Download AP - 02 - Atoms Molecules and Ions

AP / College Chem
Chapter 2 – Atoms, Molecules & Ions
- Atomic Theory of Matter:
o Democritus (400 B.C.) was a Greek philosopher who came up with the
first concept of the atom (atomos – indivisible)
o Plato and Aristotle did not believe in atoms (no limit to how small you
could divide matter)
o John Dalton (1803 – 1807) came up with the modern atomic theory
 Each element is composed of extremely small indivisible particles
called ATOMS
 All atoms of the same element are identical but atoms of different
elements are different from each other
 Atoms of one element cannot be changed into atoms of a
different element by chemical reactions (ATOMS are neither
created or destroyed!)
 Compounds are formed when atoms of a more than one element
combine and a given compound always has the same relative
number and kind of atoms (law of definite composition)
o Explains the law of conservation of matter (the total mass of the
reactants must equal the total mass of the products)
o Dalton also used his theory to develop the law of multiple proportions
(if two elements A and B combine to form more than one compound,
the masses of B that combine with a given mass of A are in a ratio of
small whole numbers)
- Discovery of Atomic Structure:
o We know today that atoms are not indivisible like Dalton thought but
are composed of subatomic particles (protons, neutrons and electrons)
o Particles with the SAME charge repel each other and particles with
OPPOSITE charges attract each other
o Cathode Rays and Electrons
 mid 1800s scientist were studying cathode ray tubes (glass tubes
filled with a gas that had a cathode (-) and anode (+) end)
 when high voltage was applied to the electrodes, radiation known
as CATHODE RAYS were produced that would cause certain
materials to emit light (FLUORESCE)
 cathode rays discovered by Johann Hittorf and Eugen Goldstein
named cathode rays in 1876
 experiments showed that cathode rays were deflected by electric
or magnetic fields in a way that was consistent with their being a
stream of negative electrical charge
 In 1897, J.J. Thompson describe cathode rays as a stream of
negatively charged particles that behaved exactly the same no
matter what material was used for the cathode
 Credited with the discovery of the electron
Cathode rays (electrons) begin at the
cathode and are accelerated toward
the anode which has a hole in the
center. This produces a beam of
electrons that travel toward a
fluorescent screen. The strengths of
the electric and magnetic fields are
adjusted so their effects cancel each
other and allow the beam to travel
along a straight path (to calculate the
charge to mass ratio
 Thompson also found the electrical charge to mass ratio of an
electron to be 1.76 x 108 Coulombs/gram
 Robert Millikan (1909) measured the charge of an electron via the
Oil-Drop Experiment
Small drops of oil fall between
electrically charged plates and gain
electrons after exposure to X-rays.
Millikan varied the voltage between
the plates and measured the rate of
fall. From this data he calculated the
charge of a single electron to be
1.602 x 10-19 Coulombs (C).
 Combining data from these two experiments, the mass of an
electron was calculated
1.602 x 10-19 C
Electron mass = ------------------------ = 9.10 x 10-28 g
1.76 x 108 C/g
 This result means the electron is almost 2000 times smaller in
mass than a Hydrogen atom!!
o Radioactivity
 Henri Becquerel (1896) discovered that Uranium spontaneously
emits high-energy radiation called RADIOACTIVITY
 Marie & Pierre Curie worked on isolating the radioactive
components of that compound that Becquerel worked on
 Ernest Rutherford determined 3 different types of radiation
(alpha, beta and gamma) and showed that alpha and beta rays
consisted of fast moving particles
 Alpha (He nucleus and has a +2 charge), Beta is fast moving
electrons and are (-) charged and gamma have no charge and are
similar to X-rays
o Nuclear Model of the Atom:
 J.J. Thompson came up with the Plum Pudding Model of the
atom (atom was a mass of (+) charge with (-) electrons dispersed
throughout like plum pudding)
 Ernest Rutherford (1910) studied angles at which alpha particles
were deflected or scattered through a thin sheet of gold foil
 He had his assistant (Marsden) look for greater angles and
surprisingly found some particles scattered backwards which was
totally inconsistent with Thompson’s model
 Rutherford came up with the NUCLEAR MODEL in which the
protons are located in a small, dense mass called the nucleus and
the negative electrons located in the mostly empty space outside
of the nucleus
 James Chadwick (1932) discovered the neutron (neutral particle)
- The Modern View of Atomic Structure:
o Every atom is made of protons and neutrons in the nucleus of the atom
and electrons located outside the nucleus in mostly empty space
o The charge of an electron is assigned a value of (-1) and a proton (+1)
making them equal in charge but opposite in sign (even though protons
are MUCH more massive than electrons are the magnitude of the
charges are the same!!)
