Download Introduction to atoms

Introduction to atoms and
molecules
Chapter 2-1 – 2-5
Chapter 5-7 and 5-9
Chapter 4-5 – 4-6
Key concepts in this unit
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Dalton’s atomic theory
The electron, proton, and neutron
Isotope, atomic number, mass number
Molecules and molecular formulas
Introduction to ions
Nomenclature for inorganic compounds
Concept of the atom
• Democritus (ancient Greek) -- 
(atomos); indivisible particle of a
substance
• Democritus’ ideas were promptly lost for
the next ~ 2000 years or so.
Dalton—birth of the atomic theory
(1803)
Dalton’s postulates:
1. Each element is composed of extremely
small, indivisible particles called atoms.
2. All atoms of a given element are identical.
Atoms of different elements are different,
and have different properties, including
mass.
Dalton’s atomic theory (con’t)
3. Atoms of an element are not changed into
different types of atoms in a chemical reaction
(i.e., you can’t turn lead into gold). Atoms are
not created or destroyed in chemical reactions.
4. Compounds are formed when atoms of more
than one element combine in small, whole
number ratios.
5. A given compound always has the same
relative number and kind of atoms.
Dalton’s theory explains laws
• law of constant composition
• law of conservation of matter
• law of multiple proportions
Law of multiple proportions
• Elements combining in more than one
proportion will do so in multiples of whole
numbers.
• Examples:
– H2O vs. H2O2.
– NO
NO2
N2O
Fundamental particles
• Masses of fundamental particles:
Particle
Mass (g)
Mass (amu)
Proton
1.67262158  10-24
1.00727646
Neutron
1.67492716  10-24
1.00866492
Electron
9.10938188  10-28
5.485779911  10-4
• Atomic mass unit:
1 amu = 1.66053873  10-24 g
Atoms
• Atoms vary by atomic number (Z), which
indicates ________________________.
• Atoms contain the same number of electrons
and protons (they have no charge).
• Atomic number increases from left to right on
periodic table.
• The atomic weight of the atom is indicated below
the symbol (in amu).
• Why are these weights not whole numbers?
Isotopes
(chap 5-7, 5-9)
• Atoms of an element have same number of
protons in the nucleus.
• Isotopes of an element vary in number of
neutrons, but not in number of protons.
• Atomic symbols: 126 C
(“carbon-12”)
– C: carbon
– Subscript number: atomic number; number of
protons
– Superscript number: mass number, number of
protons + neutrons
• What is the atomic symbol for carbon-14?
Isotope symbols
Atomic
symbol
Isotope
name
Atomic
number
Mass
number
No. of
protons
No. of
neutrons
Chlorine37
24
_
Na
197
118
Average atomic weight
• The average atomic weight (the one in the
periodic table) is a weighted average.
• Not all isotopes of an atom exist in the same
amounts, and this must be taken into account.
AWavg   (isotope _ mass)(isotope _ abundance)
• Example 1: Let’s assume carbon has only
2 isotopes, 13C and 12C. The natural
abundance of 13C is 1.0 %. What is the
average atomic mass?
• Example 2: Cl has 2 isotopes, 35C and
37C. The average atomic mass of Cl is
35.5 amu. What is the natural abundance
of the two isotopes?
Molecules
• A molecule is
__________________________________
_________________________________.
– In some cases, an atom is also a molecule.
Example: He, Ne, Ar.
Some elements exist in nature as
molecules.
– Diatomic molecular
elements:
– Polyatomic molecular
elements:
– Allotropes:
Formulas
• “shorthand” for writing and describing
molecular composition.
– Chemical formula:
– Structural formula:
Ionic compounds
• Ionic compounds are composed of
_________________________________.
– Cation:
– Anion:
• Ionic compounds do not form molecular units,
but are _______________ containing a defined
ratio of cations and anions, forming a
_________.
• Formula unit:
Nomenclature for inorganic
compounds
• The charge on an ion is designated by a
number written above and to the right of
the symbol
• Examples:
– Na1+ (or Na+)
– Al3+
– F– O2-
Periodic groups and ions
• Normally (but not always) atoms in the following groups
form ions of the designated charge.
