Download Atom Review

Atomos (Democritus c. 460 – c. 370 B.C.)
The first suggestion that atoms make up all matter
The Law of Conservation of Matter (Lavoisier 1743-1794)
Matter is neither created nor destroyed in a chemical reaction.
The Law of Constant Composition (Joseph Proust 1754-1826)
A given compound always contains the same elements in the same
proportions by mass.
Example: H2O is always 88.9% oxygen and 11.1% hydrogen
Dalton’s Atomic Theory of Matter (John Dalton 1766-1844)
Postulates 1-4 (pg. 92 in book)
1. Each element is composed of extremely small 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. They are
dissociated, combined or recombined.
4. Compounds are formed when atoms of different elements combine with one
another. A given compound always has the same relative numbers and kinds
of atoms.
Dalton’s Model
Described atom as a
Hard Dense, Indivisible sphere
Thomson’s Model
Credited with the discovery of the electron
using a cathode ray tube under the influence of
an electric field. Described atom using his
“Plum Pudding Model” where
Pudding = + positively charged
Plums = - negatively charged
Rutherford’s Model
First discovered alpha, beta and gamma radiation.
Used alpha radiation in the Gold Foil
Experiment to discover the NUCLEUS.
(This led to the proton discovery)
Bohr (Orbital) Model
“Electrons exist in energy levels
outside the nucleus. The further
away the electron is from the
nucleus, the more energy it
requires to exist there.
+
Chadwick (Rutherford’s student) then discovered the neutron.
The History of Atomic Structure.
J.J. Thomson’s (1856-1940)
Discovery of the ELECTRON
Stream of particles
Cathode Ray Tube
Battery or other source of electricity
The History of Atomic Structure.
J.J. Thomson’s (1856-1940)
Discovery of the ELECTRON
Electric Field
Stream of particles
(+)
Cathode Ray Tube
(-)
Battery or other source of electricity
The History of Atomic Structure.
Earnest Rutherford (1871-1937)
Discovery Alpha, Beta & Gamma radiation
Magnetic Field
Lead Block
(+)
b
g
a
(-)
Radioactive material
The History of Atomic Structure.
Earnest Rutherford (1871-1937)
Gold Foil Experimentation
Lead Block
Radioactive material
The History of Atomic Structure.
Earnest Rutherford (1871-1937)
Gold Foil Experimentation
Lead Block
Radioactive material
The History of Atomic Structure.
Earnest Rutherford (1871-1937)
Gold Foil Experimentation
Lead Block
Radioactive material
The nucleus is a very small positively charged core containing
protons and neutrons.
Negatively charged electrons are extremely tiny and occupy the vast
majority of the atom’s volume.
• The smallest particle of an element that retains the chemical identity
• composed of three main parts.
Proton:
•Positively charged
•Found in the nucleus
•Atomic number (z) =
protons
Electrons:
Neutron:
•Neutral (no charge)
•Found in the nucleus
•Atomic mass – atomic
number = neutrons.
Protons and neutrons give
the atom its mass.
Particle
Proton
Neutron
Electron
Charge (C)
+1.602x10-19
0
-1. 602x10-19
•Negatively charged
•Exist in electron cloud
•Number of electrons =
number of protons.
(in a neutral atom)
Electron Cloud gives the
atom its volume.
Mass (g)
1.673x10-24
1.675x10-24
9.109x10-28
Mass (amu)
1.0073 = 1
1.0087 = 1
0.0006 = 0.0
The atomic number (z) describes the number of protons
in an element.
(The number of protons is equal to the number of
electrons in a neutral atom)
19
K
39.098
Potassium
The atomic number is the smallest of the two numbers
seen in each box on the periodic table.
The mass number describes the average mass of the
element isotopes.
19
K
39.098
Potassium
The mass number (m) is the largest number
in each box on the periodic table.
The mass number – the atomic number is equal to the number of
neutrons in an atom.
m – z = neutrons
•Atoms that have the same number of protons
but different numbers of neutrons.
•Isotopes of an element exhibit identical
chemical behavior.
•This is why H-1 and H-2 both bond with
oxygen to form water.
•Carbon-14 and carbon-12 bond with oxygen to
form CO2
Isotope symbols  Carbon-12, C-12 or
P
N
E
Protium
Hydrogen’s 3 isotopes
Deuterium
Tritium
2
1H
1
1H
H-1
12
C
6
H-2
3
1H
H-3
Ions: electrically charged atoms.
o Atoms that gain electrons are negatively charged
 Nonmetals gain electrons
o Atoms that lose electrons are positively charged
 Metals lose electrons
Potassium
19 protons
19 electrons
20 neutrons (m – z: 39 – 19)
If potassium loses an electron:
Potassium: + 19 protons (+ charges)
- 18 electrons (- charges)
+1 overall charge of the potassium ion
(also called the oxidation number or state)
Ion representations
Cs+1
Cs+
133
+1
Cs
55
All three of these indicate that cesium has
given 1 electron away.
Isoelectric: Indicates the number of electrons in
an ion matches the number of electrons in the
noble gas it wants to become to increase stability.
So what Noble Gas is Cs+1 isoelectric with?
Xenon
Isotope symbols
Cesium-133
Cs-133
133
55
Cs
How many protons, neutrons, and electrons
are present in the
56
+2
Fe
26
P = 26
N=m–z
56 – 26 = 30
E = 24
+2 indicates two
more protons
than electrons.
What is the chemical symbol for the ion
with 15 protons and 18 electrons?
P-3
31
15
P-3
Atomic Mass
Scientist’s chose to define an atomic mass unit in terms of
an arbitrary standard—a carbon-12 atom.
This means that 1 atomic mass unit is 1/12 of the mass of a
carbon-12 atom.
