Class 2: Bonding, Intermolecular Forces, Thermodynamics

Standard State

- A substance has a standard heat of formation of 0. This is best explained by the fact that the substance is an element in its standard state. - For elements in their standard state, ΔHf° is equal to zero by definition. - The standard states of elements are the forms that they adopt at a temperature of 25°C and pressure of 1 atmosphere (1 atm)

Co-60m is an unstable isotope that releases energy in the form of gamma radiation. What is the most likely electron configuration for Co-60m? A. [Ar] 3d9 B. [Ar] 4s2 3d7 C. [Ar] 4s1 3d8 D. [Ar] 4s2 3d6 4p

- B. [Ar] 4s2 3d7 - Gamma radiation consists of very high-energy photons released from a nucleus in an excited state. It is not the result of excited electrons returning back to lower energy states, although the answer choices lead you to believe this may be the case. Since the metastable state of Co-60m is due to nucleon excitation and not electron excitation, the electron configuration is not expected to be any different from the electron ground state.

CF2Cl2 is a common freon used in refrigerators. The strongest intermolecular forces holding these molecules together are: A. ionic forces. B. hydrogen bonding. C. London dispersion forces. D. dipole-dipole forces.

- CF2Cl2 is not an ion, so it cannot experience ionic forces. - In order to experience hydrogen bonding, the molecule must have at least one hydrogen atom bonded to a highly electronegative element (F, O, N), which CF2Cl2 does not. - While all molecules have some degree of London dispersion forces, CF2Cl2 is a slightly polar molecule since fluorine is more electronegative than chlorine. Therefore, dipole forces is the best answer since they are stronger than London dispersion forces.

Given the same number of moles, which of the following solids will have the lowest melting point? A. SiO2 B. Cu C. NaI D. H2Se

- Hydrogen selenide is a molecular solid. Its intermolecular forces consist of dipole-dipole interactions. Though a relatively strong intermolecular force, they are much weaker than the covalent bonds between the atoms in a network solid (choice A), the covalent bonds between the atoms in a metallic solid (choice B), and the ionic forces between the ions in an ionic solid (choice C). Thus, the weaker intermolecular forces will allow hydrogen selenide to melt more easily than the others. A. SiO2 = network solid (covalent bonds) B. Cu = metallic solid (covalent bonds) C. NaI = ionic solid (ionic bonds) D. H2Se = molecular solid (dipole-dipole) *could watch khan video

A reaction occurs that results in a set of products with more stable bonds and more orderly arrangement than were present in the reactants. Which of the following is true of this reaction? A. The enthalpy and entropy changes are negative. B. The enthalpy change is positive, and the entropy change is negative. Your Answer C. The enthalpy change is negative, and the entropy change is positive. D. The enthalpy and entropy changes are positive.

- In order for the products to be more stable, they must have lost energy; thus, ΔH is negative (=exothermic) - If the products have a more orderly arrangement than the reactants, then the entropy decreased; thus ΔS is also negative, and the answer is "the enthalpy and entropy changes are negative."

what is the enthalpy of the combustion of methane?

- In the combustion of methane (CH4 + 2 O2 → CO2 + 2 H2O) four H-C bonds and two O=O bonds are broken while two C=O bonds and four H-O bonds are formed. To find the enthalpy of the overall reaction, we must use the following equation: H = Hbondsbroken - Hbondsformed

Ion-Dipole Forces

- Ion-dipole forces are produced between ions and polar molecules (i.e. ionic molecule + polar solvent) (ex: NaCl - H2O) the larger the ionic charge and the larger the dipole = the larger the force - a more highly charged ion generates a stronger force - a more polar molecule generates a stronger force ex: Na+ Ca2+ = larger charge and creates stronger ion forces bonded to H2O (negative O2 is attracted to positive Ca+) * full charges (cations and anion) are STRONGER than H-bonding (partial charges)

NP and P bonding

- NP = equal sharing - P = unequal sharing - electron density is always higher around the more EN element

