4/14- Chemical Bonding and Structure Questions

List out the name and bond angles of the molecular geometry of molecules with 6 electron domains.

1. Octahedral, no lone pairs, 90 degrees

2. Square pyrimidal, one lone pair, less than 90 degrees.

3. Square planar, two lone pairs, 90 degrees

List out the name and bond angles of the molecular geometry of molecules with 5 electron domains.

1. trigonal bipyrimidal, no lone pairs, 90 degrees, 120 degrees.

2. see saw, one lone pair, less than 90 and 120 degrees.

3. T shaped, two lone pairs, 90 degrees

4. linear, 3 lone pairs, 180 degrees

Explain in detail how covalent bond is formed.

The electrostatic attraction between a shared pair of electrons and a positive nuclei.

Covalent bonds result from the overlap of atomic orbitals. A sigma bond (σ) is formed by the direct head-on/end-to-end overlap of atomic orbitals, resulting in electron density concentrated between the nuclei of the bonding atoms.

A pi bond (π) is formed by the sideways overlap of atomic orbitals, resulting in electron density above and below the plane of the nuclei of the bonding atoms.

State the formula for formal charge, and how it is used.

Formal charge (FC) can be used to decide which Lewis (electron dot) structure is preferred from several. The FC is the charge an atom would have if all atoms in the molecule had the same electronegativity. FC = (Number of valence electrons)-1⁄2(Number of bonding electrons)-(Number of non-bonding electrons). The Lewis (electron dot) structure with the atoms having FC values closest to zero is preferred.

What are some exceptions to the octet rule?

Exceptions to the octet rule include some species having incomplete octets and expanded octets.

Hydrogen: 2 electrons

Boron: 6 electrons

Beryllium: 4 electrons

Sulphur: 6 e domains, 12 electrons

some atoms that can hold 5 electron domains

Explain the wavelength of light required to dissociate oxygen and ozone.

wavelength required to dissociate oxygen: 241nm

-because of double bond, stronger

wavelength required to dissociate ozone: 330nm

-1.5 bond order, weaker

Describe the mechanism of the catalysis of ozone depletion when catalysed by CFCs and NOx.

Check notes!!!

Explain how a hybrid orbital is formed.

A hybrid orbital results from the mixing of different types of atomic orbitals on the same atom.

Identify and explain the relationships between Lewis (electron dot) structures, electron domains, molecular geometries and types of hybridization.

sp = linear = 2 electron domains = 180 degrees

sp2 = trigonal planar = 3 electron domains = 120 degrees

sp3 = tetrahedral, trigonal pyrimidal, bent = 4 electron domains = 109.5 degrees, 107 degrees, 104.5 degrees

List atoms with 6 and 5 bonding domains.

6 bonding domains:

-S

-Xe

-Br

5 bonding domains:

-S

-I

-P

-Cl

Explain the meaning of electron delocalization.

pi/π--π_-electrons shared by more than two atoms/nuclei / a pi/π--π_-bond/overlapping p-orbitals that extends over more than two atoms/nuclei;

Describe the structure and bonding of an ionic compound, and how it is formed.

The ionic bond is due to electrostatic attraction between oppositely charged ions. Under normal conditions, ionic compounds are usually solids with crystal lattice structure.

Ionic compounds are ormed when electrons are transerred rom one atom to another to orm ions with complete outer shells o electrons.

Positive ions (cations) form by metals losing valence electrons.

Negative ions (anions) form by non-metals gaining electrons.

The number of electrons lost or gained is determined by the electron configuration of the atom.

List the properties of an ionic compound, and use the bonding models to explain its properties.

1. Solid at room temperature

2. Have high melting and boiling point, low votality (do not form gas easily): to separate particles into liquid or gas would require high energy to break the strong ionic bonds.

3. Do not conduct electricity as a solid, does conduct electricity when liquid or in an aqueous solution: no freely moving charged particles.

4. brittle, hard: strong ionic compound do not allow the structure to bend or deflect.

5. all dissolves somewhat in water bur solubility varies discriminately

What is the octet rule, and what are the exceptions to it?

The rule states that all atoms should have in total of 8 electrons after covalently bonding with other atoms.

The exceptions to the octet rule include:

1. Hydrogen (2 electrons)

2. Beryllium (4 electrons)

3. Boron (6 electrons)

Describe how the formulas of ionic compounds are written.

The ionic compounds are written based on the charges (2+ and 2-, 3+ and 3-).

Usually the cations come first and then the anion is written.

There are some exceptions to this:

CH3COONa

CH3COOH

The Na and H are written at the end to indicate that they are actually bonded to the O.

Define covalent bonding.

The electrostatic attraction between a shared pair of electrons and positive nuclei.

