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Chapter 2: Atoms, Molecules, and Ions
Chemistry 1061: Principles of Chemistry I
Andy Aspaas, Instructor
Atomic theory of matter
• Dalton’s atomic theory
– All matter composed of indivisible atoms
• Atom: small particle, retains identity in reactions
– Element: type of matter composed of only one kind of
atom
– Compound: type of matter composed of fixed proportion
of 2 or more elements
– Chemical reaction: rearrangement of atoms to give new
chemical combinations
Deductions from Dalton’s Atomic Theory
• Law of conservation of mass
• Law of definite proportions
– Example: 1.00 g C reacted with oxygen
– 2 compounds formed
• 1.3321 g O : 1.00 g C
• 2.5542 g O : 1.00 g C
Structure of the Atom
• Dalton said atoms were indivisible, but experiments
starting around 1900 showed that atoms themselves
consist of particles
– Nucleus: atom’s central core
• Positively charged
• Contains most of atom’s mass
– Electron
• Very light, negatively charged particles
Discovery of the electron
• 1897, British physicist J.J. Thomson
– Experiments showed atoms are not indivisible
– Cathode ray tube
– Calculated ratio of electron’s mass to its charge
• 1909, American Robert Millikan
– Charge on electron: 1.602 x 10-19 coulomb
– Mass: 9.109 x 10-34 kg
Nuclear model of the atom
• 1911, British physicist Ernest Rutherford
– Gold foil experiment
– Most alpha particles pass straight through gold
foil
– Most of atom = empty space
– If a golf ball was the nucleus, the atom would be
about 3 miles in diameter
Gold foil experiment
Nuclear isotopes
• Nucleus comprised of 2 kinds or particles
– Protons
• Have positive charge equal in magnitude to
electron’s negative charge
• More than 1800 times more massive than
electron
– Neutrons
• Mass nearly identical to proton, no charge
Atomic number and mass number
• Atomic number (Z): total number of protons in the
nucleus of an atom
– New element definition: substance whose atoms
all have the same atomic number
• Mass number (A): total number of protons and
neutrons in a nucleus
• Nuclide: atom characterized by certain mass
number and certain atomic number
• Nuclide symbol: Z subscript, A superscript, element
Isotopes
• Isotopes: atoms whose nuclei have the same atomic
number but different mass numbers
• Some elements have only one naturally occurring
isotope (sodium-23)
• Some have several naturally occurring isotopes:
– Oxygen contains the following isotopes
• 99.769% oxygen-16
• 0.037% oxygen-17
• 0.204% oxygen-18
Atomic weights
• Dalton: an atom of a certain element has a
characteristic mass
• But, naturally occurring elements may be a mixture
of isotopes
– Isotope percentages are generally constant
• Dalton actually calculated average atomic masses
– They were relative to the mass of hydrogen, the
smallest element
Atomic mass units and atomic weight
• Carbon-12 isotope chosen arbitrarily as standard
– 12 atomic mass units (amu)
• Atomic weight: average atomic mass for a naturally
occuring element, expressed in amu
– Appears on periodic table along with atomic
number
Calculating atomic weight
• Fractional abundance: fraction of a total number of
atoms, which consists of a particular isotope
• Isotopic mass is not exactly equal to mass number
– Neon-20, mass = 19.992 amu, abund = 0.9051
– Neon-21, mass = 20.994 amu, abund = 0.0027
– Neon-22, mass = 21.991 amu, abund = 0.0922
• Multiply isotopic mass by fractional abundance, and
sum total to get atomic weight
• Gives a “weighted average”
Periodic table of the elements
• 1869, Dimitri Mendeleev (and J. Lothar Meyer,
independently)
– Arranged elements in order of atomic number
– Placed elements in horizontal rows so that
vertical columns of elements with similar
properties formed
Periods
• Period: elements in any one horizontal row of the
periodic table
– First has only 2
– 2nd and 3rd have 8
– 4th and 5th have 18
– 6th has 32
– 7th is incomplete
Groups
• Group: elements in any one vertical column of the
periodic table
– Are numbered with Roman numeral and letter
– (or just numbered with a number)
• Main-group elements (labeled with A)
• Transition elements (labeled with B)
• Inner transition elements (shown below rest of table)
Groups
• Elements in a single group have similar properties
– For example:
• IA: alkali metals - all react vigorously with
water (except hydrogen)
• VIIA: halogens - all react vigorously with
sodium
Metals, Nonmetals, Metalloids
• Metal: substances or mixtures that…
– Have a characteristic shine
– Generally are good conductors of heat and
electricity
– Are usually malleable (can be hammered into
sheets)
– Are usually ductile (can be drawn into wire)
– Metallic elements are solid at room temp., except
for mercury
Metals, nonmetals, metalloids
• Nonmetal: element that does not exhibit characteristics of a
metal
– Most are gases
– Solid nonmetals are usually brittle and hard
– Bromine is only liquid nonmetal