Polarity (tell what kinds of bonds are polar)

- Polarity is determined by the difference in electronegativity between the atoms sharing electrons in a bond. - Electronegativity generally increases from the bottom left to the top right corners of the periodic table. - Elements that are farther from each other on the periodic table have larger differences in electronegativity, while elements that are closer to each other have smaller differences. ex: - Even though fluorine is the most electronegative element, a F—F bond is nonpolar because there is no difference in electronegativity between the two atoms - Since the C—Si and N—O pairs of elements are right next to each other on the table, they are only slightly polar - while hydrogen and chlorine are far from each other, yielding the most uneven sharing of electrons in the bond. (H - Cl is a polar bond)

Which of the following compounds displays the lowest vapor pressure at 15˚C? RANK BASED ON MW A. NH3 B. CO2 C. SO2 D. HF

- The compound with the least vapor pressure has the strongest intermolecular forces. lower vapor pressure = lower rate of evaporation = strong IMFs (vapor pressure always increases as the temperature increases) - Both NH3 and HF can form hydrogen bonds, the strongest type of intermolecular force, so one of these compounds is the most likely correct answer. - The polarity of the H-F bond is much greater than that of an H-N bond and HF has a greater molecular weight, both of which result in HF having the lowest vapor pressure of the compounds listed. - Carbon dioxide is nonpolar, so it would only have weak London dispersion forces making it a gas at 15°C (CO2 is wrong). - Sulfur dioxide is a polar compound and would experience both dipole-dipole interactions and moderate London dispersion forces as it has a relatively large molecular weight. These forces are weaker than the hydrogen bonds in HF, so SO2 is wrong. While this question asks for the lowest vapor pressure rather than to rank all four compounds based on this property, it's good to note that the boiling point of SO2 is higher than the boiling point of NH3 due to the larger molecular weight of SO2.

Which of the following will release the most energy with the addition of an electron? A. Al B. Ga C. Si D. G

- The energy released when an atom gains an electron in its valence shell is its electron affinity. - The electron affinity periodic trend becomes more negative as you move from the bottom left to the top right of the periodic table. - As the atomic radius of an atom decreases, an electron added to an atom has farther to fall to reach the valence shell closer to the nucleus, thereby giving off more energy. - Since silicon is the smallest atom of the options given, it has the most negative electron affinity, indicating the greatest gain in stability.

Kinetics vs. Thermodynamics`

- Thermodynamics predicts the spontaneity (and the equilibrium) of reactions, not their rates - if you had a starting line and a finish line, thermo tells you how far you will go, while kinetics tells you how quickly you will get there

Is combustion spontaneous?

- Yes, exothermic so -H, +S = -G - Yes, combustion is spontaneous at all temperatures G = H - TS

Enthalpy book def

- a measure of the heat energy that is released or absorbed when bonds are broken and formed during a reaction that's run at constant pressure - bond is formed = energy is released, H<0 = G>L>S - energy must be put into a bond to break it, H>0= S>L>G

The second law of Thermodynamics

- all processes proceed toward maximum disorder or entropy - increase entropy (ex: dissolving a solid in a liquid tremendously increases the disorder of the two components as they mix, so ΔS > 0.) (ex: CO2(s) -> CO2(g) has a positive ∆Srxn - A positive ΔSrxn means that the entropy of the products is greater than that of the reactants, i.e., disorder increases)

Dipole-Induced Dipole Forces

- are produced between polar and nonpolar molecules "induced" = NP (must count e-; more e- = easier to destroy e- cloud & induce dipole) + P (creates stronger force) the force is generated when a polar molecule attracts electron density in a nonpolar molecule - a more polar molecule generates a stronger force (think FONClBrISCH) - a molecule with more electrons and larger size generates a stronger force (more polarizable) ex: I2 vs F2 I2 has (53e-)2, F2 has (9e-)2. So I2 has a huge e- cloud which makes it easier to make polar and induce a dipole moment *dipole-induced dipole forces are very easily cleaved -weak bonds due to instantaneous moment in time. A molecule with a dipole (like H2O) disrupts e- cloud of a NP molecule and shifts it (e- repel positive dipole) for a brief moment

