Chapter 6 Chemical Bonding Bonding Atoms seldom exist

Chapter 6 Chemical Bonding

Bonding Atoms seldom exist as independent particles in nature--instead they are combinations of atoms that are held together by chemical bonds.

A chemical bond is a mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together. By bonding, the atoms decrease in potential energy, thereby creating

bonding: Ionic bonding: bonding chemical bonding that results from the electrical attraction between large numbers of cations and anions (metal and nonmetal) n. Atoms completely give up electrons to other

Large numbers of oppositely charged ions come together because of electrical attraction (p+ and e -). Also note the change in the size of atoms (Atomic and Ionic Radii).

Covalent bonding chemical bonding that results from the sharing of electron pairs between two atoms (nonmetal and nonmetal). Ex.

Sharing one or more of the valence electrons. The shared electrons are simultaneously attracted to both atomic nuclei (p+ and e-).

Ionic or Covalent? ? ? Bonding is rarely purely ionic or covalent. It usually falls somewhere between, depending on how strongly the atoms of each element attract electrons. Electronegativity measure of an atom’s ability to attract

Remember… Electronegativity increases across a period and decreases down a group. Which atom has the highest electronegativity? n Flourine Which atom has the lowest?

Ionic vs. Covalent Bond Type Electronegativity % ionic difference character Nonpolar covalent 0. 0 to <0. 4 Polar covalent 0. 4 to <1. 7 Ionic 1. 7 and greater 4% ionic 96 % covalent 4 -50% ionic 50% covalent

Practice Elements H and Cl Li and Cl Na and Cl H and H Electronegativity Bond type difference

Nonpolar covalent bond a covalent bond in which the bonding electrons are shared equally by the bonding atoms, resulting in a balanced distribution of electrical

Polar covalent bonds that have an uneven distribution of charge(one end of the molecule is more electronegative that the other ones). Example: H 2 O

Examples: Nonpolar H-H Polar H-Cl What about HBr, CCl 4, CO 2, NH 3?

Notice the + and – poles.

Covalent Bonding Molecular Compounds Many chemical compounds including most of the chemicals that are in living things, are composed of molecules n Molecule a neutral group of atoms that are held together by covalent bonds n Molecular compound a chemical compound whose simplest units

The composition of a compound is given by its chemical formula. n. Chemical formula indicates the relative numbers of atoms of each kind in a chemical compound by using

Molecular formula shows the types and numbers of atoms combined in a single molecule of a molecular compound n Diatomic molecule a molecule containing only two atoms n

Covalent Bond Length the distance between two bonded atoms at their minimum potential energy, that is, the average distance between two bonded atoms Bond Energy the energy required to break a chemical bond and form neutral isolated atoms (in kj/mol) n Bond length and bond energies

Octet Rule Octet rule chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level They want to be like the noble gases, with their outer s and p

Exceptions to the Octet Rule Hydrogen-only needs 2 electrons to be happy Be can have only 4 electrons B and Al can have only 6 electrons rd 3 row or lower can sometimes have more than 8

Types of covalent bonds Single n Hydrogen (H 2) Double n Ethene (C 2 H 4) Triple n Nitrogen (N 2)

Compounds Ionic compound composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal n. The chemical formula of an ionic compound represents the simplest ratio of the compounds combined ions

Formula unit the simplest collection of atoms from which an ionic compound’s formula can be established n The ratio depends on

Bonding In order for ions to minimize their potential energy, they combine in an orderly arrangement known as a crystal lattice. Lattice energy the energy released when one mole of an ionic crystalline compound is formed from gaseous ions

Shows actual arrangement, with electron clouds just touching Expanded view clarifies the positioning of the ions in the structure.

Compounds Molecular (covalent) Ionic Low melting point High melting point Low boiling point High boiling point Volatile Non-volatile Doesn’t conduct electricity Conducts electricity in solution Hard, brittle

Ionic Compounds Remember, valence electrons are transferred from one atom to another Must have an element with low electronegativity and an element with high electronegativity 2 ions are formed-one positive and one negative (cation, anion) n They are attracted to each other by their opposite charges

Electron-Dot Notation Electron-dot notation an electron configuration in which only the valence electrons of an atom of a particular element are shown, indicated by dots placed around the elements symbol Examples: F H N

Dot Diagrams for Ionic Compounds Metals lose valence electrons and have positive charges Nonmetals gain enough electrons to have 8 valence electrons and have a negative charge Alternate positive and negative ions in the compound formula

Practice Na. Cl Mg. F 2 Ga 2 O 3 Ca. O

Lewis Structures Electron-dot notation can also be used to represent molecules by combining the notations of two individual atoms n. H : H F F n Shared pair- a pair of electrons involved in bonding n Unshared pair (lone pair)- a pair of electrons that is not involved in bonding n You can change the shared pair to a

