Atoms, Molecules, and Ions

Chapter 2: Atoms, Molecules, and Ions

Authors: William L. Masterton, Cecile N. Hurley, Edward J. Neth
Source: Cengage.com/Chemistry/Masterton
Institution: University of Connecticut

Learning a Language in Chemistry

• Chemistry has a language akin to learning a new spoken language.

  • Steps to learn:

  • Start with the alphabet (elements).

  • Form words (molecules).

  • Form sentences (chemical reactions).

  • This chapter introduces the fundamentals of the language in chemistry.

Outline of Chapter 2

1. Atoms and Atomic Theory

2. Components of the Atom

3. Quantitative Properties of the Atom

4. Introduction to the Periodic Table

5. Molecules and Ions

6. Formulas of Ionic Compounds

7. Names of Compounds

The Language of Chemistry

• The fundamental language of chemistry includes:

  • Atoms

  • Molecules

  • Ions

  • Chemical formulas

  • Chemical names

The Structure of Matter

• Atoms: Basic units composed of electrons, protons, and neutrons.

• Molecules: Combinations of two or more atoms.

• Ions: Charged particles that can be formed from atoms or molecules.

Atoms and Atomic Theory

• An element is composed of tiny particles called atoms.

• All atoms of the same element share identical chemical properties.

• In a chemical reaction, atoms are neither created nor destroyed.

• Compounds form when two or more atoms of different elements combine.

Fundamental Laws of Matter

1. Law of Conservation of Mass: Matter is conserved in chemical reactions.

2. Law of Constant Composition: Any sample of a pure substance contains the same proportion of elements.

   • Example: Pure water has the same composition everywhere.

3. Law of Multiple Proportions: When two elements form different compounds, the ratios of their masses are small whole numbers.

   • Example: Cr₂O₃ vs. CrO₃.

Components of the Atom

• The atomic theory of matter raised additional questions about the structure of atoms.

  • Could atoms be broken down into smaller particles?

  • Over 100 years after the formulation of atomic theory, new experimental insights emerged to answer these questions.

Fundamental Experiments

• Key figures:

  • J.J. Thomson at Cavendish Laboratories, Cambridge.

  • Ernest Rutherford at McGill University, Canada, plus Manchester and Cambridge Universities, England.

Electrons

• The first evidence for subatomic particles was identified through electric conduction tests on gases at low pressures by J.J. Thomson in 1897.

  • Cathode Rays:

  • These are streams of negatively charged particles, identified as electrons.

  • Electrons carry a unit negative charge: -1

  • Mass: Approximately 1/2000 of the lightest atomic mass.

The Electron and the Atom

• Every atom has at least one electron.

• Atoms can possess up to 100 or more electrons, ensuring electrical neutrality through a number of electrons equal to the positive charges in an atom.

The Plum-Pudding Model

• Proposed by J.J. Thomson:

  • Atoms are modeled as positively charged spheres with electrons embedded within.

  • Often referred to as the plum-pudding model (or raisin-bread model).

Protons and Neutrons - The Nucleus

• In 1911, Ernest Rutherford conducted pioneering experiments on gold foil bombarded with alpha particles.

  • Expected outcomes were that particles would easily pass through; however, some were deflected, leading to the discovery of a dense, positively charged nucleus at the atom's center.

Rutherford’s Model

• The results indicated:

  • A small, dense core (nucleus) with a positive charge exists at the center of the atom.

  • Electrons orbit this nucleus, meaning that most of the atom is just empty space.

Sub-atomic Particles

• Protons (p⁺):

  • Charged positively, located in the nucleus.

  • Mass nearly equivalent to a hydrogen atom (1.00728 amu).

• Neutrons (n⁰):

  • No charge, also situated in the nucleus.

  • Mass slightly higher than that of protons (1.00867 amu).

• Electrons (e⁻):

  • Carry negative charge, orbiting the nucleus with negligible mass (0.00055 amu).

Mass and the Atom

• Over 99.9% of an atom's mass is found in the nucleus; the volume of the nucleus is significantly smaller compared to the atom's overall volume.

Terminology

• Atomic Number (Z): Indicates the number of protons in the atom.

• Mass Number (A): Sum of protons and neutrons in the nucleus (A = Z + n).

Isotopes

• Isotopes are atoms of the same element (same atomic number) but differing mass numbers due to variations in neutron quantity.

Nuclear Symbolism

• A describes the mass number, Z signifies the atomic number, and X is the chemical symbol: X_{A}^{Z}.

Isotopes of Hydrogen

• Hydrogen Isotopes:

  • ^{1}H - Regular hydrogen

  • ^{2}H - Deuterium

  • ^{3}H - Tritium

  • Different in mass and applications.

