PACKET #7:Chemical Bonding

Reference Table: PT & Table S www.regentsprep.org

Chemical Bonding

• Chemical Bond: an attraction between the protons of one atom 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.

• Bonds can be classified as being either polar or non-polar.

• Polarity: tendency of a molecule, or compound, to be attracted or repelled by electrical charges because of an asymmetrical arrangement of atoms around the nucleus.

• 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.

Electronegativity DifferenceRemember . . .• Electronegativity: an atom’s

attraction for electrons in a bond.• The higher the EN, the more the atom

attracts electrons.• The lower the EN, the less the weaker

the attraction for electrons.• REFER TO TABLE S: Note which

elements have higher EN (TRENDS).

Ionic Bond

• Attraction between oppositely charged ions

• Occurs when electrons are transferred from one ion (charged particle) to another

• Electronegativity difference 1.7+ • Metals react with Nonmetals to form

ionic compounds• Always

Polar !!!

Forming Ionic Bonds

• Ionic bonds are formed when valence electrons are transferred from metals to non-metals forming ionic compounds. 

• Metals lose electrons and become cations (+).• Nonmetals gain electrons and become anions (-)

• Example: KCl:

Example: CaBr2

• An ionic bond was formed by the TRANSFER OF ELECTRONS!! (NO SHARING)!!!

• Draw the following ionic compounds: SrF2, LiI, BaCl2, Na2O, AlCl3, and Al2O3 

Polyatomic Ions• Ionic compounds containing polyatomic ions

can have both ionic and covalent bonding. • Example: KNO3

• Notice that NO3- is composed of 2 nonmetals

therefore the bonding is covalent between N and O but the bonding between K+ and NO3

-  is ionic.

Example: (NH4)3PO4

• The bonding is covalent between N & H, as well as between P & O, but ionic between the two individual polyatomic ions.

• Draw the Lewis structure of the following ionic compounds NaOH, Mg(NO3)2 and (NH4)2CO3

Possible Combinations for Ionic Compounds

Formula

(+) ( - ) Examples

M NM NaCl, KI, and CaF2

M P LiNO3 and Sr(CN)2

P NM NH4Br, Hg2S, and (H3O)2Se

P P NH4NO3 and H3OMnO4

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 Cl2 or H2O]

• Neither atom pulls strongly enough to remove an electron from the other

• The EN difference is < 1.7 • Unpaired electrons pair up in such a way that

the atoms complete their outer shells• Covalent compoundsalso referred to as molecularcompounds

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 covalent0.5 - 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:

• H2, N2, O2, F2, Cl2, Br2, I2

• 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 CH4

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 NH3

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 CCl4

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 – O2 & N2

DOUBLE BOND TRIPLE BOND

MORE . . .

• Some other covalent compounds that are important to know how to draw are:

H2, F2, Cl2, Br2, I2, O2, N2, HCl, HF, HBr, HI, CO2, CF4, CBr4, H2O, H2S,

CCI4. SiH4

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 thelower EN is partially (+)

Polar & Non-Polar Bonds vs. Polar & Non-Polar Molecules

• Just to add a little more confusion . . .• RECALL:

Polar Bond = different ENNon-Polar Bond = same (similar) EN

BUT

Polar Molecule = AsymmetricalNon-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.

Polar Molecules

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 diatomicmolecules 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!!!)

• H2O

• N2

• NH3

• CO2

• CH4

HF H2O CO2 CH4 NH3

HCl H2S CS2 CF4 PH3

HBr SiO2 CCl4HI SiS2 CBr4

CI4

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: BF3

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: PCl5 (10 valence electrons)

ICl4- (12 valence electrons)

Other Types of Covalent Bonds

Coordinate Covalent Bond:• When one atom donates both of the

electrons that are shared

• Example: NH4+ and H3O+

• 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 (SiC), and silicon dioxide (SiO2)

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 intra-molecular force (a bond within a molecule). Dipole-Dipole attractions would be a inter-molecular force (bonds found between molecules)

• The higher the degree of polarity in the bonds the stronger the intermolecular forces

Dipole-Dipole Attractions

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

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 H2O is much higher than H2S

Hydrogen Bond

London Dispersion ForcesAKA: 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 NaCl, and a molecule such as water (or any other polar covalent compound).

• When you put NaCl 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.

CCl4

Molecular Shapes

• There are five fundamental shapes.

Review Questions1) Which type of bond results when one

or more valence electrons are transferred from one atom to another? 1) a nonpolar covalent bond2) a polar covalent bond3) a hydrogen bond4) an ionic bond

2) Which compound contains ionic bonds? 1) CO2 2) CaO 3) NO2 4) NO

3) Which molecule contains a triple covalent bond? 1) N2 2) H2 3) Cl2 4) O2

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 metal2) transferred from one atom to another3) captured by the nucleus4) shared between two atoms

6) Which molecule contains a nonpolar covalent bond?

7) Which formula represents a nonpolar molecule? 1) CH4 2) H2S 3) HCl 4) NH3

8) Which pair of characteristics describes the molecule illustrated below?

1) symmetrical and polar2) asymmetrical and nonpolar3) symmetrical and nonpolar4) 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 NaCl(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: Cl2• Molecule B: CCl4• Molecule C: NH3

Explain why CCl4 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.