Matter Structure: Atoms, Bonds, and Reactions
Structure of Matter
Dalton thought that an atom could not be broken into smaller pieces.
Thomson discovered that cathode rays were negative particles, which he called electrons.
Rutherford bombarded gold foil with positively charged particles and observed that some were barely diverted, while others bounced around. His conclusions were:
- All positive charge is concentrated in the nucleus.
- The lightweight core of the hydrogen atom is formed by a positively charged particle called a proton.
- The nuclei of atoms of metals produce larger deviations, indicating a greater positive charge and a greater number of protons.
Atomic Spectra (Chadwick): Only certain long wavelengths appear as streaks or bands on a dark background.
Emission: Radiation is emitted from a source or focus.
Absorption: If radiation from a source crosses a material, the material absorbs some radiation, causing black bands to appear in the spectrum.
The radiation emitted by atoms is the propagation of electromagnetic waves, which have the same speed of propagation but differ in wavelength and frequency. The energy of this radiation is emitted and absorbed in small packets or quanta of energy, called photons.
According to the Bohr model:
- Electrons move in orbits around the nucleus without emitting energy.
- Electrons can only orbit the nucleus in specific orbits where the angular momentum of the electron is an integral multiple of h (Planck’s constant).
- When an electron moves from an outer orbit to an inner one, the energy difference is emitted as electromagnetic radiation.
- Electrons do not radiate energy while moving in an orbit, only when changing orbits.
- The energy absorbed or emitted is the sum of the energy levels.
- Electrons are arranged in various circular orbits that determine different energy levels.
Limitations of Bohr’s Model
The circular orbit model was very restricted and only applied to the hydrogen atom and other simple atoms. Spectrographs showed that many rays were doublets or triplets. The Zeeman effect showed that when an atom is subjected to a magnetic field, its spectral lines split. This model was a mixture of classical and quantum mechanics.
OJ Bernier showed that some elements had similar properties, such as triads, where the atomic weight of the central element is approximately the arithmetic mean of the atomic weights of the other two.
Newlands arranged elements in order of increasing atomic weight, noting that every eighth element repeated properties (law of octaves).
Mendeleev created the first periodic table, ordering elements by atomic weight and establishing a relationship between their masses and properties. Meyer developed a similar system, plotting properties as a function of atomic masses, observing periodic variations. Mendeleev’s table was more complete, ordering elements by increasing atomic mass and matching them according to their valences. He also placed elements according to their electronic configurations in rows and columns. Moseley determined the atomic number and found that if elements were placed in order of increasing atomic number, they fit better in the table.
Groups or Families of Elements
- Alkali metals: ns1
- Alkaline earth metals: ns2
- Transition metals: (n-1)d1-10 ns2
- Inner transition elements (lanthanides and actinides): (n-2)f1-14 (n-1)d0-1 ns2
- Earth metals: ns2p1
- Carbon group: ns2p2
- Nitrogen group: ns2p3
- Oxygen group: ns2p4
- Halogens: ns2p5
- Noble gases: ns2p6
Atomic Radius: Half the distance between the centers of two adjacent atoms.
Ionic Radius: When an atom is ionized, its volume decreases when losing electrons and increases when gaining electrons.
Ionization Energy: The energy needed to remove an electron from an atom.
Electron Affinity: The energy change when an electron is added to a gaseous atom. The sign indicates whether energy is released or absorbed.
Oxidizing Character: The tendency to gain electrons.
Electronegativity: The tendency of an atom to attract electrons when combined with other atoms.
Chemical Bonding
Atoms tend to achieve a stable state by completing their outer electron shell, usually with eight electrons (octet rule).
The stability of a bond is related to the energy of the system. The lower the energy, the more stable the bond.
Morse Curve: Represents the potential energy of two atoms or ions as they approach each other. It shows the balance between repulsive and attractive forces.
Ionic Bond
Ions are united by electrostatic attraction between opposite charges. They form a crystal lattice. The type of crystal lattice depends on the size of the ions. The coordination number is the number of ions of opposite sign surrounding a given ion.
Types of Ionic Networks
- Cubic network
- Fluorite network
- Tetragonal network
Lattice Energy: The energy released when gaseous ions combine to form a crystal.
Born-Haber Cycle: A thermochemical cycle used to determine the experimental lattice energy.
Properties of Ionic Compounds
- High melting and boiling points: Due to strong electrostatic forces.
- Solubility: Generally soluble in polar solvents like water, but not in non-polar solvents.
- Hardness: They resist scratching due to strong bonds.
- Low electrical conductivity: In solid state, ions are fixed in the lattice.
- Fragility: When a force is applied, ions of the same charge come into contact, causing repulsion and breaking the crystal.
Metallic Bond
A metallic crystal is formed by metal ions surrounded by a sea of delocalized valence electrons. The stability of the crystal is due to the interaction between the metal ions and the delocalized electrons.