Skill: Practice drawing the Lewis Dot Diagram!

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Define bond energy (bond enthalpy), and provide some examples (using equations).

The energy required to break one mol of a covalent bond in gaseous molecules.

Cl2 (g) -> 2Cl (g)

I2 (g) -> 2I (g)

Outline the energy required in forming and breaking covalent bonds.

Breaking bonds requires energy, so the reaction is endothermic.

Forming bonds always releases energy so the reaction is exothermic.

Describe the relationship between bond length and bond strength.

Average bond enthalpy decreases as bond length increases.

Although there is considerable variation in the bond lengths and strengths of single bonds in different compounds, double bonds are generally much stronger and shorter than single bonds, because the bond length is shorter. The strongest covalent bonds are shown by triple bonds, with the shortest bond length.

Define resonance bonding.

A double bond could appear in 2 (or more) locations.

When writing the Lewis structures for some molecules it is possible to write more than one correct structure.

Skill: Draw the resonance bonding for Benzene (C6H6), Ozone (O3), Carbonate (CO3 2-), and hydrogen carbonate (HCO3 -).

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State the valence shell electron pair repulsion (VSEPR) theory.

The theory states that pairs of electrons arrange themselves around the central atom so that they are as far apart from

each other as possible. There will be greater repulsion between non-bonded pairs of electrons than between bonded pairs.

List our what counts as an electron domain.

-Lone pair

-Single bond

-Double bond

-Triple bond

List out the different shapes of molecules and its corresponding bond angles.

1. linear, 180 degrees

2. trigonal planar, 120 degrees

3. tetrahedral, 109.5 degrees

4. trigonal pyramidal, 107 degrees

5. bent/ VEE shape, 104.5 degrees

Explain how polarity in covalent bond.

Bond polarity results from the difference in electronegativities of the bonded atoms.

The difference in electronegativity results in the formation of an ionic bond, while a SMALLER electronegativity (0.5-1.7) results in the formation of a polar covalent bond.

One end of the molecule will thus be more electron rich than the other end, resulting in a polar bond. This relatively small difference in charge is represented by + and -. The bigger the difference in electronegativities the more polar the bond. This is called a bond dipole.

Covalent bonds can also be non-polar, where the bonds occur between atoms that have no different in electronegativity (the "gens"), they are also called pure covalent bonds. (0.1-0.4)

List out giant covalent/network covalent structures of carbon and silicon and explain their properties in terms of their structure.

Giant covalent/network covalent structures are covalent substances that don't exist as discrete molecules. Some examples of it include: diamond, silicon, and silicon dioxide.

Diamond:

-made up of purely carbon atoms bonded by strong covalent bond; each carbon atom is covalently bonded to four other carbon atoms to form a giant covalent structure.

as a result, it has a high melting and boiling point.

-because all the electrons are localized it does not conduct electricity.

Silicon:

-each silicon atom is bonded to 4 other silicon atoms in a tetrahedral arrangement (like carbon)

-again, because all electrons are held in fixed position, it is a poor conductor at low temperature.

Silicon Dioxide (an empirical formula!):

-Silicon dioxide has a bent structure caused by the lone pairs of electrons on the oxygen atom

-it is a strong covalent bond, a hard substance with high melting and boiling point and is a poor conductor of electricity.

List out the allotropes of carbon and explain their properties in terms of their structure.

Allotropes are different forms of the same element in the same physical state.

Carbon exists as 3 allotropes: diamond, graphite, graphene, and fullereness.

Graphite:

-each carbon atom has very strong bonds to three other carbon atoms

-this gives layers of hexagonal rings consisting of carbon rings in fused hexagonal rings.

-composed of planar sheets of hexagonally arranged C atoms stacked on top of each other

-layers are held together by relatively low LDF

-because only bonded to 3 other C atoms, each C atom has an electron which becomes delocalized across the plane; so it conducts electricity

-trigonal planar structure; sp2

-soft and slippery because layers are able to slide over one another due to weak IMF

Graphene:

-a single layer of hexagonally arranged carbon atoms, i.e. it is essentially a form of graphite which is just one atom thick; have high tensile strength (thin)

-extremely light, functions as a semiconductor and is 200 times stronger than steel

-trigonal planar

-it is the most chemically reactive carbon allotrope

C60 fullerene:

-each C atom is covalently bonded to three other C atoms

-carbons bonded together in 20 hexagons (6C rings) and 12 pentagon (5C ring); gives a geodesic spherical structure

-presence of delocalized electrons = conducts electricity

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Explain coordinate bonding.

A coordinate covalent bond is where the pair of shared electrons only come from one atom.

The most important thing to note that is when drawn into a Lewis dot diagram, coordinate bond looks exactly like a regular bond.

