CHEM 1315: Unit 1 - Matter, Atoms, Elements and Particles

Proton, Neutron, Electron

  • Fundamental subatomic particles in atoms:

    • Proton

    • Neutron (n)

    • Electron

CHEM 1315

  • Course Title: Chemistry 1315

  • Unit 1: Matter: Atoms, Elements, and Particles

  • Instructor: Dr. Oikeh

  • Lecture notes available at canvas.ou.edu

Learning Objectives for Unit 1

  • 1.1 Explain the structure, properties, and interactions of matter and elements; represent matter from macro to submicroscopic views.

  • 1.2 Explain experiments leading to the Modern Atomic Theory.

  • 1.3 Characterize subatomic particles, isotopes, and ions.

  • 1.4 Determine atomic mass in the Periodic Table based on isotope abundances.

  • 1.5 Characterize the Periodic Table of elements.

  • 1.6 Name elements and monatomic ions, including symbols.

  • 1.7 Relate the mole, mass, and number of particles through Avogadro's number, including diatomic elements.

Chemistry as the Central Science

  • Chemistry integrates knowledge from various fields of science:

    • Biology

    • Molecular Biology

    • Biochemistry

    • Organic Chemistry

    • Medicine

    • Toxicology

    • Materials Science

    • Nanotechnology

    • Inorganic Chemistry

    • Analytical Chemistry

    • Environmental Science

    • Physical Chemistry

    • Chemical Engineering

    • Chemical Physics

    • Nuclear Chemistry

    • Geochemistry

    • Physics

    • Geology

    • Earth Sciences

Definition of Chemistry

  • Chemistry is the science that investigates materials of the universe and the changes these materials undergo.

  • Study of matter includes:

    • Composition: What matter is made of.

    • Structure: Arrangement of components.

    • Properties: Physical and chemical characteristics.

    • Transformation: Changes in states of matter and associated energy.

  • Matter: Defined as anything with mass and volume, constituting the 'stuff' of the universe.

Macro and Micro Views of Matter

  • Macro: Large-scale view of matter.

  • Micro: Small-scale or submicroscopic components.

Classification of Matter

  • Matter can be classified into two main categories:

    • Pure Substances

    • Have definite composition and distinct properties.

    • Further classified into:

      • Elements: Consist of only one type of atom (examples:

      • Carbon (C)

      • Iron (Fe)

      • Diatomic elements: Oxygen (O2), Hydrogen (H2)).

      • Compounds: Made of two or more different types of atoms in fixed ratios (examples:

      • Water (H2O)

      • Carbon dioxide (CO2)

      • Sodium chloride (NaCl)).

    • Mixtures

    • Physical combinations of two or more substances which can be separated without changing their identities.

    • Types:

      • Homogeneous Mixtures: Uniform composition (e.g., tea).

      • Heterogeneous Mixtures: Non-uniform composition (e.g., wet sand).

States of Matter

  • All substances can exist in three states:

    • Solid

    • Liquid

    • Gas

  • Transitions among states can occur through physical methods without altering chemical composition:

    • Melt: Solid to liquid.

    • Freeze: Liquid to solid.

Kinetic Molecular Theory

  • This theory explains physical states of matter in terms of kinetic and potential energy.

  • Two critical factors dictate matter's state:

    • Strength of attraction between particles (potential energy).

    • Amount of kinetic energy of particles.

Properties of Matter

  • Include physical properties (e.g., boiling point of water at 100°C) and chemical properties (e.g., nitroglycerine's explosiveness).

Changes in Matter

  • Physical Changes: Do not alter chemical composition; examples include boiling and melting.

  • Chemical Changes: Alter the chemical composition, forming new substances through reactions (e.g., Sodium + Chlorine → Sodium Chloride).

Modern Atomic Theory

  • Developed through key laws:

    • Law of Conservation of Mass: Matter is neither created nor destroyed in chemical reactions (Antoine Lavoisier).

    • Law of Definite Proportions: Compounds have specific ratios of constituent elements (Joseph Proust).

    • Law of Multiple Proportions: Ratios of masses of one element that combine with a fixed mass of another can be expressed in small whole number ratios (John Dalton).

John Dalton’s Atomic Theory

  • Fundamental tenets:

    • Matter is composed of indivisible atoms.

    • Atoms of an element are identical in mass and properties.

    • Compounds are combinations of different atoms in simple ratios.

    • Chemical reactions involve the rearrangement of atoms.

  • Modifications to Dalton’s theory include:

    • Atoms consist of subatomic particles (protons, neutrons, electrons).

    • Atomic mass may vary due to isotopes.

Discovery of Atomic Structure

  • Atoms consist of three primary subatomic particles:

    • Protons: Positively charged particles within the nucleus.

