
# Chapter 2 Atoms, Molecules, and Ions

Shaun Williams, PhD

## The Early History of Chemistry

### Early History of Chemistry

• Greeks were the first to attempt to explain why chemical changes occur.
• Alchemy dominated for 2000 years.
• Several elements discovered.
• Mineral acids prepared.
• Robert Boyle was the first "chemist".
• Performed quantitative experiments.
• Developed first experimental definition of an element.

## Fundamental Chemical Laws

### Three Important Laws

• Law of conservation of mass (Lavoisier):
• Mass is neither created nor destroyed in a chemical reaction.
• Law of definite proportion (Proust):
• A given compound always contains exactly the same proportion of elements by mass.
• Law of multiple proportions (Dalton):
• When two elements form a series of compounds, the ratios of the masses of the second element that combine with 1 gram of the first element can always be reduced to small whole numbers.

## Dalton's Atomic Theory

### Dalton's Atomic Theory (1808)

• Each element is made up of tiny particles called atoms.
• The atoms of a given element are identical; the atoms of different elements are different in some fundamental way or ways.
• Chemical compounds are formed when atoms of different elements combine with each other. A given compound always has the same relative numbers and types of atoms.
• Chemical reactions involve reorganization of the atoms - changes in the way they are bound together.
• The atoms themselves are not changed in a chemical reaction.

• Gay—Lussac
• Measured (under same conditions of T and P) the volumes of gases that reacted with each other.
• At the same T and P, equal volumes of different gases contain the same number of particles.
• Volume of a gas is determined by the number, not the size, of molecules.

## Early Experiments to Characterize the Atom

### J. J. Thomson (1898-1903)

• Postulated the existence of negatively charged particles, that we now call electrons, using cathode-ray tubes.
• Determined the charge-to-mass ratio of an electron.
• The atom must also contain positive particles that balance exactly the negative charge carried by electrons.

Carthode-Ray Tube

### Robert Millikan (1909)

• Performed experiments involving charged oil drops.
• Determined the magnitude of the charge on a single electron.
• Calculated the mass of the electron
• $$(9.11 \times 10^{-31}\,\chem{kg})$$.

Millikan Oil Drop Experiment

### Henri Becquerel (1896)

• Discovered radioactivity by observing the spontaneous emission of radiation by uranium.
• Three types of radioactive emission exist:
• Gamma rays ($$\gamma$$) – high energy light
• Beta particles ($$\beta$$) – a high speed electron
• Alpha particles ($$\alpha$$) – a particle with a 2+ charge

### Ernest Rutherford (1911)

• Explained the nuclear atom.
• The atom has a dense center of positive charge called the nucleus.
• Electrons travel around the nucleus at a large distance relative to the nucleus.

Rutherford's Gold Foil Experiment

## The Modern View of Atomic Structure: An Introduction

### The Modern Atom

• The atom contains:
• Electrons – found outside the nucleus; negatively charged.
• Protons – found in the nucleus; positive charge equal in magnitude to the electron’s negative charge.
• Neutrons – found in the nucleus; no charge; virtually same mass as a proton.
• The nucleus is:
• Small compared with the overall size of the atom.
• Extremely dense; accounts for almost all of the atom’s mass.

### Isotopes

• Atoms with the same number of protons but different numbers of neutrons.
• Show almost identical chemical properties; chemistry of atom is due to its electrons.
• In nature most elements contain mixtures of isotopes.

Two isotopes of sodium

### Isotope Symbols

• Isotopes are identified by:
• Atomic Number (Z) – number of protons
• Mass Number (A) – number of protons plus number of neutrons

## Molecules and Ions

### Chemical Bonds

• Covalent Bonds
• Bonds form between atoms by sharing electrons.
• Resulting collection of atoms is called a molecule.
• Ionic Bonds
• Bonds form due to force of attraction between oppositely charged ions.
• Ion – atom or group of atoms that has a net positive or negative charge.
• Cation – positive ion; lost electron(s).
• Anion – negative ion; gained electron(s).