List and explain the 3 types of intermolecular forces.

London Dispersion Force:

-refers to instantaneous induced dipole forces that exist between any atoms or groups of atoms and should be used for non-polar entities

-made up of both instantaneous dipoles and induced dipoles (caused by changes in electron density and movement of electrons within an atom or molecule)

-instantaneous dipoles induce nearby atoms to make it into an induced dipole

-ALL atoms and molecules have LDF (both polar and non-polar)

-larger molar mass = larger surface area = increase in LDF

Dipole dipole forces:

-occur only between polar molecules (molecules that have a net dipole movement)

-dipole dipole force is the electrostatic attraction between a partial positive charge on one molecule, and a partial negative charge on the other.

-they ALSO have LDF!

-The strength of the dipole-dipole forces depends on the overall polarity of a molecule—the more polar the molecule, the stronger the dipole-dipole force.

-lower electronegativity = partial positive charge

-higher electronegativity = partial negative charge

Hydrogen Bonding:

-occurs between molecules that have an electronegative nitrogen, oxygen, or fluorine atom directly bonded to a hydrogen atom

-The strongly electronegative atom pulls the (only) electron away from the hydrogen nucleus, "exposing" the nucleus.

-this exposed hydrogen nucleus is in turn strongly attracted to the lone pairs on nearby molecules.

Define a metallic bond and explain their properties and strength.

A metallic bond is the electrostatic attraction between a lattice of positive ions and delocalized electrons.

The positive ions are arranged into a lattice structure, and there are valence electrons that create a sea of delocalized electrons which are not bound to one atom but can float around and are attracted to any nearby nucleus.

This creates non-directional bonds between metal atoms that explains why metals are malleable. (Since there are no strong attractions "planes" of atoms can slide over each other easily.); good conductors of electricity

Their properties include:

-good conductors

-ductile (can be made into wires)

-malleable (can be bent into shape)

-shiny when polished

The strength of metallic bond depends on the charge of the ions and the radius of metal ion.

greater ionic charge (more delocalized electrons) = smaller ionic radius = greater strength = hight melting and boiling points

So towards the left of the periodic table, the melting and boiling point increases.

Moving down the periodic table, the melting and boiling point decreases as it increases in atomic size.

Describe what alloys are and explain their uses.

Alloys usually contain more than one metal (homogenous mixtures composed of 2 or more metals or a metal and a non-metal) and have enhanced properties.

Alloys can be created with the addition of carbon or phosphorus (non-metals) in some cases.

**They are often stronger and have more resistance to corrosion.

The addition of different sized atoms in the alloy means that they layers cannot slide over each other as easily (originally due to non-directional bonding. This results in the alloy being harder than its component metals.

Examples!:

Brass is an alloy of copper and zinc.

Steel is an alloy of iron with carbon and some other elements such as nickel, tungsten, or molybdenum.

Different grades and types of steel are created by changing the amount and type of alloying agents.

Explain why hydrogen bonding is important.

Hydrogen bonding is important because it makes water makes life possible.

For e.g., proteins are chains of amino acids that contain both NH2 and COOH groups, and thus have hydrogen bonding.

DNA consists of a double helix structure that is held together by hydrogen bonds. Without the hydrogen bonds the DNA would unravel; but if the helix structures were held together with covalent bonds they would be too strong to pull apart and replicate.

Water is the most dense at 4°C, and expands as it becomes a solid. This is because the water molecules hydrogen-bond to each other in a six-sided shape (ref. diagrams on p. 112). The end result is that, unlike most substances, solid water (ice) floats on top of liquid water. (Which is a good thing for all the fish.)

high latent heat of vaporization

high specific heat capacity

Define the term average bond enthalpy.

Energy required to break one mol of covalent bond in gaseous state averaged over similar compounds.

State the formula for the following polyatomic ions:

Sulphate

Ammonium

Hydroxide

Carbonate

Hydrogen carbonate/bicarbonate

Nitrate

Phosphate

Acetate/ethanoate

Sulphate:

SO4 2-

Ammonium:

NH4+

Hydroxide:

OH-

Carbonate:

CO3 2-

Hydrogen carbonate/bicarbonate:

HCO3-

Nitrate:

NO3-

Phosphate:

PO4 3-

Acetate/ethanoate:

CH3COO-

Define dipole dipole forces.

dipole dipole force is the electrostatic attraction between a partial positive charge on one molecule, and a partial negative charge on the other.

Explain why sodium conducts electricity but phosphorus does not.

Na: delocalized electrons / mobile sea of electrons / sea of electrons free to move;

No mark for just "mobile electrons".

Draw a structure of BrO3- that does not follow the octet rule.

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State the level of ozone in the atmosphere.

abour 10 ppm