Network solids

- atoms are connected in a lattice of covalent bonds, meaning that all interactions between atoms are covalent bonds - only intramolecular forces, making them very strong and tend to be hard solids at RT (diamonds and quartz) if the EN = 0 or VERY SMALL, the bond is NONPOLAR and e- are shared equally = covalent

Intermolecular Forces (IMFs)

- between 2 molecules - attractive - intermolecular attractions are produced when particles of opposite charge attract each other (similar to an ionic bond) the strength of an IMF is dependent on: - larger charges = stronger attractive force - particle size is negligible due to large distances * trying to break IMFs = going through a phase change (S - L - G); means you are looking at physical traits/properties NOT a chemical reaction - stronger IMF = particles held together more tightly - pulling apart the particles is always ENDOTHERMIC

Metallic solids

- covalently bound lattice of nuclei and their inner shell electrons, surrounded by a sea of electrons - the freely roaming valence electrons are called conduction electrons - metals = excellent conductors of electricity and heat and are malleable and ductile - most metals are solid at RT

Is the dephosphorylation of ATP spontaneous?

- endothermic process because breaking bonds, +H = requires energy - increasing entropy, +S = increasing randomness - G at high temperatures +G at low temperatures - so yes spontaneous, but only at high temperatures

bond dissociation energy

- energy required to break a bond homolytically - homolytic bond cleavage, one electron of the bond being broken goes to each fragment of the molecule = two radicals will form (Not the same thing as heterolytic bond cleavage - aka - dissociation - where electrons of the electron pair that make up the bond end up on the same atom; this forms both a cation and an anion) - the higher the bond order (single = 1, double = 2, triple =3), the shorter and stronger the bond r, measures bond length in Angstroms where 1 = 10x10^-10m

Enthalpy (H)

- energy stored within chemical bonds or any attractive force - the change in enthalpy of a reaction is the difference between what is stored in the reactants vs. the products if reactants stored E is greater than products -- > products = exothermic, H<0 if reactants -- > products stored E is greater than reactants = endothermic, H>0 (BREAKING BONDS IS ALWAYS ENDOTHERMIC, must put in energy to break the bonds)

nonpolar bonds

- equal sharing of electrons; no difference in electronegativity - symmetrical; bond dipoles are symmetrically oriented in a molecule - EN = 0

Is freezing water spontaneous? (L>S) (G>L>S = endo or exo?)

- exothermic, -H but disorder is decreasing because we are going from the liquid to solid state so -S (disorder decreases) +G = -H - T(-S) G will be nonspontaneous (+) at high temperatures G will be spontaneous (-) at low temperatures -G= -H - smallT(-S)

Gibbs free energy equation

- free energies of formation of standard state elements are 0 Grx = n(Gfproducts) - n(Gfreactants) ΔG = ΔH - TΔS -G -H +S = spontaneous (most favored) H<0 exo H>0 endo S<0 nonspontaneous S>0 spontaneous

ionic solids

- held together by electrostatic attraction between cations and anions in a lattice structure - ionic bonds are strong and solid at RT - strength is dependent on magnitude of ion charges (greater is stronger) and size of ions (smaller size is stronger) *the greater the charge, the stronger the force of attraction between the ions *the smaller the ions, the more they are attracted to each other

Intermolecular hydrogen bonding

- in order to participate in intermolecular hydrogen bonding, a molecule must be able to act as both a H bond donor and aceptor

Types of Solids

- ionic solids - network solids - metallic solids - molecular solids

Dipole-Dipole Forces

- most common - these forces are produced between polar molecules (think vs. NP molecules/symmetrical) the force is aligned along the permanent MOLECULAR dipole - a more polar molecule generates a stronger force: FONClBrISCH tells you polarity (distance between these atoms tells you polarity, N-Cl is not a polar bond but N-I is) i.e O is more polar than S so stronger D-D forces - dipole forces are easily cleaved - two water molecules bonded have dipole-dipole forces because positive end (H side) is attracted to the other water molecules negative dipole (O side)

