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CHAPTER 2 ATOMS, MOLECULES AND IONS
6/2015
ATOMIC THEORY OF MATTER:
1. Democritus 400BC Greek
Through logical deduction theorized that matter
is made up of small particles that cannot be
divided.
Called these particles “atomos”
His idea was dismissed until the 1700’s
2. John Dalton’s Atomic Theory, 1805
- Based on his reading and observations in the
lab
Dalton’s Atomic Theory of Matter ( 1803)
a. Each element is composed of very small,
indivisible particles called atoms (Democritus)
b. All atoms of an element are identical, but
they differ from the atoms of other elements.
c. Atoms are neither created or destroyed in
chemical reactions (Law of Conservation of
Matter /Mass– Antoine Lavoisier)
d. A given compound always contains the same
elements in the same proportions, by mass.
(Law of Constant Composition , or
Law of Definite Proportions, John Proust,
1799)
Dalton used his theory to propose the
Law of Multiple Proportion: If two elements, (ex
A & B) form more than one compound, then the
ratio of the second element that reacts with a
fixed amount of the first element will always be
in ratios of small whole numbers.
Ex H2O and H2O2
H2O For every 2 g H : 16 g O
H2O2 For every 2 g H : 32 g O
The ratio of oxygen in each compound is 1:2
More advanced technology and application of
knowledge of electrical charges were used to explore
the structure of atoms.
 Two types of electrical charges , + and –
 Like charges repel
 Opposite charges attract
a. Cathode Ray Tubes: closed glass tubes with
electrons at each end, and containing little
air.
* when electricity is applied , radiation was
produced that travelled from the
cathode (- end) to the anode (+ end), causing
the gas in the tube to glow.
3. JJ Thomson, 1897
Found he could move a cathode ray in a cathode
ray tube using a magnet or electrical field
Concluded that the cathode ray was made up of
a stream of neg charged particles with a mass
smaller than hydrogen, he named them
electrons.
This was a breakthrough in that it was realized
that subatomic particles exist.
Thompson published a mass: charge ratio for the
electron of 1.76 x 10-28 C/g
He developed the Plum Pudding model of atom
– the atom was a large, positively charged
sphere, with negative charges spread
throughout .
4. Robert Millikan, 1909
Determined charge and mass of an electron
using an oil drop experiment
1.60 X 10-19 coulomb, 9.11 x 10-28 g
5. Henri Becquerel, 1896
Discovered radioactivity – the spontaneous
emission of radiation from the nucleus of an
atom. (He observed Uranium)
6. Later , in the early 1900’s Ernst Rutherford was
able to separate radiation into 3 types using a
focused stream of radiation, charged electrical
plates, and a detection screen.
Alpha particles ( α ) with a charge of +2
Beta particles (β) with a charge of -1
Gamma radiation (α), no charge; no mass
7. Rutherford and Associates, 1909
Gold foil experiment
http://phet.colorado.edu/en/simulation/rutherfordscattering
http://www.youtube.com/watch?v=XBqHkraf8iE
Evidence
most alpha particles
went straight through foil
Conclusion
the atom is mostly
empty space
some alpha particles
bounced nearly straight
back
the atom contains a
dense, positively
charged nucleus
Rutherford’s model of an atom: Rutherford and his
associates concluded that (1) most of an atom’s
mass and all of its positive charge is in a small core
(the nucleus) and (2) most of the atom is empty
space.
8. Rutherford discovered protons in 1919
9. James Chadwick discovered neutrons in 1932.
Although subatomic particles smaller than protons,
neutrons, and electrons exist, chemical behaviour
can be explained by considering these three
particles.
STRUCTURE OF ATOMS – THE BASICS
ATOM: the smallest particle of an element that
shows the chemical properties of that
element.)
Three main sub particles make up an atom
Particle
Proton
Charge
+1
Mass
1 amu
Location
nucleus
Neutron
0
1 amu
nucleus
Electron
-1
.0006 amu
(~0)
orbitals
around
nucleus
Atomic mass units (amu)
* a unit of mass equal to 1/12 the mass of a
Carbon-12 atom (which has a mass of
12 amu)
1 amu ~ 1.66 x 10-24 g
protons & neutrons have masses of ~ 1 amu
Atomic number: the # protons in its nucleus
The number of protons in the nucleus
determines the element.
