CHEM 102 (Pre-Exam 2) by Christelle Seri

The Three States

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The phase in which a substance exists depends on the relative extents of its intermolecular forces (IMFs) and the kinetic energies (KE) of its molecules.

Particles in a solid are tightly packed together and often arranged in a regular pattern; in a liquid, they are close together with no regular arrangement; in a gas, they are far apart with no regular arrangement.
Particles in a solid vibrate about fixed positions and do not generally move in relation to one another; in a liquid, they move past each other but remain in essentially constant contact; in a gas, they move independently of one another except when they collide.

IMF and KE

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IMFs are the various forces of attraction that may exist between the atoms and molecules of a substance due to electrostatic phenomena, as will be detailed in this module. These forces serve to hold particles close together, whereas the particles’ KE provides the energy required to overcome the attractive forces and thus increase the distance between particles.

IMF and KE (illustrated)

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Intermolecular Force

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All of the attractive forces between neutral atoms and molecules are known as van der Waals forces

Dispersion Force

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At any moment in time, an atom or molecule can develop a temporary, instantaneous dipole if its electrons are distributed asymmetrically --> can distort neighboring electrons and create an induced dipole--> result in a relatively weak electrostatic attraction between the species

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Why does melting and boiling points increase with increasing atomic/molecular size?

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In a larger atom, the valence electrons are, on average, farther from the nuclei than in a smaller atom. Thus, they are less tightly held and can more easily form the temporary dipoles that produce the attraction.

Shape and Dispersion Force

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The greater the surface area, the greater the dispersion forces.

Think Velcro

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Dipole-Dipole Attraction

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The electrostatic force between the partially positive end of one polar molecule and the partially negative end of another

Hydrogen Bonds

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The large difference between the boiling points is due to a particularly strong dipole-dipole attraction that may occur when a molecule contains a hydrogen atom bonded to a fluorine, oxygen, or nitrogen atom (the three most electronegative elements).

Hydrogen Bonds (cont.)

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Despite use of the word “bond,” keep in mind that hydrogen bonds are intermolecular attractive forces, not intramolecular attractive forces (covalent bonds). Hydrogen bonds are much weaker than covalent bonds, only about 5 to 10% as strong, but are generally much stronger than other dipole-dipole attractions and dispersion forces.

Intermolecular Forces*****

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1. London Dispersion Forces (LDF)
2. Dipole‐Dipole Interactions
3.Ion Dipole Interactions (salt dissolving in solution)
4. Hydrogen Bonding

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LDF

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Weak and short-lived but can be significant in large atoms
usually increase with molar mass
are present in all compounds.
can occur between atoms or molecules.
are due to electron movement, not to differences
in electronegativity

Polarizability

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The ease with which the electron distribution in the atom
or molecule can be distorted, inducing a temporary dipole.

Polarizability Trends

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Smaller atoms or ions are less polarizable than larger ones because their electrons are closer to the nucleus and therefore held more tightly.
Polarizability increases down a group as larger electron clouds are
easier to distort. Polarizability decreases across a period dues to increasing Zeff making the atoms smaller.
Cations are less polarizable than their parent atoms and anions are more polarizable than their parent atoms.

Pi Bonds and Polarizability

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Pi bonds allow for the delocalization of electrons (they are free roaming) making them easier to distort.

Ion-Dipole Interactions

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The larger the charge, the stronger the force

Hydrogen Bonding

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Involves Hydrogen in a bond with F, O, or N (FON Times), interacting with a lone pair in another F, O, N molecules.

Some Quick Facts

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High Adhesion = High Surface Tension = High IMF

Polar Liquids = High IMF = Capillary Action

Adhesion vs Cohesion

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Co-hesion = between like molecules
Ad-hesion = between unlike molecules

Cohesive forces within water (polar) try to minimize the surface area.

Vaporization

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Vaporization is endothermic (heat is absorbed) because energy is required to overcome the relatively strong intermolecular forces in the liquid.
High temperatures are required to vaporize water due to hydrogen
bonding among its molecules in the liquid state.

Vapor Pressure

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The pressure of the vapor (gas) present at equilibrium is called the
vapor pressure of the liquid.

Definition of vapor pressure: The pressure of the vapor resulting from evaporation of a liquid above a sample of the liquid in a closed container.

Volatility

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A substance with a high vapor
pressure is considered to be
volatile (changes from liquid to gas easily)

Normal Boiling Point

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The normal boiling point of a liquid
is the temperature at which the
vapor pressure of the liquid is
exactly 760 torr.

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Lower boiling points correspond to
higher vapor pressures and weaker
IMF.

A Quick Comment

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Be sure to check the polarity before determining hydrogen bonds. Also watch out for ionic-dipole bonds

Mass

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The mass of an element shown in a periodic table or listed in a table of atomic masses is a weighted, average mass of all the isotopes present in a naturally occurring sample of that element. This is equal to the sum of each individual isotope’s mass multiplied by its fractional abundance.

