Unit 2: Molecules

Anhydride

A compound derived from another compound by the removal of water; it will combine with water to form an acid (acidic anhydride) or base (basic anhydride).

Acetate

C2H3O2-

Valence Bond Theory & Hybridization of Atomic Orbitals summary

Covalent bonds form by overlap of atomic orbits (s,p,d...) and hybrid orbitals (sp, sp^2, sp^3...). Forms sigma and pi bonds. The degree of overlap determines the strength of the bond.

Mo Theory drawbacks

Easily used for diatomic molecules of the first and second period atoms; becomes more complex as the size of the molecules increases.

H-bonding

H bound to F, N, or O

(h) SO3(g) + H2O(l) (i) NaHSO4(aq) + NaOH(aq)

H2SO4(aq). The production of sulfuric acid (or the occurrence of acid rain). Na2SO4(aq) + H2O(l). An acid-base neutralization reaction.

Phosphorus acid

H3PO3

(b) Increasing acidity: ammonia, water, hydrogen fluoride

In these binary acids (NH3, H2O, HF), the acidic proton is bonded directly to a nonmetal atom X (N, O, or F). The acid strength is related to the ease of breaking this H-X bond.Across a period, the dominant factor is the polarity of H-X bond, which is proportional to the difference in electronegativity. Thus, the ranking is: NH3 < H2O < HF (strongest acid).

Strength of Ionic bonding

Ionic radius The greater the size of the ion, the larger the distance between the two charges, so the smaller the force Ionic charge Larger the charge, the stronger the bond

Rank in the order specified, and explain your reasoning with the aid of Lewis structures: (a) Increasing basicity: hypochlorite ion, chlorite ion, chlorate ion, perchlorate ion

Lewis structures for hypochlorite (ClO-), chlorite (ClO2-), chlorate (ClO3-), and perchlorate (ClO4-) ions are shown below. In the corresponding conjugate acids (HClO, HClO2, HClO3, and HClO4), a greater number of terminal O atoms (which are highly electronegative) tends to cause more electron density to be withdrawn away from the O-H bond, polarizing it and thus making it easier to break this bond heterolytically into H+ and the conjugate base. Thus, HClO is the weakest acid and HClO4 is the strongest acid. Correspondingly, ClO- is the strongest base and ClO4- is the weakest base, or in order of increasing base strength: ClO4- < ClO3- < ClO2- < ClO- (strongest base).

True/False: Lewis Structures with nonzero formal charges are incorrect.

Lewis structures with formal charges are not wrong. Formal charges are used to pick which lewis structure i the best for a neutral molecule. The sum of the lewis structures in a molecule must equal zero. When there are more than one lewis structures possible, the formal charge helps us understand the arrangement of the atoms.

(a) Which gives a more basic solution in H2O: (i) CaO or SO3; (b) BeO or BaO; (c) P4O10 or K2O ?

More basic solution: (i) CaO vs. SO3; (ii) BeO vs. BaO; (iii) P4O10 vs. K2O. Remember that metal oxides are basic anhydrides, whereas nonmetal oxides are acid anhydrides.

(b) Which has more covalent character in its bonds: (i) LiCl or KCl; (ii) AlCl3 or PCl3; (iii) NCl3 or AsCl3 ?

More covalent character: (i) LiCl vs. KCl; (ii) AlCl3 vs. PCl3; (iii) NCl3 vs. AsCl3. Compare the difference in electronegativity (between element E and Cl) and remember the trends in EN.

(g) Br2(l) + 2 Cl-(aq)

NR. Note that Cl2 is a stronger oxidizing agent than Br2.

Which of the following species are amphiprotic (or amphoteric) in aqueous solution? For such a species, -+---write equations showing it acting as an acid and as a base. OH , NH4 , H2O, HS , NO2 , HCO3 , HBr.

Of the listed species, H2O, HS-, and HCO3- are commonly considered to be amphiprotic (or amphoteric). Reactions with water are shown below:H2O acting as an acid: H2O(l) + H2O(l) = OH-(aq) + H3O+(aq)H2O acting as a base: H2O(l) + H2O(l) = H3O+(aq) + OH-(aq) HS- acting as an acid: HS-(aq) + H2O(l) = S2-(aq) + H3O+(aq)HS- acting as a base: HS-(aq) + H2O(l) = H2S(aq) + OH-(aq)HCO3- acting as an acid: HCO3-(aq) + H2O(l) = CO32-(aq) + H3O+(aq) HCO3- acting as a base: HCO3-(aq) + H2O(l) = H2CO3(aq) + OH-(aq)

Lead

Pb

VSPER Theory uses

Predicts molecular shapes. Polarity of bonds can be used to determine the polarity of the molecule.

