Transcript The Periodic Table, Atomic Structure, Isotopes, Ions and Nomenclature

Atomic Structure, Isotopes, And Ions
Lecture 3
Ch.2
Suggested HW: 15, 19, 23, 28,
38, 49, 52, 58
Understanding the Nature of Atoms
• If you cut a piece of graphite from the tip of a pencil into
smaller and smaller pieces, how far could you go?
• You would eventually end up with atoms (translates to
“indivisible” in greek) of pure carbon.
• You can not divide a carbon atom into smaller
pieces and still have carbon
Matter
• An atom is the smallest identifiable unit of an element
• The theory that all matter is composed of atoms grew out
of two primary laws
1. Law of constant composition
2. Law of conservation of mass
Law of Constant Composition
The relative amounts of each element in a
given substance are always the same,
regardless of how the substance was
made.
Law of Constant Composition
• For a molecule, AB:
• mass A + mass B = total mass AB
• %A + %B = 100
• %A =
𝑚𝑎𝑠𝑠 𝐴
𝑡𝑜𝑡𝑎𝑙 𝑚𝑎𝑠𝑠 𝐴𝐵
𝑥 100%
• Example: We analyze 1.630 g of CaS and find that it’s 0.906 g
Ca. Find the mass of S? Find the mass% of Ca and S?
* 55.6% Ca, 44.4% S
* This means that all pure CaS in the universe has the same
composition as calculated above, regardless of how it was made
or where it was found.
Law of Conservation of Mass
In a chemical reaction, atoms are not
created or destroyed, only rearranged.
The total mass of substances present
before a reaction is equal to the total
mass after.
Atomic Structure
• We have established that matter is comprised of atoms.
But what are atoms made of?
• In the 1800’s, physicists conducted numerous
experiments which revealed that the atom itself is made
up of even smaller, more fundamental particles.
• The three types of sub-atomic particles that make up the
atom are known as:
• electrons
• protons
• neutrons
Discovery of the Electron
J.J. Thomson’s Cathode Ray Experiment (late 1800’s)
• No Electric Field
• With Electric Field
Applying voltage to a metal cathode produces a beam of particles. The
beam can be deflected by electric fields towards a positive pole. Mass of
cathode plate does not change during this process. What does this mean?
Plum Pudding Model
• Atoms are charge neutral. If electrons
reside within an atom, then an
equivalent number of positive charges
must also exist, appropriately named
protons.
• How do all these charges coexist?
• Thomson proposed the very first
theoretical model of the atom, the socalled plum pudding model (PPM)
shown to the right.
• Electrons reside in a sea of uniform
positive charge
Protons and The Nucleus
• Ernest Rutherford sought to test
the PPM using the gold foil
experiment (below)
• A beam of positively charged
α-particles were focused on a
very thin sheet of gold
• Based on the PM model, this
beam would pass right through
the gold foil. In actuality, the
beam was deflected at odd
angles, with some α-particles
bouncing directly back!!
transmitted
through cloud
Electron cloud
THE ATOM
α particles
Nucleus
scattered,
repulsed
particles
True model of the atom is a dense,
positively charged, proton-loaded
nucleus surrounded by a sparse
electron cloud ! The vast majority
of an atom’s mass is contained
within the nucleus.
Neutrons
• Rutherford’s model was incomplete. For example, a hydrogen
atom has one proton and one electron, but is only ¼th the
mass of a helium atom which has two electrons and two
protons.
• If all of the mass of an atom comes from its sub-atomic
particles, how do we explain the unaccounted for mass?
• The answer is neutrons, particles that are equal in mass to
protons, but with no electrical charge.
Subatomic Particles and Their Relative Masses & Charges
Particle
Relative Charge
Mass (amu)
Charges shown in table are relative to the
charge of a proton. A proton has an actual
charge of 1.602 x 10-19 Coulombs (C), an
electron has a charge of -1.602 x 10-19 C.
Opposite charges attract! Like charges repel!!
Elemental Symbols
6
Atomic #
C
Carbon
12.0107
• The number of protons in an atom is called the atomic number.
An element is defined by its atomic number. (ex. only carbon has
6 protons)
• For a given element, the number of protons DOES NOT CHANGE
• In a neutral atom, the number of protons is equal to the number
of electrons.
Elemental Symbols
6
C
Carbon
12.0107
Mass #
• The mass number of an element is the sum of its protons
and neutrons.
• The mass #’s listed on the periodic table are averages (the
unit of atomic mass is the amu, or atomic mass unit).
• These averages are used because numerous variations of
elements called isotopes exist in nature.
Isotopes
• Isotopes are variations of elements with the same number
of protons but different numbers of neutrons. Isotope
symbols are shown below for the three isotopes of
nitrogen with their % abundances in nature. The 14N and
15N isotopes have 7 and 8 neutrons, respectively.
mass number
atomic number
𝟏𝟒
𝟕𝑵
(99.636%)
𝟏𝟓
𝟕𝑵
(0.346%)
Avg. atomic mass is obtained using the % abundance and the isotope mass.