o Every atom has the SAME number of protons and electrons making the
atom electrically NEUTRAL
o Atomic mass unit (AMU) – the mass unit of the atom
o 1 amu = 1.66 x 10-24 grams
o 1 proton is 1.0073 amu
o 1 neutron is 1.0087 amu
o 1 electron is 5.486 x 10-4 amu (so that means that it takes 1,836
electrons to equal the mass of a proton!)
o Atomic size is usually measured in ANGSTROMS (Å) which is 1 x 10-10 m
o Atoms have diameters anywhere from 1 – 5 Å
o Electromagnetic Forces are determined by COULOMB’S LAW
F is the electromagnetic force of
attraction (-) or repulsion (+) between
2 charged particles. It is proportional
to the magnitude of the charges and
inversely proportional to the square
of the distance (r) between the
centers of the charged particles
o Atomic Numbers, Mass Numbers & Isotopes
 Atomic number (Z) – the total number of protons in the nucleus of
the atom
 The atomic number identifies an element
 Mass number (A) – the total number of protons AND neutrons in
the nucleus of the atom
 A = # protons + # neutrons
 # neutrons = A – Z
 The mass number identifies an ISOTOPE (different form of the
same element that has a different mass because of a different
number of neutrons)
 There are at least 2 isotopes (sometimes more) for each element
- Atomic Weights
o Atomic Mass Scale
 Atoms of different elements have different masses
 Early scales based on H having a mass of 1, then O having a mass
of 16
 After discovering isotopes it became clear that an isotope needed
to be the standard for atomic mass so C-12 was used and given a
mass of 12.0000 amu
 Average atomic mass is calculated as a weighted average
Calculate the average atomic mass of Chlorine based on the following data
35
Cl is 75.78% abundance with a mass of 34.969 amu
37
Cl is 24.22% abundance with a mass of 36.966 amu
Average mass = (0.7578)(34.969 amu) + (0.2422)(36.966 amu)
= 26.50 amu + 8.953 amu
= 35.45 amu
 Relative abundance and masses of isotopes are determined by a
MASS SPECTROMETER
 Produces a MASS SPECTRUM that shows mass and abundance
 Used to identify chemical compounds and analyze mixtures of
substances
- The Periodic Table:
o Most significant tool that chemists use for organizing and remembering
chemical facts
o When elements are arranged in order of increasing atomic number,
their chemical and physical properties show a repeating (periodic)
pattern
o PERIODS – horizontal rows (7) correspond to the 7 main energy levels of
electrons in the atom
o GROUPS – vertical columns (18) correspond to a family of elements that
share similar chemical and physical properties mostly because their
outer energy level electron structure is the SAME as all other members
of that group or family
o Metals – located at the left side and middle of the periodic table
o Nonmetals – located at the right side of the periodic table
o Metalloids – semiconductors or hybrids that are located along a
stepwise dividing line in the periodic table (B, Si, Ge, As, Sb, Te)
o Metalloids have some properties of metals and some of nonmetals
- Molecules and Molecular Compounds:
o Molecules have 2 or more atoms combined together by covalent bonds
(sharing pairs of electrons)
o Molecules may be elements (DIATOMIC ELEMENTS – H2, O2, N2, Cl2, Br2,
I2, F2)
o Molecules may also be molecular compounds that are sharing electrons
between atoms of different elements (H2O, NH3, CO2, etc)
o Molecules are almost always made of ONLY NONMETALS (and
sometimes metalloids)
o Molecules are always identified by their MOLECULAR FORMULA (shows
the exact number of each atom in one molecule) that is NEVER reduced
as a formula
o Empirical formula – shows a chemical formula with the SMALLEST
WHOLE NUMBER RATIO of elements in the formula (may be reduced)
o Empirical formulas are NOT typically used for molecules
- Ionic Compounds:
o Ionic compounds are made from the attraction of (+) and (-) ions which
are formed by a TRANSFER of electrons from a metal to a nonmetal
atom
o Results in FULL charges and is a much stronger bond than a covalent
bond
o Periodic table is very helpful in determining the CHARGE on a
monatomic ion (can figure out from what group the element is in)
o Ions in ionic compounds are arranged in 3-D structures called a crystal
lattice and therefore only EMPIRICAL formulas are used to show the
smallest whole # ratio of ions in the compound
o Each Na+ in sodium chloride is surrounded by 6 Cl- and each Cl- is
surrounded by 6 Na+ so we actually have a 1:1 ratio  NaCl
o Although ionic compounds are made up of (+) and (-) ions, the
compounds themselves have a total charge of ZERO and are NEUTRAL
o That means the (+) and (-) charge has to balance so when figuring out
the empirical formula for an ionic compound we CRISS-CROSS the
charges to become the subscripts on the opposite ion and then reduce
the formula (if needed)
o If you have more than one unit of a polyatomic ion it must be placed in
PARENTHESIS with the subscript after the parenthesis
o No parenthesis are needed if your subscript is ONE on a polyatomic ion
- Naming Inorganic Compounds:
o Chemical nomenclature is used for naming compounds
o System of assigning a unique name to every individual compound
o Some compounds are known for their common names (NH3 – ammonia)
o Organic molecules have their own system for naming
o Inorganic based on the following system and depend on the type of
compound that is involved
o Names and Formulas of IONIC compounds:
 Naming Ionic Compounds is easy – they are all made of (+) and (-)
ions so to name any ionic compound--Name the (+) ion first and
the (-) ion last
 The trick is knowing HOW to name the IONS!!!
