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5.4: A Molecular View of Elements and Compounds

CHEM 114: Introductory Chemistry · LibreTexts

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    From LibreTexts, under CC BY-NC-SA 4.0, as the source states it; its equations are written in TeX, which the PDF shows typeset. This page at LibreTexts.

    Chapter 1
    1. Chapter 1: The Chemical World
    2. 1.1: The Scope of Chemistry
    3. 1.2: Chemicals Compose Ordinary Things
    4. 1.3: Hypothesis, Theories, and Laws
    5. 1.4: The Scientific Method: How Chemists Think
    6. 1.5: A Beginning Chemist: How to Succeed
     •  Chapter 2
    1. Chapter 2: Measurement and Problem Solving
    2. 2.1: Taking Measurements
    3. 2.2: Scientific Notation: Writing Large and Small Numbers
    4. 2.3: Significant Figures: Writing Numbers to Reflect Precision
    5. 2.4: Significant Figures in Calculations
    6. 2.5: The Basic Units of Measurement
    7. 2.6: Problem Solving and Unit Conversions
    8. 2.7: Solving Multistep Conversion Problems
    9. 2.8: Units Raised to a Power
    10. 2.9: Density
    11. 2.10: Numerical Problem-Solving Strategies and the Solution Map
    12. 2.E: Measurement and Problem Solving (Exercises)
     •  Chapter 3
    1. Chapter 3: Matter and Energy
    2. 3.1: In Your Room
    3. 3.2: What is Matter?
    4. 3.3: Classifying Matter According to Its State: Solid, Liquid, and Gas
    5. 3.4: Classifying Matter According to Its Composition
    6. 3.5: Differences in Matter: Physical and Chemical Properties
    7. 3.6: Changes in Matter: Physical and Chemical Changes
    8. 3.7: Conservation of Mass: There is No New Matter
    9. 3.8: Energy
    10. 3.9: Energy and Chemical and Physical Change
    11. 3.10: Temperature: Random Motion of Molecules and Atoms
    12. 3.11: Temperature Changes: Heat Capacity
    13. 3.12: Energy and Heat Capacity Calculations
    14. 3.E: Exercises
     •  Chapter 4
    1. Chapter 4: Atoms and Elements
    2. 4.1: Experiencing Atoms at Tiburon
    3. 4.2: Indivisible: The Atomic Theory
    4. 4.3: The Nuclear Atom
    5. 4.4: The Properties of Protons, Neutrons, and Electrons
    6. 4.5: Elements: Defined by Their Numbers of Protons
    7. 4.6: Looking for Patterns: The Periodic Law and the Periodic Table
    8. 4.7: Ions: Losing and Gaining Electrons
    9. 4.8: Isotopes: When the Number of Neutrons Varies
    10. 4.9: Atomic Mass: The Average Mass of an Element’s Atoms
      •  Chapter 5
    1. Chapter 5: Molecules and Compounds
    2. 5.1: Sugar and Salt
    3. 5.2: Compounds Display Constant Composition
    4. 5.3: Chemical Formulas: How to Represent Compounds
    5. 5.4: A Molecular View of Elements and Compounds
    6. 5.5: Writing Formulas for Ionic Compounds
    7. 5.6: Nomenclature: Naming Compounds
    8. 5.7: Naming Ionic Compounds
    9. 5.8: Naming Molecular Compounds
    10. 5.9: Naming Acids
    11. 5.10: Nomenclature Summary
    12. 5.11: Formula Mass: The Mass of a Molecule or Formula Unit
      •  Chapter 6
    1. Chapter 6: Chemical Composition
    2. 6.1: How Much Sodium?
    3. 6.2: Counting Nails by the Pound
    4. 6.3: Counting Atoms by the Gram
    5. 6.4: Counting Molecules by the Gram
    6. 6.5: Chemical Formulas as Conversion Factors
    7. 6.6: Mass Percent Composition of Compounds
    8. 6.7: Mass Percent Composition from a Chemical Formula
    9. 6.8: Calculating Empirical Formulas for Compounds
    10. 6.9: Calculating Molecular Formulas for Compounds
      •  Chapter 7
    1. Chapter 7: Chemical Reactions
    2. 7.1: Grade School Volcanoes, Automobiles, and Laundry Detergents
    3. 7.2: Evidence of a Chemical Reaction
    4. 7.3: The Chemical Equation
    5. 7.4: How to Write Balanced Chemical Equations
    6. 7.5: Aqueous Solutions and Solubility: Compounds Dissolved in Water
    7. 7.6: Precipitation Reactions: Reactions in Aqueous Solution That Form a Solid
    8. 7.7: Writing Chemical Equations for Reactions in Solution: Molecular, Complete Ionic, and Net Ionic Equations
    9. 7.8: Acid–Base and Gas Evolution Reactions
    10. 7.9: Oxidation–Reduction Reactions
    11. 7.10: Classifying Chemical Reactions
    12. 7.11: The Activity Series
     •  Chapter 8