Example 2.1

• a) Isotope of cobalt (Co, Z = 27) with 33 neutrons:

  • Nuclear symbol: ^{60}Co (A = p + n = 27 + 33)

• b) Radioactive isotope of strontium (^{90}Sr):

  • Protons: 38

  • Neutrons: 52

  • Nuclear symbol of gallium with 31 protons and 38 neutrons: ^{69}Ga.

Atomic Masses: The Carbon-12 Scale

• Atomic mass measures how heavy, on average, an atom of an element is.

• Unit: atomic mass unit (amu).

• Mass of one ^{12}C atom: 12 amu (exactly).

Determining Atomic Masses

• Atomic masses can be measured with high precision using mass spectrometry.

• The process separates substances based on their mass and charge; the results show abundance against mass on a graph.

Mass Spectrometry

• Atoms are ionized at low pressure enabling them to accelerate towards a magnetic field.

• The extent of cation deflection in this field is inversely related to mass.

Fluorine

• Atomic fluorine has a single isotope with an exact mass of 19.00 amu.

Carbon

• Atomic mass of ^{12}C: exactly 12.000 amu

• Carbon on the periodic table shows an average mass of 12.011 amu.

• Most atomic masses are not whole numbers due to isotopic prevalence.

Isotopic Abundance

• To determine an element's mass, one must know both the mass of each isotope and their relative abundances (isotopic abundance).

• Mass spectrometry determines isotopic abundance.

Example 2.2

• For bromine with an average atomic mass of 79.90 amu and two isotopes:

  • ^{79}Br and ^{81}Br.

• The abundance of the heavier isotope can be calculated through equations correlating atomic masses and percentages.

Relative Atomic Masses

• Representation of the mass of an element shown below its symbol in the periodic table (e.g., He = 4.003 amu).

Masses of Individual Atoms

• Knowing the mass in grams for practical purposes can help determine quantities using Avogadro's number (6.022 x 10²³).

Example 2.3

• For arsenic, calculations based on Avogadro's number enabled you to determine:

  1. Mass of a single arsenic atom: 1.244 imes 10^{-22}g.

  2. The number of atoms in a 10-gram sample.

  3. Number of protons in a 0.1500 lb sample.

Introduction to the Periodic Table

• Periodic Table Structure:

  • Rows are periods; columns are groups.

  • IUPAC numbering from 1 to 18.

Blocks in the Periodic Table

• Main group elements: Group 1, 2, 13-18; Transition metals: Groups 3-12; Post-transition metals: Right of transition metals.

Families in the Periodic Table

• Alkali Metals (Group 1): Soft and reactive.

• Alkaline Earth Metals (Group 2): Reactive but less so than alkali metals.

• Halogens (Group 17): Very reactive nonmetals.

• Noble Gases (Group 18): Generally unreactive.

Importance of Families

• Elements within the same family exhibit similar chemical properties.

• Example: Alkali metals are highly reactive; noble gases are inert.

Mendeleev

• Dmitri Mendeleev: Arranged elements by their properties, anticipated unknown elements.

• By 1886, elements discovered matched his predicted properties (Sc, Ga, Ge).

Metals and Nonmetals

• A diagonal line starting with B separates metals from nonmetals.

• Metalloids exhibit properties of both categories.

Biological View of the Periodic Table

• Essential Elements: Essential to life (C, H, O, S)

• Toxic Elements: Can become lethal at higher concentrations (e.g., Selenium).

Molecules

• Formed through the combination of two or more atoms, typically nonmetals.

• Covalent Bonds: Hold molecules together.

• Molecular Formulas: Indicate the number of each atom, as in water (H_{2}O) and ammonia (NH_{3}).

Structural Formulas

• Illustrates bonding patterns in molecules (e.g., H-O-H for water).

Example 2.4

• Molecular Formulas:

  • Ethyl Alcohol: C_{2}H_{6}O

  • Ethylamine: C_{2}H_{7}N

  • Recognize that molecular formulas do not convey atomic arrangement.

Molecular Elements

• Some elements exist as molecules:

  • Oxygen (O_{2}), Hydrogen (H_{2}), and Nitrogen (N_{2}).

Ions

• Ions are formed when atoms or molecules gain or lose electrons.

  • Cations: Positively charged (e.g., Na^{+}).

  • Anions: Negatively charged (e.g., O^{2-}).

  • No change in proton count during ion formation.

Example 2.5

• Questions investigating protons, neutrons, and electrons in ions:

  • Aluminum Ion (Al^{3+}): 13 protons, 14 neutrons, 10 electrons.

  • Sulfur Ion: 16 protons (S), 16 neutrons, 18 electrons (S^{2-}): ^{32}S^{2-}.

Polyatomic Ions

• Groups of atoms can also carry a charge (e.g., Hydroxide OH^{-}, Ammonium NH_{4}^{+}).

Ionic Compounds

• Compounds form between cations and anions (e.g., Sodium Chloride NaCl).