Atomic mass = # protons + # neutrons
(the number of particles in the nucleus)
Most of the mass of an atom is in the nucleus
Single, unreacted atoms are neutral, meaning
the number of protons equals the number
of electrons, so the opposite charge
values cancel out:
Ex. Carbon 6 + (protons)
6 – (electrons)
0 no overall charge
The number of neutrons in atoms of one element
can vary.
Isotopes : naturally occurring versions of an element
that vary in the numbers of neutrons in their
nucleus.
 Isotopes of an element will vary in mass
because some have more or fewer
neutrons.
 Isotopes of an element have the same
chemical properties since these are
determined by the number of protons and
electrons.
 symbols for isotopes may be written 2
ways:
 name – mass Ex: Chlorine – 35
symbol - mass
Cl-35
atomic mass
atomic #
Symbol
14
6C
Average atomic mass: the “weighted” average of all
naturally occurring isotopes of an element
periodic table lists average atomic masses
avg atomic mass is close to the most common
isotope
Calculating Average Atomic Mass:
Steps to solve the problem:
1. Read problem: What is the average atomic
mass of chlorine, if 75% of chlorine atoms
have a mass of 35 and 25% have a mass of
37?
2. Organize your data into a table, matching
mass and % abundance of each isotope
Isotope mass
35
37
% abundance
75%
25%
decimal
.75
.25
3. Change % to a decimal by moving decimal
point two places to the left (put in table)
4. Plug numbers into the equation
AAM = (mass isotope1) (abundance of 1) +
(mass isotope 2)(abundance of isotope 2)
+… etc
AAM = (35)(.75) + (37)(.25) = 35.45 amu
Practice: Calculate the average atomic mass of
Silicon. 92.21% of Si atoms have a mass of 27.977
amu, 4.70% have a mass of 28.976 amu and 3.09%
have a mass of 29.974 amu.
FOUR BASIC FORCES
A basic understanding of the four basic forces is
helpful in understanding the structure and reactivity
of atoms.
1. Gravitational force: the attraction between
two objects.
 It is directly related to mass and inversely
proportional to distance between centers
of the two objects.
 Subatomic particles are too small for
gravity to have a significant effect on
chemical behaviour.
2. Electromagnetic forces: attractive and
repulsive forces that act between charged
particles or magnetic objects.
 Like charges (poles) repel, opposite attract
 Similar to gravity in that it is directly
proportional to size of the charge of
magnetic force and inversely proportional
to distance.
3. Strong nuclear force: inward force on the
nucleus of an atom that works to keep the
particles together. (Remember protons are all
repelling each other).
4. Weak nuclear force: interacts with quarks,
bosons, and fermions. Involved in beta decay.
THE PERIODIC TABLE – or Why are the elements
arranged like that?
1. Periodic Law: when elements are
arranged by increasing atomic number,
they show repeating patterns of chemical
and physical properties.
2. Groups (aka Chemical family): are elements
found in a column of the periodic table
that share similar chemical (and often
physical) properties.
- They have similar chemical properties b/c
they have the same number of electrons in
their outermost energy level.
- Groups may be labelled several ways
Ex 1 – 18 or 1A – 8A and 1B – 8B
3. Some groups have special names
1 (1A) Alkali metals
2 (2A) Alkaline earth metals
17 (7A) Halogens
18 (8A) Noble gases
4. Periods: horizontal rows in the periodic
table. These rows represent energy levels
in the atoms.
- An element in the 3rd period has electrons
in the 1st, 2nd, and 3rd energy levels.
5. Metals: elements located on the left and
middle section of the periodic table are
metals
- Exception is H, a non-metal
Characteristics of metals include luster,
ability to conduct heat; ability to conduct
electricity. Most are solids at room
temperature (exception being Hg,which is a
liquid at RT)
6. Non-metals: located on the right side of
the table.
- Non-metals vary more than metals in their
properties and characteristics
- Many are gases, they are biologically
important (C, H, O )most common element
in living things.
7. Metalloids: are located along the stair-step
line of the periodic table
- they have characteristics between metals
and non-metals
- B, Si, Ge, As, Sb, Te
8. Information listed for each element
includes
 Atomic number
 Symbol
 Name
 Atomic mass (average)
 Density
 Oxidation states
Atoms react with each other to form molecules and
ionic compounds in order to obtain more
favourable, lower energy states.
MOLECULES AND MOLECULAR COMPOUNDS
Molecule: general name used for two or more
non-metal atoms bonded together by sharing
of electrons.
 Bonds formed by the sharing of electrons
are called covalent bonds
 Several elements are found in nature in
molecular form , meaning 2 or more of the
same type of atom bonded together.
