Atomic Models, Periodic Trends, and Chemical Bonding

Classified in Chemistry

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Scientific models used to explain why and how atoms form molecules:

  • Lewis dot structure
  • Valence bond theory

Valence bond theory describes a covalent bond as the overlap of half-filled atomic orbitals (each containing a single electron) that yield a pair of electrons shared between the two bonded atoms. According to valence bond theory, a covalent bond results when two conditions are met: (1) an orbital on one atom overlaps an orbital on a second atom and (2) the single electrons in each orbital combine to form an electron pair.

  • Molecular orbital theory

model that describes the behavior of electrons delocalized throughout a molecule in terms of the combination of atomic wave functions. It describes the distribution of electrons in molecules in much the same way that the distribution of electrons in atoms is described using atomic orbitals. Using quantum mechanics, the behavior of an electron in a molecule is still described by a wave function, Ψ, analogous to the behavior in an atom. Just like Electrons around isolated atoms, electrons around atoms in molecules are limited to discrete (quantized) energies. The region of space in which a valence electron in a molecule is likely to be found is called a molecular orbital (Ψ2). Like an atomic orbital, a molecular orbital is full when it contains two electrons with opposite spin.

Meaning and periodic trends of:

  • Electron configurations allow us to understand many periodic trends. Covalent radius increases as we move down a group because the n level (orbital size) increases. Covalent radius mostly decreases as we move left to right across a period because the effective nuclear charge experienced by the electrons increases, and the electrons are pulled in tighter to the nucleus. Anionic radii are larger than the parent atom, while cationic radii are smaller, because the number of valence electrons has changed while the nuclear charge has remained constant. Ionization energy (the energy associated with forming a cation) decreases down a group and mostly increases across a period because it is easier to remove an electron from a larger, higher energy orbital. Electron affinity (the energy associated with forming an anion) is more favorable (exothermic) when electrons are placed into lower energy orbitals, closer to the nucleus. Therefore, electron affinity becomes increasingly negative as we move left to right across the periodic table and decreases as we move down a group. For both IE and electron affinity data, there are exceptions to the trends when dealing with completely filled or half-filled subshells.
  • covalent and ionic atomic sizes:covalent radius
  • Ionic radius
  • ionization energy
  • Electron affinity

The chemical basis for the formation of:

  • Cations
  • Anions
  • ionic compounds
  • covalent bonds
  • molecular compounds.

Electronegativity and the different types of chemical bonds.

a measure of the tendency of an atom to attract electrons (or electron density) towards itself. It determines how the shared electrons are distributed between the two atoms in a bond. The more strongly an atom attracts the electrons in its bonds, the larger its electronegativity.

The influence exerted by the molecular structure on the way the presence of polar bonds affect the molecular polarity of covalent compounds.

The valence shell electron pair repulsion model and the difference between the concepts of electron pair geometry and molecular structure.

Valence shell electron-pair repulsion theory (VSEPR theory) enables us to predict the molecular structure, including approximate bond angles around a central atom, of a molecule from an examination of the number of bonds and lone electron pairs in its Lewis structure. The VSEPR model assumes that electron pairs in the valence shell of a central atom will adopt an arrangement that minimizes repulsions between these electron pairs by maximizing the distance between them. The electrons in the valence shell of a central atom form either bonding pairs of electrons, located primarily between bonded atoms, or lone pairs. The electrostatic repulsion of these electrons is reduced when the various regions of high electron density assume positions as far from each other as possible. We differentiate between these two situations by naming the geometry that includes all electron pairs the electron pair geometry. The structure that includes only the placement of the atoms in the molecule is called the molecular structure. The electron-pair geometries will be the same as the molecular structures when there are no lone electron pairs around the central atom, but they will be different when there are lone pairs present on the central atom.

The interpretation given by the valence bond theory for the hybridization of atomic orbitals, and the formation of sigma and pi bonds:

Valence bond theory describes a covalent bond as the overlap of half-filled atomic orbitals (each containing a single electron) that yield a pair of electrons shared between the two bonded atoms. We say that orbitals on two different atoms overlap when a portion of one orbital and a portion of a second orbital occupy the same region of space.

The overlap of two s orbitals (as in H2), the overlap of an s orbital and a p orbital (as in HCl), and the end-to-end overlap of two p orbitals (as in Cl2) all produce sigma bonds (σ bonds).

A pi bond (π bond) is a type of covalent bond that results from the side-by-side overlap of two p orbitals.

The formation of bonding and antibonding molecular orbitals:

  • Adding electrons to these orbitals creates a force that holds the two nuclei together, so we call these orbitals bonding orbitals.
  • The attractive force between the nuclei and these electrons pulls the two nuclei apart. Hence, these orbitals are called antibonding orbitals.

The side-by-side overlap of two p orbitals gives rise to a pi (π) bonding molecular orbital and a π* antibonding molecular orbital

The explanation of why the diatomic molecule of oxygen is paramagnetic.

when we pour liquid oxygen past a strong magnet, it collects between the poles of the magnet and defies gravity, as in Figure8.1. Such attraction to a magnetic field is called paramagnetism, and it arises in molecules that have unpaired electrons. And yet, the Lewis structure of O2 indicates that all electrons are paired. How do we account for this discrepancy?

The numerical part consists of three operational exercises focused on:

  1. calculations of the formal charge and the drawing of feasible resonance structures.

formal charge = # valence shell electrons (free atom)− # lone pair electrons − # bonds

1. A molecular structure in which all formal charges are zero is preferable to one in which some formal charges are not zero.

2. If the Lewis structure must have nonzero formal charges, the arrangement with the smallest nonzero formal charges is preferable.

3. Lewis structures are preferable when adjacent formal charges are zero or of the opposite sign.

4. When we must choose among several Lewis structures with similar distributions of formal charges, the structure with the negative formal charges on the more electronegative atoms is preferable.

2. The identification of the hybridization state of central atoms and the number of sigma and pi bonds in simple molecular structures.

3. The use of the octet rule and Lewis model to draw the Lewis dot structure

  1. Count up valence electrons
  2. Skeleton around least electronegative
  3. Distribute excess around outside atoms
  4. Remaining e- go in the center
  5. Multiple bonds for octets

A multiple-choice section of 10 points designed to assess your overall comprehension of:

The chemical bonding when you are confronted with the analysis of questions for which you must decide from options that differ in only one or two critical factors.

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