• Ionic compounds are constituted by strong bonds, represented as a 3D arrangement, primarily in solid-state with high melting points.

Ionic Compounds in Solution

• Ionic compounds dissociate into ions when dissolved in water, which conduces electricity (strong electrolytes).

• Conversely, molecular compounds that do not form ions do not conduct electricity (nonelectrolytes).

Example 2.6

• Construct a structural representation of a solution with Potassium Sulfate (K_{2}SO_{4}).

Formulas of Ionic Compounds

• Charge balance is crucial for ionic formulas; equal positive and negative charges must be present.

• Example: Calcium Chloride (CaCl_{2}): Ca^{2+} requires two Cl^{-} for neutrality.

Noble Gas Connections

• Ions derived from elements in groups 1, 2, 16, and 17 tend to mimic the electron configuration of the nearest noble gas.

Cations of Transition and Post-Transition Metals

• Example: Iron commonly forms Fe^{2+} and Fe^{3+}.

Polyatomic Ions Summary

• Common cations include:

  • Ammonium NH_{4}^{+}

  • Mercury Hg_{2}^{2+}

• Most other polyatomic ions are negatively charged.

Example 2.7

• Predict the ionic compound formulas based on the charge of involved ions.

  • Barium with iodine: BaI_{2}

  • A transition metal compound with oxide: Cu_{2}O

  • Alkaline earth with nitrogen: Sr_{3}N_{2}

  • Ammonium and phosphate: (NH_{4}){3}PO{4}$.

Naming Compounds - Cations

• Monatomic Cations: Named from their respective metals:

  • Sodium: Na^{+} - Sodium ion.

  • Iron can vary in charge: Fe^{2+} (iron(II)), Fe^{3+} (iron(III)).

Naming Compounds - Anions

• Monatomic Anions: Named by adding -ide to the stem of the element name:

  • Oxygen becomes oxide: O^{2-} converts to oxide.

  • Polyatomic ions have unique names (e.g., NO_{3}^{-}: nitrate).

Oxoanions

• Nonmetals forming two types of oxoanions follow specific naming conventions:

  • -ate used for the one with more oxygen, -ite for less oxygen.

• Additional prefixes (per- for the largest, hypo- for the smallest) can denote different oxoanions comprising a nonmetal.

Names of Ionic Compounds

• Combine cation and anion names:

  • Cr(NO_{3})_{3} is chromium(III) nitrate.

  • SnCl_{2} is tin(II) chloride.

Example 2.8

• Names of specified ionic compounds:

  • CaS: Calcium sulfide.

  • Al(NO_{3})_{3}: Aluminum nitrate.

  • FeCl_{2}: Iron(II) chloride.

Binary Molecular Compounds

• Naming involves Greek prefixes to distinguish between the number of atoms present:

  • E.g., N_{2}O_{5} is dinitrogen pentoxide; PCl_{3} is phosphorus trichloride.

Example 2.9

• Provide names for specific molecular compounds:

  • SF_{4}: Sulfur tetrafluoride.

  • PCl_{3}: Phosphorus trichloride.

  • N_{2}O_{3}: Dinitrogen trioxide.

  • Cl_{2}O_{7}: Dichlorine heptoxide.

Acids

• Acids ionize in solution to release hydrogen ions, H^{+}.

  • HCl in gas form is hydrogen chloride; in water, it forms hydrochloric acid.

Common Acids

• Hydrochloric acid from HCl(g) is hydrogen chloride in solution.

• Hydrobromic acid originates from HBr(g), while hydriodic acid comes from HI(g).

Oxoacids

• Two prevalent oxoacids:

  • HNO_{3} (nitric acid)

  • H_{2}SO_{4} (sulfuric acid).

• Oxoacid names derive from corresponding oxoanions, modifying suffixes (-ate to -ic and -ite to -ous) while retaining prefixes.

Example 2.10

• Naming different acids based on their structures:

  • HCl(g): Hydrogen chloride

  • HNO_{2}(aq): Nitrous acid (via modification of nitrite).

• For H_{2}SO_{4}(aq) and HIO(aq)$$, respective names are sulfuric acid and hypoiodous acid.

Naming Flow Charts

• Flow charts outline how to name both molecular and ionic compounds based on characteristics such as gas state and presence of hydrogen or oxoanions.

Key Concepts

1. Relation of nuclear symbols to protons and neutrons in the nucleus.

2. Correlation of atomic mass, isotopic abundance, and average mass of an element.

3. Atomic mass relation to Avogadro’s Number.

4. Connection between elements and their arrangement on the periodic table.

5. Understanding structural, condensed, and molecular formulas.

6. The relationship between ionic charge and electron count.

7. Predicting ionic compound formulas based on ion charges.

8. Relating compound names to their formulas, including ionic, molecular, and oxoacids.