Chemical Bonding Properties

The Covalent Bond

• Atoms gain stability by sharing electrons to form covalent bonds.

Why do atoms bond?

• Understanding bonding helps in developing new chemicals and technologies.
• Noble gases, with stable electron arrangements, rarely form compounds.
• Stability is linked to lower energy states; atoms gain stability by achieving noble-gas electron configurations.
• Atoms can acquire stable electron configurations by sharing valence electrons.
• Water molecules form when hydrogen and oxygen atoms share electrons.

What is a covalent bond?

• Atoms in nonionic compounds share electrons to form covalent bonds.
• A molecule forms when two or more atoms bond covalently.
• Shared electrons are part of the outer energy levels of both atoms in a covalent bond.
• Covalent bonding typically occurs between elements near each other on the periodic table.
• Most covalent bonds form between atoms of nonmetallic elements.

Covalent bond formation

• Diatomic molecules like hydrogen (H2), nitrogen (N2), oxygen (O2), fluorine (F2), chlorine (Cl2), bromine (Br2), and iodine (I2) form when atoms share electrons.
• Diatomic molecules are more stable than individual atoms.
• Fluorine, with seven valence electrons, needs one more to form an octet.
• Forces of attraction and repulsion act between two approaching fluorine atoms.
• Attraction increases until a point of maximum net attraction is achieved and covalent bond and molecule forms.
• In a covalent bond, the most stable arrangement of atoms exists at an optimal internuclear distance, with net attraction exceeding net repulsion.
• Fluorine exists as a diatomic molecule because sharing one pair of electrons provides each fluorine atom with a stable noble-gas configuration.
• In a fluorine molecule, each atom has one shared electron pair (covalently bonded) and three unshared pairs, also known as lone pairs

Single Covalent Bonds

• Sharing one pair of electrons results in a single covalent bond, like in hydrogen molecule.
• Shared electron pair is often referred to as the bonding pair.
• In a hydrogen molecule, covalently bonded atoms equally attracts the pair of shared electrons.
• The two shared electrons belong to each participating atom, resulting in the helium configuration and lower energy.
• Lewis structures illustrate the arrangement of electrons in a molecule.
• A line or a pair of vertical dots between element symbols represents a single covalent bond in a Lewis structure.

Group 17 and single bonds

• Halogens (Group 17), like fluorine, have seven valence electrons and need one more to complete octet.
• Halogens form single covalent bonds with other nonmetals like carbon.
• Some Group 17 elements form covalent bonds with identical atoms like Fluorine (F2), and Chlorine (Cl2).

Group 16 and single bonds

• Group 16 elements can share two electrons and form two covalent bonds.
• Oxygen has electron configuration of 1s22s22p4, forms water molecule with two hydrogen atoms.
• Each hydrogen atom has the helium configuration by sharing one electron with oxygen.
• Oxygen has the neon configuration by sharing one electron with each hydrogen atom.
• The oxygen atom in the water molecule has two single covalent bonds and two unshared pairs of electrons

Group 15 and single bonds

• Group 15 elements form three covalent bonds with nonmetal atoms.
• Nitrogen has electron configuration of 1s22s22p³, forms ammonia (NH3) with three single covalent bonds.
• Three nitrogen electrons each bond with one hydrogen atom, leaving one unshared electron pair.
• Nitrogen forms similar compounds with Group 17 elements like nitrogen trifluoride (NF3), nitrogen trichloride (NCl3), and nitrogen tribromide (NBr3), with each atom sharing an electron pair.

Group 14 and single bonds

• Group 14 elements form four covalent bonds.
• Methane (CH4) forms when a carbon atoms bonds with four hydrogen atoms
• Carbon atom with electron configuration of 1s22s22p2 has four valence electrons and needs four more for noble gas configuration and forms four bonds
• Because hydrogen atom, a group 1 element, has one valence electron, four hydrogen atoms are used to provide required electrons for a carbon atom
• Carbon also forms single covalent bonds with other nonmetal atoms, including those in group 17.

The sigma bond

• Single covalent bonds are also called sigma bonds (σ).
• Sigma bonds have shared electrons in area centered between two atoms.
• Valence atomic orbitals overlap end-to-end, concentrating electrons in bonding orbital, a localized region where bonding electrons are most likely to be.
• Sigma bonds can form from overlapping s-orbitals, overlapping s- and p-orbitals or overlapping p-orbitals
• Water (H2O), ammonia (NH3), and methane (CH4) have sigma bonds.

Multiple Covalent Bonds

• Atoms can achieve noble gas configurations by sharing more than one pair of electrons, forming multiple covalent bonds.
• Double and triple covalent bonds are examples of multiple bonds.
• Carbon, nitrogen, oxygen, and sulfur atoms often form multiple bonds with other nonmetals.
• Number of covalent bonds corresponds to number of valence electrons needed to complete an octet.