    • Neutrons: Neutral particles within the nucleus.

    • Electrons: Negatively charged particles surrounding the nucleus, forming a cloud of negative charge.

  • J.J. Thompson: Discovered the electron through cathode rays.

  • Robert Millikan: Measured the charge of the electron.

  • Ernest Rutherford: Proposed that most of the atom's mass is concentrated in the nucleus, with electrons surrounding it.

  • James Chadwick: Discovered neutrons in the nucleus.

Atomic Particles: Mass and Charge

  • Data Table:

    • Proton:

    • Mass: 1.67262×10−24 g1.67262 \times 10^{-24} \text{ g}; 1.00727 amu; Charge: +1.

    • Neutron:

    • Mass: 1.67493×10−24 g1.67493 \times 10^{-24} \text{ g}; 1.00866 amu; Charge: 0.

    • Electron:

    • Mass: 9.10938×10−28 g9.10938 \times 10^{-28} \text{ g}; 0.00055 amu; Charge: -1.

Atomic Number and Mass

  • Atomic Number (Z): Number of protons in the nucleus, characteristic of each element. Neutral atoms possess equal numbers of protons and electrons.

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

Isotopes

  • Atoms of the same element with the same number of protons but different numbers of neutrons.

  • Isotopes typically share similar chemical properties but differ in physical properties (e.g., Carbon-12, Carbon-13, Carbon-14).

  • Isotopes are represented symbolically using the elemental symbol alongside mass and atomic numbers (e.g., 614extC^{14}_{6} ext{C} for Carbon-14).

Ions

  • Charged species formed when atoms gain or lose electrons:

    • Cations: Positively charged ions (lose electrons).

    • Anions: Negatively charged ions (gain electrons).

  • Example Representation:

    • Sodium ion: Na+

    • Chloride ion: Cl-

The Mole

  • Definition: A mole is defined as the amount of substance containing 6.022imes10236.022 imes 10^{23} particles (Avogadro's number).

  • Example:

    • 1 mole of electrons = 6.022×10236.022 \times 10^{23} electrons.

Molar Mass

  • Molar mass refers to the mass of one mole of a substance expressed in g/mol. For example:

    • Carbon = 12.01 g/mol

    • Water (H2O) = (2 * 1.008 g/mol) + (15.999 g/mol) = 18.015 g/mol

  • Molar masses for compounds can be calculated based on their constituent elements.

Calculating Atomic Mass and Example

  • Example Calculation of Atomic Mass:

    • Consider Carbon Isotopes:

    • 12extC^{12} ext{C}: 12.000 amu (by definition)

    • 13extC^{13} ext{C}: 13.003355 amu

    • ext{Atomic weight} = ext{(% abundance of }^{12} ext{C)} \times \text{mass of }^{12} ext{C} + ext{(% abundance of }^{13} ext{C)} \times \text{mass of }^{13} ext{C}

    • Example: extAtomicweight=0.9889×12+0.0111×13.003355=12.0111extamuext{Atomic weight} = 0.9889 \times 12 + 0.0111 \times 13.003355 = 12.0111 ext{ amu}

Key Points on Calculation of Atomic Weight

  • Total mass represented in atomic mass is a significant means of determining elemental behavior in reactions.

  • Individual atomic masses of atoms often not whole numbers due to averaging based on isotope distribution.

The Periodic Table

  • Periodic Law: When arranged by increasing atomic number, properties recur periodically.

  • Groups and Periods:

    • Groups: Vertical columns representing elements with similar properties

    • Periods: Horizontal rows; elements behave similarly but not as consistently as those in the same group.

Categorizing Elements

  • Metals: Tend to lose electrons.

  • Nonmetals: Tend to gain electrons.

  • Metalloids: Have properties intermediate between metals and nonmetals.

Common Element Properties

  • Metals: High thermal and electrical conductivity, malleability, ductility, metallic luster.

  • Nonmetals: Brittle solids or gases, variable properties.

Naming Ions and Chemical Symbols

  • Monoatomic Ions: Named based on the element root and charge notation.

  • Examples:

    • Sodium (Na+): Sodium ion

    • Oxide (O2-): Oxide ion

  • Main Group Ions: Typically follow common charge patterns based on group placement.

Moles, Atoms, and Conversion

  • Conversion requires recognizing the relationship between mass, moles, and number of particles using molar mass and Avogadro’s number.

  • Example Problems:

    • Calculate moles of a compound from its empirical formulas.

    • Derive the number of atoms present in a specific sample mass.

Final Thoughts

  • Mastery of topics in this unit builds a solid foundation for understanding chemistry concepts, properties of elements, and reactions as we advance in the coursework.