Diatomic Molecules and Friend

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Most samples of the elements hydrogen, oxygen, and nitrogen are composed of molecules that contain two atoms each (called diatomic molecules) and thus have the molecular formulas H2, O2, and N2, respectively. Other elements commonly found as diatomic molecules are fluorine (F2), chlorine (Cl2), bromine (Br2), and iodine (I2). The most common form of the element sulfur is composed of molecules that consist of eight atoms of sulfur; its molecular formula is S8

Molecular Formula

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Note that a molecular formula is always a whole-number multiple of an empirical formula.

MOLES

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1 mole of any element contains the same number of atoms as 1 mole of any other element. The masses of 1 mole of different elements, however, are different, since the masses of the individual atoms are drastically different. The molar mass of an element (or compound) is the mass in grams of 1 mole of that substance, a property expressed in units of grams per mole (g/mol)

Avogadro's Number

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The number of entities composing a mole has been experimentally determined to be 6.02214179 x 10²³

Formula Mass

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Keep in mind, however, that the formula for an ionic compound does not represent the composition of a discrete molecule, so it may not correctly be referred to as the “molecular mass.”

Empirical Formulas

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In summary, empirical formulas are derived from experimentally measured element masses by:

Deriving the number of moles of each element from its mass
Dividing each element’s molar amount by the smallest molar amount to yield subscripts for a tentative empirical formula
Multiplying all coefficients by an integer, if necessary, to ensure that the smallest whole-number ratio of subscripts is obtained

I don't even understand wtf is going on....

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Fundamental Aspects of Any Chemical Equation

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The substances undergoing reaction are called reactants, and their formulas are placed on the left side of the equation.
The substances generated by the reaction are called products, and their formulas are placed on the right sight of the equation.
Plus signs (+) separate individual reactant and product formulas, and an arrow separates the reactant and product (left and right) sides of the equation.
The relative numbers of reactant and product species are represented by coefficients (numbers placed immediately to the left of each formula). A coefficient of 1 is typically omitted.

Formula Units????

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Stoichiometry

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This is a big topic

ION DIPOLE CHRISTELLE GADAMMIT

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Solubility

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For solubility, "like dissolves like" means that polar solutes dissolve
in polar solvents and nonpolar solutes dissolve in nonpolar solvents.

IOOOOONNNNNN DIPOOOOOOLLLLLEEEEE

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Number of Atoms?

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Empirical Formulas

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Molecular to Empirical

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Percents Based on Mass

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Formula UNITS

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ALL possible uses of Avogadro's Number

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Finding Limiting Reagents When The Masses are Different

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Remember Molar Ratios When Calculating Limiting Reagents

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When ionic compounds dissolve in water, they may dissociate into their constituent ions, which are subsequently dispersed homogenously throughout the resulting solution (a thorough discussion of this important process is provided in the chapter on solutions). Ionic compounds dissolved in water are, therefore, more realistically represented as dissociated ions, in this case:

Dissolution of Ionic Compounds

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Complete Ionic Equation

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Explicitly representing all dissolved ions

S'more

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These spectator ions—ions whose presence is required to maintain charge neutrality—are neither chemically nor physically changed by the process, and so they may be eliminated from the equation to yield a more succinct representation called a net ionic equation

Precipitation

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A precipitation reaction is one in which dissolved substances react to form one (or more) solid products. Many reactions of this type involve the exchange of ions between ionic compounds in aqueous solution and are sometimes referred to as double displacement, double replacement, or metathesis reactions

Solubility

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The extent to which a substance may be dissolved in water, or any solvent, is quantitatively expressed as its solubility, defined as the maximum concentration of a substance that can be achieved under specified conditions. Substances with relatively large solubilities are said to be soluble. A substance will precipitate when solution conditions are such that its concentration exceeds its solubility. Substances with relatively low solubilities are said to be insoluble, and these are the substances that readily precipitate from solution.

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Electrolytes

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When some substances are dissolved in water, they undergo either a physical or a chemical change that yields ions in solution. These substances constitute an important class of compounds called electrolytes. Substances that do not yield ions when dissolved are called nonelectrolytes. If the physical or chemical process that generates the ions is essentially 100% efficient (all of the dissolved compound yields ions), then the substance is known as a strong electrolyte. If only a relatively small fraction of the dissolved substance undergoes the ion-producing process, it is called a weak electrolyte.

Conducting Solutions

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Solutions may also conduct electricity if they contain dissolved ions, with conductivity increasing as ion concentration increases

Dissociation

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When ionic compounds dissolve in water, the ions in the solid separate and disperse uniformly throughout the solution because water molecules surround and solvate the ions, reducing the strong electrostatic forces between them. This process represents a physical change known as dissociation. Under most conditions, ionic compounds will dissociate nearly completely when dissolved, and so they are classified as strong electrolytes.

Exceptions to Dissociation

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In other cases, the electrostatic attractions between the ions in a crystal are so large, or the ion-dipole attractive forces between the ions and water molecules are so weak, that the increase in disorder cannot compensate for the energy required to separate the ions, and the crystal is insoluble. Such is the case for compounds such as calcium carbonate (limestone), calcium phosphate (the inorganic component of bone), and iron oxide (rust)

Water as a Conductor

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Pure water is an extremely poor conductor of electricity because it is only very slightly ionized

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The Struggle is Real

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When performing calculations stepwise, as in Example 6.11, it is important to refrain from rounding any intermediate calculation results, which can lead to rounding errors in the final result.