MO Theory uses

Predicts the arrangements of electrons and their energy levels in molecules. Explains paramagnetism/diamagnetism and makes it possible to predict the properties of hypothetical molecules and ions.

Antimony

Sb

HCOOH is an example of a carboxylic acid, a class of weak organic acids that you can recognize by the -COOH group, with the rest of the molecule symbolized by "R".

Since it is a weak acid, so the solution is weakly acidic.

Lewis Theory uses

Structures show relative bond strengths as single, double, and triple bonds.

BrF3

T-shaped sp3d1 nonpolar

VSEPR Theory summary

The number of electron domains around the central atom in Lewis structure determines the three-dimensional shape of the molecule or polyatomic ion. Considers bonds as localized between one pair of atoms.

NH4I is a salt containing a cation originating from a weak base (NH3) and an anion originating from a strong acid (HI).

Thus, NH4+ is a weak acid, whereas I- has no acid-base properties. The overall solution is weakly acidic.

NaNO2 is a salt containing a cation originating from a strong base (NaOH) and an anion originating from a weak acid (HNO2).

Thus, Na+ has no acid-base properties, whereas NO2- is a weak base. The overall solution is weakly basic.

Lewis Theory drawbacks

Two dimensional model of three dimensional molecules; no distinction between bonds in different compounds. Does not explain why or how bonds form.

dipole moment

a property of a molecule whose charge distribution can be represented by a center of positive charge and a center of negative charge

Amphoteric

a substance that can act as both an acid and a base

CIO2

bent sp3 polar

ionic bonding

form an extended structure, hard and rigid, brittle

group 1 metal reacting with water to produce

h2 and metal hydroxide

HBr is a strong acid. The solution is

highly acidic

BeCl2

linear sp nonpolar

SCN-

linear sp polar

covalent bonding

sharing of e-, form a molecule, strong intramolecular forces, but weak intermolecular forces between molecules.

RbOH is a strong base. The solution is

strongly basic

metallic bonding

the chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons, cations can move without disrupting their attraction to the e-. Mobile, delocalized e-. High conductivity, because the delocalized e- can disperse heat more easily.

lattice energy

the energy required to separate one mole of the ions of an ionic compound, which is directly related to the size of the ions bonded and is also affected by the charge of the ions

atomic number

the number of protons in the nucleus of an atom

Melting points of ionic solids

the smaller ions can pack closer together, resulting in stronger bonds, and the MP higher. Vice versa.

mass number

the sum of the number of neutrons and protons in an atomic nucleus

CIO3-

trigonal bypyramidal, sp3 polar

CO3 2-

trigonal planar sp2 nonpolar

explain why magnesium metal is deformed by an applied force, but magnesium fluoride is shattered.

The magnesium metal has a sea of electrons which allow the e- to move around and make the metal malleable, this is because it has metallic bonding. Magnesium fluoride has ionic bonding, which makes the positive and negatively charged ions very attracted to each other and difficult to separate. When stress is applied, like-charged particles are pushed closer, causing the structure to break down. (electrostatic repulsions)

Write balanced equations. (If there is no reaction, write "NR") (a) H2(g) + 2 Na(s) = (b) H2(g) + Cl2(g)= (c) BeO(s) + H2O(l)

2 NaH(s). Here, H2 is reduced and acts as an oxidizing agent. Here, H2 is oxidized and acts as a reducing agent. NR. Unlike the other alkaline-earth oxides, BeO is not a basic anhydride (in fact, it is amphoteric) and does not react with water.

Triatomic Molecules

All triatomic molecules do not have a planar shape. They can be bent, v shaped, linear.

Lewis Theory summary

Atoms in molecules and polyatomic ions have completed valence shells following the octet rule.

Isotope

Atoms of the same element that have different numbers of neutrons

CaO(s) + H2O(l) --> (e) Br2(l) + 2 Rb(s)= (f) Br2(l) + 2 I-(aq)

Ca(OH)2(aq).Think of this as a reaction of "O2-", a very strong base, with water. 2 RbBr(s). A simple reaction between an alkali metal and a halogen. I2(aq) + 2 Br-(aq). A halogen displacement reaction.

VSPER Theory drawbacks

Does not explain how/why bonds form.

electrostatic repulsion

Describes a force between particles with opposite electrical charges that causes them to push apart from one another

Valence Bond Theory & Hybridization of Atomic Orbitals uses

Description of pi bonds helps to explain greater the reactivity of compounds with double bonds and lack of rotation around double bonds. Hybrid atomic orbitals explain some molecular geometries, but not all.

Pyrophoric

Descriptive of any substance that ignites spontaneously when exposed to air.

Allotropes

Different forms of the same element

Valence Bond Theory & Hybridization of Atomic Orbitals

Does not explain certain properties such as paramagnetism. Not easily used to predict properties of hypothetical molecules or ions. Needs multiple structures to describe resonance.

