Atoms, Molecules and Ions
Dalton’s Atomic
Theory (1808)
All matter is composed of extremely
small particles called atoms
Atoms of a given element are identical
in size, mass, and other properties;
atoms of different elements differ in
John Dalton
size, mass, and other properties
 Atoms cannot be subdivided, created, or destroyed
(Law of conservation of mass)
 Atoms of different elements combine in simple
whole-number ratios to form chemical compounds
(Law of constant composition)
 In chemical reactions, atoms are combined,
separated, or rearranged
Discovery of the Electron
In 1897, J.J. Thomson used a cathode ray tube
to deduce the presence of a negatively charged
particle.
Cathode ray tubes pass electricity through a gas
that is contained at a very low pressure.
Thomson’s Atomic Model
Thomson believed that the electrons were like
plums embedded in a positively charged
“pudding,” thus it was called the “plum
pudding” model.
Mass of the Electron
1909 – Robert
Millikan determines
the mass of the
electron.
The oil drop apparatus
Mass of the
electron is
9.109 x 10-31
kg
Conclusions from the Study
of the Electron
Cathode rays have identical properties
regardless of the element used to
produce them. All elements must contain
identically charged electrons.
Atoms are neutral, so there must be
positive particles in the atom to balance
the negative charge of the electrons
Electrons have so little mass that atoms
must contain other particles that
account for most of the mass
Rutherford’s Gold Foil
Experiment
Alpha particles are helium nuclei
Particles were fired at a thin sheet of
gold foil
Particle hits on the detecting screen
(film) are recorded
Rutherford’s Findings
 Most of the particles passed right through
 A few particles were deflected
 VERY FEW were greatly deflected
“Like howitzer shells
bounding off of tissue
paper!”
Conclusions
The nucleus is small
The nucleus is dense
The nucleus is positively
charged
Atomic Particles
Particle
Electron
Proton
Neutron
Charge
Mass (kg)
Location
-1
9.109 x 10-31 Electron
(1/1836 amu) cloud
+1
1.673 x 10-27
(1 amu)
Nucleus
0
1.675 x 10-27
(1 amu)
Nucleus
The Atomic Scale
 Most of the mass of the atom is in
the nucleus (protons and neutrons)
 Electrons are found outside of the
nucleus (the electron cloud)
 Most of the volume of the atom is
empty space
Modern Atomic Theory
Several changes have been made to
Dalton’s theory.
Dalton said:
Atoms of a given element are identical in size,
mass, and other properties; atoms of
different elements differ in size, mass, and
other properties
Modern theory states:
Atoms of an element have a characteristic
average mass which is unique to that
element.
Modern Atomic Theory #2
Dalton said:
Atoms cannot be subdivided, created, or
destroyed
Modern theory states:
Atoms cannot be subdivided, created, or
destroyed in ordinary chemical reactions.
However, these changes CAN occur in
nuclear reactions!
Atomic Number
Atomic number (Z) of an element is the
number of protons in the nucleus of each
atom of that element.
Element
Carbon
# of protons
Atomic # (Z)
6
6
Phosphorous
15
15
Gold
79
79
Mass Number
Mass number is the number of protons and
neutrons in the nucleus of an isotope.
Mass # = p+ + n0
Nuclide
p+
n0
e-
Mass #
Oxygen – 18
8
10
8
18
Arsenic – 75
33
Phosphorus –
31
15
75
16
Isotopes
Isotopes are atoms of the same element having
different masses due to varying numbers of neutrons.
Isotope
Protons Electrons Neutrons
Hydrogen-1
(protium)
1
1
0
Hydrogen-2
(deuterium)
1
1
1
1
1
2
Hydrogen-3
(tritium)
Nucleus
Atomic Masses
Atomic mass is the average of all the naturally
isotopes of that element. Carbon
Carbon== 12.011
12.011
Isotope
Symbol
Composition of
the nucleus
% in nature
Carbon-12
12C
6 protons
6 neutrons
98.89%
Carbon-13
13C
6 protons
7 neutrons
1.11%
Carbon-14
14C
6 protons
8 neutrons
<0.01%
Calculating Average Atomic
Mass
• Average atomic mass = S(% of each isotope)(mass of each isotope)
100
• Example: If 69Ga and 71Ga occur in the
percents 62.1 and 37.9, calculate the
average atomic mass of gallium atoms.
Molecules
Two or more atoms of the same or different
elements, ___________ bonded together.
Molecules are discrete structures, and their
formulas represent each atom present in the
molecule.