Properties of Metals
- High melting and boiling points
- High electrical conductivity in solid state
- High thermal conductivity
- Good mechanical properties, malleability, and ductility
Covalent Bond
Atoms share one or more pairs of electrons. A single covalent bond is formed when two atoms share one pair of electrons.
Exceptions
- Boron trifluoride: The boron atom has only six electrons around it.
- Phosphorus pentafluoride: The phosphorus atom has ten electrons around it.
Valence Bond Theory
Atoms maintain their atomic orbitals when forming molecules. Covalent bonds are formed by the overlapping of atomic orbitals, creating a region of high electron density. Overlapping orbitals must be semi-occupied. Frontal overlap forms sigma bonds, while lateral overlap forms pi bonds.
Multiple bonds are formed when atoms share more than one pair of electrons.
Covalent Bond: When two atoms approach, their orbitals overlap. The shared region contains electrons from both atoms. The electrons must have opposite spins. This forms a sigma bond.
Hybridization of Orbitals
sp3 Hybridization: In ammonia (NH3), one of the four hybrid orbitals of nitrogen is non-bonding, while the other three form bonds with hydrogen. The bond angles are approximately 107º. The molecule has a pyramidal shape with a triangular base.
H2O: The 2s and 2p orbitals of oxygen hybridize to form four sp3 hybrid orbitals. Two of these orbitals contain lone pairs, while the other two form bonds with hydrogen. The bond angle is 104.5º due to repulsion between lone pairs. The molecule is angular.
Multiple Bonds: In ethene (C2H4), each carbon atom has sp2 hybridization. Two of the sp2 orbitals form a sigma bond between the carbon atoms, while the remaining sp2 orbitals form sigma bonds with hydrogen. The unhybridized p orbitals form a pi bond. In ethyne (C2H2), each carbon atom has sp hybridization. The two sp orbitals form a sigma bond between the carbon atoms and sigma bonds with hydrogen. The two unhybridized p orbitals form two pi bonds. The molecule is linear.
Chemical Kinetics and Thermochemistry
Chemical kinetics studies the rate of reactions, the factors that influence them, and the mechanisms by which reactants become products.
Rate of Reaction
The rate of reaction is a positive quantity that expresses the change in concentration of a reactant or product over time.
Rate Law: V = k[A]x[B]y, where V is the instantaneous rate, k is the rate constant, [A] and [B] are the molar concentrations of reactants, and x and y are experimentally determined exponents. The rate constant depends on temperature but is independent of concentrations.
Factors Affecting Reaction Rate
- Nature of Reactants: The rate varies depending on the nature of the reactants.
- Contact Area: Increasing the surface area of solid reactants increases the contact area and the reaction rate.
- Concentration: Increasing the concentration of reactants increases the frequency of collisions and the reaction rate.
- Temperature: Increasing the temperature increases the kinetic energy of molecules, the frequency of collisions, and the reaction rate.
- Presence of a Catalyst: Catalysts increase the reaction rate without being consumed in the reaction.
Collision Theory
The rate of a reaction is proportional to the frequency of effective collisions between reactant molecules. Effective collisions require proper orientation and sufficient energy.
Activation Energy: The minimum energy required for reactants to reach the transition state.
Catalysis
A catalyst is a substance that increases the rate of a chemical reaction. An inhibitor decreases the rate. Catalysts change the reaction mechanism, decreasing the activation energy. They do not affect the overall change in Gibbs free energy.
Reaction Mechanism
A reaction mechanism is a series of elementary reactions that describe the overall reaction. The molecularity of an elementary reaction is the number of molecules involved.
Thermochemistry
Standard Enthalpy of Reaction (ΔH°)
The change in enthalpy when reactants in their standard states are transformed into products in their standard states.
Standard Enthalpy of Formation (ΔH°f)
The enthalpy change when one mole of a substance is formed from its elements in their standard states.
Standard Enthalpy of Combustion (ΔH°c)
The enthalpy change when one mole of a substance is completely burned under standard conditions.
Hess’s Law: If a reaction can occur in several steps, the overall enthalpy change is the sum of the enthalpy changes of the individual steps.
Calculation of ΔH°
ΔH° = ΣΔH°f(products) – ΣΔH°f(reactants), where the enthalpies of formation are multiplied by their stoichiometric coefficients.
Enthalpy of Bond
The energy required to break a bond. Bond formation releases energy. The enthalpy of reaction depends on the energy consumed to break bonds and the energy released when new bonds are formed.
Gibbs Free Energy (ΔG)
A thermodynamic potential that combines enthalpy and entropy to determine the spontaneity of a process. ΔG = ΔH – TΔS
Standard Free Energy of Formation (ΔG°f)
The change in free energy when one mole of a compound is formed from its elements in their standard states. By convention, the free energy of elements in their standard states is zero.
Standard Free Energy of Reaction (ΔG°r)
The change in free energy when reactants in their standard states are transformed into products in their standard states. ΔG°r can be determined from the free energies of formation of the reactants and products. The sign of ΔG determines the spontaneity of a reaction at constant temperature and pressure.