Hydrogen Bonding

- produced between VERY polar molecules (FON) the force is aligned along the permanent BOND dipole - more paired donors and acceptors generate a stronger force H-bonding with identical molecules is possible if they contain (think about how amino acids H-bond): - a donor (N-H, O-H, F-H) AND - an acceptor (N with LP, O with LP, F with LP) H-bonding with water is possible if they contain: - a donor (N-H, O-H, F-H) OR - an acceptor (N with LP, O with LP, F with LP) acceptors are FON only and must have LP to accept H atoms

hybridization

- sets of electrons determine this - can be used to determine geometry/shape - each bond (single, double, triple) = 1 group of electrons - each LP = 1 group of electrons

Which mixture of molecules will have the strongest interactions? A. F2/Cl2 B. Cl2/NH3 C. CH3I/H2O D. NH3/H2O

- strongest to weakest intermolecular force is: hydrogen bonding > dipole-dipole forces > London dispersion forces. - CH3I has dipole-dipole forces while H2O exhibits hydrogen bonding; the two associate via dipole-dipole forces. - Cl2 has London dispersion forces while NH3 exhibits hydrogen bonding; the two interact with London dispersion forces. - Both F2 and Cl2 exhibit only London dispersion forces (both individually and with each other). - both NH3/H2O undergo hydrogen bonding (both individually and with each other), and thus would be the mixture with the strongest interaction.

The third law of Thermodynamics

- that a perfect crystal possesses no entropy at absolute zero

The zeroth law of Thermodynamics

- that two systems at equilibrium with a third are in equilibrium with each other - thermal equilibrium - energy will flow from a body at a higher temperature to a body at a lower temperature until both bodies have the same temperature - energy flow into a system has a (+) sign - energy flow out of a system has a (-) sign

polar bonds

- unequal sharing of electrons; dipoles are asymmetrical - difference in EN - dipole always starts with positive atom and points toward negative atom ex: O-H *can have a NP molecule with polar bonds (ex: CCl4)

Thermodynamics

- when energy flows into a system from the surroundings, the energy of the system increases and the energy of the surroundings decrease - when energy flows out of a system into the surroundings, the energy of the system decreases and the energy of the surroundings increases

Which of the following molecules has the lowest boiling point? A. F2 B. Cl2 C. H2O D. HCN

- which of the choices has the weakest intermolecular forces? - strongest to weakest intermolecular force is: hydrogen bonding > dipole-dipole forces > London dispersion forces. - H2O (water) can hydrogen bond, HCN has dipole-dipole forces, and F2 and Cl2 only have London dispersion forces. Thus, H2O and HCN may immediately be eliminated. - To differentiate the two molecules that solely exhibit London dispersion forces, the Cl2 with a larger molecular weight will have the stronger force because it is more polarizable and has a larger surface area of interaction. Thus, F2 will have the lowest boiling point among the choices.

Lewis Dot Structures

1. Count valence e- (look to group on p-table) (-) charges = extra e- (+) charges = removed e- 2. Arrange the atoms with the least electronegative atom in the center - there is only one central atom - carbon always goes in center - hydrogen never goes in center 3. Connect each of the outer atoms to the central atom with on line (can add more later) 4. Add pairs of electrons as pairs of dots to all non-hydrogen outer atoms until each has 8 e- - start with more EN * resonance structures are possible if there are equivalent locations for the multiple bonds

