Transcript Masterton and Hurley Chapter 2

Chapter 2
Atoms, Molecules, and Ions
Learning a Language
• When learning a new language:
• Start with the alphabet
• Then, form words
• Finally, form more complex structures such as
sentences
• Chemistry has an alphabet and a language; in this
chapter, the fundamentals of the language of
chemistry will be introduced
Outline
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Atoms and Atomic Theory
Components of the Atom
Introduction to the Periodic Table
Molecules and Ions
Formulas of Ionic Compounds
Names of Compounds
The Language of Chemistry
• This chapter introduces the fundamental language of
chemistry
• Atoms, molecules and ions
• Formulas
• Names
The Structure of Matter
• Atoms
• Composed of electrons, protons and neutrons
• Molecules
• Combinations of atoms
• Ions
• Charged particles
Laws of Chemical Composition
Conservation of Mass - The total mass remains
constant during a chemical reaction.
EX: A bicycle rusting
Law of Definite Proportions
Law of Definite Proportions : All samples of a
compound have the same composition; that is, all
samples have the same proportions, by mass, of the
elements present
Ex: Water always contains:
~89% oxygen
~11% hydrogen
John Dalton
Atoms and Atomic Theory
• An element is composed of tiny particles called
atoms
• All atoms of the same element have the same
chemical properties
• In an ordinary chemical reaction
• There is a change in the way atoms are combined
with each other
• Atoms are not created or destroyed
• Compounds are formed when two or more atoms of
different element combine
• Atoms are indivisible and indestructible
Figure 2.1 - John Dalton and Atomic Theory
Law of Multiple Proportions
Law of Multiple Proportion: When two or more different
compounds of the same two elements are compared,
the masses of one element that combine with the
fixed mass of the second element are in the ratio of
small whole numbers.
Figure A – The Law of Multiple Proportions
Two different oxides of chromium
Law of Multiple Proportions
• Atomic theory raised more questions than it
answered
• Could atoms be broken down into smaller
particles
• 100 years after atomic theory was proposed, the
answers were provided by experiment
Fundamental Experiments
• J.J. Thomson, Cavendish Laboratories, Cambridge,
England
• Ernest Rutherford
• McGill University, Canada
• Manchester and Cambridge Universities, England
Figure 2.2 – J.J. Thomson and Ernest Rutherford
Electrons
J.J. Thomson, 1897
• First evidence for subatomic particles came from the
study of the conduction of electricity by gases at low
pressures
• Rays emitted were called cathode rays
• Rays are composed of negatively charged
particles called electrons
• Electrons carry unit negative charge (-1) and have
a very small mass (1/2000 the lightest atomic
mass)
The Electron and the Atom
• Every atom has at least one electron
• Atoms are known that have one hundred or more
electrons
• There is one electron for each positive charge in an
atom
• Electrical neutrality is maintained
Figure 2.3 – Cathode Ray Apparatus
Robert Millikan and the Oil Drop Experiment,
1909
• Produced small oil drops, acquire an electric
charge, measure velocity of a falling droplet
with and without an electric field
• Obtained the charge on an electron, which coupled
with Thomson’s work, allowed the calculation of the
mass of an electron.
• Calculated the charge of an electron:
-1.602 * 10-19 coulomb (C)
• Calculated the mass of an electron:
9.109 * 10-31 kg
Protons and Neutrons – The Nucleus
• Ernest Rutherford, 1911
• Bombardment of gold foil with alpha particles (helium
atoms minus their electrons
• Expected to see the particles pass through the foil
• Found that some of the alpha particles were
deflected by the foil
• Led to the discovery of a region of heavy mass at
the center of the atom
Rutherford’s Gold Foil
Experiment
• to prove or disprove Thomson’s model
• alpha particles (Helium atoms, 2 protons and 2
neutrons) shot at gold foil
• most particles went through, but some came back at
different angles
• “It is about as incredible as if you had fired a 15-inch
shell at a piece of tissue paper and it came back and
hit you.” -ER
• discovered the nucleus (positive core of an atom)
Figure 2.4 – Rutherford Backscattering
Thomson’s Plum Pudding Model
• aka chocolate chip cookie model
• positive charge uniformly distributed in a sphere
• electrons imbedded in the sphere
YUMMY!!!!
Rutherford's Model
• Nucleus in the center of the atom containing protons
• Electrons in the space surrounding the nucleus
Nuclear Particles
1. Protons
• Mass nearly equal to the H atom
• Positive charge
2. Neutrons
• Mass slightly greater than that of the proton
• No charge
Mass and the Atom
• More than 99.9% of the atomic mass is concentrated
in the nucleus
• The volume of the nucleus is much smaller than the
volume of the atom
Table 2.1 – Subatomic Particles
Protons and neutrons are located at the center of an atom
(at the nucleus).
