William L MastertonCecile N. Hurleyhttp://academic.cengage.com/chemistry/masterton

Edward J. Neth • University of Connecticut

Chapter 2 Atoms, Molecules, and Ions

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 is the mass number• Z is the atomic number• X is the chemical symbolXA

Z

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’sP.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)

• 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.

• 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

• 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

• 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:

• Hydrogen chloride (hydrochloric acid)• Nitric acid• Sulfuric acid• Hypobromous acid• Nitrous acid• Phosphoric acid