Checking Lewis Dot Structures

1st: check total number of e- - check that atoms obey octet - do formal charges add up to overall charge - look for the smallest set of formal charges - (+) formal charges should be on the LESS EN atom - (-) formal charges should be on the more EN atom * a resonance hybrid maximizes stability and minimizes charge

sp

2 groups = linear = 180 degrees = most stable

sp2

3 groups = trigonal planar = 120 degrees

sp3

4 groups = tetrahedral = 109.5 degrees = AX4

sp3 - bent

<< 109.5 2 LPs = AX2E2

Which has the strongest dipole moment? PBr3O PF5 CCl4 SF6

= PBrO3, because the 2 atoms involved in the bond differ in EN. However, an entire molecule can only have a dipole if it contains bond dipoles and is asymmetrical the remaining choices are trigonal bipyramidal, tetrahedral, and octahedral - so all have identical substituents, are symmetrical and have no net dipole

Intramolecular Forces

= chemical bonds = within molecules chemical bonds are formed when e- are shared between two atoms as their orbitals overlap the strength of a chemical bond is dependent upon: - more electrons shared = stronger bond (i.e. triple vs single bond) - shorter distance between atoms = stronger bond (i.e triple bond vs double bond) - stronger bonds = higher bond dissociation energies - breaking a bond is always an ENDOTHERMIC process (+ energy)

van der waals

= dipole forces, hydrogen bonding, london forces collectively - occur between neutral molecules or atoms (i.e. NOT ion-dipole)

How many orbitals make up the 5f subshell?

A total of 7 orbitals make up the 5f subshell. - There is one s orbital (2e-) - three p orbitals (6e-) - five d orbitals (10e-) - seven f orbitals. (14e-) (Note: Do not confuse the number of orbitals in a subshell with the number of electrons the subshell can hold. Each orbital can hold two electrons, so the capacity of an nf subshell is 7 × 2 = 14 electrons.)

Which two atomic orbitals interact to form the D—D bond in D2? A. s and s B. p and p C. sp and sp D. sp3 and sp3

A. The two atomic orbitals that interact to form the D—D bond in D2 are s and s. Deuterium (an isotope of hydrogen) only has a 1s orbital, and is therefore not hybridized ("sp and sp" and "sp3 and sp3" can be eliminated). The D—D bond, then, is formed by an s orbital from each deuterium.

Ionic Bonds (S=I, L=C; NM+M)

AKA an electrostatic bond 1 from s-block (M) and 1 from p-block (NM) - formed between particles of opposite charges (cations with anions); formed between elements with large differences in EN; larger charges and smaller ions male the strongest ionic bonds - the electrons are LOCALIZED on the ions - no electrons are shared in ionic bond - the ions dissociate in aqueous solution as an electrolyte - an aqueous solution of an ionic compound is a CONDUCTOR - solid compounds with ionic bonds are INSULATORS and brittle

periodic table

Alkali metals? Alkaline earth metals? Transition metals? NonmetalsMetalloids? - includes H Noble gases?

Rank the average C—O bond length from shortest to longest for CO, CO2, and CO32-. A. CO32- < CO2 < CO B. CO < CO2 < CO32- C. CO32- < CO < CO2 D. CO < CO32- < CO2

B. The rank of the average C—O bond length from shortest to longest for CO, CO2, and CO32- is CO < CO2 < CO32-. Bond length can be attributed to the hybrid orbitals which form the bond (sp orbitals form shorter bonds than sp2 orbitals, which form shorter bonds than sp3 orbitals). CO and CO2 have shorter bonds lengths than CO32-, since their carbon atom is sp hybridized, while that of CO32- is sp2 hybridized ("CO32- < CO2 < CO" and "CO32- < CO < CO2" can be eliminated). CO has a shorter bond length than CO2, since its oxygen orbitals are sp hybridized, while they are sp2 hybridized in CO2.

How many electrons are shared in the bond between sodium and chlorine in a molecule of NaCl? A. 1 B. 0 C. 2 D. 3

B. There are 0 electrons shared in the bond between sodium and chlorine in a molecule of NaCl. NaCl is an ionic compound, and no electrons are shared in an ionic bond.

