Transcript chapter2 - AlvarezHChem

William L Masterton
Cecile N. Hurley
http://academic.cengage.com/chemistry/masterton
Chapter 2
Atoms, Molecules, and Ions
Edward J. Neth • University of Connecticut
Introduction
• Atoms
• Composed of electrons, protons and neutrons
• Molecules
• Combinations of atoms
• Ions
• Charged particles
Greeks: Empedocles and Democritus
• Suggested the concept of atoms but were not taken
seriously or credited with an atomic theory
John Dalton: credited with the first atomic model
Figure 2.1 - John Dalton and Atomic Theory
Atomic Theory
1. An element is composed of tiny particles called
atoms
2. All atoms of the same element have the same
chemical properties
3. In an ordinary chemical reaction, atoms rearrange
their bonds but atoms are not created or destroyed
4. Compounds are formed when two or more atoms of
different element combine
Fundamental Laws of Matter
Law of Conservation of Mass
Matter is conserved in chemical reactions
This applies to all chemical reactions but DOES NOT
include nuclear reactions
Law of Constant Composition
Compound always contains the same elements in the
same proportions by mass.
Pure water has the same composition everywhere.
Law of Multiple Proportions
• The masses of one element that combine with a
fixed mass of the second element are in a ratio of
small whole numbers.
Compare CO and CO2
Figure A – The Law of Multiple Proportions
Two different oxides of chromium
Components of the Atom
• 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
• Finding the
Electrons:
Protons:
Neutrons:
J.J. Thomson
• Discovered the electron
Figure 2.2 – J.J. Thomson and Ernest Rutherford
Figure 2.3 – Cathode Ray Apparatus
Electrons
• First evidence for subatomic particles
• J.J. Thomson in 1897
• 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)
J.J. Thomson’s Model
• 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
Ernest Rutherford:
Discovered the nucleus of the atom
Gold Foil Experiment:
• Bombardment of gold foil with α 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 = nucleus
Figure 2.4 – Rutherford Backscattering
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
Atomic Mass
• The average mass of all of the isotopes of an
element accounting for their relative abundances
Table 2.1 – Subatomic Particles
Terminology
• Atomic number, Z
• Number of protons in the atom
• Mass number, A
• Number of protons plus number of neutrons
• Mass # = p+ + n0
Nuclear symbolism
A
Z
X
• A is the mass number
• Z is the atomic number
• X is the chemical symbol
Isotopes
• Isotopes are two atoms of the same element
• Same atomic number but differ in number of
neutrons
• Different mass numbers
• Mass # = p+ + n0
Example 2.1
Radioactivity
• Radioactive isotopes are unstable (Radioactive
decay is not a chemical process)
1. These isotopes decay over time
2. Emit other particles and are transformed into
other elements
• Particles emitted
1. Beta (β) particles: High speed electrons
2. Alpha (α) particles: helium nuclei
3. Gamma (γ) rays: high energy light
Nuclear Stability
• depends on the neutron/proton ratio
• For light elements, n/p is approximately 1/1
• For heavier elements, n/p is approximately 1.4/1
Figure 2.5 – The Nuclear Belt of Stability
2.3 Introduction to the Periodic Table
• Dmitri Mendeleev: 1836-1907
• Arranged elements by chemical properties
• Left space for elements unknown at the time
• Predicted detailed properties for several
undiscovered elements:
• Sc, Ga, Ge
• By 1886, all these elements had been discovered, and
with properties similar to those he predicted
Mendeleev’s
P.T.