    1. Chapter 8: Quantities in Chemical Reactions
    2. 8.1: Climate Change: Too Much Carbon Dioxide
    3. 8.2: Stoichiometry
    4. 8.3: Making Molecules: Mole-to-Mole Conversions
    5. 8.4: Making Molecules: Mass-to-Mass Conversions
    6. 8.5: Limiting Reactant, Theoretical Yield, and Percent Yield
    7. 8.6: Limiting Reactant, Theoretical Yield, and Percent Yield from Initial Masses of Reactants
    8. 8.7: Enthalpy: A Measure of the Heat Evolved or Absorbed in a Reaction
    Chapter 9
    1. Chapter 9: Electrons in Atoms and the Periodic Table
    2. 9.1: Blimps, Balloons, and Models of the Atom
    3. 9.2: Light: Electromagnetic Radiation
    4. 9.3: The Electromagnetic Spectrum
    5. 9.4: The Bohr Model: Atoms with Orbits
    6. 9.5: The Quantum-Mechanical Model: Atoms with Orbitals
    7. 9.6: Quantum-Mechanical Orbitals and Electron Configurations
    8. 9.7: Electron Configurations and the Periodic Table
    9. 9.8: The Explanatory Power of the Quantum-Mechanical Model
    10. 9.9: Periodic Trends: Atomic Size, Ionization Energy, and Metallic Character
     •  Chapter 10
    1. Chapter 10: Chemical Bonding
    2. 10.1: Bonding Models and AIDS Drugs
    3. 10.2: Representing Valence Electrons with Dots
    4. 10.3: Lewis Structures of Ionic Compounds: Electrons Transferred
    5. 10.4: Covalent Lewis Structures: Electrons Shared
    6. 10.5: Writing Lewis Structures for Covalent Compounds
    7. 10.6: Resonance: Equivalent Lewis Structures for the Same Molecule
    8. 10.7: Predicting the Shapes of Molecules
    9. 10.8: Electronegativity and Polarity: Why Oil and Water Don’t Mix
     •  Chapter 11
    1. Chapter 11: Gases
    2. 11.1: Extra-Long Straws
    3. 11.2: Kinetic Molecular Theory: A Model for Gases
    4. 11.3: Pressure: The Result of Constant Molecular Collisions
    5. 11.4: Boyle’s Law: Pressure and Volume
    6. 11.5: Charles’s Law: Volume and Temperature
    7. 11.6: Gay-Lussac's Law: Temperature and Pressure
    8. 11.7: The Combined Gas Law: Pressure, Volume, and Temperature
    9. 11.8: Avogadro’s Law: Volume and Moles
    10. 11.9: The Ideal Gas Law: Pressure, Volume, Temperature, and Moles
    11. 11.10: Mixtures of Gases: Why Deep-Sea Divers Breathe a Mixture of Helium and Oxygen
    12. 11.11: Gases in Chemical Reactions
     •  Chapter 12
    1. Chapter 12: Liquids, Solids, and Intermolecular Forces
    2. 12.1: Interactions between Molecules
    3. 12.2: Properties of Liquids and Solids
    4. 12.3: Intermolecular Forces in Action: Surface Tension and Viscosity
    5. 12.4: Evaporation and Condensation
    6. 12.5: Melting, Freezing, and Sublimation
    7. 12.6: Types of Intermolecular Forces: Dispersion, Dipole–Dipole, Hydrogen Bonding, and Ion-Dipole
    8. 12.7: Types of Crystalline Solids: Molecular, Ionic, and Atomic
    9. 12.8: Water: A Remarkable Molecule
     •  Chapter 13
    1. Chapter 13: Solutions
    2. 13.1: Prelude - Tragedy in Cameroon
    3. 13.2: Solutions: Homogeneous Mixtures
    4. 13.3: Solutions of Solids Dissolved in Water: How to Make Rock Candy
    5. 13.4: Solutions of Gases in Water: How Soda Pop Gets Its Fizz
    6. 13.5: Solution Concentration: Mass Percent
    7. 13.6: Solution Concentration: Molarity
    8. 13.7: Solution Dilution
    9. 13.8: Solution Stoichiometry
    10. 13.9: Freezing Point Depression and Boiling Point Elevation: Making Water Freeze Colder and Boil Hotter
    11. 13.10: Osmosis: Why Drinking Salt Water Causes Dehydration
      •  Chapter 14
    1. Chapter 14: Acids and Bases
    2. 14.1: Sour Patch Kids and International Spy Movies
    3. 14.2: Acids: Properties and Examples
    4. 14.3: Bases: Properties and Examples
    5. 14.4: Molecular Definitions of Acids and Bases
    6. 14.5: Reactions of Acids and Bases
    7. 14.6: Acid–Base Titration: A Way to Quantify the Amount of Acid or Base in a Solution
    8. 14.7: Strong and Weak Acids and Bases
    9. 14.8: Water: Acid and Base in One
    10. 14.9: The pH and pOH Scales: Ways to Express Acidity and Basicity
    11. 14.10: Buffers: Solutions That Resist pH Change
     

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    Learning Objectives
    • Classify substances as atomic elements, molecular elements, molecular compounds, or ionic compounds.