Among the following, which molecular geometry CANNOT result in a nonpolar structure? A. Bent B. Diatomic covalent C. Square planar D. Trigonal planar

Bent - Only molecules with complete symmetry around the central atom can be nonpolar. - Bent molecules lack this symmetry, because both noncentral atoms lie on the same side of the central atom. (bc < 180 degrees, LPs push central atoms down on same side)

How many hybrid orbitals does each fluorine atom in XeF4 have? A. Five B. Six C. Four Correct Answer D. Seven

C. There are four hybrid orbitals on each fluorine atom in XeF4. Each fluorine atom has three lone pairs and a single bond, giving a total of four electron groups. These require four orbitals: one s and three p to make four sp3 orbitals. - Xe can have more than 8e-

The intermolecular forces that exist among the molecules of NH3 gas are: A. dipole-dipole forces only. B. London dispersion forces only. C. both dipole-dipole and London dispersion forces. D. neither dipole-dipole nor London dispersion forces.

C. both dipole-dipole and London dispersion forces - All "real" molecules and atoms will exhibit London dispersion forces. It is not possible to exhibit dipole-dipole forces only. - Strictly nonpolar molecules will exhibit London dispersion forces. Since NH3 has a permanent dipole, it will also exhibit dipole-dipole forces of attraction. - Since NH3 is a permanent dipole, it will exhibit dipole-dipole intermolecular forces in addition to the London dispersion forces exhibited by all molecules.

Chemical bonds (intramolecular forces)

Chemical bonds can be either ionic or covalent - determined by the difference in electronegativity of the atoms FON - ClBrISC - H - if the EN = 0 or VERY SMALL, the bond is NONPOLAR and e- are shared equally - if the EN>0 and moderate in value, the bond is POLAR and e- are NOT shared equally = IONIC i.e. F will have a polar and ionic bond H will have a covalent bond

Which of the following types of orbitals of the central atom are involved in bonding in octahedral compounds? A.sp B.sp3 C.p D.d2 sp3

D.d2 sp3 octahedral compounds have six σ (sigma) bonds and no lone pairs.

The first law of thermodynamics

Energy can be transferred and transformed, but it cannot be created or destroyed - the total energy of the universe is constant - energy is always conserved and in this example, it is simply changing form from chemical energy to thermal energy - the internal energy of an object is proportional to its temperature *an isolated system has constant energy - no transformation of the energy is possible. When systems are in contact, however, energy is allow to flow, and thermal equilibrium can be attained. **work can be put into a system to increase its overall energy - this may or may not occur with a corresponding change in temperature

Hess's Law of Heat Summation

Enthalpy is a state function, so it is independent of pathway Hess's Law problems require you to combine 2+ reactions and their enthalpies to find the change in enthalpy for the overall reaction - reversing the direction of a reaction changes the sign of H - if you change the stoichiometric coefficients you must scale the value of H for the reaction - add reactions together to cancel out intermediate species

Coordinate Covalent - Hemoglobin

Fe = electrophile = lewis acid; it is ligated and chelated 4 Nitrogens all have LP to donate to Fe

Gibbs Free Energy

Free Energy (G) - energy available to do work in a chemical process - G is not stored energy like H a spontaneous process = exergonic -G (products are lower in energy than reactants) a non-spontaneous process = endergonic +G (products are higher in energy than reactants)

At room temperature, H2S is a gas, but H2O is a liquid. Which of the following plays an important role in this observation? A. The electronegativity of oxygen is greater than that of sulfur. B. Sulfur has a greater ionization energy than oxygen has. C. Sulfur has a smaller atomic radius than oxygen. Your Answer D. Stronger hydrogen bonds form between molecules of hydrogen sulfide than between molecules of water.

