Atomic Structure

Atoms, Molecules and Ions

Outline

• Atomic Theory

• Atomic Structure

• Periodic Table

• Molecules and Molecular Compounds

• Ions and Ionic Compounds

• Naming and Formulas for Ionic Compounds

• Naming and Formulas for Acids

• Naming and Formulas for Molecular Compounds

Atomic Theory

• Aristotle (384-322 BC)

  • Proposed that matter was continuous.

• Democritus (460-370 BC)

  • Suggested that the world was made up of tiny indivisible particles called atomos.

Development of Atomic Theory

• Antoine Lavoisier (1743-1794)

  • Defined elements as materials made of a fundamental substance that cannot be broken down into anything else.

  • Transformed water into hydrogen and oxygen, proving that it was not an element.

  • Introduced the Law of Conservation of Mass: Mass before chemical reaction equals mass after.

• John Dalton (~1800) - Formulated the Atomic Theory of Matter:

  1. All elements are composed of extremely small particles called atoms, which are the smallest fundamental units of matter.

  2. All atoms of a given element are identical.

  3. In chemical reactions, atoms are not created, destroyed, or changed.

  4. Compounds are formed when atoms combine.

Dalton’s Theory Explained

• Law of Constant Composition (Definite Proportions) (Proust)

  • In a given compound, the relative numbers and kinds of atoms are constant.

• Law of Conservation of Mass (Lavoisier)

  • Mass before = Mass after.

• Law of Multiple Proportions

  • Atoms join in different whole number ratios to form different compounds.

Atomic Structure - Discovery of Subatomic Particles

• Discovery of the Electron

  • JJ Thomson conducted experiments with cathode rays.

  • Cathode rays were found to be composed of negatively charged particles known as electrons.

  • The mass to charge ratio of electrons was calculated as $1.76 \times 10^8 \text{ coulombs/g}$.

• Oil Drop Experiment by Robert Millikan

  • Measured the charge of the electron as $1.60 \times 10^{-19} \text{ C}$.

  • Calculated the mass of the electron as $9.11 \times 10^{-28} \text{ g}$.

Additional Subatomic Particles

• Atoms have additional subatomic particles discovered through radioactivity.

  • Observations by Henri Becquerel, Marie and Francis Curie, Ernest Rutherford indicated that uranium spontaneously emits high-energy radiation of three different types, which respond differently to an electric field.

JJ Thomson's Plum Pudding Model

• Proposed that the atom consisted of a uniform positive sphere with electrons embedded throughout.

• This model was later proven incorrect.

Ernest Rutherford - Discovery of the Atomic Nucleus

• Conducted experiments leading to the discovery of the atomic nucleus.

  • Identified that positively charged particles were densely packed in the center of the atom, thus not scattered throughout.

  • Most of the atom is considered empty space.

Radioactivity and Subatomic Particles

• The observation of radioactivity contributed significantly to the understanding of atomic structure.

• Types emitted from radioactive sources like uranium include:
  a. Helium nuclei (alpha particles)
  b. High energy electrons (beta particles)
  c. Gamma rays

• Silverman's gold foil experiment showed significant insights about atomic structure:

  1. The atom consists of protons, electrons, and neutrons.

  2. Electrons have a charge equal and opposite to that of the proton.

  3. Most of the atom is predominantly empty space.

  4. Electrons orbit the nucleus, though not in defined motions.

Modern Atomic Structure - Subatomic Particles

• Electrons:

  • Found outside the nucleus, negatively charged (-1), with a mass of $9.109 \times 10^{-28} \text{ g}$.

• Protons:

  • Located in the nucleus, positively charged (+1), with a mass of $1.673 \times 10^{-24} \text{ g}$.

• Neutrons:

  • Also in the nucleus, neutral with a mass of approximately same as protons, $1.675 \times 10^{-24} \text{ g}$.

• The forces that hold the nucleus together include:

  • Gravitational force

  • Weak nuclear force

  • Electrostatic force

  • Strong nuclear force (most significant).

Atomic Identity

• Atomic Number (Z):

  • Defines the identity of an atom. It is equal to the number of protons in the nucleus.

  • As per periodic table, elements are arranged in order of increasing atomic number.

  • For neutral atoms, number of protons equals the number of electrons.

• Isotopes:

  • Atoms of the same element can have differing numbers of neutrons, leading to various isotopes.

  • Mass Number: The total number of protons and neutrons.

• Element Symbols:

  • Conventionally represented as:

  • X = symbol of element

  • Z = atomic number (number of protons)

  • A = mass number (number of protons + neutrons)

  • Example of Isotope Notation: Hydrogen represented as ${}^{1}H, {}^{2}H, {}^{3}H$ denotes hydrogen-1, hydrogen-2 (deuterium), and hydrogen-3 (tritium).

