Download Chemistry 127 – Chapter 2 Atoms, Molecules and Ions

Chapter 2: Atoms, Molecules
and Ions
Joseph DePasquale
Building Blocks of Chemistry
2.1: Early Ideas in Atomic Theory
2.2: Evolution of Atomic Theory
2.3: Atomic Structure and Symbolism
2.4: Chemical Formulas
2.5: The Periodic Table
2.6: Molecular and Ionic Compounds
2.7: Chemical Nomenclature
2.1 Atomic Theory (1807)
John Dalton
1) Matter is composed of small particles
called atoms. An atom is the smallest unit
of an element that can participate in a
chemical change.
2) An element consists of only
one type of atom.
3) Atoms of one element differ
in properties from atoms of
all other elements.
2.1 Atomic Theory (1807)
John Dalton
4) A Compound consists of atoms of two or more
elements combined in a small, whole-number ratio. In a
given compound, the number of atoms of each of its
elements are always present in the same ratio.
2.1 Atomic Theory (1807)
John Dalton
5) Atoms are neither created nor destroyed during a
chemical change, but instead rearrange to yield a different
type(s) of matter.
Fundamental Laws of Matter
• Dalton’s Atomic Theory is the basis for
the three fundamental laws of matter.
1) Law of Conservation of Mass
• There is no change in mass in an ordinary
chemical reaction.
• If atoms are conserved in a reaction, then
mass must also be conserved.
• Matter can be neither created nor destroyed
Fundamental Laws of Matter
2) Law of Definite Proportions (or Constant
Composition)
• All samples of a pure compound always contains
the same elements in the same proportions by
mass.
• Example: Pure water has the same composition
everywhere.
Fundamental Laws of Matter
3) Law of Multiple Proportions
• Applies to situations where 2 elements form more
than one compound.
• The law states that when two elements react to form
more than one compound, a fixed mass of one
element will react with masses of the other element in
a ratio of small, whole number.
The Law of Multiple Proportions: Two
different chromium-oxygen compounds
Green: 2.167 g Cr/1 g O
Red: 1.083 g Cr/1 g O
2.2 Evolution of Atomic Theory
• Atomic theory raised more questions than it
answered.
• What are atoms composed of?
• Is there something smaller than an atom?
• Almost 100 years after atomic theory was proposed,
subatomic particles were discovered.
Discovery of the Electron
• J.J. Thomson (1897)
• Cathode Rays Experiment
• Cathode rays are emitted by all
materials.
• These cathode ray particles turned
out to be electrons.
• Electrons have a charge of -1
• Electrons have a very small mass.
Nobel Laureates:
J. J. Thompson &
Ernest Rutherford
Discovery of the Electron
Evidence that cathode rays are composed of
negatively charged particles, which we now refer
to as electrons:
1) Repelled by the negative end of a magnet.
2) Attracted to the positive end of a magnet.
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Thomson’s Model of the Atom
• “Plum Pudding Model”
• Atoms consisted of a positively
charged mass with negatively
charged electrons embedded
in that mass.
• Raisins = electrons
• Bread = Positively charge
Discovery of the Atomic
Nucleus
• Ernest Rutherford (1911)
• Gold Foil Scattering experiment
• Discovered the Atomic Nucleus and
determined that it is was positively
charged.
Nobel Laureates:
J. J. Thompson &
Ernest Rutherford
Rutherford’s Gold Foil Experiment
• Alpha particles (positively charged) were fired at gold foil.
• Most passed through the foil
• Some were reflected at sharp angles
Conclusions:
1) The center of an atom contains a small positively
charged center that contains most of the atom’s mass.
2) Most of the atom is “empty space”.
Structure/Components of an Atom
• Nucleus: Positively charged center. Composed of two
main parts
1) Protons: Charge = +1
2) Neutrons (Discovered by James Chadwick,
1932): Charge = 0
• Protons and neutrons are much heavier than
electrons. Over 99.9% of an atom’s mass is
concentrated in the nucleus.
• Electron Cloud: Negatively charged outer region.
Accounts for most of the atom’s volume, composed of
only electrons.
• Electrons have a charge = -1, very small mass, and
spread far apart.
2.3: Atomic Structure and Symbolism
Atomic Structure
• Diameter of an atom ~ 10-10 m
• Diameter of a nucleus is 100,000 times smaller ~ 10-15 m
• The nucleus accounts for most of the atom’s mass, but very
little of it’s volume.