H2O molecules are much more polar than H2S molecules because O is more electronegative than S. As a result, the hydrogen bonds between water molecules are much stronger than the dipole-dipole interactions between H2S molecules

Orbital geometry vs Molecular geometry

Orbital geometry - will either be linear (sp), trigonal planar (sp2), or tetrahedral (sp3) Shape or Molecular geometry - accounts for how lone pairs change the shape of a molecule ex: H2O orbital geometry = tetrahedral (sp3 hybridization) shape/molecular geometry = bent (2LPs, still sp3)

Entropy (S)

POTENTIAL randomness - increasing the number of particles increases entropy - increasing the volume (of the container) increases entropy - changing phase from s to l to g = increases entropy -increasing the temperature increases entropy (makes things move faster) - use absolute entropies for equation - otherwise use Hess's Law of Summation to calculate entropy

Strength of Intermolecular Forces (IMFs)

STRONGEST: ion dipole hydrogen bonding dipole dipole induced dipole LDF

Strength of Chemical bonds (intramolecular forces)

STRONGEST: ionic covalent coordinate covalent metallic

Molecular solids

Solids that are composed of molecules held together by their intermolecular forces (3): - hydrogen bonds - dipole-dipole - LDFs - weaker than ionic, network, or metallic bonds so they have much lower melting and BP - often liquids or gases at RT; more likely to be solids as the strength of their intermolecular forces increase

The Lewis structure of the azide, N3-, ion is shown below. How many unhybridized p-orbitals are there on the central nitrogen? A. 0 B. 1 C. 2 D. 3

The central N has two bonding partners and no lone pairs. As such it has an sp type hybridization. This leaves two unhybridized orbitals to participate in the two π-bonds in the molecule.

Which of the following represents the ground state electron configuration of a Co(III) ion? [Ar] 4s2 3d10 or [Ar] 3d6

The ground state electron configuration of a neutral cobalt atom is [Ar] 4s2 3d7. Co(III) is a cation as metals do not gain electrons

When a sample of solid NH4NO3 is dissolved in water, the reaction flask becomes cold to the touch. What can be concluded about the given thermodynamic values for the solvation process? A. ΔG < 0, ΔH < 0, ΔS > 0 B. ΔG > 0, ΔH < 0, ΔS > 0 C. ΔG < 0, ΔH > 0, ΔS > 0 D. ΔG < 0, ΔH > 0, ΔS < 0

When a sample of solid N4NO3 is dissolved in water, the reaction flask becomes cold to the touch. In can be concluded that ΔG < 0, ΔH > 0, ΔS > 0 for the solvation process. Since the solvation process is observed to occur, it is spontaneous, so ΔG < 0. If the reaction flask is cold to the touch, then the enthalpy change of solvation is positive (ΔH > 0) as this is an endothermic process. Finally, dissolving a solid in a liquid tremendously increases the disorder of the two components as they mix, so ΔS > 0.

London dispersion forces

are produced between all molecules (for NP molecules - this is all they have) - aka induced dipole-induced dipole forces - the force produces a random molecular collision, which then produces a temporary and transient dipole the force is produced by collisions that produce temporary but small dipoles by deforming the electron cloud - a molecule with more electrons and larger size (larger cloud) generates a stronger force (more polarizable) ex: CCl4 (NP) at first, even distributions of e- cloud. Then, e- shift to the left = induced dipole moment * will affect neighboring molecules (i.e another CCl4) and cause them to shift left too = temporary and non-permanent. Constantly changing & poles shift. ** the LDFs are between the 2 CCL4 molecules** (or any NP molecule I2, Ch4, Br2) * dispersion forces are very weak and easily cleaved

Energy is needed to break a bond

endothermic, +H (if a reaction occurred and flask becomes cold to touch, +H) ** when bonds are broken, energy is absorbed

isentropic

entropy remains constant

Energy is released in making a bond

exothermic, -H = releases heat (a molecule's bonds becoming more stable, losing energy) (reactants have more energy than products, energy is released to get to products) ** energy is released when bonds are formed

Which of the following lists hydrogen halides in terms of increasing standard heats of formation? A. HF < HBr < HCl < HI B. HI < HBr < HCl < HF C. HBr < HF < HCl < HI D. HF < HI < HCl < HBr

he hydrogen halides in terms of increasing standard heats of formation is: HF < HBr < HCl < HI. The general formula for hydrogen halide formation is: 1/2H2+ 1/2X2 -> HX - The reaction that forms the most stable HX will be the most exothermic and have the most negative standard heat of formation. - Fluorine is the smallest, most electronegative halogen and forms the only hydrogen halide of the group that is not a strong acid. HF is the most stable hydrogen halide and will have the most negative heat of formation. - the heat of formation of HI is actually endothermic. This eliminates choice HF < HI < HCl < HBr.