Lewis Structures formulas in which atomic symbols represent nuclei and inner-shell electrons, dot-pairs or dashes between two atomic symbols represent electron pairs in covalent bonds, and dots adjacent to only one atomic symbol represent unshared pairs Structural formula indicates that kind, number, arrangement and bonds, but not the unshared pairs

Structures 1. Determine the type and number of atoms in the molecule 2. Write the electron dot notation for each type of atom in the molecule 3. Determine the total number of valence electrons in the atoms to be combined 4. Arrange the atoms to form a skeleton structure for the

5. Add unshared pairs of electrons so that each hydrogen atom shares a pair of electrons and each other nonmetal is surrounded by 8 electrons. 6 a. Count the electrons in the structure to be sure that the number of valence electrons used equals the number available 6 b. If too many electrons have been used, subtract one or more lone pairs until the total number of valence electrons is correct. Then move lone pairs until all have 8.

Practice NH 3 H 2 S CH 4 H 2 O

Multiple Covalent Bonds Some atoms can share more that one electron pair n Double bond- a covalent bond produced by the sharing of 2 pairs of electrons between two atoms n Triple bond- a covalent bond produced by the sharing of 3 pairs of electrons between two atoms n Examples: C 2 H 4, C 2 H 2, N 2

More examples… CH 2 O CO 2 HCN

Resonance Structures Some molecules and ions cannot be represented by a single Lewis structure Resonance- refers to bonding in molecules or ions that cannot be correctly represented by a single Lewis structure n Example: O 3 (ozone), SO 3

Polyatomic Ions Polyatomic ions- a charged group of covalently bonded atoms n Polyatomic ions combine with ions of opposite charge to form ionic compounds n The charge of the ion results from an excess of electrons (charge) or a shortage of electrons (+ charge) + - -2

Molecular Geometry The properties of molecules depend not only on the bonding of the atoms but also on molecular geometry Molecular polarity the uneven distribution of

VSEPR Theory VSEPR (valence shell, electron-pair repulsion) states that repulsion between the sets of valence-level electrons surrounding an atom causes these sets to be oriented as far as possible away from each other

Linear Trigonal planar Tetrahedr al

Practice (use page 184 & 186) Give the molecular geometry and bond angles of the following: n HI n CBr 4 n Al. Br 3 n CH 2 Cl 2

VSEPR and Unshared Pairs The lone pairs occupy space around the central atom, but the actual shape of the molecule is determined by the positions of the atoms only n Examples w. NH 3 w H 2 O -

Hybridization To explain how the orbitals of an atom become rearranged when the atom forms covalent bonds, a different model is used Hybridization the mixing of 2 or more atomic orbitals of similar energies of the same atom to produce new orbitals of equal energies

Hybrid orbitals of equal energy produced by the combination of 2 or more orbitals on the same atom n Look at carbon, 2 of carbon’s valence electrons occupy the 2 s and 2 occupy the 2 p. To achieve four equivalent bonds, carbon’s 2 s and 2 p orbitals hybridize

Metallic Bonding Metals have unique properties 1. Excellent electrical properties 2. Shiny appearance 3. Malleability- the ability of a substance to be hammered into thin sheets 4. Ductility- the ability of a substance to be pulled into

Model The highest energy levels of most metal atoms are occupied by very few electrons The vacant orbitals overlap

The electrons become delocalized- they move freely in the empty orbitals (called a sea of electrons) The chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons

Intermolecular Forces The forces of attraction between molecules are known as intermolecular forces These forces vary in strength but are generally weaker than bonds that join atoms in molecules, ionic compounds, or metal atoms in solid metals. In other words, it generally takes far less energy to separate molecules from one another than it does to take

Dipole-Dipole Forces The strongest intermolecular forces exist between polar molecules and only polar molecules Dipole created by equal but opposite charges that are separated by a short distance The direction of a dipole is from the dipole’s positive pole to its negative pole Polar covalent molecules act as little magnets, they have positive ends and

A dipole is represented by an arrow with a head pointing toward the negative pole and a crossed tail situated at the positive pole n H Cl n The forces of attraction between polar molecules are known as dipole-dipole forces n Examples: NH 3, CO 2

Hydrogen Bonding Hydrogen bonding the intermolecular force in which a hydrogen atom that is bonded to a highly electronegative atom is attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule (usually N, O,

Forces Even noble-gas atoms and molecules that are nonpolar experience a weak intermolecular attraction London dispersion forces the intermolecular attractions resulting from the constant motion of electrons and the creation of an instantaneous dipole

London forces increase with increasing atomic or molar massthe bigger the molecule, the stronger the force! These forces arise from the weak attractive force of the electrons on one molecule for the nuclei of another molecule. This is the only way the noble gases stay together.