Ex: Oxygen gas is O2
Ozone is O3
Diatomic molecules: molecules made up of
two of the same type of atom bonded
together. There are 7 elements that occur
as diatomic molecules, all non-metals
H2 N2 O2 F2 Cl2 Br2 I2
 Molecular compounds : substances made
of more than one type of atom. Most
molecular substances are made up of only
non-metals.
Molecular formula: formula written using chemical
symbols followed by subscripts to indicate
the actual number of each type of atom in
one molecule.
 element symbols to represent atoms
 subscripts follow symbols to show how many of
those atoms are in the molecule
 no subscript means one atom
C6H12O6 -- this molecule is composed of 6 C,
12 H and 6 O atoms.
H2O2
-- One molecule of hydrogen
peroxide contains 2 H & 2 O
atoms
 Molecular formulas are always used for
molecules.
Empirical formulas: Formulas that show the lowest
whole number ratio of elements in a
compound.
 The empirical formula for glucose : CH2O
 The empirical formula for hydrogen
peroxide is HO
 Empirical formulas are usually used for
ionic compounds.
Structural formulas: show the relative arrangement
of atoms in a compound in a diagram in which
chemical symbols indicate the element and dashes
represent bonds. Used for molecules.
Writing Formulas and Names for Molecular
Compounds:
Some molecules have special names. (memorize )
H2O is water
NH3 is ammonia
Rules for Binary Molecular Compounds
This naming system is for compounds composed of
two nonmetallic elements.
1. The first element keeps its name and gets a
prefix if it has a subscript in the formula
2. The second element gets the –ide suffix
(ending) and ALWAYS gets a prefix
Details….details….
1. Prefixes are used in molecular names to indicate the
number of each type of atom in the compound. :
Mono = 1
Hexa = 6
Di = 2
Hepta = 7
Tri =3
Octa = 8
Tetra = 4
Nona = 9
Penta = 5
Deca = 10
 If the second element’s name begins with an
“a” or “o”, drop the last vowel off the prefix
name
- Ex Carbon monoxide CO
Carbon dioxide
CO2
Diphosphorus pentoxide
www.sciencegeek.net/Chemistry
P2O5
IONS AND IONIC COMPOUNDS
Metals react with non-metals to form ionic
compounds.
 Metals give away their electrons to the
non-metals, resulting in positively charged
metal ions and negatively charged nonmetal ions. The oppositely charged ions
are attracted to each other to form an ionic
compound.
Ions: Atoms that have lost or gained electrons by
reacting with another atom, and , as a result ,
have an net (overall) charge.
Cation = a positively charged ion caused by
losing e Generally metals form cations
Anion = a negatively charged ion caused by
gaining e Generally non-metals form anions
The charge of an ion is equal to the total # of
positive (proton) charges plus the total # of
negative (electron) charges
Ion symbol : written using the chemical symbol
followed by its charge written as a superscript
Ex: Cl1- Mg2+
or
atomic mass
atomic #
Symbol charge
Ex:
An unreacted oxygen atom has
8 protons (+8)
8 electrons (-8)
Total charge
0 neutral, uncharged
But… oxygen reacts with other atoms by gaining
two electrons…so
8 protons (+8)
10 electrons (-10)
Total charge -2
 the oxygen ion has a -2 charge
General chemical family trends for ion formation:
Alkali metals lose 1 e- and form +1 cations
Alkaline earth metals lose 2 e- and form +2 cations
Group 3A lose 3 e- and form +3 cations
Group 4A can lose or gain 4 e- to form either +4 or -4
charged ions
Group 5A gain 3 e- and form -3 anions
Group 6A gain 2 e- and form -2 anions
Halogens gain 1 e- and form -1 anions
Noble gases do not generally form ions – they are
unreactive or inert.
Polyatomic ion: group of generally non-metal atoms
bonded together to form a molecule with an overall
charge.
Ex: SO42- sulfate - see reference sheet
The majority of polyatomic ions are oxoanions –
polyatomic anions made up of an element
covalently bonded to oxygen
(covalent bonds are the result of atoms
sharing electrons)
You will always have the reference sheet for these,
but the following is the rationale behind their
names…
a. The “most important” form is named by
shortening the element name and adding “ate”
(_____ ate) Ex ClO3- Chorate
b. Other forms are named relative to the ___ate
form
* if it has 1 less oxygen = _______ ite
ClO2- = Chlorite
* if it has 2 less oxygens = hypo____ ite
ClO- = Hypochlorite
* if it has 1 more oxygen = per ____ ate
ClO4- = Perchlorate
Ionic compounds are electrically neutral, meaning
total positive charges on the cations = total charges
the anions. This is the key to writing formula
correctly.