Subatomic Particles
• Small units are needed to describe the properties of
subatomic particles.
• Atomic mass unit (amu)
• Fundamental unit of charge (e)
• A proton has a mass about 1800 times greater than
that of an electron.
• Neutrons are just slightly larger than protons.
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Atomic Number (Z)
• All atoms of a particular element have the same
number of protons in the nucleus.
• The number of protons in the nucleus of an atom is
its atomic number (Z).
• Therefore, all atoms of the same element have the
same atomic number.
• In a neutral atom:
Number of protons = Number of electrons
Ions
• When the number of protons and electrons are not equal,
the atom is electrically charged and is called an ion.
• Atomic charge =
• Atoms acquire charge by losing or gaining electrons.
• There is no change in the number of protons in the nucleus
when an ion forms. Only the number of electrons increases
or decreases.
Cations and Anions
• Positively charged ions are called cations.
• Cations form by loss of electrons.
• Negatively charged ions are called anions.
• Anions form by gain of electrons.
Mass Number (A)
• Mass number (A) – The total number of
protons and neutrons in an atom.
A = number of protons + number of neutrons
• Atoms of the same element do not always
have the same mass number.
• Atoms must have the same number of
protons, but the number of neutrons may vary.
Isotopes
• Atoms that contain the same number of
protons but a different number of neutrons are
called isotopes.
• Isotopes are atoms of the same element that
differ in mass.
Chemical Symbols
• Each element is assigned a symbol.
• Consists of one or two letters. Some newer ones contain
three letters.
• First letter is always capitalized. Second and third letters
are always lower case.
• The symbols are typically derived from the English name.
• Sometimes the symbol is derived from other languages.
Chemical Symbols
• The composition of an atom can be represented
by a chemical symbol.
Mass number
(number of protons + neutrons)
Atomic number
(number of protons)
A
Z
X
Elemental symbol
(from periodic table)
Three Isotopes of Hydrogen
Isotopes of Hydrogen
• Notice the ice on top
of the water.
• Notice some of the
ice on the bottom of
the glass.
Atomic Mass
• Each proton and each neutron has a mass of ~ 1 amu
• Each electron weighs far less.
• Therefore the atomic mass of a single atom in amu is
approx. equal to its mass number.
Atomic Mass
• However, most elements exist naturally as a mixture of
two or more isotopes.
• The periodic table shows the average atomic mass of
each element, which represents the average atomic
mass of the naturally occurring mixture of isotopes of
that element.
• Therefore, the atomic masses in the periodic table are
not whole numbers.
• Exception:
Average Atomic Mass of Carbon
• Carbon has two main isotopes, C-12 and C-13, which account
for basically 100% of naturally occurring carbon.
Isotope
Isotopic mass (amu)
Natural
abundance (%)
12C
12.00000
98.93
13C
13.003355
1.070
• Average atomic masses are calculated as weighted averages.
%Y1
%Y2
atomic mass Y = (atomic mass Y1 )
+ (atomic mass Y2 )
+...
100
100
Determination of Atomic Mass and
Isotopic Abundance
• An atom’s atomic mass can be determined to a
highly precise value by using a high tech
instrument known as a mass spectrometer.
• The mass spectrometer separates matter based
on its mass and charge.
• This data can be used to determine the
abundance and mass of each isotope in a
naturally occurring sample of that element.
Mass Spectrometer
(instrument)
Mass spectrum
(data)
2.4: Chemical Formulas
• Molecular Formula – A representation of a
molecule or compound which consists of the
following:
1) Chemical symbols to indicate the types of atoms.
2) Subscripts after the symbol to indicate the number of
each type of atom.
Molecular Formulas
• Examples:
• Water, H2O
• Ammonia, NH3
• Methane, CH4
Structural Formulas
• A Structural Formula shows the same information as
a molecular formula but also how the atoms are
connected.
Structural Formulas Showing the
Molecule’s Geometry
Space-Filling
Model
Ball-and-Stick
Model
More on geometries in Ch. 7…
Structural Formula: Caffeine
Molecular Elements
• Some elements exist as molecules:
• Diatomic molecules:
H2, N2, O2, F2, Cl2, Br2, I2,
• Others: P4, S8
Empirical Formulas
• Empirical Formula – Indicates the simplest whole
number ratio of the number of atoms in a
compound.
• Molecular formula – Indicates the actual number
of atoms in a molecule or compound.
