Download Chapter 5: Atomic Structure & The Periodic Table

Chapter 5: Atomic Structure & The
Periodic Table
Democritus– 4th century B.C., teacher in Greece,
first suggested the existence of atoms, lacked
experimental support because scientific testing
was unknown at the time.
2000 years after Democritus, the real nature of
atoms and observable changes at the atomic
level were established.
John Dalton (1766-1844)—English school teacher,
performed experiments to test and correct his
atomic theory.
Dalton’s Atomic Theory
• All elements are composed of tiny
indivisible particles called atoms. (*Now
known to be divisible—broken down into
subatomic particles)
• Atoms of the same element are identical.
Atoms of any one element are different
from those of any other element.
• Atoms of different elements can physically
mix together or can chemically combine
with one another in simple whole-number
ratios to form compounds
• Chemical Rxs. occur when atoms are
separated, joined, or rearranged. Atoms
of one element, however, are never
changed into atoms of another element as
a result of a chemical rx.
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Atom—smallest particle of an element that
retains the properties of that element
( a scanning tunneling microscope can be
used to view the surface of individual
atoms. Ex: pg 108 with gold atoms)
Subatomic Particles
Electrons—negatively charged subatomic
particles.
J.J. Thomson (1856-1940) discovered
electrons in 1897
• performed experiments that involved passing
electric current through gases at low
pressure. The gases were sealed in glass
tubes fitted at both ends with metal disks
called electrodes. (pg 109 apparatus used)
• Cathode-ray tube, the electrons travel as a
ray from the cathode(-) to the anode(+)
Robert A. Millikan (1868-1953)
• Mass of e- is 1/1840 the mass of a hydrogen
atom (“proton”)
• Electrons are negatively charged (1-)
E. Goldstein (1886) observed a cathode-ray
tube and found rays traveling in the direction
opposite to that of the cathode rays.
(+ charge)
Protons—positivley charged particle
mass= @ 1840 x that of an electron
James Chadwick (1891-1974) English
physicist, 1932 discovered/confirmed
neutrons
Neutrons—neutral particle
mass= to that of a proton
Properties of Subatomic Particles
Particle
Symbol
Relative
mass
Actual
mass (g)
e-
Relative
electric
charge
1-
Electron
1/1840
9.11 x
10-28
Proton
p+
1+
1
1.67 x
10-24
Neutron
n0
0
1
1.67 x
10-24
Atomic Nucleus
How are subatomic particles arranged in an
atom:
1911- Ernest Rutherford (1871-1937) tested
the theory of Atomic Structure
Used massive alpha particles—helium atoms
that have lost their 2e- and have a double +
charge because of the two remaining protons
Rutherford’s Gold-foil Experiment
• Rutherford directed a narrow beam of alpha
particles at a very thin sheet of gold foil.
• Alpha particles passed straight through the
gold atoms without deflection
• Some of the alpha particles bounced off the
gold foil at very large angles.
Rutherford’s Theory of the Atom
• Atom is mostly empty space (explaining
the lack of deflection of most of the alpha
particles)
• All the positive charge and almost all the
mass is concentrated in a small region
(nucleus). (accounts for the great
deflection of some of the alpha particles)
• Nucleus- the central core of an atom
composed of protons and neutrons. Tiny
compared to the atom as a whole.
Contains most of the atoms mass.
Atomic Number
• Atomic number= the number of protons in
the nucleus
• Identifies and element
• Elements are different because they
contain different number of protons
• Number protons=number electrons (atoms
electrically neutral
Mass Number
• Mass number= number of protons +
number of neutrons
• If you know the atomic number and the
mass number of an atom of any element,
you can determine the atom’s composition
• Number of neutrons= mass numberatomic number
Isotopes
• Atoms that have the same number of
protons but different number of neutrons
• Different mass numbers
• Chemically alike because they have the
same number of protons and electrons,
which are the subatomic particles
responsible for chemical behavior
• Ex: Carbon-12, Carbon-14
• Neon-20, Neon-21, Neon-22
Atomic Mass
• Atomic mass unit (amu)– 1/12 the mass of
a Carbon-12 atom
• Compare the relative masses of atoms
using a reference isotope as a standard.
• C-12 was assigned a mass of exactly 12
atomic mass units
• He-4 atom with a mass of 4.0026 amu,
has @ 1/3 the mass of a C-12 atom
• Ex: how many C-12 atoms would have @
the same mass as a Nickel-60 atom?
Answer: 5 C-12 atoms = the mass of 1
Nickel-60 atom
• Atomic masses are not in whole numbers
Ex: Cl amu= 35.453
relative abundance of the naturally
occurring isotopes of the element.
• In nature, most elements occur as a
mixture of two or more isotopes. Each
isotope of an element has a fixed mass
and a natural percent abundance
• Atomic mass—a weighted average mass of the
atoms in a naturally occurring sample of the
element
• Weighted average mass reflects both the mass
and the relative abundance of the isotopes as
they occur in nature
Ex: Element X has two natural isotopes. The
isotope with a mass of 10.012 amu has a relative
abundance of 19.91%. The isotope with a mass
of 11.009 amu has a relative abundance of
80.09%. Calculate the atomic mass of this
element.
Solve: (10.012 amu x 0.1991) + (11.009 amu x
0.8009) = 10.810 amu
Development of the Periodic Table
Dmitri Mendeleev
• Russian Chemist
• Constructed the first periodic table
• Periodic table- an arrangement of the
elements according to similarities in their
properties.
• Listed in columns by increasing atomic
mass
Henry Moseley
• British physicist
• Determined the atomic number of the
atoms of the elements
• Arranged elements in a table by order of
atomic number instead of atomic mass
• Today’s periodic table
• Each element is identified by its symbol
placed in a square
Periods- horizontal rows of periodic table
7 periods
number of elements/period ranges
from 2 to 32
Properties of elements within a period
change as you move across from element
to element
Periodic Law– when the elements are
arranged in order of increasing atomic
number, there is a periodic repetition of
their physical and chemical properties
• Group or Family– each vertical column of
elements
• Elements in any group or family have similar
physical and chemical properties
• Groups 1A—8A elements are called the
representative elements (have a wide range of
physical and chemical properties)
• Groups 1B—8B elements in the middle of the
table are called transition elements (metals)
• Bottom rows of elements under the main table
are called the inner transition elements (metals)
(aka rare-earth elements)
80% of all elements are metals, which are solids at
room temp. except for Hg (Mercury), a liquid
•
•
•
•
Group 1A- alkali metals
Group 2A- alkaline earth metals
Groups 3A-8A are nonmetals
Bordering the black line that divides
metals/nonmetals are metalloids
• Metalloids are elements with properties
that are intermediate between those of
metals and nonmetals
• Left side of the periodic table, except for
Hydrogen are metals
• Upper right corner of periodic table are
nonmetals
Metals:
• High electrical
conductivity
• High luster when
clean
• Ductile (can be drawn
into wires)
• Malleable (able to be
beaten into thin
sheets)
Nonmetals:
• Poor conductors of
electricity
• Non lustrous
• Brittle
Periodic Table