Atoms combine to achieve a more stable electron configuration. Maximum stability results when an atom is isoelectronic with a noble gas.
Valence electrons are the outer shell electrons of an atom. The valence electrons are the electrons that participate in chemical bonding.
A Lewis dot symbol consists of the symbol of an element and one dot for each valence electron in an atom of the element.
As a rule, the elements most likely to form cations in ionic compounds are the alkali metals and alkaline earth metals, and the elements most likely to form anions are the halogens and oxygen.
An ionic bond is the electrostatic force that holds ions together in an ionic compound. The compound itself is electrically neutral.
Examples (remember to balance the equation):
A quantitative measure of the stability of any ionic solid is its lattice energy, defined as the energy required to completely separate one mole of a solid ionic compound into gaseous ions.
There is a rough correlation between lattice energy and melting point. The larger the lattice energy, the more stable the solid and the more tightly held the ions. It takes more energy to melt such a solid, and so the solid has a higher melting point than one with a smaller lattice energy. Note that $\ce{MgCl2}$, $\ce{MgO}$, and $\ce{CaO}$ have unusually high lattice energies. The first of these ionic compounds has a doubly charged cation ($\ce{Mg^{2+}}$) and in the second and third compounds there is an interaction between two doubly charged species ($\ce{Mg^{2+}}$ or $\ce{Ca^{2+}}$ and $\ce{O^{2−}}$).
Covalent bond is a bond in which two electrons are shared by two atoms. Covalent compounds are compounds that contain only covalent bonds.
For the sake of simplicity, the shared pair of electrons is often represented by a single line. Thus, the covalent bond in the hydrogen molecule can be written as $\ce{H-H}$. In a covalent bond, each electron in a shared pair is attracted to the nuclei of both atoms. This attraction holds the two atoms in $\ce{H2}$ together and is responsible for the formation of covalent bonds in other molecules.
lone pairs: pairs of valence electrons that are not involved in covalent bond formation.
A Lewis structure is a representation of covalent bonding in which shared electron pairs are shown either as lines or as pairs of dots between two atoms, and lone pairs are shown as pairs of dots on individual atoms.
Octet rule: An atom other than hydrogen tends to form bonds until it is surrounded by eight valence electrons.
A single bond, two atoms are held together by one electron pair. If two atoms share two pairs of electrons, the covalent bond is called a double bond. A triple bond arises when two atoms share three pairs of electrons, as in the nitrogen molecule ($\ce{N2}$).
For the same pair of atoms, multiple bonds are shorter than single covalent bonds. Bond length is defined as the distance between the nuclei of two covalently bonded atoms in a molecule. Generally, shorter means stronger.
In a Polar covalent bond, or simply a polar bond, the electrons spend more time in the vicinity of one atom than the other.
Electronegativity is the ability of an atom to attract toward itself the electrons in a chemical bond.
Electronegativity is related to electron affinity and ionization energy. Thus, an atom such as fluorine, which has a high electron affinity (tends to pick up electrons easily) and a high ionization energy (does not lose electrons easily), has a high electronegativity. On the other hand, sodium has a low electron affinity, a low ionization energy, and a low electronegativity.
In general, electronegativity increases from left to right across a period in the periodic table, as the metallic character of the elements decreases. Within each group, electronegativity decreases with increasing atomic number, and increasing metallic character.
Electronegativity and electron affinity are related but different concepts. Both indicate the tendency of an atom to attract electrons. However, electron affinity refers to an isolated atom's attraction for an additional electron, whereas electronegativity signifies the ability of an atom in a chemical bond (with another atom) to attract the shared electrons.
Oxidation number refers to the number of charges an atom would have if electrons were transferred completely to the more electronegative of the bonded atoms in a molecule.
Requirements¶
- Understand the concepts
- Understand how to distinguish ionic band and covalent bond based on the electronegtivities.
Requirements¶
- Remember the Lewis structures of the example compounds in the textbook.
An atom's formal charge is the electrical charge difference between the valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure.
To assign the number of electrons on an atom in a Lewis structure, we proceed as follows:
Example
Requirements¶
- Learn how to count formal charge.
A resonance structure is one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure.
The term resonance itself means the use of two or more Lewis structures to represent a particular molecule.
A common misconception about resonance is the notion that a molecule such as ozone somehow shifts quickly back and forth from one resonance structure to the other. Keep in mind that neither resonance structure adequately represents the actual molecule, which has its own unique, stable structure. “Resonance” is a human invention, designed to address the limitations in these simple bonding models.
Examples
Requirements¶
- Understand what is resonance structure and remember the examples.
The octet rule applies mainly to the second-period elements. Exceptions to the octet rule fall into three categories characterized by an incomplete octet, an odd number of electrons, or more than eight valence electrons around the central atom.
In some compounds, the number of electrons surrounding the central atom in a stable molecule is fewer than eight. Consider, for example, beryllium, and boron.
Example, in boron trifluoride there are only six electrons around the boron atom
Although boron trifluoride is stable, it readily reacts with ammonia.
The bond between $\ce{N}$ and $\ce{B}$ is called a coordinate covalent bond, defined as a covalent bond in which one of the atoms donates both electrons.
Some molecules contain an odd number of electrons. Among them are nitric oxide ($\ce{NO}$) and nitrogen dioxide ($\ce{NO2}$).
Odd-electron molecules are sometimes called radicals. Many radicals are highly reactive. For example, when two nitrogen dioxide molecules collide, they can form dinitrogen tetroxide in which the octet rule is satisfied for both the N and O atoms:
Atoms of the second-period elements cannot have more than eight valence electrons around the central atom, but atoms of elements in and beyond the third period of the periodic table form some compounds in which more than eight electrons surround the central atom. In addition to the $3s$ and $3p$ orbitals, elements in the third period also have $3d$ orbitals that can be used in bonding. These orbitals enable an atom to form an expanded octet.
One compound in which there is an expanded octet is sulfur hexafluoride, a very stable compound. The electron configuration of sulfur is $[\ce{Ne}]3s^23p^4$. In $\ce{SF6}$, each of sulfur's six valence electrons forms a covalent bond with a fluorine atom, so there are twelve electrons around the central sulfur atom:
Requirements¶
- Understand the concepts.
- Remember the examples.
Bond enthalpy is the enthalpy change required to break a particular bond in 1 mole of gaseous molecules.
The bond enthalpy of the same bond, for example $\ce{O-H}$, in two different molecules, such as methanol ($\ce{CH3OH}$) and water ($\ce{H2O}$), will not be the same because their environments are different. Thus, for polyatomic molecules we speak of the average bond enthalpy of a particular bond. Table 9.3 in the textbook.
The enthalpy of reaction in the gas phase is given by
$$\Delta H=\Sigma D_{\text { bonds broken }}-\Sigma D_{\text { bonds formed }}$$The reason is
Requirements¶
- Understand the concepts.
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