IB Chemistry HL

Combination of all Syndey Wong quizlet sets

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

ionic bonding

the electrostatic attraction between two oppositely charged ions

cation

A positively charged ion

anion

A negatively charged ion

multivalent

describing the ability of an element to form ions in more than one way, depending on the chemical reaction it undergoes

votality

tendency to vaporize

covalent bonding

the electrostatic attraction between a positive nuclei and a shared pair of electrons

bond energy

energy required to break one mole of a covalent bond in gaseous molecules

delocalized electrons

electrons that are free to move in metals; they come from resonance bonding

resonance bonding

a double bond could appear in 2 (or more) locations

coordinate covalent bond

where a pair of shared electrons only come from one atom

bond order

the number of bonds between pairs of electrons

expanded octets

Occur for non metals in rows 4-7 where the central atoms can hold more than 8 electrons b/c the extra electrons expand/fill into the d-orbital

pure covalent bonds

bond that occurs between atoms that have no difference in electronegativity

polar covalent bonds

bond that occurs between atoms that have a different in electronegativity between 0.5-1.7 units

spectrum of bonding

a gradual shift from non-polar covalent bond to polar ones, and then to ionic

net dipole moment

a measure of its overall polarity, the sum of all the bond dipoles in a molecule

intramolecular bonds

bonds within molecules

intermolecular forces

forces of attraction between molecules

London dispersion forces

the intermolecular attraction resulting from the uneven distribution of electrons and the creation of temporary dipoles

instantaneous dipole

temporary dipole that occurs for a brief moment in time when the electrons of an atom or molecule are distributed asymmetrically

induced dipole

a dipole temporarily created in an otherwise nonpolar molecule, induced by a neighboring charge

Dipole-dipole forces

intermolecular forces that exist between polar molecules; the strengths this intermolecular attractions increase when polarity increases

Hydrogen bonding

strong type of intermolecular dipole-dipole attraction, occurs between hydrogen and F, O or N

giant covalent structure

covalent substances that don't exist as discrete molecules

allotropes

two or more different molecular forms of the same element in the same physical state

geodesic dome

a spherical structure formed by 20 hexagons and 12 pentagon shapes

alloy

a combination; a mixture of two or more metals/non-metals

ethanoic acid (acetic acid)

CH3COOH / C₂H₄O₂

hydrochloric acid

HCl

calcium carbonate

CaCO3

molecular elements(diatomic molecules)

Have No Fear Of Ice Cold Beer!

Hydrogen, Nitrogen, Fluorine, Oxygen, Iodine, Chlorine and Bromine

hydrogen(H2), nitrogen (N2), oxygen (O2), fluorine (F2), chlorine (Cl2), bromine (br2), iodine(I2)

molecular elements

elements with diatomic molecules (two atoms of that element bonded together)

molecular compound

compounds formed from two or more nonmetals

hydroxide

OH-

phosphate

PO4 3-

nitrate

NO3 1-

sulfate

SO4 2-

carbonate

CO3 2-

ammonium

NH4 1+

acetate

C2H3O2 1-

aqueous

dissolved in water

sublimation

a change directly from the solid to the gaseous state without becoming liquid

homogeneous mixture

a mixture in which substances are evenly distributed throughout the mixture, it is a solution

heterogeneous mixture

a mixture in which different materials can be distinguished easily, particles disperse with the substance

solute

A substance that is dissolved in a solution.

solvent

A liquid substance capable of dissolving other substances (e.g. water)

empirical formula

a chemical formula showing the ratio of elements in a compound rather than the total number of atoms

molecular formula

a chemical formula of a molecular compound that shows the kinds and numbers of atoms present in a molecule of a compound

meniscus

The curved upper surface of a liquid in a tube

precipitate

a solid that forms and settles out of a liquid mixture

formula unit

the lowest whole-number ratio of ions in an ionic compound

relative molecular mass

the weighted mean mass of a molecule compared with one-twelfth of the mass of an atom of carbon-12.

atomic mass unit

one twelfth the mass of a carbon-12 atom

relationship of protons, electrons, and neutrons in an atom

the number of protons equal the number of electrons

isotope

an atom with the same number of protons and a different number of neutrons from other atoms of the same element.

electrolyte

An ionic compound whose aqueous solution conducts an electric current, highly soluble in water

oxide

a binary compound of oxygen with another element or group.

dichromate

Cr2O7 2-

miscible

liquids that dissolve freely in one another in any proportion (in any concentration)

solubility

a measure of how much solute can dissolve in a given solvent at a given temperature and specific solvent

immiscible

liquids that are not soluble in each other