Chemical Bonding

Chemical Bonding

Chemical Bonds • Form when atoms or ions are strongly attached to one another • Defined as forces of attraction that hold two atoms together and allows them to function as a unit • Three main types: • Ionic • Covalent • Metallic

The Octet Rule • Atoms tend to gain, share, or lose electrons in order to obtain a full set of valence electrons (in most cases this equals 8) • An octet of electrons consists of full s and p sublevels on an atom. • Exceptions: transition elements and rare earth elements

Determining Types of Bonds • Determined by electronegativity differences - difference determines the percentage ionic or percentage covalent character • Type of bonds are determined by which percentage is prevalent

Ionic Bonding • Occurs when electrons are transferred from one atom to another, forming two ions • Cation – positive ion Anion – negative ion • The ions stay together because of electrostatic attractions – ionic bond • Ionic bonds NEVER form molecules • Ionic bonds form easily between alkali metals and halogens • All ionic compounds are electrically neutral

Example - + Na Cl Na + Cl +

Properties of Ionic Compounds • Ionic compounds do not form molecules; they form a crystal lattice • Formula unit – the simplest collection of atoms from which an ionic compound’s formula can be established • The ions lower their potential energy forming orderly, 3-D array in which the cations and anions are balanced • Formula units arrange themselves in repeating patterns

This is a crystal of CaCl2. Each ion is held rigidly in place by strong electrostatic forces that bond it to several oppositely charged ions

Other Properties • Normally form between metals and nonmetals • Ionic compounds have ions that form very strong bonds, which makes them hard and brittle • They have high melting points and high boiling points • Most are solids at room temperature

Properties continued • When dissolved in water, the solution will conduct electricity • Conduct electricity in the molten state, but will not conduct electricity in the solid state • Tend to be soluble in water • Crystallize as sharply defined particles • One atom has a low electronegativity and a low electron affinity • The other is vise versa

Types of Ions • There are two types of ions • Monatomic: cation or anion that consists of a single atom. Examples: Na+ and Cl- • Polyatomic: two or more atoms that act as a single ion (or particle). Examples: (CO3)2- and (OH)-

Monatomic Ions

Common Polyatomic Ions

Metallic Bonds • Metal members of the representative groups have some, if not all, vacant p orbitals • Many of the transition metals contain vacant d orbitals - allow electrons to roam freely throughout the metal • Electrons are delocalized - they do not belong to any one atom (sea of mobile electrons) • Metallic bonding is the result of the attraction between metal atoms and the surrounding sea of electrons • Responsible for metallic properties such as conductivity, malleability, ductility, and luster

Forming Covalent Bonds • covalent (co – sharing; valent – outermost shell) - when electrons are shared between two nuclei • AKA: molecular bonds • molecules – a group of atoms held together by covalent bonds

Electron Pairs in Covalent Bonds • Unshared pairs – pairs of electrons that do not participate in bonding and belong to only one atom - also called lone pairs • Bonding pairs – pairs of electrons being shared between two atoms thus creating a covalent bond • Electrons are not always equally shared • Unequal/equal sharing is determined by electronegativity differences

Covalent Bonds • Electronegativity - tendency to attract electrons in a chemical bond • polar – “having opposite ends” • covalent bonds in which the bonding electrons are more strongly attracted by one of the bonding atoms • nonpolar – not “having opposite ends” • covalent bonds in which the bonding electrons are shared equally between the 2 bonding atoms

Properties of Covalent Compounds • Typically low melting points • Most are gases, liquids, or very soft solids at room temperature • Do not conduct electricity • Are brittle when solids • Typically form between nonmetals

Types of Covalent Bonds • Covalent bonds, unlike ionic bonds, can form multiple bonds • Single bonds – two atoms share two electrons (one pair) • Double bonds – two atoms share four electrons (two pair) • Triple bonds – two atoms share six electrons (three pair)

Types of Covalent Bonds • Coordinate covalent bonds – a covalent bond in which a single atom contributes both of the electrons to a shared pair • Covalent bonds are separated into two types of bonds • Sigma bonds (s) • pi bonds (p)

Sigma and Pi Bonds • Sigma bonds • Formed along the horizontal axis between two atoms • The primary bonds • Pi bonds • formed above and below the horizontal axis between two atoms • The secondary bonds

Bond Composition • Single bonds - consist of only one sigma bond • Double bonds - consist of one sigma bond and one pi bond • Triple bonds - consist of one sigma bond and two pi bonds • http://Sigma and Pi Bonds

Bond Guide

Polarity • Remember: Polar vs Nonpolar are also determined by electronegativity differences • Polar Bonds: 0.4 < x < 1.7 ( x represents electronegativity difference) • Nonpolar Bonds: x < 0.4 • Areas of partial charge build up because a shift in electron charge density occurs • Shift itself indicated by use of arrows along the bond • Partial charge indicated by uses of

Bond Length • Bond length – average distance between bonded atoms • Measured from center of one nucleus to center of the neighboring nucleus • Different pairs of atoms form bonds of different length • Atomic radius of each atom participating in the bond therefore directly affects bond length • Not fixed because atoms vibrate through the bond in a spring-like fashion • Multiple bonds are shorter than single bonds

Bond Angle • Bond angle – the angle between the two bond axes • Bond angles are also not fixed because of the atom vibration

Bond Energy • Bond energy – energy required to break a bond • indicates bond strength • usually reported in units of kilojoules/mole • the closer the atoms the greater the bond energy required to separate them • more energy is required to break multiple bonds than single bonds • also called bond dissociation energy

Chemical Bonding