Early Theories of Matter

Early Theories of Matter

• Democritus: Believed matter was composed of indivisible particles called atoms. Proposed that atoms were uniform, solid, hard, incompressible, and indestructible, and that they moved through empty space.
• Aristotle: Rejected Democritus's theory, believing instead that all materials on Earth were made of four elements: earth, fire, water, and air.
• Dalton's Atomic Theory: Proposed that all matter is made up of atoms, which are the smallest unit of an element. Key points:
  • All matter is composed of extremely small particles called atoms.
  • Atoms of a given element are identical in size, mass, and other properties.
  • Atoms of different elements differ in size, mass, and other properties.
  • Atoms cannot be subdivided, created, or destroyed.
• Thomson: Discovered the electron, a negatively charged subatomic particle found within atoms.
• Nagaoka: Proposed a model of the atom with a small, positively charged nucleus surrounded by a ring of electrons.
• Rutherford: Through his gold foil experiment, determined that the positive charge of an atom is concentrated in a tiny, dense nucleus, while electrons occupy the space surrounding it.

Subatomic Particles

• Three main subatomic particles:
  • Protons: Positively charged, reside in the nucleus.
  • Neutrons: Neutral charge, reside in the nucleus.
  • Electrons: Negatively charged, orbit the nucleus.
• Atomic number (Z): The number of protons in an atom's nucleus. It defines the element.
• Mass number (A): The sum of protons and neutrons in an atom's nucleus.
• Isotopes: Atoms of the same element that have the same number of protons but different numbers of neutrons. This results in different mass numbers.
• Relative Atomic Mass (Aᵣ): The weighted average of the masses of all isotopes of an element.
• Unified Atomic Mass Unit (u): A unit of mass used to express atomic masses. It is approximately equal to the mass of a proton or neutron.
• Atomic Number - Atomic Mass = Number of Neutrons*

Atomic Radius

• Atomic radius: The distance from an atom's nucleus to the furthest electron orbital.
• Atomic radius decreases across a period: As you move across a period from left to right, the number of protons increases in the nucleus, increasing the positive charge and attracting the electrons more strongly. This pulls the electron cloud closer to the nucleus, decreasing the radius.
• Atomic radius increases down a group: As you move down a group, the number of electron shells increases, resulting in electrons further from the nucleus. This weakens the attraction between the nucleus and electrons.

Nuclear Force, Shielding Effect, and Electron Configuration

• Nuclear Force: The strong force holding protons and neutrons together in the nucleus.
• Shielding Effect: The decrease in the attraction between the nucleus and valence electrons due to the presence of inner electrons, which partially shield the valence electrons from the positive nuclear charge.
• Shielding effect is more prominent in elements at the bottom of a group: This is because the nucleus is further from the valence electrons, and the inner electrons are more effective at shielding the valence electrons from the nucleus's attraction.

Ionic Radius

• Ionic radius: The distance from the nucleus of an ion to the outermost edge of its electron cloud.
• Cations (positive ions) are smaller than their parent atoms: Because they lose electrons, thus reducing the size of the electron cloud.
• Anions (negative ions) are larger than their parent atoms: Because they gain electrons, increasing the size of the electron cloud.

Ionization Energy

• Ionization energy: The amount of energy required to remove an electron from a gaseous atom or ion.
• First ionization energy: The energy needed to remove the least tightly bound electron from a neutral atom in the gaseous phase.
• Second ionization energy: The energy required to remove the next least tightly bound electron from a 1+ ion of the element.
• Ionization energy increases across a period: Because the atomic radius decreases, the electrons are held more tightly by the increasing positive nuclear charge, making it harder to remove them.
• Ionization energy decreases down a group: Because the atomic radius increases, the electrons are further from the nucleus and less tightly held, making it easier to remove them.

Electron Affinity

• Electron affinity: The change in energy (in kJ/mol) when an electron is added to a neutral atom in the gaseous phase to form a negative ion.
• Electron affinity can be positive or negative. Negative values mean the addition of an electron releases energy, making the process energetically favorable. Positive values mean the addition of an electron requires energy, making the process unfavorable.
• Electron affinity generally increases across a period: Because the smaller atomic radius increases the attraction between the nucleus and the incoming electron, making it more favorable to add an electron.
• Electron affinity generally decreases down a group: Because the larger atomic radius decreases the attraction between the nucleus and the incoming electron, making it less favorable to add an electron.

