Transcript ch3 - Otterville R-VI School District

Ch. 3: Atoms
3.1 Foundations
History

Democritus
 named the most basic
particle
 atom- means “indivisible

Aristotle
 didn’t believe in atoms
 thought matter was
continuous
History
 by


1700s, all chemists agreed:
on the existence of atoms
that atoms combined to make
compounds
 Still
did not agree on whether
elements combined in the same ratio
when making a compound
Law of Conservation of Mass
 mass
is neither created or destroyed
during regular chemical or physical
changes
Law of Definite Proportions
 any
amount of a compound contains
the same element in the same
proportions by mass
No matter
where the
copper
carbonate is
used, it still
has the same
composition
Law of Multiple Proportions
applies when 2 or more elements combine
to make more than one type of compound
 the mass ratios of the second element
simplify to small whole numbers

Law of Multiple Proportions
Dalton’s Atomic Theory
1.
2.
3.
4.
5.
All mass is made of atoms
Atoms of same element have the same
size, mass, and properties
Atoms can’t be subdivided, created or
destroyed
Atoms of different element combine in
whole number ratios to make compounds
In chemical reactions, atoms are
combined, separated, and rearranged.
Modern Atomic Theory

Some parts of Dalton’s theory were wrong:



atoms are divisible into smaller particles
(subatomic particles)
atoms of the same element can have different
masses (isotopes)
Most important parts of atomic theory:


all matter is made of atoms
atoms of different elements have different
properties
Ch. 3: Atoms
3.2 Structure of Atom
Structure of Atom

Nucleus:



contains protons
and neutrons
takes up very little
space
Electron Cloud:


contains electrons
takes up most of
space
Subatomic Particles
 includes



all particles inside atom
proton
electron
neutron
 charge
on protons and electrons are
equal but opposite
 to make an atom neutral, need equal
numbers of protons and electrons
Subatomic Particles
number of protons identifies the atom as a
certain element
 protons and neutrons are about same size
 electrons are much smaller
 nuclear force- when particles in the nucleus
get very close, they have a strong attraction




proton + proton
proton + neutron
neutron + neutron
Atomic Radius




size of atom
measured from center of nucleus to outside of
electron cloud
expressed in picometers (1012 pm = 1 m)
usually 40-270 pm
Example
 An
atom has a radius of 140 pm.
How large is that in meters?
1m
12
10
140 pm  12
 140 10  1.4 10 m
10 pm
Ch. 3 Atoms
3.2b: Important Discoveries
Discovery of Electron
resulted from scientists passing electric
current through gases to test conductivity
 used cathode-ray tubes
 noticed that when current was passed
through a glow (or “ray”) was produced

Discovery of
Electron
Noted Qualities of Ray Produced:
1. existed- there was a shadow on the
glass when an object was placed
inside
2. had mass- the paddle wheel placed
inside, moved from one end to the
other so something must have been
“pushing” it
Discovery of Electron
Noted Qualities of Ray Produced:
3.
negatively charged- the rays behaved the same
way around a magnetic field as a conducting wire
4.
negatively charged- were repelled by a negatively
charged object
Discovery of Electron

all of these led scientists to believe
there were negatively charged
particles inside the cathode ray

Milliken found the mass of the
electron
Discovery of Electron
 J.J.
Thomson (English 1897) did more
experiments to actually make the
discovery
 he found ratio of charge of this
particle to this mass of the particle
 since the ratio stayed constant for any
metal that contained it, it must be the
same in all of the metals
Are electrons the only particles?
 since
atoms are neutral, something
must balance the negative charge
 since an atom’s mass is so much
larger than the mass of its electrons,
there must be other matter inside an
atom
Discovery of Nucleus
Rutherford discovered the nucleus by
shooting alpha particles (have positive
charge) at a very thin piece of gold foil
 he predicted that the particles would go
right through the foil at some small angle

Discovery of Electron
Discovery of Nucleus
 some
particles (1/8000) bounced back
from the foil
 this meant there must be a “powerful
force” in the foil to hit particle back
Predicted Results
Actual Results
Discovery of Nucleus
Characteristics of
“Powerful Force”:
1. dense- since it was strong
enough to deflect particle
2. small- only 1/8000 hit the
force dead on and bounced
back
3. positively charged- since
there was a repulsion
between force and alpha
particles
 Find
the element Sodium on your
periodic table. What do you know
about atoms of sodium from the
information on the table?
Ch. 3 Atoms
3.3 Counting Atoms
Atomic Number
 (Z)
number of protons
 All atoms of the same element have
the same atomic number
 located above the symbol in the
periodic table
 order of the elements in the periodic
table
Isotopes
atoms of the same element with different
numbers of neutrons
 most elements exist as a mixture of
isotopes
 What do the Carbon isotopes below have in
common? What is different about them?

