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Unit 2: Atomic Structure
Democritus
ATOMS
Democritus (4th century BC)
1. IDEA OF THE ATOM
a. Matter consists of tiny, indivisible particles (ATOMOS)
b. Atoms of elements differ from one another in size and shape
- no experiments to support this!
ATOMIC THEORY – John Dalton - 1803
a. All matter is made up of tiny indivisible particles called atoms
1. Based on EXPERIMENTAL evidence
ATOM (modern definition)
Smallest particle of an element that still retains the
properties of that element
John Dalton
1766-1844
Atomic Theory of Matter based
on the following postulates:
1) Each element is composed of
indivisible particles called atoms.
2) All atoms of a given element are
identical but they differ from those
of any other element.
3) Atoms are neither created nor
destroyed in any chemical reaction.
4) A given compound always has the
same relative numbers and kinds of
atoms.
SEEING ATOMS
Law of Definite (Constant) Composition
A compound always has the same composition by
mass.
CO2 : 27% C
CO: 43% C
73% O
57% O
COMPOSITION OF THE ATOM
1.ELECTRON: J.J. THOMSON – 1897
a. Cathode rays are composed of small, negatively charged
particles – ELECTRONS
b. These particles are present in ALL atoms.
JJ Thompson
•Cathode Ray Tube (CRT)
•Discovered the electron
Millikan
c.
MILLIKAN – Oil drop experiment
1. Determined the SIZE of charge on the electron
(-1). This allowed calculation of the electron’s
mass, 9.1 x 10-28 g or ~ 1/1840 mass of H atom.
Models of the atom
GOLD FOIL EXPERIMENT

2. MODEL OF ATOM
a. Gold foil experiment- Rutherford
(1911)
1. Bombarded thin gold foil with alpha
particles (+ charge)
2. Most went straight through:
“Most of the atom is empty space.”
3. A few were strongly deflected or actually
BOUNCED BACK
GOLD FOIL EXPERIMENT
4. Conclusion
Discovered two main parts of the atom:
a. Atoms have a very small, positively
charged center containing almost all of
the mass of the atom (1/10,000 diameter of
the atom) – the NUCLEUS
b. Negatively charged electrons occupy the
space around the nucleus
c. Positive charge on the nucleus = negative
charge outside the nucleus
From Thomson to Rutherford
MODERN CONCEPT:
Electrons occupy “cloud” outside the
nucleus
Orbit
Electron Cloud
Progression of the Atom
5. OTHER PARTICLES:
a. PROTON:
CHARGE: +1
MASS: 1.673 X 10-23 g or 1 amu
LOCATION: nucleus
b. NEUTRON
1. CHARGE: 0
2. MASS: 1.675 X 10-23 g or 1 amu
3. LOCATION: nucleus
4. Discovered by Chadwick (1932)
Remember for the ELECTRON
1. CHARGE: -1
2. MASS: 9.1 X 10-28 g or 0 amu
3. LOCATION: electron cloud
ATOMIC MASS of an element
a. average mass of an atom of an element,
compared to a standard
b. Standard: Carbon-12 (12.000)
c. Atomic mass unit (amu):
1/12 mass of a carbon-12 atom
d. Atomic mass of an element on the
periodic table is an AVERAGE of the
naturally occurring isotopes
Isotopes of Hydrogen
Hydrogen – 1
(protium)
1p
1
H
1e
1
0n
Hydrogen – 2
(deuterium)
1p
2
H 1e
1
1n
in “heavy water”
Hydrogen – 3
(tritium)
1p
3
H 1e
1
2n
Radioactive
Note: Use the mass number on the periodic table, unless
it is specifically given.
Isotopes of Carbon
C-12
C-14
or
or
12
6
C
6p
6e
6n
14
6
C
6p
6e
8n
ATOMIC MASS
1. There are two isotopes of copper, Cu-63 and Cu-65.
Which is more abundant?
ATOMIC MASS
2. Calculate the atomic mass of Chlorine, given that 75.77%
of its atoms are Cl-35 (34.97 amu) and 24.23% are Cl-37
(36.97 amu).
ATOMIC MASS
3. The atomic mass of aluminum is 26.98. The atomic mass
of nitrogen is 14.01. An aluminum atom is how many
times heavier than a nitrogen atom?