Double bonds

• A double covalent bond occurs when two atoms share two pairs of electrons.
• Oxygen atoms only exist as diatomic molecules where each oxygen atom has six valence electrons and needs two more, as shown in Figure 8.8a.
• Double covalent bond forms when each oxygen atom shares two electrons to achieve a noble-gas configuration.

Triple bonds

• A triple covalent bond is created when three pairs of electrons between 2 atoms are shared.
• Nitrogen molecules (N2) have a triple covalent bond.
• Each nitrogen atom shares three electron pairs with other nitrogen atom, as shown in Figure 8.8b.

The pi bond

• Multiple covalent bond consists of one sigma bond and at least one pi bond. Pi bond (π) forms with parallel orbitals overlapping and sharing electrons where the shared electron pair occupies space above and below the bond line.
• Molecules with multiple covalent bonds contain both sigma and pi bonds.
• Double covalent bond consists of one pi bond and one sigma bond.
• Triple covalent bond consists of two pi bonds and one sigma bond.

The Strength of Covalent Bonds

• Covalent bonds involve attractive and repulsive forces, with balanced nuclei/electron attraction.
• Nuclei balance repels other nuclei, and electron repels other electrons.
• Covalent bonds differ in strength, and are influenced by multiple factors.

Bond length

• Covalent bond strength depends on distance between bonded nuclei.
• Bond length is between the two bonded nuclei at maximum attraction and determined by atom sizes and number of shared electron pairs.
• Increasing number of shared pairs shortens bond length
• Shorter bond length makes stronger bond.
• Single bonds (e.g., in F2) are weaker than double bond (e.g., in O2), and double bond (e.g., in O2) is weaker than triple bond (e.g., in N2)

Bonds and energy

• Bond formation and breakage involves energy changes.
• Energy is released when bond forms, but energy input is needed to break bond.
• Bond-dissociation energy is required energy to break specific covalent bond (always positive value).
• Lower bond length is associated to greater bond-dissociation energy.
• Bond energy and bond length have an inverse relationship.
• Sum of bond-dissociation energy values represents total chemical potential energy in molecule.
• Total energy change is defined by energy used to both break and form bonds.
• Endothermic reaction requires greater energy to break existing bonds than released during new bond formation.
• Exothermic reaction occurs when more energy released from product's bond formation exceeds that required for reactants

Naming Binary Molecular Compounds

• Some compounds are commonly named, requiring scientific names to decipher composition.
• Follows similar processes described in Chapter 7 for ionic compound
• Composed of two nonmetal atoms, unlike metal atoms or ions.
• N2O (dinitrogen monoxide or nitrous oxide/laughing gas) explains rules

Naming Binary Molecular Compounds’ rules

• The first element in the formula is always named first, using the entire element name, like N for nitrogen.
• The second element in the formula is named using its roots, and adding the suffix -ide, like O for oxygen in oxide.
• Prefixes are used to indicate the number of atoms of each element present in the compound.
• Two atoms of Nitrogen would use the prefix di and become dinitrogen, One atom of oxygen uses the prefix mono and becomes monoxide
• To avoid awkward proncuation, the mono- prefix is never utilized on the first element, i.e. CO is carbon monoxide, not monocarbon monoxide
• Furthermore, vowels are removed to also alleviate awkwardness, for example, monooxide becomes monoxide

Common names for some molecular compounds

• Many Binary molecular compounds are known by common, everyday names and the scientific, formalized names, i.e. H2O is water
• For example, baking soda is sodium hydrogen carbonate and common table salt is sodium chloride
• Discovered before naming conventions, nitrous oxide and water is given a common name
• Other binary covalent that are named by common terms are ammonia (NH3), hydrazine (N4H4), and nitric oxide (NO).

Naming Acids

• Water solutions of some molecules form acids and are named as such.
• Acids are important due to their specific unique chemical properties.
• Any compound that produces hydrogen ions (H+) in solution, is acidic, as HCi
• Two acid types exist, called binary acids and oxyacids.

Naming binary acids

• A Binary acids are acids containing hydrogen and one other element.
• To name Binary Acids, the prefix hydro- is added, i.e. HCi is known as "hydrochloric" acid because hydrogen and chlorine become hydrochloric.
• Add suffix -ic after the root of second element
• Always add the term acid at end. Thus HCi is known as hydrochloric acid
• Although acids are binary and indicate two elements, two exception exist.
• Even with more than two acids, they are named by name rule. The acid is named the same way, with no oxygen.
• The root of second part name is the root of polyatomic ion the acid contains and called hydrocyanic acid in soulution.