Periodic Table Overview

• Arranged in order of increasing atomic number with repeating properties defined as periodic.

• Horizontal rows = periods;

• Vertical columns = groups, indicating similar properties.

• Metals, nonmetals, and metalloids are organized in distinct regions, with use cases or significance defined for each group.

Molecules and Molecular Compounds (Covalent Compounds)

• Molecule: Defined as two or more nonmetal atoms bound together.

  • Can consist of:

  • Atoms of a single element - Example: O₂

  • Atoms of more than one type of element - Example: H₂O.

• Molecular compounds have distinct properties dissimilar from the properties of the constituent elements.

• Common diatomic molecules include: H₂, N₂, O₂, F₂, I₂, Cl₂, Br₂.

Chemical Formulas

• Chemical formulas indicate the type and number of atoms in a compound:

  • Two types of chemical formulas exist:

  • Molecular formulas: Express actual number of atoms present.
    Example: glucose C₆H₁₂O₆ (molecular) vs CH₂O (empirical).

  • Empirical formulas: Express the simplest ratio of the constituent atoms.
    Example: Water (H₂O) maintains same empirical and molecular formula.

Ions and Ionic Compounds

• Ions: Atoms that have gained or lost electrons:

  • Anions: Atoms that have gained electrons (negatively charged).

  • Cations: Atoms that have lost electrons (positively charged).

• The number of protons in an ion remains unchanged.

• Trends in ion formation:

  • Metal atoms typically lose electrons to form cations.

  • Nonmetals usually gain electrons to form anions.

  • Aim for stability by attaining electron configuration similar to that of the nearest noble gas.

• The properties of ions diverge from the neutral atoms from which they derive.

Composition of Ionic Compounds

• Composed typically of a metal and a nonmetal, ensuring the presence of a cation and an anion.

• Formed through electrostatic attractive forces known as ionic bonds.

• Always represented by empirical formulas that represent the simplest integer ratio of ions in the compound.

Formulas for Ionic Compounds

• To generate the empirical formula:

  • Write the cation symbol first followed by the anion.

  • Charges of cations and anions must balance to achieve a neutral compound.

  • Common strategy: “Swap and Drop” charges to determine subscripts.

Polyatomic Ions

• Contain multiple atoms bonded covalently but carry a net charge (positive or negative). Examples:

  • Nitrate (NO₃)

  • Hydroxide (OH^-)

  • Sulfate (SO₄^2-)

  • Phosphate (PO₄^3-)

Naming Ionic Compounds

• Monatomic Cations: Name of the element + “ion” at the end. Example: Ca²⁺ is "calcium ion".

• Transition metals forming multiple charges require Roman numerals in parentheses to indicate their charge, e.g., Fe²⁺ is "iron (II) ion".

• Monatomic Anions: Element name with “-ide” suffix (S^2- -> sulfide).

• Naming Polyatomic Ions: Retain the ion’s name as is.

  • E.g. NH₄⁺ as ammonium.

Naming Acids

• Definition: A substance yielding hydrogen ions (H⁺) when dissolved in water.

• Naming Binary Acids: For acids containing hydrogen and a monatomic ion:

  1. Add prefix “hydro-”

  2. Change “-ide” ending to “-ic”

  3. Add “acid” at the end of the name.

  • Example: HCl -> hydrochloric acid.

• Oxyacids Formation: 1. Change “-ite” to “-ous” and “-ate” to “-ic.” Retain “hypo-” or “per-” prefixes.

Molecular Compounds and Naming Rules

• Molecular compounds are primarily made of nonmetals.

1. Name the leftmost element on the periodic table first (Oxygen, unless it’s with Fluorine).

2. Higher atomic number element to be prioritized when in the same column.

3. Change second element's name ending to “-ide.”

4. Use prefixes to denote counts, except for the first element if only one.

• Common prefixes:

  • mono- = 1

  • di- = 2

  • tri- = 3

  • tetra- = 4

  • penta- = 5

  • …

  • deca- = 10.

Example Molecular Naming

• CO: carbon monoxide

• N2O3: dinitrogen trioxide

Note: This extensive format should allow students to utilize these study notes as a complete reference for the concepts presented in the transcript.

Atoms, Molecules and Ions

Outline

• Atomic Theory

• Atomic Structure

• Periodic Table

• Molecules and Molecular Compounds

• Ions and Ionic Compounds

• Naming and Formulas for Ionic Compounds

• Naming and Formulas for Acids

• Naming and Formulas for Molecular Compounds

Atomic Theory

• Aristotle (384-322 BC): Proposed that matter was continuous.

• Democritus (460-370 BC): Suggested that the world was made up of tiny indivisible particles called atomos.