NaCl
Na has a 1+ charge, and Cl has a 1- charge
Na2S
2 x Na1+ = +2 total One S2- = -2 total
 ionic compounds form crystalline solids ( the
ions are arranged in a repeating pattern called a
crystal lattice)
 ionic compounds have no distinct molecule
because of the lattice arrangement, so they
are represented by an empirical formula:
the simplest, whole number ratio of atoms
or ions in a compound
WRITING FORMULAS FOR IONIC COMPOUNDS
1. Write the chemical symbol for the cation
first, followed by the symbol for the anion
2. Write subscripts after the chemical symbols to
indicate the ratio of cations to anions for a
neutral compound
How to determine the ratio of cations to anions in
an ionic compound…..
Remember ionic compounds are neutral, meaning
total charges must add up to zero!
…this is because all electrons being lost
MUST be gained by another atom, so total positive
charges on cations MUST equal the total negative
charge on anions involved.
i. Look up charges of ions
ii. Figure out how many you will need of each ion,
so when you add up the total positive charges
from the cations, it equals the total negative
charges of the anions (and is then neutral!)
Note – do not write a subscript of 1
Ex: Fe2O3 is a solid made up of a ration of
2 Fe atoms for every 3 O atoms.
3. Polyatomic ions must be put in parentheses if
they need a subscript added.
Ex Fe(NO3)2 indicates 1Fe2+ : 2 NO31-
Ex: Write the correct formula for
Sodium sulfide
Magneisum hydroxide
Calcium nitrate
Barium oxide
Ammonium nitrate
Sodium fluoride
Aluminum oxide
http://www.chemfiles.com/flash/formulas.html
NAMING IONIC COMPOUNDS
A little background info is necessary to know:
Cation names (three types of cations):
a. Metals that form only one kind of cation are
named with just their element name.
Alkali metals +1
Alkaline earth +2
Boron Group +3
Some transition metal as such as Zn and Ag
b. Transition Metals form more than type of
cation, so their name includes the name of the
element followed by their charge in roman
numerals in parentheses.
Ex Iron2+ and Iron3are named Iron (II) and Iron (III)
c. Polyatomic cations – learn to recognize these from
your reference sheet - and just write their name.
(most polyatomic ions are anions)
Look up name on reference sheet
NH41+ Ammonium
* Note if your ionic compound formula includes 3 or
more elements, it includes a polyatomic ion. They are
not always surrounded by parentheses. They only need
to be surrounded by parentheses when the compound
contains more than one of them.
NaNO3 Sodium nitrate NH4NO3 Ammonium nitrate
Fe(CN)2 Iron (II) cyanide
Anion names:
* there are two types of anions
a. Monatomic non-metal anions are
named by shortening element name and
adding an “ide” ending
F1O2Fluoride Oxide
b. Polyatomic anions – look the names up on
your reference sheet.
Naming Ionic Compounds
1. If the compound you are naming has more than 2
elements in a formula, then you have a polyatomic
ion.
Ex: KOH is K & OH Potassium hydroxide
NaNO3 is Na & NO3 Sodium nitrate
NH4F is NH4 & F Ammonium flouride
Practice – name the following:
NaBr
Ca(OH)2
(NH4)2SO4
AlCl3
KNO3
K2S
Na2SO3
KCN
2. If the metal is a transition metal, you must put ( )
after the name and show the charge as Roman
numerals
Ex Lead ( IV )
You must figure out the charge on the metal
by using your knowledge that ionic
compounds are neutral!!!
a. Figure out charge on transition metals by
looking up the charge on the anion, then
multiply that by the number of anions to get the
total negative charge.
b. Remember that total charges on cations must
be equal (but opposite charge) to total neg
charges so that the ionic compound is neutral.
c. divide total positive charge by the number of
cations present in the formula to find the charge
on one of them. Write this value in the () as a
Roman numeral.
Ex: Write the name for CoCl2:
Cobalt ( ) chloride
 start with the anion charge Cl has a
charge of -1
…there are two Cl, so total neg
charge is -2.
 Total neg charges MUST equal the total
positive charges, so the charge on the one
Cobalt cation must be +2
 Cobalt (II) chloride
Practice naming the following:
SnO2
CuO
Cu2O
Co(NO3)2
NiSO4
Latin System for naming ionic substances that
involve a metal that forms two types of ions
The metal is named using a latin root followed
by a suffix that indicates the higher or lower of
the two possible charges.