## An Introduction to the Periodic Table

### The Periodic Table

• Metals vs. Nonmetals
• Groups or Families – elements in the same vertical columns; have similar chemical properties
• Periods – horizontal rows of elements

### Groups or Families

• Table of common charges formed when creating ionic compounds.
• Group or Family Charge
Alkali Metals (1) 1+
Alkaline Earth Metals (2) 2+
Halogens (17) 1-
Noble Gases (18) 0

## Naming Simple Compounds

### Naming Compounds

• Binary Compounds
• Composed of two elements
• Ionic and covalent compounds included
• Binary Ionic Compounds
• Metal—nonmetal
• Binary Covalent Compounds
• Nonmetal—nonmetal

### Binary Ionic Compounds (Type I)

1. The cation is always named first and the anion second.
2. A monatomic cation takes its name from the name of the parent element.
3. A monatomic anion is named by taking the root of the element name and adding –ide.
• Examples
• $$\chem{KCl}$$ - Potassium chloride
• $$\chem{MgBr_2}$$ - Magnesium bromide
• $$\chem{CaO}$$ - Calcium oxide

### Binary Ionic Compounds (Type II)

• Metals in these compounds form more than one type of positive ion.
• Charge on the metal ion must be specified.
• Roman numeral indicates the charge of the metal cation.
• Transition metal cations usually require a Roman numeral.
• Elements that form only one cation do not need to be identified by a roman numeral.
• Examples
• $$\chem{CuBr}$$ - Copper(I) bromide
• $$\chem{FeS}$$ - Iron(II) sulfide
• $$\chem{PbO_2}$$ - Lead(IV) oxide

### Polyatomic Ions

• Must be memorized (see Table 2.5 on pg. 65 in text).
• Examples of compounds containing polyatomic ions:
• $$\chem{NaOH}$$ - Sodium hydroxide
• $$\chem{Mg(NO_3)_2}$$ - Magnesium nitrate
• $$\chem{(NH_4)_2SO_4}$$ - Ammonium sulfate

### Binary Covalent Compounds (Type III)

• Formed between two nonmetals.
1. The first element in the formula is named first, using the full element name.
2. The second element is named as if it were an anion.
3. Prefixes are used to denote the numbers of atoms present.
4. The prefix mono- is never used for naming the first element.
Prefix Number Prefix Number
mono- 1 di- 2
tri- 3 tetra- 4
penta- 5 hexa- 6
hepta- 7 octa- 8
nona- 9 deca- 10

### Binary Covalent Compounds (Type III) - Examples

• $$\chem{CO_2}$$ - Carbon dioxide
• $$\chem{SF_6}$$ - Sulfur hexafluoride
• $$\chem{N_2O_4}$$ - Dinitrogen tetroxide

### Acids

• Acids can be recognized by the hydrogen that appears first in the formula -$$\chem{HCl}$$.
• Molecule with one or more $$\chem{H^+}$$ ions attached to an anion.
• If the anion does not contain oxygen, the acid is named with the prefix hydro– and the suffix –ic.
• Examples
• $$\chem{HCl}$$ - Hydrochloric acid
• $$\chem{HCN}$$ - Hydrocyanic acid
• $$\chem{H_2S}$$ - Hydrosulfuric acid

### Acids (cont.)

• If the anion does contain oxygen:
• The suffix –ic is added to the root name if the anion name ends in –ate.
• Examples:
• $$\chem{HNO_3}$$ - Nitric acid
• $$\chem{H_2SO_4}$$ - Sulfuric acid
• $$\chem{HC_2H_3O_2}$$ - Acetic acid
• If the anion does contain oxygen:
• The suffix –ous is added to the root name if the anion name ends in –ite.
• Examples:
• $$\chem{HNO_2}$$ - Nitrous acid
• $$\chem{H_2SO_3}$$ - Sulfurous acid
• $$\chem{HClO_2}$$ - Chlorous acid

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