Electrons are dispersed around the nucleus.
Terminology
• Mass number, A =
• Number of protons plus number of neutrons
• Atomic number, Z=
• Number of protons in the atom
Nuclear symbolism
A
Z
X
• A is the mass number
• Z is the atomic number
• X is the chemical symbol
19 1
F
9
Isotopes
• Isotopes are two atoms of
the same element
• Same atomic number
• Different mass numbers
• Number of neutrons is A-Z
• Number of neutrons differs
between isotopes
• Isotopes & Their
Uses
• Heart scans with
radioactive
technetium-99.
•
99
43Tc
• Emits gamma rays
Isotopes of hydrogen
• 1H, 2H, 3H
• Hydrogen, deuterium, tritium
• Different masses
1 H
1
1 proton and 0 neutrons,
protium
2 H 1 proton and 1 neutron,
1
deuterium
3 H 1 proton and 2 neutrons,
1
tritium radioactive
Note that some of
the ice is at the
bottom of the glass
– this is 2H2O
Isotopes
Atoms can be represented using the element’s symbol
and the mass number (A) and atomic number (Z):
A
E
Z
35
Cl
17
37
Cl
17
Isotopes
• Example 2.1 Write the atomic symbols for the
following species:
• a. the isotope of carbon with a mass of 13
• b. the nuclear symbol when Z= 92 and then number
of neutrons = 146
Answers:
13
C
6
238 U
92
Ions
• An electrically charged particle comprised of one or
more atoms
Example 2.2
• Write the atomic symbols for the following:
• a. a species having 16 protons, 16 neutrons and 18
electrons
• b. the phosphide ion (symbol = P) with an overall
charge of -3
Answers:
32
S 2-
16
31 P 315
An atomic mass unit (amu) is defined as exactly onetwelfth the mass of a carbon-12 atom
• 1 u = 1.66054 × 10–24 g
• The atomic mass of an element is the weighted
average of all of the isotopes an element has
• Weighted average is the addition of the contributions
from each isotope
• percent abundance is the percent or fraction of each
isotope found in nature.
• Example 2.3 Determine the average atomic mass of
magnesium which has three isotopes with the
following masses: 23.98 (78.6%), 24.98 (10.1%),
25.98 (11.3%).
Radioactivity
• Radioactive isotopes are unstable
• These isotopes decay over time
• Emit other particles and are transformed into other
elements
• Radioactive decay is not a chemical process!
• Particles emitted
• High speed electrons: β (beta) particles
• Alpha (α) particles: helium nuclei
• Gamma (γ) rays: high energy light
Nuclear Stability
• Nuclear stability depends on the neutron/proton ratio
• For light elements, n/p is approximately 1
• For heavier elements, n/p is approximately 1.4/1
• The belt of stability
Figure 2.5 – The Nuclear Belt of Stability
Introduction to the Periodic Table
The Periodic Table: Elements Organized
• Know location and description of:
• groups or families
• periods or series
• metals, metalloids, nonmetals and their properties
• main group elements
• transition metals
• lanthanides and actinides
Periods and Groups
• Vertical columns are groups
• IUPAC convention: use numbers 1-18
Families with Common Names
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Alkali Metals, Group 1
Alkaline Earth Metals, Group 2
Halogens, Group 17
Noble Gases, Group 18
Importance of Families
• Elements within a family have similar chemical
properties
• Alkali metals are all soft, reactive metals
• Noble gases are all relatively unreactive gases;
He, Ne and Ar do not form compounds
Group 1A: Alkali Metals
Li, Na, K, Rb, Cs
Reaction of potassium + H2O
Cutting sodium metal
Group 2A: Alkaline Earth Metals
Be, Mg, Ca, Sr, Ba, Ra
Magnesium
Magnesium
oxide
Group 3A: B, Al, Ga, In, Tl
• Al resists corrosion
Gallium is one of the few metals
that can be liquid at room temp.
Group 4A: C, Si, Ge, Sn, Pb
Quartz, SiO2
Diamond
Group 5A: N, P, As, Sb, Bi
Ammonia, NH3
White and red
phosphorus
Group 6A: O, S, Se, Te, Po
Sulfuric acid dripping from snot-tite
in cave in Mexico
Elemental S has a ring structure.
