Download Ch. 4 Slides

Ch. 4: Atoms and Elements
Dr. Namphol Sinkaset
Chem 152: Introduction to
General Chemistry
I. Chapter Outline
I. Introduction
II. Atomic Theory
III. The Nuclear Atom
IV. Elements
V. The Periodic Table
VI. Ions
VII. Isotopes
VIII. Atomic Mass
I. Introduction
• Atoms are the
building blocks of
everything we
experience.
• What we smell, what
we feel, what we see.
• In this chapter, we
trace the history of
the atom and learn
about its makeup.
II. The Greeks
• From out perspective, matter can be
infinitely divided.
• However, Leucippus and Democritus
(5th century B.C.) believed there was a
limit.
 Eventually, you will reach something that
was “atomos” or “indivisible.”
• Unfortunately, their idea was not
accepted.
II. Revival of the Atom
• The idea of the atom lay dormant for
over 2000 years.
• John Dalton revived the idea in order to
explain 3 natural laws that puzzled
everyone at the time.
• Dalton’s Atomic Theory worked so well
that it was quickly accepted.
II. Postulates of Dalton’s Theory
1. Each element is composed of tiny,
indestructible particles called atoms.
2. All atoms of a given element have the
same mass and other properties that
distinguish them from the atoms of
other elements.
3. Atoms combine in simple, wholenumber ratios to form compounds.
II. Atoms
• Today, overwhelming evidence points
towards the existence of atoms.
• Atoms can be imaged and arranged!
III. Not “Atomos”
• Dalton’s theory treated atoms as
permanent, indestructible building
blocks that composed everything.
• However, J.J. Thomson discovered
electrons, which were much smaller
than an atom and negatively charged!
• Since atoms are neutral, where’s the
positive charge?
III. Plum Pudding
• J.J. Thomson
proposed the
plum pudding
model of the
atom.
 Electron “raisins”
 “Pudding” of
positive charge
III. Rutherford likes Plum
Pudding
• Ernest Rutherford was a student of J. J.
Thomson.
• He tried to prove the plum pudding
model by shooting a-particles at gold
foil.
• Note that a-particles are 7000x more
massive than an electron and have a
positive charge.
III. Rutherford’s Expectation
III. Rutherford’s a-Particle
Experiment
III. Conclusions from
Rutherford’s Experiment
• Most of an atom’s mass and all of its
positive charge exist in a nucleus.
• Most of an atom is empty space,
throughout which tiny electrons are
dispersed.
• By having equal numbers of positivelycharged particles (protons) and electrons,
an atom remains electrically neutral.
III. Rutherford’s Interpretation
III. The Nuclear Atom
• Surprisingly, an atom is mostly empty space!
• The nucleus holds 99.9% of the atom’s mass.
III. Components of an Atom
• Protons. Positively-charged particles in
the nucleus. Mass of 1.67262 x 10-27 kg
or 1.0073 amu.
• Neutrons. Neutral particles in the
nucleus. Mass of 1.67493 x 10-27 kg or
1.0087 amu.
• Electrons. Negatively-charged particles.
Mass of 9.1 x 10-31 kg or 0.00055 amu.
III. Charge
• Charge is a
fundamental property.
• To designate charge,
the sign GOES
AFTER the
magnitude, e.g. 2+.
• Matter is charge
neutral.
IV. An Atom’s Identity
• The number of protons in an atom
determines its elemental identity.
IV. Referring to Elements
• Since each element has a unique # of
protons, we could refer to elements
using Z, the atomic number, which
equals the # of protons in an atom.
 e.g. The Z = 2 element.
• More commonly, we use an element’s
name or chemical symbol.
 e.g. The element helium, or He.
IV. Chemical Symbols
• Chemical symbols are a one or two
letter abbreviation of an element’s
name.
• First letter always capitalized; second
letter is LOWERCASE.
• Some symbols are based on historical
names: e.g. Au from aurum.
IV. The Periodic Table
IV. Sample Problem
• Find the name and atomic number of
the following elements.
a)
b)
c)
d)
e)
V
N
Hg
Rh
Mo
V. Organizing Chemical Info
• Dmitri Mendeleev was the first to organize
information of elements according to periodic
law, i.e. when arranged properly, elements
show repeating properties.