Latin roots include:
Plumb = Lead
Stann = Tin
Ferr = Iron
Cupr = Copper
ous suffix – used for the lower of the two possible
charges
ic suffix – used for the higher of the two possible
charges
2+
Ex: Fe = Ferrous and Fe3+ = Ferric
Practice naming the following using the Latin system
FeCl2
FeO
Fe2O3
ACIDS: a molecular substance that produces
hydrogen ions (H+) when dissolved in water.
1. Acids are covalently bonded
2. Acids dissociate (break apart) when
dissolved in water to produce H+ and an
anion
Ex: HCl dissociates to form H+ and ClFormulas for acids can be written by adding one H
for every negative charge on the anion. (Like ionic
compounds they are electrically neutral)
Ex: SO42-
Acid formula: H2SO4
Sulfate will join with 2 H to produce Sulfuric acid
Ex: Cl1Acid formula: HCl
Chloride will join with 1 H to form Hydrochloric
acid
Acids are named depending on the anion they form
when dissolved in water.
Name of anion
________ ide
Hypo____ite
________ite
________ate
Per_____ate
Name of the acid
Hydro________ ic acid
Hypo____ous acid
_______ous acid
_______ic acid
Per_____ic acid
So, how do you go about naming an acid:
1. Recognize the compound is an acid from the H in the
beginning of the formula.
2. Look at the second half of the formula and identify
(look up the name) of the anion.
3. Find the anion name in the list above, and name the
acid accordingly.
Ex:
HCl
HClO
H3PO3
H3PO4
HClO4
chloride anion
hypochlorite ion
phosphite ion
phosphate ion
perchlorate ion
Hydrochloric acid
Hypochlorous acid
Phosphorous acid
Phosphoric acid
Perchloric acid
http://www.kentchemistry.com/links/naming/acids.htm
Writing Formulas for Acids:
- work backwards from the name and identify the
anion, and the charge on the anion.
- Formula will have 1 H for every neg charge on
anion
Ex: Hydrosulfuric acid
the name tells you it is an acid containing a S2anion
Formula will be H2S
Naming pure molecular compounds that contain H:
HCl in the pure state (not dissolved in water) is named
using Hydrogen + anion name with ide ending
Hydrogen chloride
It is Hydrochloric acid only when it is dissolved in
water. This is usually designated by an (aq)
subscript.
Acid Salt – formed when a metal is substituted for some of
the Hydrogens in an acid. The remaining H become part of a
polyatomic ion.
Ex: Hydrosulfuric acid H2S
Sodium replaces one of the H’s = NaHS
Acid Salt Naming Rule #1
metal name + hydrogen (if one H) + polyatomic or anion
name
dihydrogen (if 2 H’s)
NaHS = Sodium hydrogensulfide
NaH2PO4 = Sodium dihydrogen phosphate
Alternate naming of acid salts:
when only one of the H is substituted for by a metal, the
prefix “bi” is inserted between the metal name anion
NaHSO4 = Sodium bisulfate
SIMPLE ORGANIC COMPOUNDS
1. Compounds of C and H, often containing O,
N and other non-metals
2. The carbon atoms in organic compounds
are usually bonded to each other in the
center of the molecule forming a “carbon
backbone”.
3. Hydrocarbons are usually mostly non-polar
(they may have polar ends or side chains),
and generally have weak attractions among
neighboring particles, causing low melting
and boiling points.
4. FYI : Common organic substances used in a
lab include benzene, toluene, hexane,
propanol, methanol, ethanol. They are
generally used as solvents.
Functional group: specific groupings of atoms that
give compounds similar properties.
Hydrocarbons: compounds containing only C and H
Alkanes: a sub category of hydrocarbon compounds
in which each C is bonded to 4 other atoms ,
meaning 4 single bonds around the Carbons.
- They have an “ane” ending to name
3. Alcohols: derivative of alkanes in which an
H is replaced by an OH (hydroxide) group.
(aka and OH group is attached to a carbon )
- Named by removing “e” from the alkane
name and adding “ol”.
Methanol:
Ethanol
1-Propanol
the “1” indicates the
hyroxyl group is on the first or #1 carbon.
2-Propanol
the “2” indicates the
hyroxyl (OH) group is on the #2 carbon
Isomers: compounds with the same molecular
formula but different configurations of atoms,
such as the 1-propanol and 2-propanol.