Transitional Page
𝐴𝑣𝑒𝑟𝑎𝑔𝑒 𝑎𝑡𝑜𝑚𝑖𝑐 𝑚𝑎𝑠𝑠 =
𝑖𝑠𝑜𝑡𝑜𝑝𝑒 𝑚𝑎𝑠𝑠 𝑥 (% 𝑎𝑏𝑢𝑛𝑑𝑎𝑛𝑐𝑒)
Group Work
• For the table below, fill in the numbers of protons, neutrons,
and electrons for each isotope of carbon.
• Then, using the given abundances and isotope masses,
calculate the average atomic mass of C. Does it match the
reported value?
ISOTOPE
P
N
E
%A
Mass (amu)
𝟏𝟐
𝟔𝑪
98.93
12
𝟏𝟑
𝟔𝑪
1.07
13.003 354 8378
𝟏𝟒
𝟔𝑪
~0
14.003 2420
Group Work
• Boron has two isotopes, 10B and 11B. Using the given isotope
masses, determine the % abundances of each isotope.
ISOTOPE
%A
Mass (amu)
𝟏𝟎
𝟓𝑩
10.013
𝟏𝟏
𝟓𝑩
11.009
Proton-Neutron Ratio and Radioactivity
• The nuclei of most naturally occurring isotopes are very
stable, despite the massive repulsive forces that exist between
the protons in the nucleus.
• A strong force of attraction between neutrons and protons
known as the nuclear force counteracts this repulsion.
• As the number of protons increases, more neutrons are
required to stabilize the atom. Stable nuclei up to atomic
number 20 have equal numbers of protons and neutrons.
• For nuclei with atomic number above 20, the number of
neutrons exceeds the protons to create a stable nucleus.
Proton-Neutron Ratio and Radioactivity
• Radioactive isotopes are unstable (high in energy). This
instability is attributed to a neutron/proton ratio that is either
too high or too low.
• To become stable, they spontaneously release particles or
radiation to lower their energy.
• This release of energy is called radioactive decay.
Radioactivity
• The three most common types of radioactive decay are alpha,
beta, and gamma
Property
α
β
γ
Charge
2+
1-
0
Mass
6.64 x 10-24 g
9.11 x 10-28 g
0
Emitted Radiation
Type
2 protons and 2
neutrons ( 42𝐻𝑒)
High energy
electron.
Pure energy
(Radiation)
Penetrating Power
Low. Stopped by
paper. Blocked by
skin.
Moderate.
Stopped by
aluminum foil.
(10α)
High. Can
penetrate several
inches of lead.
(10000α)
Radioactivity
• For example, the 238
92𝑈 isotope undergoes alpha decay to
decrease its n/p ratio:
238
92𝑈
→
234
90𝑇ℎ
+ 42𝐻𝑒
• The Thorium-234 isotope then undergoes beta decay which
lowers the ratio even more:
234
90𝑇ℎ
→
234
91𝑃𝑎
+
0
−1𝑒
– In beta decay, a neutron is converted to a proton and an
electron. This causes the proton count to increase:
1
0𝑛
→ 11𝑝 + −10𝑒
• Gamma (γ) decay usually accompanies α or β decay to release
residual excess energy. γ is not shown in equations.
Applications of Radiochemistry: Carbon Dating
𝟏𝟒
𝟕𝑵
+ 𝟏𝟎𝒏 →
𝟏𝟓
𝟕𝑵
𝟏𝟓
𝟕𝑵
→
𝟏𝟒
𝟔𝑪
+ 𝟏𝟏𝒑
𝟏𝟒
𝟔𝑪
→
𝟏𝟒
𝟕𝑵
+ −𝟏𝟎𝒆
Beta decay!
Half life = 5700 yrs
Applications of Radiochemistry: Radiomedicine
131I
treatment for thyroid cancer
Radioactive tracers can be linked to chemical compounds to allow doctors
to monitor physiological processes. These compounds ‘glow’ upon decay
via γ-decay
• Organ malfunction can be indicated if a radioisotope is taken
up either too little or too much.
• Can monitor blood flow
• Intestinal blockages can be detected by accumulation of tracer.
• Tumor detection
• 18F, 99Tc
Ions
• Thus far, we’ve learned that each element has an exact number
of protons.
– For example, Hydrogen has only one proton. If you force a
second proton onto the atom, you no longer have hydrogen…
you now have Helium.
• We have also learned that atoms can have variable numbers of
neutrons (isotopes).
• Next, we will discuss ions.
Ions
• Ions are electrically charged atoms, resulting from the gain
or loss of electrons.
• Positively charged ions are called cations. You form cations
when electrons are lost
• Negatively charged ions are called anions. You form anions
when electrons are gained
Ion Nomenclature
• A cation is named by adding the word “ion” to the end of the
element name
• Anions are named by adding the suffix –ide to the end of an
element
𝑳𝒊+
Lithium ion
𝑪𝒍−
Chloride
𝑵𝒂+
Sodium ion
𝑺𝟐−
Sulfide
Magnesium ion
𝑶𝟐−
Oxide
Aluminum ion
𝑷𝟑−
Phosphide
𝑴𝒈𝟐+
𝑨𝒍𝟑+
Group Work
• Fill in the missing information below
ISOTOPE
P
N
E
13
14
10
32
16𝑆
32 216𝑆
??
?? 4+
??𝑃𝑡
95