Chapter 2
Notes
AP Chemistry
Unit 1
I.

A.




Atomic Theory
Dalton’s Theory
1. An element is composed of tiny particles called
atoms
2. The atoms of an element are identical, while atoms of
different elements have different properties
3. Atoms of one element cannot be changed into
atoms of a different element by chemical means; atoms
can neither be created or destroyed in chemical
reactions.
4. Compounds are formed when atoms of 2 or more
elements combine in a fixed ratio: a given compound
always contains the same number and kind of atoms
(Law of Constant Composition)
Atomic Theory
 B.

Law of Conservation of Mass
There is no change in mass during a
chemical reaction.
 Law

of Multiple Proportions
applies when 2 elements can form 2 or
more different compounds. * For a given
mass of one element, the masses of the
second element in the compound are in
small whole number ratios.
II.
The Discovery of Atomic
Structure
 A.


J.J. Thomson
Used cathode ray tubes (see page ) to detect
negatively charged particles, which he called
electrons
Developed the “plum pudding”
model of the atom in 1897
 Robert
Millikan
Used the “oil drop” experiment
to determine the charge and
mass of an electron in 1909

 Henri


Becquerel
Discovered “radioactivity” in 1896
This led to further study of alpha, beta, and
gamma particles.
 Eugene


Rutherford
Used gold foil and alpha
particles to discover the
existence of the nucleus
in 1911
Later discovered proton
In 1919
 James

Discovered neutrons in 1932
 Niels

Chadwick
Bohr
Developed the common model of the atom,
referred to as the “planetary” model where
electrons orbit the nucleus
III. The Modern View of the Atom
 A.



Three main particles
Proton = positive, in nucleus,
Neutron – neutral, in nucleus
Electron – negative, cloud, negligible mass
 Characteristics



of Atoms
Atomic Number, Z: equals the number of protons
in atom
Mass Number, A: equals the number of protons
and neutrons in the atom
Isotopes: different forms of an element that have
different numbers of neutrons and therefore
different mass numbers.
Practice
 Using
the nuclide symbols for these
elements, identify the number of protons,
neutrons, and electrons
a.
13
p=
e=
n=
𝐶
3
b. 𝐻
p=
e=
n=
+
c.
19 -1
𝐹
p=
e=
n=
C. Four Basic Forces Known in
Nature

1. Gravitational



Occur between all objects
Since masses of atoms and subatomic particles
are so small, this force is negligible
2. Electromagnetic


These are attractive and repulsive forces that
exist between electrically charged or magnetic
objects.
Often these forces are referred to as Coulombic
forces, since Coulomb’s Law quantifies the
force between 2 charged objects: F = Kq1q2/d2
 3.

Strong Nuclear Force
This is the force that holds the nuclei of atoms
with more than one proton together
 Weak


Nuclear Force
This force is weaker than electromagnetic
forces but stronger than gravitational forces
The weak interaction is responsible for the
existence and structure of atomic nuclei and is
responsible for both the radioactive decay
and nuclear fusion of subatomic particles
IV. Masses of Atoms
 A.

Atomic Mass: The C-12 scale
Amu – atomic mass unit
 B.
Atomic Masses and Isotopic
Abundance or Weighted Atomic Mass


The average mass of an element
accounting for the abundance of all
isotopes
Calculated by multiplying each
isotope’smass by its % abundance, then
add all isotopes.
Example:
 Given
the percentages and masses of different
isotopes, find the weighted atomic mass of the
element and identify the element.
Isotope
mass(amu)
% abundance
1
6.01513
7.494
2
7.01601
92.506
(.07494 * 6.01513) + (0.92506 * 7.01601) =
0.4508 + 6.4902 = 6.9410 amu
Element:
Lithium
 C.
Mass Spectrometer
Read the description on page ___ about how it
works
 This instrument provides data about the masses of
the isotopes of an atom: data collected in this
way proved that one of Dalton’s ideas was not
true – not all atoms of an element are identical –
they have variable mass depending on the
number of neutrons
What element is this?