Mass Number
sum of particles in nucleus
 A = #p + #n
Hydrogen isotopes have special names:





protium
deuterium
tritium
What do the
prefixes
in their names
come from?
Designating Isotopes
 Hyphen


notation:
Name - mass number
ex. Carbon – 13
 Nuclear
pn
p
Symbol notation:
Symbol
Ex : 136C
Examples
1.
7 protons, 8 neutrons
Nitrogen-15
2.
15
7
N
17 electrons, 19 neutrons
Chlorine- 36
36
17
Cl
Examples
3.
Z=5, 6 neutrons
Boron- 11
3.
11
5
B
A=75, 42 neutrons
Arsenic- 75
75
33
As
Examples
#p
#n
#e
Z
A
Carbon- 13
6
7
6
6
13
Xenon-131
54
77
54
54
131
Sodium-24
11
13
11
11
24
Oxygen- 15
8
7
8
8
15
Examples
#p
#n
#e
Z
A
Carbon- 13
6
7
6
6
13
Xenon-131
54
77
54
54
131
Sodium-24
11
13
11
11
24
Oxygen- 15
8
7
8
8
15
A neutral atom contains 34
electrons and has an A of
59. Write the nuclear
symbol notation and
hyphen notation for this
isotope.
Ch. 3 Atoms
3.3 Counting Atoms
Relative Atomic Mass
 since
masses of atoms are so small, it
is more convenient to use relative
atomic masses instead of real masses
 to set up a scale, we have to pick one
atom to be the standard
 since 1961, the carbon-12 nuclide is
the standard and is assigned a mass of
exactly 12 amu
Relative Atomic Mass
 atomic
mass unit (amu)- one is exactly
1/12th of the mass of a carbon-12
atom
 mass
of proton= 1.007276 amu
 mass of neutron= 1.008665 amu
 mass of electron= 0.0005486 amu
Relative Atomic Mass
 the
mass number (A) and the relative
atomic mass are very close but not the
same because


relative atomic mass includes electrons
the proton and neutron masses aren’t
exactly 1 amu
Average Atomic Mass
 weighted
relative atomic masses of the
isotopes of each element
 each isotope has a known natural
occurrence (percentage of that
elements’ atoms)
Calculating Average Atomic Mass
 Naturally


occurring copper consists of:
69.71% copper-63 (62.929598 amu)
30.83% copper-65 (64.927793 amu)
(0.6971 x 62.929598)+(0.3083 x 64.927793)
=63.55 amu
Calculating Average Atomic Mass
 An
element has three main isotopes
with the following percent occurances:



#1: 19.99244 amu, 90.51%
#2: 20.99395 amu, 0.27%
#3: 21.99138 amu, 9.22%
 Find
the average atomic mass and
determine the element.
Calculating Average Atomic
Mass
The mole
a
unit for measuring a very large
amount- like number of atoms or
molecules in a sample
 like one dozen (1 dozen = 12 things)
 except bigger:
1 mole = 6.022x1023 things
6.022x1023 ?
 6.022x1023 is the number of atoms
in exactly 12 g of carbon-12
 Why
The mole
 6.022x1023
is called
Avogadro’s Number in
honor of all of his
contributions to chemistry
 can be used as a conversion
factor between a number of
things and mole
6.022 10 atoms
1 mole
OR
23
1 mole
6.022 10 atoms
23
Molar Mass
 the
mass of one mole of pure
substance in grams per mole
 numerically equal to average atomic
mass
 under the symbol on the periodic
table
 can be used as a conversion factor
between moles and grams
15.9994 g
1 mole
for Oxygen :
OR
1 mole
15.9994 g
Conversion Factors
Grams
#
Atoms
Moles
Use
Molar
Mass:
Use
Avog.’s
Number:
grams
per
mole
atoms
per mole
Gram  Moles
use molar mass
 Ex. 32.3 g Na = ? mol Na

1 mol Na
32.3 g Na 
 1.40 mol Na
22.99 g Na

Ex. 0.56 mol Fe = ? g Fe
55.85 g Fe
0.56 mol Fe 
 31 g Fe
1 mol Fe
1. What is the mass in grams of 3.5 mol of Cu?
2. Grams of Fe in 2.25 mol of Fe?
3. Grams of K in 0.375 mol of K?
4. How many moles are in 11.9g Al?
5. Number of moles in 5g of Ca?
6. Number of moles in 3.60 x 10-16 g Au?
# Atoms  Moles
use Avogadro’s Number
 Ex: 1.40 mol Na = ? Na atoms

6.022 10 atoms
23
1.40 mol Na 
 8.43 10 Na atoms
1 mol
23

Ex: 3.4x1023 atoms Fe = ? mol Fe
1 mol
3.4 10 atoms Fe 
 0.56 mol Fe
23
6.022 10 atoms
23
1. # of moles of Ag in 3.01 x 1023 atoms
of Ag?
2. Moles of Pb in 1.50 x 1012 atoms Pb?
3. Moles of Sn in 2500 atoms of Sn?
1. 3.01x1023 Ag atoms ( 1 mol Ag) =
6.022x1023
0.500 mol Ag
2. 1.50x1012 atoms Pb ( 1 mol Pb) =
6.022x1023
2.49x10-12 mol Pb
3. 2500 Sn atoms ( 1 mol Sn) = 4.2x10-21
6.022x1023
Grams 
# Atoms

use both: Avogadro’s # and molar mass

Ex: 0.0326 g N = ? atoms of N
1 mol N 6.022 1023 atoms
0.0326 g N 

 1.40 1021 atoms N
14.01 g N
1 mol

Ex: 2.01x1041 atoms of H = ? g H
1 mol
1.01 g H
17
2.0110 atoms H 


3.37

10
gH
23
6.022 10 atoms
1 mol
41
1. Mass of 1.20x108 atoms of Cu?
2. Mass of 7.7x1015 atoms of Ni?
3. # of S atoms in 4g of S?