PACKET 7 Chemical Bonding Reference Table PT Table

PACKET #7: Chemical Bonding Reference Table: PT & Table S

Chemical Bonding • Chemical Bond: an attraction between the protons of one atom (nucleus) and the electrons of the next atom that attaches the atom together. • Formed by transferring or sharing of electrons. • A chemical bond has stored or potential energy. • After a chemical bond is formed, the atoms have a complete outer shell they are stable (8 valence electrons)

Polarity: tendency of a molecule, or compound, to be attracted or repelled by electrical charges because of its particle arrangement of atoms around the nucleus. POLAR = PULL (UNEQUAL) NON-POLAR = NO PULL (EQUAL) • Think of it like a game of tug of war, if one end of the compound is pulling on the electrons more than the other, there is an unequal pull, and therefore, the substance is polar. If there is an equal pull, then the substance is non-polar. • This concept of polarity is determined by ELECTRONEGATIVITY DIFFERENCE

When the teams pull equally it’s like they’re not pulling at all (NONPOLAR) When the teams pull unequally, then there’s a pull (POLAR)

Electronegativity Difference Remember. . . Table S • Electronegativity: an atom’s attraction for electrons in a bond. • The greater the EN difference is the MORE POLAR. • The lower the EN difference is the LESS POLAR. • When the EN difference equals 0, then it’s NON-POLAR.

Ionic Bond • Attraction between opposite charged ions (+)(-) • Occurs when electrons are transferred from one ion (charged particle) to another • Electronegativity difference is greater than 1. 9 • Always Polar!! (so much pull that electrons are pulled

Ionic Bonds – Lewis Dot Structure • Example: Na. Cl:

Example: Ca. Cl 2

Polyatomic Ions • Ionic compounds containing polyatomic ions have both ionic and covalent bonding. • Example: KNO 3 • Notice that NO 3 - is composed of 2 nonmetals therefore the bonding is covalent between N and O but the bonding between K+ and NO 3 - is ionic.

Example: (NH 4)3 PO 4 • TRY (NH 4)3 PO 4 • Another example of a compound with both ionic and covalent bonds

Properties of Ionic Compounds • Hard • Good conductors of electricity in liquid or aqueous form only, because ions can move in solution and in liquid form, but not in solid form. • High melting and boiling points • Solid at room temperature • Dissolve in polar substances: like water. (Polar – opposite charges).

Covalent Bonds Formed when 2 atoms (both nonmetals) share electrons. [Example Cl 2 or H 2 O] • Neither atom pulls strongly enough to remove an electron from the other. Electrons are shared. • The EN difference is < 1. 9 • Unpaired electrons pair up in such a way that the atoms complete their outer shells • Covalent compounds also referred to as molecular compounds

Properties of Covalent Bonds • • Gases, liquids or solids Soft Nonmetals Poor conductors of heat and electricity because they are not charged particles. (No ions or mobile electrons) • Low melting and boiling points because of weak attraction between molecules.

Polar vs. Non-Polar Covalent Bonds • Unlike an ionic compound, a covalent compound can be classified as either a polar covalent bond, or a non-polar covalent bond. • If the EN of the atoms are different then it is a polar covalent bond. • If the EN of the atoms are the same or very similar then it is a non-polar covalent bond. 0. 0 - 0. 4 = non-polar covalent 0. 5 - 1. 6 = polar covalent

Polar Covalent Bonds • There will always be an unequal sharing of electrons due to the EN difference. • Example: HCl • EN of H = 2. 1 EN of Cl = 3. 2 • Difference is 1. 1, which is less than 1. 7, but greater than 0. 4

Non-Polar Covalent Bonds • EN difference 0. 0 – 0. 4 • The non-polar covalent bonds you must commit to memory are the diatomic molecules: • H 2, N 2, O 2, F 2, Cl 2, Br 2, I 2 • These are considered covalent bonds because they are two non-metals sharing electron, and are considered non-polar because since they are the same element, they have the same EN, and therefore the difference is 0. • Since “likes dissolve in likes” non-polar covalent compounds will only dissolve in non-polar solvents.

Number of Covalent Bonds • Single covalent bond: one pair of shared electrons; 2 electrons total • -Double covalent bond: two pairs of shared electrons; 4 electrons total • -Triple covalent bonds: three pairs of shared electrons; 6 electrons total

A Little Review. . . The Octet Rule: • Atoms seek to have eight valence electrons when in a bond. • All Noble Gases have a full octet. • Exceptions to the octet rule: H and He will only hold 2 electrons each. He already has 2 electrons and will not bond with other atoms.