Benzene, C6H6
Examples of Molecules
Examples of Molecules
– Substance and
Formula
Element or
Compound
Description
Hydrogen (H2); Oxygen
(O2); Nitrogen (N2);
Fluorine (F2); Chlorine
(Cl2); Iodine (I2); Bromine
(Br2)
Elements
Diatomic
Water (H2O)
Compound
Polyatomic
Ammonia (NH3)
Compound
Polyatomic
Covalent Network
Substances
Covalent network substances have covalently
bonded atoms, but do not have discrete
formulas. Groups of these substances are
called __________.
Why Not??
Graphite
Diamond
Ions
• Atoms have equal numbers of protons
and electrons so they have ________
• When atoms gain or lose electrons,
the proton:electron numbers are
uneven causing a charged particle
called an ____
Ions
Cation: A ________ ion
Mg2+ (monatomic), NH4+ (polyatomic)
Anion: A _______ ion
Cl- (monatomic), SO42- (polyatomic)
Ionic Bonding: Force of attraction
between oppositely charged ions.
Ionic compounds form crystals, so their
formulas are written _________ (lowest
whole number ratio of ions).
Predicting Ionic Charges
Group 1: Lose 1 electron to form 1+ ions
H+ Li+ Na+ K+
Group 2: Loses 2 electrons to form 2+
ions
Be2+ Mg2+
Ca2+ Sr2+ Ba2+
Predicting Ionic Charges
Group 13: Loses 3 electrons to form 3+
ions
B3+ Al3+ Ga3+
Group 14: Loses 4 electrons or gains 4
electrons
Caution! C22- and C4- are both called
carbide
Predicting Ionic Charges
Group 15: Gains 3 electrons to form 3- ions
N3P3As3Nitride
Phosphide Arsenide
Group 16: Gains 2 electrons to form 2- ions
O2S2Se2Oxide
Sulfide
Selenide
Predicting Ionic Charges
Group 17: Gains 1 electron to form 1ions
F1Cl1Br1I1Fluoride Chloride Bromide Iodide
Group 18: Stable Noble gases do not
form ions!
Predicting Ionic Charges
Groups 3-12: Many transition elements have
more than one possible oxidation state
(charge).
Iron(II) = Fe2+ Iron(III) = Fe3+
Some transition elements have only one
possible oxidation state.
Zinc = Zn2+
Silver = Ag+
Binary Compounds of Nonmetals
and Metals (Ionic Compounds)
To write the formula from the name
• Write the formulas for the _____ and
_____, including CHARGES!
2. Check to see if charges are balanced.
3. Balance charges , if necessary, using
________. Use parentheses if you
need more than one of a polyatomic
ion.
Writing Ionic Compound
Formulas
Example: Barium nitrate
Example: Ammonium sulfate
Example: Iron(III) chloride
Example:
Example:
Example:
Example:
Aluminum sulfide
Magnesium carbonate
Zinc hydroxide
Aluminum phosphate
Naming Ionic Compounds
1. Cation first, then anion
2. Monatomic cation = name of the element
Ca2+ = calcium ion
3. Monatomic anion = root + -ide
Cl- = chloride
CaCl2 = calcium chloride
Naming Ionic Compounds
Metals with multiple oxidation states
 Some metal forms more than one cation
 Use Roman numeral in name
PbCl2
Pb2+ is the lead(II) cation
PbCl2 = lead(II) chloride
Binary Acids
• An acid is a compound that produces
hydrogen ions in water
• Formed when hydrogen ions combine
with monatomic anions
• To name an acid
– Use the prefix hydro followed by the
non-metal name modified to an –ic ending
– Add the word acid
– Example: HCl and HI
Polyatomic Ions and Oxyanions
• Polyatomic ion – more than one element are
combined to create a species with a charge
– Have to memorize these
• Oxyanions – a polyatomic ion that contains
oxygen
• To name a compound containing a
polyatomic ion
– Positive ion written first
– Polyatomic ion written second, unmodified
– Example: K2CO3
Oxyacids
• Formed when hydrogen ions combine
with polyatomic oxyanions
• To name an oxyacid
– Use the name of the oxyanion and
replace the –ite ending with –ous or the
–ate ending with –ic
– Add the word acid
– Example: HClO4 and HNO3
Naming Binary Compounds of
Two Nonmetals
 Two nonmetals form molecules
 First element in the formula is named first.
 Second element is named as if it were an
anion.
 Use prefixes – mono, di, tri, tetra, penta,
hexa, hepta and octa
 Only use mono on second element P2O5 = diphosphorus pentoxide
CO2 = carbon dioxide
CO =carbon monoxide
N2O = dinitrogen monoxide
Hydrates
• Ionic formula units with water associated
with them
• Water molecules are incorporated into the
solid structure of the ions
• To name a hydrate
– Use the naming of an ionic compound followed
by the term hydrate with an appropriate prefix
for the hydrate
– Example: CuSO4.5H2O