Group 7A: Halogens
F, Cl, Br, I, At
Group 8A: Noble Gases
He, Ne, Ar, Kr, Xe, Rn
• Horizontal rows are periods
• First period is H and He
• Second period is Li-Ne
• Third Period is Na-Ar
Arrangement of Elements
• Periods
• Arranged by increasing atomic number
• Families
• Arranged by chemical properties
Metals and Nonmetals
• Diagonal line starting with B separates the metals
from the nonmetals
• Elements along this diagonal have some of the
properties of metals and some of the properties of
nonmetals
• Metalloids
• B, Si, Ge, As, Sb, Te
Blocks in the Periodic Table
• Main group elements
• 1, 2, 13-18
• Transition Metals
• 3-12
• Post-transition metals
• Elements in groups 13-15 to the right of the transition
metals
• Ga, In, Tl, Sn, Pb, Bi
• Lanthanides and Actinides
• Below the main part of the table
A Look at the Sulfur Group
• Sulfur (nonmetal), antimony (metalloid) and silver
(metal)
Biological View of the Periodic Table
• “Good guys”
• Essential to life
• Carbon, hydrogen, oxygen, sulfur and others
• “Bad guys”
• Toxic or lethal
• Some elements are essential but become toxic at
higher concentrations
• Selenium
Figure 2.8 – Biologically Important and Toxic
Elements
Mendeleev
• Dmitri Mendeleev, 1836-1907
• Arranged elements by chemical properties
• Left space for elements unknown at the time
• Predicted detailed properties for elements as yet
unknown
• Sc, Ga, Ge
• By 1886, all these elements had been discovered, and
with properties similar to those he predicted
Moseley
• Arranged the periodic table according to atomic
number
Molecule
• Two or more atoms may combine to form a molecule
• Atoms involved are often nonmetals
• Covalent bonds are strong forces that hold the
atoms together
• Molecular formulas
• Number of each atom is indicated by a subscript
• Examples
• Water, H2O
• Ammonia, NH3
Empirical and Molecular Formulas
Empirical formula: the simplest whole number ratio
of elements in a compound
• A molecular formula : gives the number of each
kind of atom in a molecule.
• Example:
Molecular formula of glucose – C6H12O6
The elemental ratio C:H:O is 1:2:1, so the empirical
formula is CH2O
Structural Formulas
• Structural formulas show the bonding patterns within
the molecule
Structural Formulas
• Condensed structural formulas suggest the bonding
pattern and highlight specific parts of a molecule,
such as the reactive group of atoms
Ball and Stick Models
Ions
• When atoms or molecules lose or gain electrons,
they form charged particles called ions
• Na  Na+ + e• O + 2e-  O2• Positively charged ions are called cations
• Negatively charged ions are called anions
• There is no change in the number of protons in
the nucleus when an ion forms.
Example 2.3
Polyatomic Ions
• Groups of atoms may carry a charge; these are the
polyatomic ions
• OH• NH4+
Ionic Compounds
• Compounds can form between anions and cations
• Sodium chloride, NaCl
• Sodium cations and chloride ions associate into a
continuous network
Forces Between Ions
• Ionic compounds are held together by strong forces
• Electrostatic attraction of + and – for each other
• Compounds are usually solids at room
temperature
• High melting points
• May be water-soluble
Solutions of Ionic Compounds
• When an ionic compound dissolves in water, the ions
are released from each other
• Strong electrolytesPresence of ions in the solution leads to electrical
conductivity
• When molecular compounds dissolve in water, no
ions are formed
• NonelectrolytesWithout ions, solution does not conduct electricity
Figure 2.12 – Electrical Conductivity
Formulas of Ionic Compounds
• Charge balance
• Each positive charge must have a negative
charge to balance it
• Calcium chloride, CaCl2
• Ca2+
• Two Cl- ions are required for charge balance
Noble Gas Connections
• Atoms that are close to a noble gas (group 18) form
ions that contain the same number of electrons as
the neighboring noble gas atom
• Applies to Groups 1, 2, 16 and 17, plus Al (Al3+) and
N (N3-)
Cations of Transition and Post-Transition Metals
• Iron
• Commonly forms Fe2+ and Fe3+
• Lead
• Commonly forms Pb2+ and Pb4+
Copper forms either
copper(I) or copper(II) ions.
Titanium forms both
titanium(II) and
titanium(IV) ions.
What is the charge on a
zirconium(IV) ion?
