending a narrow beam of alpha
particles (Helium atoms – 2 electrons giving a double + charge)
through a thin sheet of gold foil and observing what the particles
did.
RESULT: majority of particles passed through, some deflected at
large angles, and some bounced right back.
Nuclear Atom (Rutherford model)
--electrons distributed around the nucleus and occupy almost all the
volume of the atom.
--nucleus very small in comparison with rest of atom
EX: nucleus would be the size of a marble on a football field ,
which would be the rest of the atom.
--Model turns out to be incomplete.
Draw table 4.1 on page 106 into notes
Note guide 2-3 Distinguishing among atoms
•Atomic Number  tells the # of protons in the atom.
+each element (atom) has its own unique atomic #, which
distinguishes it from others.
EX: Hydrogen has 1 proton and Helium has 2. Hydrogen
will always have 1 and can never have 2 or more.
•Mass number  total # of protons + neutrons in an atom
--we can calculate mass number by adding protons or neutrons in
an atom.
--can show mass number in shorthand notation
197 -------------- mass number
Au
79 -------------- atomic number
•Isotopes  atoms with different numbers of neutrons compared with the
protons of the atom
--KEY: different number of neutrons DOES NOT change the
atom (gold isotope is still gold)
EX: Neon
--atom = 10p+ 10n0 10e--isotope = 10p+ 12n0 10ewe name this Neon-22 (10p+ + 12n0 )
•Atomic Mass  1/12 the mass of a carbon-12 atom.
--reference isotope used to compare one mass to another mass
--because we use carbon-12 (6 p+ and 6n0) as the standard, we say
that a single p+ and n0 is 1 AMU (atomic mass unit)
--However the atomic mass of an element is not usually a whole #?
Why?
Because most elements occur as a mixture of 2 or more
Isotopes (p+  n0)
--because isotopes are in the picture, whole number ratios rare
--atomic mass takes the weighted average mass of the atoms in a
naturally occuring sample of the element.
this includes mass and how much of the element there is.
CALCULATING ATOMIC MASS
--need to know 3 things
1. Number of stable isotopes of the of the element
2. Mass of each isotope
3. Natural percent abundance of each isotope
--Atomic mass = mass of each isotope X natural abundance
(decimal) + the products
•Periodic Table Preview
--arrangement of elements in which the elements are
separated into groups based on a set of repeating
properties.
--allows comparison of properties of one element (or
groups of elements) to another element (or group of
elements)
--broken into:
1. Rows (horizontal) called Periods
--7 periods
--properties vary as you move across
2. Columns (vertical) called groups (families)
+ properties vary as you move down
+ characteristics and properties are similar within groups
CONCEPTUAL PROBLEM 4.2 PG. 113
-- DO 19 AND 20
Note guide 2-4 Nuclear Radiation
•Radioactivity: process by which nuclei emit particles and rays (rays
called radiation)
--caused by nuclear reactions, which differ from chemical reactions
How?
1. In chem. Reactions, atoms tend to look to gain stability (full
electron shell by losing or gaining electrons. In nuclear
reactions, the nuclei of unstable isotopes (radioisotopes)
gain stability by undergoing changes.
--these changes emit large amounts of energy
2. Nuclear reactions are not affected by changes in temp.,
pressure, or the prescence of catalysts. Chemical
reactions affected by these.
--discovery of radioactivity blows away Dalton’s idea of atoms
being indivisible
--radioisotopes have unstable nuclei
stability of a nucleus depends on the ratio of neutrons to
protons and the size of the nucleus. The bigger the spread,
the more unstable the nucleus.
--unstable nuclei emit radiation during radioactive decay
3 types of radiation:
1. Alpha Radiation: consists of helium nuclei that have been emitted.
--particles called alpha particles
2. Beta Radiation: electron resulting from the breaking apart of a
neutron, turning into a proton which remains in the
nucleus and a fast moving electron is released.
--WAIT…WAIT…TIME OUT. You said a new proton cannot
be created, you can never add or subtract to change the atomic #
--Yes but, unstable radioisotopes are the only exception
to the never rule
3. Gamma Radiation: high energy photon emitted by a radioisotope
(gamma ray)
--emission of gamma radiation does not alter the atomic
number or mass
--very dangerous rays.
Note guide 2-5 Nuclear transformations
•Nuclear Force = attractive force; acts between ALL nuclear particles that
are extremely close together (protons and neutrons)
--if no force in the nucleus, the like charges would repel, then break
apart.
•Stability = 264 nuclei are stable, more than 1230 are not, therefore they
decay over time
•Type of decay caused by neutron to proton ratio
--if too many neutrons (to protons), these decay by turning a neutron
into a proton beta particle = Beta emission
--if not enough neutrons (to protons), these increase stability by
turning protons into neutrons
--all nuclei that have an atomic number greater than 83 are radioactive.
--if too many protons and neutrons, they undergo radioactive decay.
Most emit alpha particles = alpha emission
•Half life = time required for ½ of the nuclei of a radio isotope to decay to
products.
--after each half life, half of the existing radioactive atoms have
decayed into atoms of a new element.
--half lives can be short (fraction of a second) to very long term (billions
of years.