sp2 - bent

i.e. O3 (ozone) < 120 degrees = bent and has one LP

Coordinate Covalent Bonds (NM+TSM)

needs d-orbital (NM + TSM) - are formed between atoms with lone pairs and electron deficient species (one atom is donating all of the electrons) = a lewis base (e- pair donor) and lewis acis (e- pair acceptor); electrons are shared - the electrons are LOCALIZED directly between the atoms - electrons are donated from the nucleophile (= lewis base, ligand, chelate) - some ligands (=lewis base/nuclophile) may donate more than one lone pair of electrons - a ligand that donates multiple LPs is a chelate - compounds with coordinate covalent bonds are easily dissociated (coordinate complexes)

+G -H +S

nonspontaneous

at high temperatures: +G -H -S

nonspontaneous

at low temperatures: +G +H +S

nonspontaneous

Covalent bonds (I; NM+NM)

p - block elements (NM + NM) - are formed between atoms with high electronegativity (nonmetals with nonmetals) - formed between elements with similar electronegativites - the e- are LOCALIZED directly between the atoms - the e- in the bond are donated from both atoms (equally shared) - compounds with covalent bonds are INSULATORS and rigid (Covalent compounds are insulators as they are strong bonds compounds and there are no free electrons available to conduct electricity, so they are insulators.) - the molecular dipole is found by adding all the bond dipoles

Metallic Bonds (C) (M+M)

s - block elements (can never coordinate covalent bond because no d-orbital) - bonds are formed between atoms with low electronegativity (metals with metals) i.e. Na (s), Li (s) - the electrons are DELOCALIZED among ALL the atoms - the electrons in the bond are donated from ALL the atoms = "sea of electrons" - valence electrons are donated in a sea that neighboring atoms share - compounds with metallic bonds are CONDUCTORS and malleable (ductile) (The availability of "free" electrons contributes to metals being excellent conductors)

sigma bonds vs pi bonds

sigma - 2 e- localized between 2 nuclei. It is formed by the end-to-end overlap of one hybridized orbital from each of the two atoms participating in the bond in C2H6, sigma bonds formed between C-H (sp3-s bond) and between C-C (p3-sp3) pi - composed of 2 e- that are localized to the region that lies on opposite sides of the plane formed by the two bonded nuclei and immediately adjacent atoms, not directly between the nuclei as with the sigma bond. Formed by proper, parallel, side-to-side alignment of two unhybridized p orbitals on adjacent atoms **In any multiple bond, there is only one sigma bond; the remainder are pi bonds: - single bond = 1 sigma bond - double bond = 1 sigma + 1 pi bond - triple bond = 1 sigma + 2 pi bonds

at high temperatures: -G +H +S

spontaneous

at low temperaturese: -G -H -S

spontaneous

Enthalpy change of a reaction

the enthalpy change of a reaction can be estimated by summing the total enthalpy required to break the bonds in the reactants with the energy released by forming the bonds in the products: sum of bonds broken (reactants) + sum of bonds formed (products) breaking bonds is ALWAYS ENDOTHERMIC (+)H - reactants are lower on graph than products (must add energy into the process) forming bonds is ALWAYS EXOTHERMIC (-)H because you are releasing heat - reactants are higher on graph than products

Enthalpies of Formation

the heats of formation (Hf) is the amount of energy associated with formed ONE MOLE of a compound from its constitutive elements in their standard states - the Hf of any element in its standard state = 0 - bond dissociation energies (BDE) are enthalpy changes required to break bonds: Hf = n(BDEbroken)-n(BDEformed) - must account for the stoichiometry of the reaction in the equation

sp3 (1LP)

trigonal pyramidal = <109.5 degrees = AX3E