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A mass-volume percent is a ratio of a solute’s mass to the solution’s volume expressed as a percentage.

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Many acids release hydrogen ions, H+, when dissolved in water. To denote this distinct chemical property, a mixture of water with an acid is given a name derived from the compound’s name. If the compound is a binary acid (comprised of hydrogen and one other nonmetallic element):
1. The word “hydrogen” is changed to the prefix hydro-
2. The other nonmetallic element name is modified by adding the suffix -ic
3. The word “acid” is added as a second word
For example, when the gas HCl (hydrogen chloride) is dissolved in water, the solution is called hydrochloric acid.

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However, we will briefly discuss the important compounds known as oxyacids, compounds that contain hydrogen, oxygen, and at least one other element, and are bonded in such a way as to impart acidic properties to the compound (you will learn the details of this in a later chapter). Typical oxyacids consist of hydrogen combined with a polyatomic, oxygen-containing ion. To name oxyacids: 1. Omit “hydrogen” 2. Start with the root name of the anion 3. Replace –ate with –ic, or –ite with –ous 4. Add “acid” For example, consider H2CO3 (which you might be tempted to call “hydrogen carbonate”). To name this correctly, “hydrogen” is omitted; the –ate of carbonate is replace with –ic; and acid is added—so its name is carbonic acid.

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An acid-base reaction is one in which a hydrogen ion, H+, is transferred from one chemical species to another.

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The nature of HCl is such that its reaction with water as just described is essentially 100% efficient: Virtually every HCl molecule that dissolves in water will undergo this reaction. Acids that completely react in this fashion are called strong acids, and HCl is one among just a handful of common acid compounds that are classified as strong (Table 7.2). A far greater number of compounds behave as weak acids and only partially react with water, leaving a large majority of dissolved molecules in their original form and generating a relatively small amount of hydronium ions. Weak acids are commonly encountered in nature, being the substances partly responsible for the tangy taste of citrus fruits, the stinging sensation of insect bites, and the unpleasant smells associated with body odor.

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A base is a substance that will dissolve in water to yield hydroxide ions, OH−. The most common bases are ionic compounds composed of alkali or alkaline earth metal cations (groups 1 and 2) combined with the hydroxide ion—for example, NaOH and Ca(OH)2. When these compounds dissolve in water, hydroxide ions are released directly into the solution. For example, KOH and Ba(OH)2 dissolve in water and dissociate completely to produce cations (K+ and Ba2+, respectively) and hydroxide ions, OH−. These bases, along with other hydroxides that completely dissociate in water, are considered strong bases.

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Solids (s), liquids (l), and gases (g) do not get separated into ions.

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We find that temperature and pressure are linearly related, and if the temperature is on the kelvin scale, then P and T are directly proportional (again, when volume and moles of gas are held constant)

Gay-Lussac's/Amonton's Law

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the pressure of a given amount of gas is directly proportional to its temperature on the kelvin scale when the volume is held constant.

For a confined, constant volume of gas, the ratio of P/T is therefore constant.

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The volume increases as the temperature increases, and decreases as the temperature decreases.

Charles's Law

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Charles’s law states that the volume of a given amount of gas is directly proportional to its temperature on the kelvin scale when the pressure is held constant.

For a confined, constant pressure gas sample, V/T is constant

Boyle's Law

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The volume of a given amount of gas held at constant temperature is inversely proportional to the pressure under which it is measured.

P and V exhibit inverse proportionality: Increasing the pressure results in a decrease of the volume of the gas.

Avogadro's Law

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For a confined gas, the volume (V) and number of moles (n) are directly proportional if the pressure and temperature both remain constant.

The Four Gas Laws

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The Combined Gas Law

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Standard Units

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Chemists sometimes make comparisons against a standard temperature and pressure (STP) for reporting properties of gases: 273.15 K and 1 atm (101.325 kPa). At STP, an ideal gas has a volume of about 22.4 L—this is referred to as the standard molar volume

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Mean Free Path

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The mean free path is the average distance a molecule travels between collisions. The mean free path increases with decreasing pressure; in general, the mean free path for a gaseous molecule will be hundreds of times the diameter of the molecule

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Ratios of effusion and diffusion are the same

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Graham’s law of effusion: The rate of effusion of a gas is inversely proportional to the square root of the mass of its particles:

Kinetic Molecular Theory (KMT)

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1. Gases are composed of molecules that are in continuous motion, traveling in straight lines and changing direction only when they collide with other molecules or with the walls of a container.

2. The molecules composing the gas are negligibly small compared to the distances between them.

3. The pressure exerted by a gas in a container results from collisions between the gas molecules and the container walls.

4, Gas molecules exert no attractive or repulsive forces on each other or the container walls; therefore, their collisions are elastic (do not involve a loss of energy).

5. The average kinetic energy of the gas molecules is proportional to the kelvin temperature of the gas.

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Ideal Gas Law Constraints

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Only works for low pressures