• Metalloid (or semimetal) - both metallic and nonmetallic
properties
– Many are semiconductors (silicon, germanium)
– Don’t conduct when pure, do when doped or at very high
temperature
Chemical formulas
• Chemical formula: atomic symbols with numeric
subscripts that show relative proportions of atoms of
different elements in a substance
• Al2O3: ratio of Al atoms to O atoms is 2 : 3
• No subscript: ratio of 1 (NaCl, 1:1)
Molecular substances
• Molecule: group of atoms that are chemically
bonded together - tightly connected by attractive
forces
• Molecular substance: composed of all the same
molecules
• Molecular formula: shows number of different atoms
of an element in a molecule
• Structural formula: shows how atoms are bonded
together, a line indicates a chemical bond
Molecular substances
• Some elements are molecular substances
– Cl2, O2, N2, S8
• Some exist as individual atoms
– He, Ne
• Some exist as a very large but indefinite number of atoms
bonded together
–C
• Polymer: extremely large molecules made of small molecules
repeatedly bonded together (monomers)
– Natural (wool, cotton), synthetic (plastics, Nylon,
polyester, Kevlar)
Ionic substances
• Unlike molecular substances, ionic substances are formed
from charged atoms or groups of atoms called ions
– Metal atoms tend to lose electrons: result is a positive
charge (cation)
– Nonmetals tend to gain electrons: result is a negative
charge (anion)
• Sodium loses one electron to become Na+
• Calcium loses two electrons to become Ca2+
• May consist of a group of bonded atoms that have a
deficiency or surplus of electrons (SO42+)
Ionic compounds
• Ionic compounds are composed of cations and
anions
• Ionic formula: given by smallest possible integer
number of different ions in substance
– Charges are omitted
– Ionic compounds are uncharged as a whole
– Ex. Na+ and Cl- = NaCl
– Fe3+ and SO42- = Fe2(SO4)3
Naming simple compounds
• Chemical nomenclature: systematic naming of
chemical compounds
• Organic compounds: carbon-containing molecular
substances
• Inorganic compounds: composed of elements other
than carbon (except CO, CO2, cyanides, and
carbonates)
Ionic compounds
• Named given by name of cation followed by name of
anion
– Potassium sulfate
• Monatomic ion: formed from a single atom
Predicting a monatomic ion’s charge
• Most main-group metallic ions have one monatomic cation
with charge equal to group number
– Aluminum, Group IIIA, Al3+
• Many high-atomic-number metals have a cation equal to
group number AND one equal to the group number minus 2
– Thallium, Group IIIA, Tl3+, Tl+
• Most transition elements can have several monatomic
cations. +2 is very common
• A nonmetal main-group ion forms an anion
– Charge = 8 - group #
– Oxygen, Group VIA, O2-
Naming monatomic ions
• Monatomic cations are named after element if
there’s only one possible cation for that element
• If more than one possible cation, use charge as
roman numeral after name
– Fe2+ = iron(II), Fe3+ = iron(III)
– Or use suffixes with latin name (-ous for lower
charge, -ic for higher)
• Cu2+ = cuprous, Cu3+ = cupric
• Monatomic anions use -ide suffix (bromide, sulfide)
Polyatomic ions
• Polyatomic ion: ion consisting of 2 or more atoms
chemically bonded together, which carry a net
electric charge
• Oxoanions: central element + some oxygens
– Suffixes indicate relative amount of oxygens
– Fewer oxygens: -ite
– More oxygens: -ate
– Ex. SO32- is sulfite, SO42- is sulfate
Polyatomic anions
• If more than 2 oxoanions, use prefixes
– hypo- and per• ClO- hypochlorite ion
• ClO2- chlorite ion
• ClO3- chlorate ion
• ClO4- perchlorate ion
Binary molecular compounds
• Name first element with its exact
name
• Suffix second element with -ide
• Use Greek prefixes to indicate
number
• Don’t use mono unless there are
more than one possible amount
– Ex. CO vs CO2
Acids and corresponding anions
• Acids have H+ as the cation
• Oxoacids contain an oxoanion
– -ate becomes -ic
– -ite becomes -ous
– End name with “acid”
• Binary compounds of hydrogen and nonmetal
– Prefix hydro– Suffix -ic acid
Hydrates
• Simply indicate how many water molecules are
present by using the Greek numerical prefix and hydrate at the end
• MgSO4 • 7H2O = magnesium sulfate heptahydrate
Writing chemical reactions
• Reactant ----> Product
• Arrow means “reacts to form” or “yields”
• Useful to indicate state or phase
– (g) = gas, (l) = liquid, (s) = solid,
– (aq) = aqueous solution
• Use a coefficient to indicate relative number of particles
involved
• ∆ over arrow means heat is applied
• A compound written over the arrow is usually a catalyst
Balancing chemical equations
• Mass must be conserved so use coefficients to
make sure the same number of each atom occurs
on each side of the equation
• Start by balancing atoms for elements that occur in
only one substance on each side
– Ex: H3PO3 ---> H3PO4 + PH3
– Start by balancing oxygen
– Practice!