• This commonly occurs such that the # of electrons in the
ion are equal to the # of electrons in the neighboring
noble gas. We will discuss why this is so in quantum
mechanics.
Group
Alkali metals (1A)
Alkaline earths (2A)
Al or B
Chalcogens (6A)
Halogens (7A)
Noble Gases (8A)
Common charge
1+
2+
3+
210
• Transition metals often form more than
one type of cation (i.e., different charges,
such as Fe+, Fe2+, Fe3+.)
• We have methods of writing formulas and
names to indicate what charge the ions
take.
Identifying ionic compounds
• Generally (but again, not always) ionic
compounds are formed by combining a metal
and a nonmetal. Molecular compounds are
generally formed by combining nonmetals.
• The sum of the charges in an ionic compound
must equal zero. The subscript of each ion
indicates how many of that ion are in the formula.
• Example: what is the formula for the ionic
compound formed when Al3+ combines with O2-?
Nomenclature rules
•
Cations
1.
Cations from metal atoms take the name of the metal.
•
2.
Examples:
Cations from transition metals have the charge indicated by a
Roman numeral.
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–
Examples:
Note: there is an older method of naming transition metal cations.
We will not use the older method in this course, but you should be
aware of it in case you ever run into it. (p.141)
Cations from nonmetal atoms end in –ium.
3.
•
Examples:
Table 2-3 (p. 55) and Table 4-11 (p. 142) contain the names and
symbols of common cations that you should know.
Anions
1. monoatomic anions are named by replacing
the ending of the element with –ide.
– Examples:
2.
polyatomic anions containing a variable number of
oxygens. These occur in ternary salts.
The common anion is named using the suffix –ate.
NO3XO 3- (X = F, Cl, Br, I)
CO32SO42PO43-
nitrate
fluorate, chlorate, bromate, iodate.
carbonate
sulfate
phosphate
• The suffix or prefix of the name changes when
the number of oxygens is increased or
decreased
# of oxygens added or subtracted to
common ion
prefix/suffix used
Add one oxygen
common ion (+0 oxygens)
subtract one oxygen
subtract two oxygens
per-___-ate
-ate
-ite
hypo-__-ite
Note that the CHARGES on a “family” of oxyanions doesn’t change when the
number of oxygens varies. The oxidation number of the central atom does change.
We’ll discuss this later.
Examples of ternary salts
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sodium sulfate
iron (II) oxide
ammonium chloride
lithium carbonite
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oxyanions with an H+ or 2 H+ added are indicated using the
word hydrogen, or dihydrogen.
examples:
(Sometimes the addition of a hydrogen is indicated by the prefix bi-. For
instance, HCO3- is the bicarbonate anion, as in NaHCO3, sodium
bicarbonate. Use of either bi or hydrogen is correct).
Table 2-3 (p. 55) and Table 4-11 (p. 142) contain the names
and symbols of common anions that you should
know. You should also be able to properly name the
relatives of the “common” oxyanions.
Acids
• For the moment, an acid is an ionic compound
that, when dissolved in water, yields a H+.
anion prefix/suffix
acid prefix/suffix
example
-ide
hydro-__-ic
per-___-ate
-ate
per-___-ic
-ic
-ite
hypo-___-ite
-ous
hypo-___-ous
HCl, hydrochloric
acid (from chloride
ion)
HClO4, perchloric
acid
HClO3, chloric acid
HClO2, chlorous acid
HClO, hypochlorous
acid
Double check: the formula for any ionic compound must = ZERO.
Binary molecular compounds
1.
2.
3.
4.
The atom farthest to the left
in the periodic table is
usually written first.
For elements in the same
group, the atom appearing
lower in the table is usually
written first.
the name of the second
element in the formula name
is given an –ide ending.
Greek prefixes indicate the
number of each type of atom
appearing in the molecule.
(p. 142) You should know
these prefixes.
one
two
three
four
five
six
seven
eight
nine
ten
monoditritetrapentahexheptoctnondec-