Briefly explain these observations exemplifying the anomalous behaviour of second-period elements. (a) Oxygen forms the compound OF2, but sulfur can form the hexafluoride SF6. (b) The bond energy in N2 is 946 kJ mol , but in P2 (stable only above 800 °C) it is only 488 kJ mol . (c) Lithium reacts with excess oxygen to give Li2O, but the analogous reaction for sodium gives Na2O2. (d) Molten BeCl2 is a poor electrical conductor, whereas molten CaCl2 is an excellent one.

Second- vs. third-period elements.The second-period element O can only form a maximum of four covalent bonds whereas the third-period element S can expand its octet through use of low-lying empty d-orbitals. Consistent with this expectation, we find that the oxygen atom is sp3-hybridized in OF2, whereas the sulfur atom is sp3d2- hybridized in SF6 Multiple bonding in these triply-bonded diatomic molecules involves p-overlap of two adjacent p orbitals. The 2p-2p p-overlaps are good in N2 but the 3p-3p p-overlaps are poorer in P2. (Note that this is not the stable form of elemental phosphorus at room temperature.) In these alkali-metal oxides, the smaller size of the second period element explains why a small Li+ ion needs to match with a small O2- ion to form the normal lithium oxide Li2O, whereas a larger Na+ ion matches better in size with a larger O2 2- (peroxide) ion to form sodium peroxide Na2O2. Beryllium is unusual among the alkaline-earth elements in that all its compounds are covalent, not ionic. Thus, BeCl2 is a molecule containing covalent Be-Cl bonds; it does not consist of Be2+ and Cl- ions. In contrast, CaCl2 is an ionic solid containing Ca2+ and Cl- ions; when CaCl2 is molten, these ions become mobile and can conduct electricity.

Classify the following as Arrhenius, Brønsted-Lowry, or Lewis acid-base reactions; identify which is the acid and which is the base. A reaction may fit all, two, one, or none of these categories. (a) Ag+ + 2 NH3 = [Ag(NH3)2]+. (b) H2SO4 + NH3 ® HSO4- + NH4+. c) 2 HCl ® H2 + Cl2. (d) AlCl3 + Cl- ® AlCl4-.

The Arrhenius definition does not apply here because the reaction does not involve H+(aq) or OH-(aq). The Brønsted-Lowry definition does not apply here because there is no exchange of protons. The Lewis definition applies here, with the Lewis acid Ag+ accepting an electron pair from the Lewis base NH3. The Arrhenius definition does not apply here because no OH- (aq) is involved. The Brønsted-Lowry definition applies here, with the acid H2SO4 donating a proton to the base NH3. The Lewis definition also automatically applies here, with the base NH3 donating an electron pair to H2SO4 (and thereby forming an N-H bond, while an H-O bond is broken). This is not an acid-base reaction. For the same reasons as in (a), the Arrhenius and Brønsted-Lowry definitions do not apply here. However, the Lewis definition does apply, with the Lewis acid AlCl3 accepting an electron pair from the Lewis base Cl- (forming an Al-Cl bond).

(c) Identify the fluorides formed with each of the chemically active Period 2 elements, draw their Lewis structures, and comment on the trend in covalent vs. ionic bonding in this series.

The fluorides of the period 2 elements are: LiF, BeF2, BF3, CF4, NF3, OF2, and F2 itself. The trend is towards decreasing ionic and increasing covalent character in the element-fluorine bonds on proceeding from LiF to F2, as the difference in electronegativity becomes reduced. Lewis structures are shown below.

KNO3 is a salt containing an cation originating from a strong base (KOH) and an anion originating from a strong acid (HNO3).

Thus, neither K+ nor NO3- has acid-base properties. The overall solution is neutral.

Molecular Orbital Theory (MO)

atomic orbitals combine to form bonding and anti bonding molecular orbitals (∂, ∂*, π, π*) that extend over the entire molecule (electrons delocalized). Describes bonding and antibusing interactions based on which orbitals are filled.

metallic oxides

basic oxides

when liquid benzene (C6H6) boils, the gas consists of:

molecules, because of the intermolecular forces between the molecules are separated, but the dispersion forces are not strong enough to separate the atoms bonded within the molecule (intramolecular forces). The bond is extremely strong, as both C and H have an EN value difference of less than 0.5.

formula for percent abundance

multiply 1 by x, plus the other by 1-x, make it equal to the avg from the periodic table and solve.

Nonmetal oxides react with water to produce

oxyacids

Sea of e-

prevents repulsions among the cations.

XeOF4

square pyramid sp3d2 polar

percent ionic character

the percentage of a bond's measured dipole moment compared to what it would be if the electrons were completely transferred