V. Mendeleev’s Breakthrough
• Mendeleev placed
elements with
similar properties in
vertical columns.
• He left blank spaces
where he thought
elements should
exist.
V. Three Types of Elements
V. Sample Problem
• Categorize the elements below as either
a metal, nonmetal, or metalloid.
a)
b)
c)
d)
e)
f)
Ru (ruthenium)
Se (selenium)
I (Iodine)
Ba (barium)
Es (einsteinium)
Kr (krypton)
V. Main Group vs. Transition
V. Families of Elements
V. Sample Problem
• To which group (new numbering
system) does each of the following
elements belong? If the group has a
name, indicate that as well.
a)
b)
c)
d)
Br (bromine)
N (nitrogen)
Cs (cesium)
Mn (manganese)
VI. Atoms Can Lose/Gain e-’s
• In chemical reactions, it’s common for
atoms to lose or gain electrons and
become ions.
• ion: a particle that has a charge
• Examples:
 Na  Na+ + e I + e-  I -
VI. Origin of the Charge
• The charge arises from the different
number of protons and electrons in the
atom.
 Ion Charge = # protons - # electrons
• A neutral Na atom has 11 protons and
11 electrons. If it loses and electron…
 Ion Charge = 11 – 10 = 1+
VI. Cations and Anions
• An ion is fundamentally different than a
neutral atom, so it needs a different
name.
• cation: a positively-charged ion
• anion: a negatively-charged ion
• Note that cations and ions have
different properties than their parent
atoms, e.g. Na vs. Na+.
VI. Sample Problem
• Determine the charges of the ions
described below.
a) A chromium atom that has lost 3
electrons.
b) A sulfur atom that has gained 2 electrons.
c) An iron atom (Fe) that has 24 electrons.
d) A phosphorus atom (P) that has 18
electrons.
VI. Ions and the Periodic Table
• The charge of an ion can be predicted
by the position of its parent element on
the periodic table IF it’s a main group
element.
• Simply count the number of spaces to
the nearest noble gas (forward or
backward).
• If you go forward, it’s an anion; if you go
backward, it’s a cation.
VI. Predicting Ion Charge
VI. Sample Problem
• What are the ions that form from atoms
of the following elements?




aluminum (Al)
tellurium (Te)
rubidium (Rb)
oxygen (O)
VII. Isotopes
• Protons are the only thing that
determines the identity of an atom.
• Therefore, it’s possible for atoms of the
same element to have different masses
due to differing number of neutrons.
• isotopes: atoms with the same number
of protons, but different numbers of
neutrons
VII. Percent Natural Abundance
• The different types and amounts of each
isotope is determined by nature.
• Note that in an isotope, the # of neutrons
varies which makes the mass number (A)
vary as well.
VII. Referring to Isotopes
• Isotopes can be represented using the
A, Z, X symbol.
VII. Referring to Isotopes
• Alternatively, the X, A notation can be
used.
VII. Sample Problem
• How many protons and neutrons are in
a potassium isotope with a mass
number of 39? What are the three ways
to represent this isotope?
VIII. What’s the Mass of an
Atom?
• It depends!
• Are we talking about the mass of a
specific atom, i.e. a given isotope?
 If so, it’s just approximately the mass
number.
• Are we talking about in general?
 Then it’s more complicated…
VIII. Atomic Mass
• Not all atoms of the same element have
the same mass, but we can calculate an
average.
• The atomic mass is the weighted
average mass of an element which
accounts for all isotopes and their
percent natural abundances.
VIII. Calculating Atomic Mass
• The equation below enables calculation
of atomic mass.
Atomic mass  (fraction isotope 1  mass isotope 1) 
(fraction isotope 2  mass isotope 2) 
(fraction isotope 3  mass isotope 3)  
VIII. Sample Problem
• Calculate the atomic mass of
magnesium using the information in the
table below.
Isotope
Mass (amu)
Natural Abundance
Mg-24
23.99
78.99%
Mg-25
24.99
10.00%
Mg-26
25.98
11.01%