Check for understanding
 Calculate
the exact molar mass for tin.
V. The Periodic Table
 A.
Most periodic tables include
information about the atomic number,
element symbol, and mass number of
each element.
 B. Periods


1. Periods are horizontal rows on PT
2. Periods represent elements that contain
approximately the same number of
electrons and energy levels.
 C.



Groups (or families)
1. Groups are the vertical columns on the
periodic table.
2. Elements in a group have very similar
physical and chemical properties because
they have the same number of valence
(outer energy level) electrons.
3. Examples of groups (name and period #)\
 A.
alkali metals – 1
 B. alkaline earth metals – 2
 C. chalcogens – 16
 D. halogens – 17
 E. noble gases - 18
 D.

Types of Elements – See page 47
1. Metals – general characteristics
 A.
high electrical conductivity: (110 x higher than
nonmetals) Ag is #1but Cu and Hg are also high.
 B. high thermal conductivity: Al, Cu and alloys like steel
are highest
 C. Ductility and Malleability: can be formed into wires
(ductile) or flattened into sheets (malleable); electrons
hold the metals’ atomic nuclei together, so metal
“crystals” can be deformed without shattering like
nonmetals.
 D. Luster: polished metal surfaces reflect light; most
have a silvery-white color because they reflect all
wavelengths of light, but gold and copper absorb
some light in the blue region of visible light
 E. All are solids at room temperature except Mercury
 F. Metals lose electrons to form cations.
 2.




 3.


Nonmetals
A. Almost always found in compound or diatomic
forms (except noble gases)
B. Most nonmetals are gases except for Br2 (liquid)
and At, I2 (solids)
C. Nonmetals gain electrons to form anions.
D. They will gain electrons to fill the valence level
octets, except H+ (which loses an electron) and
noble gases (which are typically unreactive).
Metalloids
A. Intermediates between metals and nonmetals
– located along “staircase” on table.
B. More like nonmetals chemically, but more like
metals in physical properties (conductivity and
solids)
VI. Molecules and Molecular
Compounds
 A.



Molecules
1. Definition: Compounds that contain
atoms that share their electrons to form
covalent bonds
2. Examples: H2O, C12H22O11
3. Types of molecules
 A.
Polar: electrons are distributed unevenly
throughout molecule
 B. Nonpolar: electrons are evenly distributed
in the molecule.
* We will study polarity in depth later.
 B.

Types of formulas
1. Empirical
 The
simplest formula of a compound
 Represents the smallest whole number ratio
 Examples: CO2, H2O, CH2O

2. Molecular
 The
formula that indicates the actual number of
atoms in a molcule
 Examples: C6H12O6, K2C2O4
Ways to picture molecules


Molecules occupy three
dimensional space. However, we
represent them in two dimensions.
The structural formula gives the
connectivity between individual
atoms in the molecule.


It may or may not be used to
show the three dimensional
shape of the molecule.
If the structural formula does show
the shape of the molecule, then
either a perspective drawing,
ball-and-stick model, or spacefilling model is used.
Check for Understanding
1.
Which of the following are empirical
formulas? For each molecular compound,
what is the empirical formula?

C12H22O11
C18H32O16
empirical
molec
C9H16O8
H2 O 2
molec
HO
HC2H3O2
molec
CH2O
VII. Ions and Ionic Compounds
 A.
Ions – See tables on pages 58 and 59 for list
of common ions. (See back of PT for
polyatomic ions)


1. definition: an ion is an atom that has gained or
lost electrons.
2. Types of ions
 A.
anion – negative ion formed when atom gains e B. cation – positive ion formed when atom loses e-

3. How do ions form (predicting ionic charge)
 a. Anions gain electrons to form noble gas
configuration
O + 2 e- → O2F + 1 e- → F-1
 b.
Cations
Group 1 & 2 & Al – lose electrons to form noble gas
configurations
Na → Na+ + eMg → Mg+2 + 2 eAl → Al+3 + 3 e
Transition & post-transitional metals – do not from
noble gas configurations
Fe → Fe+2 + 2 eFe → Fe+3 + 3 e
 B.