Naming oxyacids

• Oxacids contain both a hydrogen atom and oxyanion are referred.
• From chapter 7, oxyanion is polyatomic ion with one or more oxygen atoms.
• When naming nitric acid (HNO3), First is identifying the oxyanion present.
• First word consitss of root oxyanion, or -hypo is a partial name with a suffix.
• If oxyanions suffix name ends with -ate, replace suffix with -ic.
• If the name endes with -ite, replace with -ous.
• The second word is always acid, so the hydrogen in oxacid is not part of name.
• Table 8.4 displays the numerous structural forms of oxacids and how to form the proper terminology

Structural Formulas

• Models represent a molecules’s covalent compounds’s structure.
• Molecular formulas utilizes symbols and numerical subscripts to shows the amount of each type of atom in a molecule.
• Ball-and stick and space filling models utilize atomic and color spheres to represent each type of atom from element
• Letter symbols shows relative positions of atoms by bonds, in structural formula.
• Drawing Lewis structures predict structure and require other structural forms to decider atom molecule

Lewis Structures

• Albeit simple to draw lewis on nonmetal formulas, utilize the method in this problems steps
  • 1. Predict the Location of Certain Atoms.
  • 2. Determine the amount of electrons for bonding.
  • 3. Determine Amount of Bonding Pairs
  • 4. Place Bonding Pairs
  • 5. Amount of Bonding Pairs Remaining
  • 6. Determine whether central atom matches Octet Rule
• Apply strategies listed in the problem steps.

Lewis structures for polyatomic ions

• Structures for polyatomic ions is structurally simple to draw, as they covalently bond unit atoms.
• The main difference is finding total amoung of electrons available for bonding.
• If presents, added electrons are negative charged while subtracting for positive.

Resonance Structures

• Molecules with double and single bonds, can have more than one lewis structural sequence
• Structual representations results in Resonance, which is defined as multiple valid structures for the representation of ions and molecules.
• Difference include the location of lone pairs, and bonding pairs, Never amount positions
• The structures of O3 and the Polyatomic ions NO3, NO2, SO32¯, and CO32– formations use Resonance
• Each each molecules and ions goes through Resonance and behaves as if there is one structure.
• Experimentally, the bond lengths show what are bonds are identical because a single bond shorter than doubles
• Actual bond length is the bonds in actual Resonance structures
• Practice on drawing the lewis molecules for practice

Exceptions to the Octet Rule

• Typically all structures contain an octet, although some cases can and will occur.
• Odd number of valence of electrons- First, some small groups of molecule show an odd amount, resulting not on an exact electrons pairs
• For instance, the NO2 has a valence electron of 17.
  • Figure 8.15 is a valid example for uneven electron counts

Suboctets and coordinate covalent bonds

• Suboctets and coordinate covalent are stable, even with less than even electrons present around the atom
• Rare in groups, BH3 a example where 3 covalent bonds form
• Only 6 atoms are present, too few for atoms
• Coordinate can form/need energy and 2 electrons

Expanded octets

• Last atoms contains 8 or more atoms, containing d-orbital through considering
• PCL5 example and bonds, plus single orbital and 3p orbitals

Molecular Shapes

• The Valence Shell Electron Pair Repulsion model (VSEPR) determines the molecular shape based on minimizing repulsion of shared and unshared electron pairs around the central atom.
• Imagine the atoms in a molecule to be positioned at fixed angles relative, to minimize fixed angles. And the forces of attractions relative to another.
• Unshared Pairs is also important in definition, plus pushes together to minimize repulsion.
• Each amount define what shapes in molecule shown.
• i.e. two pairs of electrons create a linear force, 3 pairs create Trigonal and four shapes create tertrahedal shape.

Hybridization

• Hybridization is the process in which atomic orbitals during chemical bonding mix and form new identical hybrid orbitals.
• The Carbon atom contains four electron valence, and the mixing occurs, resulting in hybrid orbitals, called sp, because the four orbitals form
• Sp is the most element undergoes hybridization, the number tells of total electron and follows by VSEPR theory

Electronegativity and Polarity

• Chemical bonds relate each to other through attraction of electrons
• Types depend on attractions measure type

Electron Affinity, Electronegativity, and Bond Character

• Type of bond relation to attraction of electron.
• Electron tendency amount amount
• The scale tells chemists electro amount, of molecules and specific compound amounts.
• Figure tell of the atom amounts, not and never show atoms of francium.
• Higher atom amount creates hard.

Bond character

• Character depends how strong to electron attracts of a bonding element and bond.
• Type of bonds determined amount which, elements amount
• Equal bond is determined is nonpure

Nonpolar and Polar Bonds

• Polar bonds are formed because elements share electrons
• Atom bonds equal each other in strength and are nonpolar form
• Nonpolar attracts charges with electric poles
• Polar molecule, which can't
• Differences attract water to balloon and electrical
• Noncharged nonpolar has electricity in all things. Electronegativity and Polarity is polar Molecules have less water and strength, if it's not
• Weak molecules causes a force with an amount of energy to make molecules

Test your knowledge of chemical bonding types and their impact on substance properties like melting point. Explore covalent bonds, Lewis structures, and molecular stability. Identify bonding types based on melting points of unknown solids.