Rules for Drawing Lewis Diagrams of Covalent Compounds 1. Calculate the total number of valence electrons available to the molecule or ion. 2. Organize the atoms so there is a central atom (usually the least electronegative) surrounded by outer atoms. Hydrogen is NEVER the central atom. 3. Form bonds between the central atom and outer atoms with a pair of electrons. All remaining electrons should be distributed so that each atom has 8 electrons.

Example #1 - Methane CH 4 1. Determine the total number of valence electrons available: • One carbon has 4 valence electrons. Four hydrogen, each with one valence electron, totals 4. • This means there are 8 valence electrons, making 4 pairs, available.

Example #2 - Ammonia NH 3 1. Determine the total number of valence electrons available: • One nitrogen has 5 valence electrons. Three hydrogen, each with one valence electron, totals 3. • This means there are 8 valence electrons, making 4 pairs, available.

Example #3 - Carbon tetrachloride CCl 4 1. Determine the total number of valence electrons available: • One carbon has 4 valence electrons. Four chlorine, each with 7 valence electrons, totals 28. • This means there are 32 valence electrons, making 16 pairs, available.

Example #4 – O 2 & N 2 DOUBLE BOND TRIPLE BOND

MORE . . . • Some other covalent compounds that are important to know how to draw are: H 2, F 2, Cl 2, Br 2, I 2, O 2, N 2, HCl, HF, HBr, HI, CO 2, CF 4, CBr 4, H 2 O, H 2 S, CCI 4. Si. H 4

Partially Positive & Negative • In a polar covalent bond, both of the elements are non-metals, and therefore there is no “true” + or – charges; instead there are partially (+) and partially (-) charges. • The element with the higher EN is partially (-) and the one with the lower EN is partially (+)

Polar & Non-Polar Bonds vs. Polar & Non-Polar Molecules • Just to add a little more confusion . . . • RECALL: Polar Bond = different EN Non-Polar Bond = same (similar) EN BUT Polar Molecule = Asymmetrical Non-Polar Molecule = Symmetrical

Polar Molecule • Polar molecules are asymmetrical (the compound is not a mirror image if you folded it over itself) • Polar molecules result from an unequal distribution of electrons. • These molecules are also called dipoles.

Non-Polar Molecules • Non-polar molecules have an equal distribution of electrons throughout the entire compound. • The electrons are being pulled in all directions evenly. • All diatomic molecules are symmetrical

This is a SNAP! • Symmetric are • Nonpolar • Asymmetric are • Polar

Let’s Practice!! • For the following covalent bonds, determine there bond type and molecule type (EXPLAIN WHY!!!) • • • H 2 O N 2 NH 3 CO 2 CH 4

HF HCl HBr HI H 2 O H 2 S CO 2 CS 2 Si. O 2 Si. S 2 CH 4 CF 4 CCl 4 CBr 4 CI 4 NH 3 PH 3

Exceptions to the Octet Rule Less than an Octet: • When there are fewer than eight electrons around an atom in a molecule or ion. • This is a rare situation and is most often encountered in compounds of boron and beryllium • Example: BF 3

Exceptions to the Octet Rule More than an Octet: • When there are more than eight electrons in the valence shell of an atom. Much more common than having less than eight. • Examples: PCl 5 (10 valence electrons) ICl 4 - (12 valence electrons)

Other Types of Covalent Bonds Coordinate Covalent Bond: • When one atom donates both of the electrons that are shared • Example: NH 4+ and H 3 O+ • Nitrogen donates a pair of electrons to share with H+ forming a coordinate covalent bond between nitrogen and hydrogen

Other Types of Covalent Bonds Network Solids: • Solids that have covalent bonds between atoms linked in one big network or one big macromolecule with no discrete particles. This gives them some different properties from most covalent compounds. • They are hard, poor conductors of heat and electricity, and have high melting points • Examples include: Diamond (C), silicon carbide (Si. C), and silicon dioxide (Si. O 2)

Metallic Bond • Occurs only in metals (Example Copper) • Metals have low ionization energy meaning they hold onto their valence electrons very loosely • As a result the electrons in metallic substances move about very easily and are not associated with any particular atom • Therefore, the particles of a metal are usually positive ions surrounded by a mobile sea of electrons • The attraction between the positive cations and the moving electrons is what holds the metal together • Properties of Metallic Bonds are that of metals: hard, good conductors of heat & electricity, malleable, ductile, etc. . .