    Atomic Elements

    Most elements exist with individual atoms as their basic unit. It is assumed that there is only one atom in a formula if there is no numerical subscript on the right side of an element’s symbol.

    Molecular Elements

    There are many substances that exist as two or more atoms connected together so strongly that they behave as a single particle. These multi-atom combinations are called molecules. A molecule is the smallest part of a substance that has the physical and chemical properties of that substance. In some respects, a molecule is similar to an atom. A molecule, however, is composed of more than one atom.

    Table \(\PageIndex{1}\): Elements That Exist as Diatomic Molecules
    Hydrogen, H Oxygen Nitrogen Fluorine Chlorine Bromine Iodine

    Some elements exist naturally as molecules. For example, hydrogen and oxygen exist as two-atom molecules. Other elements also exist naturally as diatomic molecules—a molecule with only two atoms (Table \(\PageIndex{1}\)). As with any molecule, these elements are labeled with a molecular formula, a formal listing of what and how many atoms are in a molecule. (Sometimes only the word formula is used, and its meaning is inferred from the context.) For example, the molecular formula for elemental hydrogen is H2, with H being the symbol for hydrogen and the subscript 2 implying that there are two atoms of this element in the molecule. Other diatomic elements have similar formulas: O2, N2, and so forth. Other elements exist as molecules—for example, sulfur normally exists as an eight-atom molecule, S8, while phosphorus exists as a four-atom molecule, P4 (Figure \(\PageIndex{1}\)).

    fac00a98cce192f9de352eba3772bd51.jpg
    Figure \(\PageIndex{1}\): Molecular Art of S8 and P4 Molecules. If each green ball represents a sulfur atom, then the diagram on the left represents an S8 molecule. The molecule on the right shows that one form of elemental phosphorus exists, as a four-atom molecule.

    Figure \(\PageIndex{1}\) shows two examples of how molecules will be represented in this text. An atom is represented by a small ball or sphere, which generally indicates where the nucleus is in the molecule. A cylindrical line connecting the balls represents the connection between the atoms that make this collection of atoms a molecule. This connection is called a chemical bond.

    Ionic Compounds

    The elements in the periodic table are divided into specific groupings; the metals, the non-metals, the semi-metals, and so on. These groupings are largely based on physical properties and on the tendency of the various elements to bond with other elements by forming either an ionic or a covalent bond. As a general rule of thumb, compounds that involve a metal binding with either a non-metal or a semi-metal will display ionic bonding. Thus, the compound formed from sodium and chlorine will be ionic (a metal and a non-metal). The basic unit of ionic compounds is the formula unit.

    Molecular Compounds

    Compounds that are composed of only non-metals or semi-metals with non-metals will display covalent bonding and will be classified as molecular compounds. Nitrogen monoxide (NO) will be a covalently bound molecule (two non-metals) and silicon dioxide (SiO2) will also be a covalently bound molecule (a semi-metal and a non-metal). The basic unit of molecular compounds is the molecule.

    Example \(\PageIndex{1}\)

    Provide the classification (i.e. atomic element, molecular element, molecular compound, or ionic compound) of each substance.

    1. Fe
    2. PCl3
    3. LiBr
    4. P4
    5. oxygen gas
    Solution
    1. Fe (iron) is an element that is represented with no subscript, so it is an atomic element.
    2. PCl3 is made up of two nonmetals, so it is a molecular compound.
    3. LiBr is made up of lithium, a metal, and bromine, a nonmetal, so it is an ionic compound.
    4. P4 is a substance that is made up of four atoms of the same element, so it is a molecular element.
    5. The formula for oxygen gas is O2 so it is a molecular element.
    Exercise \(\PageIndex{1}\)

    Provide the classification (i.e. atomic element, molecular element, molecular compound, or ionic compound) of each substance.

    1. I2
    2. He
    3. H2O
    4. Al
    5. CuCl
    Answer a:
    molecular element
    Answer b:
    atomic element
    Answer c:
    molecular compound
    Answer d:
    atomic element
    Answer e:
    ionic compound

    5.4: A Molecular View of Elements and Compounds is shared under a CC BY-NC-SA 4.0 license and was authored, remixed, and/or curated by LibreTexts.