Electronegativity

• Electronegativity: A measure of an atom's tendency to attract electrons in a chemical bond.
• Electronegativity increases across a period: Due to the increasing nuclear charge, the atoms hold onto their electrons more tightly.
• Electronegativity decreases down a group: As the atom gets larger, the attraction between the nucleus and the outermost electrons weakens, making it less likely to attract electrons from other atoms.

Properties of Ionic and Covalent Compounds

• Key Differences:
  • Ionic compounds: Formed by the electrostatic attraction between positively charged ions (cations) and negatively charged ions (anions).
  • Covalent compounds: Formed by the sharing of electrons between atoms.
• Ionic Compounds:
  • Typically formed between metals and non-metals.
  • High melting and boiling points due to strong electrostatic attractions between ions.
  • Often crystalline solids at room temperature.
  • Good electrical conductors when dissolved in water or melted.
• Covalent Compounds:
  • Typically formed between non-metals.
  • Melting and boiling points vary significantly depending on their intermolecular forces.
  • Can be solids, liquids, or gases at room temperature.
  • Poor electrical conductors in water.
• Electrolyte: A substance that conducts electricity when dissolved in water. Ionic compounds are typically electrolytes because they dissociate into ions in solution.

Intermolecular Forces

• Intermolecular forces (IMFs): Attractive forces between molecules. They are weaker than the intramolecular forces (within a molecule), but they can still significantly affect a substance's physical properties.
• Types of Intermolecular Forces:
  • London dispersion force: Weakest type of IMF, present in all substances. It arises from temporary, induced dipoles in non-polar molecules.
  • Dipole-dipole force: Exists between polar molecules. It is stronger than London dispersion forces.
  • Hydrogen bonding: Strongest type of IMF, occurring between molecules containing a hydrogen atom bonded to a highly electronegative atom (like oxygen, nitrogen, or fluorine). It is responsible for many of water's unique properties.
• Stronger IMFs lead to higher melting and boiling points. This is because more energy is needed to overcome the stronger attractions between molecules to change the state of matter.

Diatomic and Polyatomic Molecules

• Diatomic molecule: A molecule composed of two atoms, either of the same element (e.g., O₂) or different elements (e.g., CO).

• Polyatomic molecule: Molecule consisting of three or more atoms (e.g., H₂O, CO₂).

  Octet Rule and Bonding Capacity

• Octet rule: Atoms tend to gain, lose, or share electrons to achieve a stable configuration of 8 valence electrons (like a noble gas).

• Bonding capacity: The number of chemical bonds an atom can form. It is often related to the number of valence electrons an atom has.

Types of Bonds

• Ionic bond: Formed by the electrostatic attraction between oppositely charged ions, usually formed between metals and non-metals.
• Covalent bond: Formed by the sharing of electrons between two atoms, typically between non-metals.

Predicting Bond Type

• Electronegativity difference (ΔEN): Can be used to predict bond type:
  • ΔEN = 0: Non-polar covalent bond (equal sharing of electrons)
  • ΔEN < 1.7: Polar covalent bond (unequal sharing of electrons)
  • ΔEN > 1.7: Ionic bond (transfer of electrons)

Summary of Key Concepts

• Understanding the arrangement of atoms and their interactions is fundamental to understanding the properties of matter.
• Atomic theory has evolved over centuries through the contributions of scientists like Democritus, Aristotle, Dalton, Thomson, Rutherford, and Chadwick. Atomic models have become increasingly refined to reflect our understanding of subatomic particles and the forces influencing their behavior.
• Atomic radius, ionization energy, electron affinity, and electronegativity are fundamental properties that influence the reactivity and bonding behavior of elements.
• Intermolecular forces are crucial for comprehending the physical properties of substances such as their melting and boiling points.
• Bonding types, including ionic and covalent, dictate the characteristics of compounds and their interactions with other substances.
• Properties of ionic and covalent compounds are determined by their bonds and intermolecular interactions, influencing their appearance, physical state, and chemical behavior.