Chapter 6 – Modern Chemistry Chemical Bonding Chapter 6 – Modern Chemistry

6.1 Types of Bonds A chemical bond is a mutual electrical attraction between the nuclei of one atom and the valence electrons of another atom Occurs in several ways Gaining and losing electrons Ionic Bonds Sharing electrons Covalent Bonds Overlapping electron clouds Metallic Bonds

Determining which is which? Covalent or Ionic? Electronegativity differences between atoms will determine whether it is ionic or covalent Covalent bonds share electrons and will have an electronegativity difference between 0 - 1.7 Ionic bonds form from ions that have gained or lost electrons and their difference is between 1.7 – 3.3

Polarity Covalent bonds can also show partial charges which is known as polarity Covalent compounds can be classified as polar or nonpolar depending on their shape and electronegativity differences Differences between 0 – 0.3 will be nonpolar (shared equally) Differences between 0.3 – 1.7 will be polar Example would be water in which we have two O-H single bonds Electronegativity of O is 3.5 and H is 2.1 Difference of 1.4 which makes the bond polar Values are found on page 161

6.2 Covalent Compounds Atoms will form covalent compounds when the valence electrons of one atom become attracted to nucleus of another atom This attraction will in return have to same affect on the second atoms valence and the first atoms nucleus

Formation of a Covalent Bond This cross attraction that occurs in covalent compounds is formed by the need of all atoms to fulfill the octet rule This rule states that all atoms are trying to fill their outer energy shell with 8 electrons The need to do this is a result of achieving their lowest potential energy state This amount of energy that is given off is the same about of energy that is required to break the bond This is known as bond energy

Electron Dot Diagrams Electron Dot Diagrams are a method of illustrating valence electrons They are created by taking the symbol of the element and treating it as if it is inside an imaginary box creating four sides Place the valence electrons around the symbol one on each side of the box before going back and pairing them up This will show how many unpaired electrons are available for bonding Covalent bonds will typically form at the unpaired electrons

Lewis Structures Lewis structures are the illustrations of covalent compounds using the dot diagrams When creating a Lewis structure, it is important that every unpaired electron is paired up in some fashion Atoms can create several types of covalent bonds in order to pair these up: single, double, triple, and coordinate covalent bonds Single – each atom shares one electron Double – each atom shares two electrons Triple – each atom shares three electrons Coordinate – one atom shares two while the other contributes none

Examples – Electron Dot Diagrams

Examples – Lewis Structures

6.3 Ionic Bonding Formation of Ionic Compounds Electron Dot Diagrams (not Lewis Structures) can be used to illustrate the formation of ionic compounds Illustrates the losing of one electron from sodium to form a positive ion and the gaining of one electron by chlorine to form a negative ion

Characteristics of Ionic Bonding Ionic compounds form a repeating pattern of negative and positive ions that will all interact with each other through the entire substance This is different than covalent compounds that share their electrons to form a single molecule This interaction forms a crystalline structure known as a crystal lattice With such a high level of interaction between ions, there is a lot of potential energy stored in the lattice structure This is known as the lattice energy It explains the high melting and boiling points of ionic compounds as well as why many of them are solids at room temperature

6.4 Metallic Bonding In metals, the vacant p orbitals of the their outer energy levels will overlap and allow the valence electrons to move freely throughout the entire substance Creates a sea of electrons Because of this overlap, metals have their unique properties Ductility Malleability Good conductors of the heat and electricity

6.5 Molecular Geometry VSEPR Theory The arrangement of shared pairs of electrons surrounding atoms must be oriented as far apart as possible If an atom has two shared pairs of electrons, then the arrangement must be 180º separation If an atom has three shared pairs of electrons, then the arrangement must be 120º separation If an atom has four shared pairs of electrons, then the arrangement must be 109.5º separation These occur without any unshared pair of electrons If an atom has unshared pairs of electrons, it will cause the angles to lessen just slightly due to the unshared pair having a greater force of repulsion than the shared pair For instance, in water, there are two shared and two unshared but instead of having angles at 109.5º, the angle is 105º

VSERP Theory Linear – 2 atoms bonded to central atom, 0 lone pairs Trigonal planar – 3 atoms bonded to the central atom, 0 lone pairs Bent/Angular – 2 atoms bonded to central atom, 1 or 2 lone pair Tetrahedral – 4 atoms bonded to central atom, 0 lone pairs Trigonal Pyramidal – 3 atoms bonded to central atom, 1 lone pair Trigonal bipyramidal – 5 atoms bonded to central atom, 0 lone pairs Octahedral – 6 atoms bonded to central atom, 0 lone pairs

Molecular Geometry Hybridization Occurs when two or more atomic orbitals of similar energies are mixed together to produce a hybrid orbital Carbon is a good example of this process Outer configuration is 2s2 2p2 but instead of only having two unpaired electrons these two orbitals blend to form an sp3 orbital at the second energy level This will also help explain why certain atoms can break the octet rule and why shape sometimes will affect polarity

Intermolecular attractions Molecules, which have atoms covalently bonded to each other, are also attracted to other molecules by various forces These forces are dipole-dipole interaction, hydrogen bonding, and London dispersion forces Dipole-dipole interaction is the attraction between the partial positive and negative ends of a polar molecule Hydrogen bonding is the attraction between an unshared pair of electrons and the hydrogen atom (when the hydrogen is connected to a highly electronegative atom) London dispersion forces is the weakest form of attraction that occurs when instantaneous dipoles occur due to the movement of the electrons