Empirical Formula vs. Molecular Formula
• Dividing the subscripts in the molecular formula by the
lowest common denominator gives the empirical formula.
Isomers
• It is possible for the same atoms to be arranged in
different ways.
• Isomers – Compounds with the same molecular
formula, but different structural formula.
Isomers
• Isomers – Compounds with the same molecular
formula, but different structural formula.
• Consider the molecular formula, C2H4O2
2.5: The Periodic Table
• The first periodic table: Mendeleev (1869) and
Lothar Meyer (1870)
• First to propose primitive
versions of the periodic table.
The First Periodic Table
• Elements listed in order of increasing atomic mass and
grouped in columns by properties.
Modern Day Periodic Table
Modern Day Periodic Table
• Elements listed in order of increasing atomic number
and grouped in columns by properties.
• Periodic Law – The properties of the elements are
periodic functions of their atomic numbers.
• Periods (rows)
• First period:
• Second period:
• Third period:
• Groups (columns)
Elements are categorized as metals, nonmetals, or metalloids.
Metals are
good conductors
of heat and
electricity.
Nonmetals are
poor conductors
of heat and
electricity.
Metalloids have
properties
intermediate between that of metals and nonmetals.
Metals, Nonmetals, and Metalloids
Nonmetals
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Blocks in the Periodic Table
• Main group elements (or representative
elements)
• Groups: 1, 2, 13-18
• Transition metals
• Groups: 3-12
• Inner transition metals
• Lanthanides
• Actinides
Common Names for Main Group Elements
• Alkali Metals, Group 1 (except hydrogen)
• All soft, reactive metals.
• Alkaline Earth Metals, Group 2
• Harder and less reactive than the alkali metals.
• Halogens, Group 17
• All reactive non-metals.
• Noble Gases, Group 18
• All relatively unreactive gases.
2.6 Molecular and Ionic Compounds
• Isolated atoms rarely appear in nature.
• Only the noble gasses consist of individual, non-reactive
atoms.
• Atoms tend to combine with each other in various ways
to either form:
1) Molecules
2) Ionic Compounds
• During the formation of molecules and compounds, only
electrons interact. The nucleus remains unchanged
Formation of Monoatomic Ions
• Monoatomic Ions –
• The periodic table can help us predict the charge of the
ion formed by many main group elements.
• Many main group elements lose or gain electrons so that
they have the same number of electrons as a near by
noble gas.
Formation of Cations
• Many main group metals lose electrons to form cations.
• Group 1  +1
• Group 2  +2
Formation of Anions
• Many main group non-metals gain electrons to form
anions.
• Group 16  -2
• Group 17  -1
Transition Metal Cations
• Several transition metals can form cations of varying
charge.
• These metals typically DO NOT form ions that have
the same number of electrons as a noble gas.
• Examples:
• Iron commonly forms Fe2+ and Fe3+
• Chromium commonly forms Cr3+ and Cr6+
Polyatomic Ions
• All ions thus far have been monoatomic.
• Many important ions in chemistry contain more than
one atom. These are known as polyatomic ions.
• OH-, hydroxide ion
• NH4+, ammonium ion
• You can think of polyatomic ions as “charged molecules”
• Most of the polyatomic anions contain one or more
oxygen atoms and are referred to as oxyanions.
Table 2.5: Polyatomic Ions
Oxyanions
• There is a system to help you remember the oxyanions.
• When a nonmetal forms two oxyanions
• -ate is the suffix used for the ion with the larger number
of oxygen atoms.
• -ite is the suffix used for the ion with the smaller number
of oxygen atoms.
• When a nonmetal forms more than two oxyanions,
prefixes are used in addition to -ate and -ite
• per- (largest number of oxygens)
• hypo- (smallest number of oxygens)
Oxyanions of Nitrogen, Sulfur and Chlorine
Ionic Bonds
• When electrons are transferred and ions
form, an ionic bond results.
• Ionic bond – Electrostatic attraction that
holds ions together.
• Ionic bonds form as a result of the
transferring of electrons
• While covalent bonds form as a result of
sharing of electrons.
Ionic Compounds
• Compounds that contains ions held together by
ionic bonds are called ionic compounds.
• Ionic compounds typically contain a metal and
nonmetal.
Formulas of Ionic Compounds
• The compound’s formula shows the simplest ratio
of cations and anions needed to produce an overall
neutral compound.