 CATIONS formed from metal atoms have the same name as the
metals plus the word “ION”
 If a metal can form CATIONS with different charges (most
transition metals and post-transition metals) then the (+) charge is
indicated by a ROMAN NUMERAL in parentheses following the
name of the metal
 Old system used suffixes (“-IC” for HIGHER charge and “-OUS” for
LOWER charge) and the Latin root name of the metal
 CATIONS formed from nonmetal atoms that end in “-IUM”
COMMON CATIONS
 ANIONS of monatomic ions are named by replacing the ending of
the element name with “-IDE”
 A few POLYATOMIC IONS also end in “-IDE”
 POLYATOMIC IONS that contain O (OXYANIONS) have the endings
“-ATE” (MORE) and “-ITE” (LESS) to indicate how many oxygen
atoms the ion contains
 When there is more than just 2 possibilities of number of O atoms
then the prefixes “PER-“ (ABOVE) and “HYPO-“ (UNDER) are used
Procedure for Naming Anions
 Anions derived by adding H+ are named by adding the prefix
hydrogen or dihydrogen. The old method uses the prefix “bi-“
Names of Common Anions
 Names of IONIC COMPOUNDS are named by naming the (+) ion
then the (-) ion
o Names and Formulas of ACIDS:
 Acid is a molecule that produces H+ when dissolved in water
 Generic formula: HA
 Acids containing anions whose names end in “-IDE” are named by
using the prefix “HYDRO-“ then the root of the anion and change
the ending to “-IC ACID”
 Acids containing anions whose names end in “-ATE” or “-ITE”
(OXYACIDS) are named by changing ATE  IC and ITE  OUS and
then adding the word ACID
o Names and Formulas of BINARY MOLECUES:
 Name the element that is farther LEFT on the Periodic Table
(except when O is combined with a HALOGEN, the halogen is
named first)
 If both elements are in the same group, the LOWER one is named
first
 The name of the second element is given the “-IDE” ending
 Greek prefixes indicated the number of atoms of each element in
ONE MOLECULE. (Mono- is never used at the start of the name!)
Prefixes used for Molecules:
- Some Simple Organic Compounds:
o Organic chemistry—study of compounds containing Carbon
o Also contain H, and sometimes O, N, or halogens
o Alkanes
 Hydrocarbons—contain only C and H
 Alkanes – simplest hydrocarbons containing only C—C single
bonds
 Alkanes are SATURATED hydrocarbons
 Names end in “-ANE”
o Some Derivatives of Alkanes
 Alcohol—has the functional group –OH in the formula
 Names end in “-OL
 Compounds with the same molecular formula but different
structural formulas are called ISOMERS
- Suggested Problems:
2.1, 2.2, 2,3 2,5, 2,7, 2.9, 2.11, 2.12, 2.13, 2.15, 2.17, 2.19, 2.20, 2.21, 2.22, 2.23,
2.25, 2.27, 2.29, 2.31, 2.35, 2.37, 2.41, 2.42, 2.43, 2.45, 2.47, 2.49, 2.50, 2.51,
2.53, 2.55, 2.57, 2.59, 2.61, 2.63, 2.65, 2.67, 2.69, 2.71, 2.74, 2.77, 2.79