1 amu = 1/12 (mass of C-12 atom) = 1.673 x 10-24 g
Isotopes of Three Common Elements
Element
Symbol
Mass
Mass (amu)
Number
Carbon
C-12
C-13
12
13
12
13.003
98.89%
1.11%
12.01
Chlorine
Cl-35
Cl-37
35
37
34.969
36.966
75.53%
24.47%
35.45
Si-28
Si-29
Si-30
28
29
30
27.977
28.976
29.974
92.21%
4.70%
3.09%
Silicon
%
Average
abundance Atomic mass
28.09
How did you get the value for the Mass?
Get the mass of the isotope and multiply it by the % abundance.
For example: To determine the actual mass, as seen of the periodic table, for
Chlorine:
1) Determine the mass in amu of each isotope and then find the total mass by
adding up the masses of each isotope.
35
17
37
17
Cl
75.53% (35)(0.7553) = 26.4355
Cl
24.47%
(37)(0.2447) = 9.0539
35.4894 amu or g
Look at the mass on the periodic table…
Do you think there are really more isotopes of Chlorine?
Compute the average atomic mass for silicon
Si-28
92.21% abundant
28 x 0.9221 = 25.8188
Silicon
Si-29
4.70% abundant
29 x 0.0470 =
1.363
Si-30
3.09% abundant
30 x 0.0309 =
0.927
28.1088 amu or g
Radioactive Decay
 When atoms emit alpha, beta or gamma radiation, it
is undergoing a radioactive decay.
 Decay occurs due to instability within the
nucleus.
 As the ratio of protons to neutrons becomes
more skewed, the nucleus becomes more
unstable.
 All isotopes with an atomic number greater
than 83 are unstable.
 Nuclear equation: an equation that keeps
track of the reaction’s components.
 Left side of equation has parent nuclide.
 Arrow means “yield”
 Right side of equation has daughter nuclide
and radiation particle
Changes in the Nucleus
In Chemical reactions, atoms interact only through
their outer electrons while their nuclei remain unchanged.
Nuclear reactions change an atom’s nucleus while the electrons
remain unchanged.
Nuclear reactions produce three kinds of radiation:
alpha radiation
Beta radiation,
gamma radiation.
4
2
0
-1
a+2
b
g
Alpha (a )Radiation-Emission
1. Alpha radiation consists of rapidly
moving helium nucleus.
a. A helium nucleus consists of 2 protons and 2-neutrons.
b. A lpha radiation is represented by the following forms:
4
2
He+2
4He+2
4
a+2
There are others, but you
get the drift!!
2. Alpha particles travel at about one-tenth the speed of light.
3. A thin sheet of paper or clothing can stop alpha particles.
It will not penetrate the skin on your body.
4. In alpha emission, the parent nuclide decays into a
daughter nuclide by emitting an alpha particle.
230
90
Th 
4 a+2
2
+
226
Ra
88
**During alpha decay the emission results
in two less protons.
** Also notice there is a change in mass; it
decreases by 4.
Beta (b- )Radiation-Emission
1. Electrons or positrons are emitted from the nucleus
at a very high speed, often approaching the speed of
light ; thus < 3.00x108 m/s
a) A positron is a positively charged particle the
same size as an electron.
2. Very high kinetic energy associated with beta or positron
particle release.
3. Beta radiation is stopped by a few millimeters of
aluminum. It will penetrate the outside of your body
only 1-5 mm. It will do more damage to the inside of
your body
a. Beta radiation is represented by the following forms:
0
-1
e
0e
0b-1
There are others, but you
get the drift!!
4. The nuclear equation to represent the radioactive
decay neon-23 by b- emission is:
23
10
Ne 
0
-1
b
+
23
Na
11
**During the decay of Neon, the emission results in
the conversion of a neutron to a proton.
** Also notice there is no change in mass.
5. Nuclide that decay by beta emission have too many
neutrons in the nucleus for the number of protons
present.
Gamma (g )Radiation-Emission
1. Gamma radiation is a form of electromagnetic radiation
similar to, but more energetic than,X-rays.
2. Gamma radiation is a pure energy form and the
most dangerous of the three radiation types.
3. Gamma radiation penetrates the farthest; several
centimeters of lead, or an even greater thickness
of concrete. It will travel right through our bodies
and thus has the potential to do much damage.
4. Gamma radiation can be
emitted along with alpha
radiation and/or along
with beta radiation.
a. Gamma radiation is
represented by: 0 g ,
Write an equation for the alpha decay of uranium-238
235
92
U
231
90
Th +
4 +2
a
2
Write the equation for the beta decay of sodium-24.
24
11
Na 
24
12
Mg +
0b
-1
The amount of time a radioactive isotope
needs to decompose to half it’s initial mass.
AE = A0• 0.5
t
t½
•AE is the amount of substance left
•A0 is the original amount of substance
•t is the elapsed time
•t1/2 is the half-life of the substance
100.0 g of isotope W has a half-life of 60 seconds. How much
of isotope W is left after 3.0 minutes. Show your work.
AE = A0• 0.5
t
t½
3.0
1.0
= (100.0g)(0.5 ) = 12.50g
Sig. figs depend on given mass.
Solve the original equation for t½
AE = A0• 0.5
T1/2 =
log 0.5
t
t½
Given 150.0g of isotope Y.
IF over a period of 10 .0days
the decay left only 12.0 g,
what would Y’s half-life be?
Show your work.
* 10.0 days
log 12.0g
150.0 g
Ans. 2.74 days
Solve the original equation for t
AE = A0• 0.5
t=
t
t½
log 20.0g
250.0 g
log 0.5
The half-life of isotope X
is 10.0 years. How many
years would it take for a
250.g sample of X to
decay to 20.0 g? Show
your work.
* 10.0 days
Ans. 36.4 years