Polyatomic Ions
• There are only two common polyatomic cations
• NH4+ and Hg22+
• All other common polyatomic ions are anions
Table 2.2 – Polyatomic ions
Oxoanions
• When a nonmetal forms two oxoanions
• -ate is used for the one with the larger number of
oxygens
• -ite is used for the one with the smaller number of
oxygens
• When a nonmetal forms more than two oxoanions,
prefixes are used
• per (largest number of oxygens)
• hypo (smallest number of oxygens)
Polyatomic Ions
• Oxoanions: the anions are composed of oxygen and
one other element
• Ex:SO42- (sulfate), NO2- (nitrite) , MnO4(permanganate)
• two oxoanions of the same element
• The anion with the smaller number of oxygens uses
the roots of the element plus “ite”
• The higher number use the root plus “ate”
• Ex: SO32- sulfite, NO2- nitrite, PO3-3
phosphite
•
SO42- sulfate,NO3- nitrate, PO4-3 phosphate
Table 2.3 – Oxoanions of Nitrogen, Sulfur and
Chlorine
2.6 Names of Compounds - Cations
Monatomic cations take the name from the metal from
which they form
• Na+, sodium ion
• K+, potassium ion
If more than one charge is possible, a Roman numeral
is used to denote the charge
• Fe2+
• Fe3+
iron(II) ion
iron(III) ion
Metals with multiple oxidation states
• Two methods: Stock method and “classical” method
• Stock system:
• metal name and the oxidation state in Roman
numbers in parenthesis
• Ex: Fe 2+ = iron(II)
• Form compound by balance charge of metal with
correct number of nonmetals
• Ex: CoCl3 = cobalt(III) chloride
Classical Method
The name of the metal ion that has the lower charge ends in “ous”, the
higher charge ends in “ic”
Names of Compounds - Anions
• Monatomic anions are named by adding –ide to the
stem of the name of the element from which they
form
• Oxygen becomes oxide, O2• Sulfur becomes sulfide, S2• Polyatomic ions are given special names (see table
2.3, p. 39)
Ionic Compounds
• Combine the name of the cation with name of the
anion
• Cr(NO3)3, chromium(III) nitrate
• SnCl2, tin(II) chloride
Write the formulas for:
• rubidium bromide
strontium hydroxide
• barium nitride
cobalt (II) sulfate
• cobalt (II) bromide
calcium phosphate
• strontium sulfide
tin (IV) phosphate
Write the formulas for:
•
•
•
•
RbBr
Ba3N
CoBr2
SrS
Sr(OH) 2
CoSO4
Ca3(PO4) 2
Sn3(PO4) 4
Example 2.6
Example:
Name the following binary ionic compounds
•
Metal nonmetal compound name
• KI
potassium iodine
potassium iodide
• Li2S lithium
sulfur
lithium sulfide
• Mg3N2 magnesium nitrogen magnesium nitride
Write the names for:
• AlCl3
(NH4)2S
K2Cr2O7
Ca3P2
Al(NO2)3
KMnO4
FeI2
PbO2
Write the names for:
• aluminum chloride
• potassium dichromate
• aluminum nitrite
ammonium sulfide
calcium phosphide
potassium permanganate
• Iron II idodide
Lead (IV) oxide
Binary Molecular Compounds
•
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
• Mono is never used on the first word
2. The second word consists of
•
•
•
The appropriate Greek prefix
The stem of the name of the second element
The suffix -ide
Writing Formulas of Binary Molecular Compounds
• Find the number associated with the prefixes and put
it after the element.
• One is never shown.
Table 2.4 - Greek Prefixes
Names of Binary Compounds
The lines trace a continuous
path from boron (B) to
fluorine (F). The element
closer to the beginning of
this path is generally written
first in the formula of a
binary molecular compound.
Some Examples
• Binary nonmetallic compounds
• N2O5, dinitrogen pentaoxide
• N2O4, dinitrogen tetraoxide
• NO2, nitrogen dioxide
• N2O3, dinitrogen trioxide
• NO, nitrogen oxide
• N2O, dinitrogen oxide
• Common names
• H2O, water
• H2O2, hydrogen peroxide
Common Molecular Compounds
Example 2.7
Write names or formulas for the following:
• B2O3
• AsO5
• As2O7
tetraphosphorus pentachloride
dihyrdogen monoxide
Answers:
• Diboron trioxide
• Arsenic pentoxide
• Arsenic heptoxide
P4Cl5
H2O
Acids
• Acids ionize to form H+ ions
• Hydrogen and chlorine
• As a molecule, HCl is hydrogen chloride
• When put in water, HCl is hydrochloric acid
Common Acids
Oxoacids
• Two common oxoacids
• HNO3, nitric acid
• H2SO4, sulfuric acid
Oxoacids of Chlorine
Example 2.8
Hydrates
A hydrate is an ionic compound in which the
formula unit includes a fixed number of water
molecules associated with cations and anions
Examples:
BaCl2 . 2 H2O
CuSO4 . 5 H2O