Chemical elements are the most basic substances in
chemistry, defined by a unique atomic number (Z), which is the number of
protons in an atom's nucleus. Every atom with the same atomic number belongs to
the same element.
Each element is identified by a name and a symbol—usually
one or two letters. The first letter is always capitalized, such as C for
carbon, O for oxygen, Ne for neon, and Si for silicon. Some symbols are derived
from Latin or other languages: Fe (iron from ferrum), Pb (lead from plumbum),
Na (sodium from natrium), K (potassium from kalium), and W
(tungsten from the German wolfram).
Elements with atomic numbers greater than 92, such as those
beyond uranium, are synthetic and created in laboratories using particle
accelerators. These elements are highly unstable and exist only for a short
time. By 2016, international scientific bodies had confirmed and named 112
elements, each with an official name and symbol.
Isotopes: Same Element, Different Mass
Atoms of the same element always have the same number of
protons but may vary in the number of neutrons. These variations are called isotopes.
The identity of an isotope is defined by its mass number (A), which is the sum
of protons and neutrons.
Example: Isotopes of Neon
- Neon-20
() — 10 protons, 10 neutrons
- Neon-21
() — 10 protons, 11 neutrons
- Neon-22
() — 10 protons, 12 neutrons
Natural abundance of these isotopes:
- Neon-20:
90.51%
- Neon-21:
0.27%
- Neon-22:
9.22%
Some elements have only one naturally occurring isotope. For
example, aluminum exists in nature only as Aluminum-27.
Ions: Gaining and Losing Electrons
Atoms can gain or lose electrons during chemical reactions,
forming charged particles called ions. The number of protons remains
unchanged, but the electron count varies:
- Loss
of electrons → Positive ion (cation)
- Gain
of electrons → Negative ion (anion)
Example:
- → 10 protons, 10
neutrons, 9 electrons
- → 10 protons, 10
neutrons, 8 electrons
Charge = Number of protons − Number of electrons
Another example is the oxygen ion , which has:
- 8
protons
- 8
neutrons
- 10
electrons
Determining Atomic Structure: Protons,
Neutrons, and Electrons
To identify the atomic structure of any atom or ion, use the
following relationships:
- Atomic
number (Z) = Number of protons
- Mass
number (A) = Protons + Neutrons
- Charge
= Protons − Electrons
Case Studies:
a) Barium-135 ()
- Z
= 56 →
56 protons
- A
= 135 →
135 − 56 = 79 neutrons
- Neutral
atom →
56 electrons
b) Selenium-80 ion ()
- Z
= 34 →
34 protons
- A
= 80 →
46 neutrons
- Charge
= −2 →
36 electrons
It’s common to simplify the notation and write 135Ba and
80Se²⁻.
Atomic Mass and Mass Spectrometry
The mass of an atom isn’t just the total of its protons and
neutrons. A small fraction of that mass is converted to energy (binding energy)
when the nucleus forms. This makes the actual atomic mass slightly less than
the sum of its parts.
To standardize atomic masses, the carbon-12 isotope
is assigned a value of 12 atomic mass units (u). Other atomic masses are
determined experimentally using a mass spectrometer.
How a Mass Spectrometer Works:
- Atoms
are ionized into charged particles.
- These
ions are accelerated through electric and magnetic fields.
- Based
on mass-to-charge ratios, ions are separated and detected.
The intensity of detection correlates with the abundance of
each isotope.
Example: Mercury Isotopes
- 196Hg
→ 0.146%
- 198Hg
→ 10.02%
- 199Hg
→ 16.84%
- 200Hg
→ 23.13%
- 201Hg
→ 13.22%
- 202Hg
→ 29.80%
- 204Hg
→ 6.85%
Although mass numbers are whole numbers, actual atomic
masses are slightly less due to energy loss. For instance, oxygen-16 has a mass
of 15.9949 u, slightly under 16.
Solving Atomic Mass Ratios
Using mass spectrometry data: If oxygen-16 is 1.06632
times heavier than nitrogen-15:
Given:
- Mass
of ¹⁶O = 15.9949 u
- Mass
ratio = 1.06632
Then: Mass of ¹⁵N = 15.9949 u ÷ 1.06632 = 15.0001
u
Key Insights to Remember
- Every
chemical element has a unique atomic number, defining its identity.
- Isotopes
differ by neutron number, not proton count.
- Ions
are atoms with unequal numbers of protons and electrons.
- Atomic
masses must be measured, not calculated, due to nuclear binding energy.
- Mass
spectrometry is a crucial tool in determining isotopic composition and
atomic weights.
- Understanding
these principles lays the groundwork for exploring chemical reactions,
molecular structures, and advanced atomic theory.
No comments:
Post a Comment