Development of Atomic Theory

• Antoine Lavoisier (1743-1794)

  • Defined elements as materials made of a fundamental substance that cannot be broken down into anything else.

  • Transformed water into hydrogen and oxygen, proving that it was not an element.

  • Introduced the Law of Conservation of Mass: Mass before chemical reaction equals mass after.

• John Dalton (~1800) - Formulated the Atomic Theory of Matter:

  1. All elements are composed of extremely small particles called atoms, which are the smallest fundamental units of matter.

  2. All atoms of a given element are identical.

  3. In chemical reactions, atoms are not created, destroyed, or changed.

  4. Compounds are formed when atoms combine.

Dalton’s Theory Explained

• Law of Constant Composition (Definite Proportions): In a given compound, the relative numbers and kinds of atoms are constant.

• Law of Conservation of Mass: Mass before = Mass after.

• Law of Multiple Proportions: Atoms join in different whole number ratios to form different compounds.

Atomic Structure - Discovery of Subatomic Particles

• Discovery of the Electron

  • JJ Thomson conducted experiments with cathode rays, finding they were negatively charged particles known as electrons.

  • The mass to charge ratio was calculated as $1.76 \times 10^8 \text{ coulombs/g}$.

• Oil Drop Experiment by Robert Millikan

  • Measured the charge of the electron as $1.60 \times 10^{-19} \text{ C}$.

  • Calculated the mass of the electron as $9.11 \times 10^{-28} \text{ g}$.

JJ Thomson's Plum Pudding Model

• Proposed that the atom consisted of a uniform positive sphere with electrons embedded throughout. This was later proven incorrect.

Ernest Rutherford - Discovery of the Atomic Nucleus

• Identified that positively charged particles were densely packed in the center of the atom (the nucleus).

• Most of the atom is considered empty space.

• The atom consists of protons, electrons, and neutrons.

Modern Atomic Structure - Subatomic Particles

• Electrons: Found outside the nucleus, negatively charged (-1), with a mass of $9.109 \times 10^{-28} \text{ g}$.

• Protons: Located in the nucleus, positively charged (+1), with a mass of $1.673 \times 10^{-24} \text{ g}$.

• Neutrons: Also in the nucleus, neutral with a mass of approximately $1.675 \times 10^{-24} \text{ g}$.

• Forces in the Nucleus: Gravitational, weak nuclear, electrostatic, and the strong nuclear force (most significant).

Atomic Identity

• Atomic Number (Z): Defines the identity of an atom; equal to the number of protons.

• Isotopes: Atoms of the same element with differing numbers of neutrons.

• Mass Number (A): The total number of protons and neutrons.

• Element Symbols: Represented as ${_{Z}^{A}X}$. Example: Hydrogen isotopes ${}^{1}H, {}^{2}H, {}^{3}H$.

Periodic Table Overview

• Arranged by increasing atomic number.

• Periods: Horizontal rows.

• Groups: Vertical columns indicating similar properties.

Molecules and Molecular Compounds

• Molecule: Two or more nonmetal atoms bound together (e.g., O2$,$H2O).

• Molecular formulas: Express the actual number of atoms present (e.g., C6H12O_6).

• Empirical formulas: Express the simplest ratio of atoms (e.g., $CH_2O$).

Ions and Ionic Compounds

• Anions: Atoms that have gained electrons (negative).

• Cations: Atoms that have lost electrons (positive).

• Ionic Bonds: Formed through electrostatic attractive forces between a metal and a nonmetal.

• Polyatomic Ions: Multiple atoms bonded covalently carrying a net charge (e.g., $SO<em>4^{2-}$, $NH</em>4^+$).

Naming Rules

• Ionic: Cation name + Anion name (ending in -ide). Transition metals use Roman numerals.

• Acids:

  • Binary: "hydro-" + element name + "-ic acid".

  • Oxyacids: "-ite" becomes "-ous", "-ate" becomes "-ic".

• Molecular: Use prefixes (mono-, di-, tri-, etc.) to denote counts of nonmetals.

Key Terms and Definitions

• Atom: The smallest fundamental unit of matter that retains the properties of an element.

• Subatomic Particles: Particles smaller than an atom, specifically protons, neutrons, and electrons.

• Nucleus: The small, dense, positively charged center of an atom containing protons and neutrons.

• Radioactivity: The spontaneous emission of high-energy radiation by an atom.

• Diatomic Molecule: A molecule composed of only two atoms (e.g., $N<em>2$, $Cl</em>2$).

• Covalent Bond: A chemical bond involving the sharing of electron pairs between nonmetal atoms.

• Electrostatic Force: The force of attraction or repulsion between particles due to their electric charge.

• Nobel Gas Configuration: A stable electron arrangement that atoms seek to achieve through ion formation.