Transcript adv ch 5 atomic structure

Ch. 4: Atomic Structure
4.1 Defining the Atom
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
Dalton’s Atomic Theory-1803
1. Matter is made of small, indivisible particles
called atoms
2. Atoms of same element have the same size,
mass, and properties
3. Atoms of different element combine in whole
number ratios to make compounds.
4. In chemical reactions, atoms are combined,
separated, and rearranged.
Consequences of Dalton’s Theory
The “billiard ball model”-the atom
is viewed as a small solid indivisible sphere.


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
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
Ch. 4: Atomic Structure
4.2 Structure of Atom
Discovery of the subatomic
particles
Discovery of Electron


resulted from scientists (JJ Thompson) passing electric
current through gases to test conductivity
used cathode-ray tubes
 noticed that when current was passed through a
tube 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 were attracted to
the positive pole (anode)- opposites attract!
Discovery of Electron

Conclusion:
there were negatively charged
particles inside the cathode ray

The particle was called ELECTRON

Cathode rays are made of electrons
Discovery of Electron
 J.J.
Thomson (English 1897) did more
experiments to actually make the
discovery
 he found ratio of charge to mass of this
particle
 since the ratio stayed constant for any
metal that contained it, it must be the
same in all of the metals
Plum Pudding Model (1897)
 proposed by Joseph John Thomson
 Nobel Prize in physics in 1906

the atom was a sphere of positive
electricity (which was diffuse) with
negative particles (electrons) imbedded
throughout
Charge and mass of electron
1916- Robert Millikan
discovered the charge and the mass of
the electron
Electron has
 a charge of -1
 a mass of 1/ 1840 of the mass of
Hydrogen atom (the smallest atom)
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
Canal rays
1886 Goldstein
 In the cathode tube experiment Goldstein
noticed another rays traveling from anode
to cathode
 He named those rays canal rays and are
made of positive particles called protons
Discovery of Nucleus

Rutherford discovered the nucleus by shooting
alpha particles (have positive charge) at a very
thin piece of gold foil
 1911-
Ernest Rutherford- the gold foil
experiment

he predicted that the particles would go right
through the foil at some small angle
Discovery of Nucleus
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
4.
“Powerful Force”= Nucleus
The Nuclear Model- proposed by Ernest
Rutherford

the atom is mostly empty space with a
dense positively charged nucleus
surrounded by negative electrons.

Rutherford received the Nobel Prize in
chemistry in 1908 for his contributions
into the structure of the atom.
The discovery of Neutrons
1932- Chadwick
 Neutrons
have NO CHARGE
 Neutrons have the mass almost equal to
the proton
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
The Subatomic Particles
Particle
Symbol
Charge
Relative
Mass
Electron
e-
-1
1/1840
Proton
p+
+1
1
Neutron
n0
0
1
Ch. 4 Atomic Structure
4.3 Distinguishing among Atoms
Atomic structure
Atomic number Z
Z= atomic number= # protons= # electrons

Indicates the position of the element in the periodic
table (the whole number by each element)

Ex: 6C , 19 K, 1H
Mass number A
A= # protons + # neutrons

To find A round up the decimal number for each
element in the Periodic table

Ex: 12C, 39K, 1H or C-12, K-39, H-1
How to calculate
 A.
electrons= atomic number Z
 B.
protons= atomic number Z
 C. neutrons= mass number – atomic number
n= A-Z
ISOTOPES

Members of the same family
 Have the same
chemical symbol
number of protons
number of electrons
atomic number
 Have different
mass numbers
numbers of neutrons

Ex: 12C, 13C, 14C, 16C- all have 6 electrons and 6
protons but 6, 7, 8, and 10 neutrons
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
 Is
the decimal number found in each box
for each element in the Periodic table
Average Atomic Mass
 To



calculate it we need to know:
number of isotopes
Mass of each isotope
Percentage (relative abundance) of each
isotope)
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
(19.99244x90.51)  (20.99395x0.27)  (21.99138  9.22)
100
 20.17945amu
Neon