ATOMIC MASS
4. Element Z has two naturally occurring isotopes, Z-203
(AM of 202.97) and Z-205 (AM of 204.97). If the average
atomic mass is 204.41, in what percentages does each
of these isotopes occur?
ATOMS FORM MOLECULES!
MOLECULE
A group of two or more atoms held together
by covalent bonds - definition
a. Basic unit of covalent compounds
Ex. H2O, CO2, C6H12O6
b. Building blocks of SOME elements
H2, O2, N2, Cl2, Br2, I2, F2
DIATOMIC ELEMENTS
MOLECULAR FORMULA
Shows the number of atoms of each element
in one MOLECULE of a substance
Ex. C6H12O6
6 C atoms
12 H atoms
6 O atoms
24 total atoms
ATOMS FORM IONIC COMPOUNDS!
IONS
An atom or group of atoms with a positive or negative
charge
a. Positive ions
Form when an atom LOSES electrons
Na → Na+ + 1 e
11 p
11p
11 e
10e
(+11 -10) = +1 or
23Na+
11
Negative ions
Form when an atom GAINS electrons
O + 2 e → O28p
8p
8e
10e
(+8 -10 = -2)
May be represented as 16O28
IONIC COMPOUNDS
a. Made up of ions held together by STRONG
electrical attractions
b. No discrete molecules
(No NaCl molecule)
EMPIRICAL FORMULAS (simplest formula)
a. Show only the LOWEST RATIO of atoms of
each element in the compound
b. Represent ALL ionic compounds
WRITING EMPIRICAL FORMULAS
PRINCIPLE: Compounds are electrically neutral
Ba2+O2- =
BaO
Ca2+F1- =
CaF2
Al3+O2- =
Al2O3
Sn4+O2- =
SnO2
PRACTICE
Write correct formulas for the following:
Li1+ and O2Sr2+ and S2Be2+ and Br1Al3+ and S2K1+ and I1-
MOLECULAR MASS(FORMULA MASS)
How heavy a molecule is compared to Carbon – 12
a. How?
b. Find the molecular mass of C4H10 (butane).
4 C + 10 H =
4(12.01) + 10( 1.01) =
58.14 amu
c. Compare the relative masses of Benzene (C6H6)
and butane.
6C+6H=
6(12.01) + 6(1.01) =
78.12 amu
78.12 = 1.344
58.14
A C6H6 molecule is 1.344 x heavier than a C4H10
molecule.
Determine the formula masses of:
Ba(NO3)2
CaSO4 * 2 H2O
3. Al2(SO4)3
4. CaCl2 * 10 H2O
THE MOLE
1 MOLE equals 6.022 x 1023 particles.
1 mole Cu contains 6.022 x 1023 atoms.
1 mole CO2 contains 6.022 x 1023 molecules.
A mole has the mass in grams which is
numerically equal to the molecular mass
(formula mass) of the substance.
This formula mass is the sum of the atomic masses
of the atoms in the substance.
63.55
1 mole Cu atoms weighs __________
grams.
1 mole CO2 molecules weighs __________
grams.
44.01
1 mole Ba(NO3)2 weighs ___________ grams.
The MASS IN GRAMS OF 1 MOLE is:
1 gram-formula mass (GFM)
1 gram-molecular mass (GMM) OR
1 gram-atomic mass (GAM)
MOLAR MASS
MOLE – GRAM CONVERSIONS
1. Find the mass in grams of 1.20 moles of propane,
C3H8.
MOLE – GRAM CONVERSIONS
2. Calculate the number of moles in 62.8 grams
of K2O, potassium oxide.
3. What is the mass of 0.785 mole of C2H5OH,
ethanol?
4. How many moles are present in 300. g of
Ca(NO3)2, calcium nitrate?
5. Given: 2.35 grams of propanol, C3H7OH. How many
moles of propanol are in the sample?
How many propanol molecules?
How many carbon atoms?
6. How many molecules are present in 100. grams of
hexane, C6H6?
How many Hydrogen atoms?
How many Carbon atoms?
6. How many molecules are present in 100. grams of
hexane, C6H14?
How many Hydrogen atoms?