Intermolecular Forces • Forces of attraction between molecules. Include: dipoles, hydrogen bonds, dispersion forces, and molecule-ion attraction. • The difference between intra- (within) and inter- (between). A covalent bond would be an intramolecular force (a bond within a molecule). Dipole-Dipole attractions would be a intermolecular force (bonds found between molecules) • The higher the degree of polarity in the bonds the stronger the intermolecular forces

Hydrogen Bonding • An intermolecular attraction between a hydrogen atom in one molecule to a nitrogen, oxygen, or fluorine atom in another molecule • The strongest intermolecular force • Substances with hydrogen bonds tend to have much higher melting and boiling points than those without hydrogen bonds • Example: The boiling point of H 2 O is much higher than H 2 S

Hydrogen Bond

Dipole-Dipole Attractions • Positive end of a polar molecule is attracted to the negative end of an adjacent polar molecule.

London Dispersion Forces AKA: van der Waals Forces • Weak intermolecular forces between non-polar molecules (like diatomic molecules) • Dispersion forces make it possible for small, non-polar molecules to exist in both liquid or solid phases under conditions of high or low temperatures. • Increases with molecular size, Ex. As you go down group 17, dispersion forces increase and boiling point increases.

Molecule-Ion Attraction • Attraction between the ions of an ionic compound such as Na. Cl, and a molecule such as water (or any other polar covalent compound). • When you put Na. Cl into water, the Na+ from the salt is attracted to the O from the water which is partially (-), and the Cl- from the salt is attracted to the H+ of the water.

Molecular Geometry • Lewis dot structures do not show us what the shape of a molecule is. • Molecular geometry allows us to learn the relationship between the two-dimensional Lewis dot structures that we learned to draw, and the three-dimensional molecular shapes that we are about to learn about. CCl 4

Molecular Shapes • There are five fundamental shapes.

Review Questions 1) Which type of bond results when one or more valence electrons are transferred from one atom to another? 1) a nonpolar covalent bond 2) a polar covalent bond 3) a hydrogen bond 4) an ionic bond 2) Which compound contains ionic bonds? 1) CO 2 2) Ca. O 3) NO 2 4) NO

3) Which molecule contains a triple covalent bond? 1) N 2 2) H 2 3) Cl 2 4) O 2 4) What is the total number of electrons shared in the bonds between the two carbon atoms in a molecule of 1) 6 2) 2 3) 8 4) 3 5) Covalent bonds are formed when electrons are 1) mobile within a metal 2) transferred from one atom to another 3) captured by the nucleus 4) shared between two atoms

6) Which molecule contains a nonpolar covalent bond? 7) Which formula represents a nonpolar molecule? 1) CH 4 2) H 2 S 3) HCl 4) NH 3

8) Which pair of characteristics describes the molecule illustrated below? 1) symmetrical and polar 2) asymmetrical and nonpolar 3) symmetrical and nonpolar 4) asymmetrical and polar 9) Which intermolecular force of attraction accounts for the relatively high boiling point of water? 1) ionic bonding 2) covalent bonding 3) metallic bonding 4) hydrogen bonding

10) In the diagram of an ammonium ion below, why is bond A considered to be coordinate covalent? 1) Nitrogen provides a pair of electrons to be shared with hydrogen. 2) Hydrogen provides a pair of electrons to be shared with nitrogen. 3) Nitrogen transfers a pair of electrons to hydrogen. 4) Hydrogen transfers a pair of electrons to nitrogen. 11) Which structural formula represents a polar molecule?

12) Which diagram best illustrates the ion-molecule attractions that occur when the ions of Na. Cl(s) are added to water? 13) Each molecule listed below is formed by sharing electrons between atoms when the atoms within the molecule are bonded together. • Molecule A: Cl 2 • Molecule B: CCl 4 • Molecule C: NH 3 Explain why CCl 4 is classified as a non-polar molecule.

14) Testing of an unknown solid shows that it has the properties listed below. • • (a) low melting point (b) nearly insoluble in water (c) nonconductor of electricity (d) relatively soft solid State the type of bonding that would be expected in the particles of this substance.