Introduction to the Periodic Table
Modern Periodic Table
• Period – a horizontal row on the periodic table
• Group – a vertical column on the periodic table
• Blocks – sections of elements with common
properties
• Families – another name for group; emphasizes the
similarity in properties within a group
Blocks in the Periodic Table
• Main group elements
• 1-2, 13-18 OR roman numeral +A groups
• Transition Metals
• 3-12 OR non roman numeral groups
• Inner Transition/Rare Earth elements
• Bottom double rows
Families with Common Names (label on PT)
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Alkali Metals, Group 1(I)
Alkaline Earth Metals, Group 2 (II)
Halogens, Group 17 (VII)
Noble Gases, Group 18 (VIII)
A Look at the Sulfur Group
• Sulfur (nonmetal), antimony (metalloid) and silver
(metal)
Example 2.3
2.4 Molecules and Ions
• Molecule: Two or more atoms chemically combined
1. Atoms involved are often nonmetals
2. 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
Structural Formulas
• Structural formulas: a formulas that shows the
bonding patterns within the molecule
Ions
• A charged particle that is the result of the loss or
gain of electrons
• Cation – a positive ion (loss)
• Anion – a negative ion (gain)
• Examples:
• Na → Na+ + e• O + 2e- → O2-
Ionic Compounds
• Compounds formed from the electrostatic attraction
of oppositely charged particles
• Sodium chloride (NaCl): Sodium cations and chloride
anions associate into a continuous network
Forces:
• Ionic compounds are held together by strong forces
• Compounds are usually solids at room
temperature
• High melting points
• often water-soluble
Solutions:
• When an ionic compound dissolves in water, the ions
are released from each other
• conductivity – the ions in a solution support the
transmission of an electric current
• Strong electrolytes – solutions that are very good conductors
• Weak electrolytes – solutions that are poor conductors
• Nonelectrolytes – solutions that do NOT conduct
Figure 2.12 – Electrical Conductivity
Formulas for 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
Transition Metals
• Polyvalent – exhibit multiple positive charges
depending on conditions
• Iron
forms Fe2+ and Fe3+
• Lead
forms Pb2+ and Pb4+
Polyatomic Ions
• Groups of atoms may carry a charge; these are the
polyatomic ions
• OH• NH4+
Noble Gas Connections
• Atoms that are close to a noble gas (group 18 or VIII)
form ions that contain the same number of electrons
as the neighboring noble gas atom
• +1, +2, +3 skip -3, -2, -1 Noble Gases
Example 2.5
2.6 Naming of Compounds
Cations: element name
• Na+, sodium
• If polyvalent, a Roman numeral is used to denote the
charge
• Fe2+
iron(II)
Names of Compounds - Anions
• Monatomic anions are named by adding –ide to the
element name
• Oxygen becomes oxide, O2• Polyatomic ions keep their names
• To name an ionic compound: name the cation first,
then, name the anion (with the word 'ion' omitted). It
is not necessary to indicate the number of cations
and anions in the compound because it is
understood that the total positive charges carried by
the cations must equal the total negative charges
carried by the anions.
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KI potassium ion + iodide ion = potassium iodide
CoCl2 cobalt(II) ion + two chloride ions = cobalt(II) chloride
CoCl3cobalt(III) ion + three chloride ions = cobalt(III) chloride
Hg2Cl2mercury(I) ion + two chloride ions = mercury(I) chloride
AgNO3silver ion + nitrate ion = silver nitrate
Oxoanions
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Per _________ate
___________ate
___________ite
Hypo_________ite
Table 2.3 – Oxoanions of Nitrogen, Sulfur and
Chlorine
Binary Molecular Compounds
• Made of 2 nonmetal elements
• Never reduce subscripts
• Covalently bonded
Mono-1 di-2 tri-3 tetra-4 penta-5 hexa-6 hepta-7
octa-8 nona-9
• Systematic naming
1. First name is the first element, with prefix to for
number of atoms (EXCEPT NO MONO)
2. Second name is prefix with element name
changed to –ide (INCLUDE MONO)
Some Examples
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Diphosphorus pentaoxide
Sulfur dioxide
Dinitrogen tetraoxide
Hydrogen dioxide
Carbon monoxide
Phosphorus trichloride
Acids
• Ionic compounds with Hydrogen as the cation
• Naming:
• Common: (strong acids)
• HBr
HI
HCl
• H2SO4
HNO3
HClO4
• Br I Cl SO NO ClO 434
Oxyacids or Oxoacids:
• Acids with and oxoanion as the anion
Acids of Chlorine (example):
Examples:
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Hydrogen chloride (hydrochloric acid)
Nitric acid
Sulfuric acid
Hypobromous acid
Nitrous acid
Phosphoric acid