Ionic Compounds
1. Ionic compounds must have electrical
neutrality – the charges in an ionic compound
must sum to zero.
2. Ionic compounds are generally combinations
of metals and nonmetals, such as NaCl or MgBr2.
Molecular compounds are generally composed of
nonmetals only, such as CO2 or H2O.
 a.


Monoatomic ions: ions that contain one atom.
i. metals – cations formed from metal atoms have the
same name as the metal.( Ex – Na and Na+1 is soldium
and sodium ion respectively). If metals can from more
than one ion we use a Roman numeral to denote the
value of the positive charge (Fe2+ is iron(II) and Fe3+ is iron
(III))
Ii. Nonmetals- anions have name formed by replacing
ending of element name with –ide. (Ex. O2- is oxide and
F1- is fluoride).
 b.
Polyatomic Ions – ions that contain more
than one atom.



See back of your PT
You may have to use subscripts to show that there
are multiple ions present.
Example: aluminum nitrate
Al+3 NO31- = Al(NO3)3
Naming Inorganic
Compounds
 Ionic






Compounds (Salts)
Composition: cation + anion
Name: Cation (space)anion name (-ide or
name polyatomic as is)
Ex: AlCl3
aluminum chloride
Mg(NO3)2
magnesium nitrate
NH4Br
ammonium bromide
Cl(CN)2
calcium cyanide
Naming Inorganic
Compounds
 ****If
the metal has more than one
possible charge you must include the
Roman numeral in parenthesis




Ex:
CuCl
FeO
CuCl2
Copper(I) chloride
Iron(II) oxide
Copper(II) chloride
B. Acids

1. Binary: hydrogen + anion ending with -ide
 Name:hydro
+ nonmetal root + icacid
 Ex: HCl (aq) hydrochloric acid H2S (aq) hydrosulfuric
acid


2. Oxyacids: hydrogen + polyatomic ion
containing oxyge
a. Oxyanions ending with –ate
 Name:
oxyanion root + ic acid
 Ex: H2SO4 sulfuric acid HNO3 nitric acid

b. Oxyanions ending with –ite
 Name:
oxyanion root + ous acid
 Ex: H2SO3 sulfurous acid HNO2 nitrous acid


b. Oxyanions ending with –ite
Name: oxyanion root + ous acid
 Ex:
H2SO3 sulfurous acid HNO2 nitrous acid
 Note:
There is a pattern to the prefixes
and suffixes given to oxyanions!
 Anion Corresponding Acid




ClO = perchlorate HClO = perchloric acid
ClO = chlorate HClO = chloric acid
ClO = chlorite HClO = chlorous acid
ClO = hypochlorite HClO = hypochlorous
acid
41-
4
31-
21-
1-
3
2
 C.


Covalent Compounds
Composition: nonmetal + nonmetal
Name: prefix nonmetal + prefix nonmetal +
-ide
 The
prefix “mono” is never attached to the
first element
 The element farthest to the left is written first. If
both are in the same group the one farther
down is written first



Example:
BH3
boron trihydride
P2O3
diphosphorous trioxide
Greek Prefixes
 1=
mono
 2= di
 3 = tri
 4 = tetra
 5 = penta
 6 = hexa
 7 = hepta
 8 = octa
9 = nona
10 = deca
 D.



Hydrates
Composition: ionic compound with water
attached in the crystal structure
Name: name of salt • prefix hydrate
Ex:
 CuSO4
• 5H2O copper(II) sulfate • pentahydrate
 CoCl2 • 7H2O
cobalt(II) chloride • heptahydrate
Writing Formulas
 A.


Molecules – all information is given
Ex: xenon hexafluoride
Dinitrogen tetroxide
XeF6
N2O4
 B.
Salts and Acids – you must look at
valence of the compounds to maintain
neutrality





Ex: nickel (II) nitrate
Barium chlorite
Phosphoric acid
Phosphorous acid
Hydrophosphoric acid
Ni(NO3)2
Ba(ClO2)2
H3PO4
H3PO3
H3 P