• Charge balance – The total positive charge of the
cations must equal the total negative charge of the
anions.
NaCl, An Ionic Compound
• There are no discrete NaCl molecules, only Na+ and Clions in a continuous network.
Ionic compounds with polyatomic ions
Properties of Ionic Compounds
• Ionic compounds have very high melting and boiling
points.
• Melting requires the strong ionic bonds to be broken.
• Oppositely charged particles need to be separated.
• The stronger the bond, the more energy required to break
that bond.
• Ionic compounds tend to have high solubility in water.
• Molten ionic compounds and water solutions of
dissolved ionic compounds can conduct electricity.
Molecular Compounds
• Many compounds consist of discrete neutral
molecules.
• Molecular compounds form when atoms share
electrons, forming covalent bonds.
• Molecular compounds usually consist of all
nonmetals.
• Properties of molecular compounds:
2.7 Chemical Nomenclature
• Compounds are identified by their formulas
and names.
• Nomenclature – A collection of rules for
naming things.
• The rules for naming a compound depends
on its type.
• Ionic compounds – Metal and nonmetal
• Molecular compounds – All nonmetals
• Simple binary molecules
• Acids
Naming Ionic Compounds
• The name of the cation is always listed first and the
name of the anion is listed second.
• Monatomic Cations - Same name as the metal they
are derived from
• Na+:
• Al+3:
• For transition metals, roman numerals in parentheses
after the name are used to indicate the ion’s charge :
Fe2+:
Fe3+:
Naming Ionic Compounds
• Monatomic anions are named by using the name of
the element, but with its ending replaced with the
suffix, –ide.
• O2- is named
• S2- is named
• Cl- is named
• Polyatomic ions (both cations and anions): Just use
the name of the ion.
Binary Molecular Compounds
•
Consist of two non-metals.
•
Unlike ionic compounds, there is no simple way to
deduce the formula of a binary molecular
compound.
•
Systematic naming
1. The first word is the name of the first element in the
formula, with a Greek prefix if necessary
2. The second word consists of
•
•
The appropriate Greek prefix
The name of the second element with its ending replaced
with the suffix, -ide
• These prefixes must be committed to memory.
• The “mono” prefix, meaning one atom, is used only
with the second word in the name but never the first
word.
Common Names for
Molecular Compounds
These names are NOT derived from the naming
rules, but are commonly used.
What is water’s actual systematic name?
Acids
• Some molecular compounds contain H atoms that
ionize in water to produce H+ ions.
• These compounds are called acids. The H atom(s)
are always listed first in the molecular formula.
• Example: HCl
• As a molecule, HCl is named hydrogen chloride.
• When put in water, HCl is named hydrochloric acid
•Special acid naming rules when in water.
. The word “acid” is added as a second word
example, when the gas HCl (hydrogen chloride) is dissolved in water, the solution is called hydrochloric
ral other examples of this nomenclature are shown in Table 2.12.
Binary Acids
Names of Some Simple Acids
Name of Gas
Name of Acid
HF(g), hydrogen fluoride
HF(aq), hydrofluoric acid
HCl(g), hydrogen chloride
HCl(aq), hydrochloric acid
HBr(g), hydrogen bromide
HBr(aq), hydrobromic acid
HI(g), hydrogen iodide
HI(aq), hydroiodic acid
H2S(g), hydrogen sulfide
H2S(aq), hydrosulfuric acid
Table 2.12
When in water:
1) Change the word hydrogen to hydroacids 2) The name of the second element with its ending
y compounds
containing
threethe
or more
elements
as organic compounds or coordination compound
replaced
with
suffix,
–ic(such
acid
ect to specialized nomenclature rules that you will learn later. However, we will briefly discuss the imp
pounds known as oxyacids, compounds that contain hydrogen, oxygen, and at least one other elemen
bonded in such a way as to impart acidic properties to the compound (you will learn the details of thi
Oxyacids
• Many acids contain oxygen in addition to hydrogen
and are referred to as oxyacids.
• Oxyacids typically consist of hydrogen combined with
an oxyanion.
• The name of the oxyacid is derived from the name of
the oxyanion.
• Never use the prefix hydro-.
Oxyacid Naming Rules
• If the oxoanion ends with –ate
• Replace –ate with –ic acid
• If the oxoanion ends with –ite
• Replace –ite with –ous acid
• The prefixes, per- and hypo- are retained in the acid
name.