Chemistry Honors: Unit 3 Test (9.1-14.3)

percent composition from chemical formula

% by mass = (mass of element in 1 mol/molar mass of compound) × 100% 1. use subscripts in formula to calculate mass of each element found in one mole of the compound 2. then that value is divided by the molar mass of the compound & multiplied by 100% *the percent composition of a given compound is always the same as long as the compound is pure

exceptions to the octet rule

(1) incomplete octet (2) odd-electron molecules (3) expanded octet

Boyle's Law

(P x V = k, where k is a constant for given sample of gas & depends only on amount of gas and temperature) states that the volume of a given mass of gas varies inversely with the pressure when the temperature is kept constant; he discovered that doubling the pressure of an enclosed sample of gas while keeping its temperature constant caused the volume of the gas to be reduced by half; constant (k) is always equal to pressure multiplied by volume; graphed like a curve asymptote (inverse relationship)

Combined Gas Law

(P x V/T = k) (P₁ x V₁/T₁ = P₂ x V₂/T₂) expresses the relationship between pressure, volume, and absolute temperature of a fixed temperature of a fixed amount of gas; only amount of gas is held constant

Gay-Lussac's Law

(P/T = k) (P₁/T₁ = P₂/T₂) states that the pressure of a given mass of gas varies directly with the absolute temperature of the gas when the volume is kept constant; discovered that when temperature increases of gas in RIGID container, kinetic energy increases resulting in increased pressure; graph illustrates direct relationship; as gas is cooled at constant volume, its pressure continually decreases until gas condenses into liquid

ideal gas law

(PV = nRT) single equation which relates the pressure, volume, temperature, and number of moles for any ideal gas; (P x V/T x n = R) (P₁ x V₁/T₁ x n₁ = P₂ x V₂/T₂ x n₂); constant (R) can be evaluated if gas is considered to be ideal ideal gas law can find molar mass of unknown gas and the density if temperature and pressure of gas are also known

Avogadro's Law

(V = k x n) (V₁/n₁ = V₂/n₂) states that volume of a gas is directly proportional to the number of moles when temperature and pressure are held constant; follows Avogadro's hypothesis (equal volumes of any gas at the same temperature and pressure contain same number of molecules); n is number of moles of gas & k is a constant; if container holding gas is rigid rather than flexible, pressure can be substituted for volume in Avogadro's Law

Charles's Law

(V/T = k; k is constant only for a given gas sample) states that the volume of a given mass of gas varies directly with the absolute temperature of the gas when the pressure is kept constant; absolute temperature is temperature measured with Kelvin scale; studied effect of temperature on volume of a gas at constant temperature; as a closed, FLEXIBLE container of gas is heated, its molecules increase in kinetic energy and push the movable piston outward, resulting in an increase in volume; constant (k) is always equal to volume divided by Kelvin temperature graphed as straight line (direct relationship); as absolute temperature of gas approaches zero, its volume approaches zero (but when gas is brought to extremely cold temperatures, its molecules eventually condense into liquid state before reaching absolute zero) law can be used to compare changing conditions for a gas (V₁/T₁ = V₂/T₂)

(AB₃E) trigonal pyramidal

(ex: ammonia NH₃ w/ H-N-H angle approx. 107°); in the tetrahedral domain geometry, four electron pairs (1 lone pair & 3 bonding pair)

(AB₂E₂) bent molecular geometry

(ex: water H₂O w/ H-O-H bond angle at 104.5°); in the tetrahedral domain geometry, four electron pairs (2 lone pairs & 2 lone pairs)

polar covalent bond

(∆EN 0.4-1.7) a covalent bond in which the atoms have an unequal attraction for electrons, so the sharing is unequal; the distribution of shared electrons within the molecule is no longer symmetrical; electron density is unevenly distributed (it is higher around the atom with greater electronegativity); ex: HF (highly polar molecule w/ electron density around fluorine higher than around hydrogen, electronegativity difference of 1.9); to illustrate the uneven electron distribution in a polar covalent bond, use the Greek letter, delta (δ) along with a positive or negative sign to indicate that an atom has a partial positive or negative charge; the atom with the greater electronegativity acquires a partial negative charge, while the atom with the lesser electronegativity acquires a partial positive charge (delta symbol is used to indicate that the quantity of charge is less than one); a crossed arrow can also be used to indicate the direction of greater electron density

nonpolar covalent bond

(∆EN<0.4) a covalent bond in which the bonding electrons are shared equally between the two atoms; the distribution of electrical charge is balanced between the two atoms & the electron density surrounding the molecule is symmetrical; any diatomic molecule in which the two atoms are the same element must be joined by a nonpolar covalent bond; molecules in which the electronegativity difference is very small (<0.4) are also considered nonpolar covalent; (ex: Cl₂)

to find molar masses of compounds

*molecular mass of a compound is the mass of one molecule of that compound *mass of one formula unit of an ionic compound is the formula mass MOLAR MASS of any compound is the mass in grams of one mole of that compound (same as the molecular mass & formula mass, but in g/mol) 1. Know the formula of the compound 2. Use the molar mass of each atom component together with the quantity of each atom in the formula to find the total molar mass

symbols used in chemical equations

+ (used to separate multiple reactants or products) → (yield sign; separates reactants form products) ↔ (replaces yield sign for reversible reactions that reach equilibrium) (s) (reactant/product in solid state) (l) (reactant/product in liquid state) (g) (reactant/product in gas state) (aq) (reactant/product in an aqueous solution-dissolved in water) Pt → (formula written above the arrow is used as a catalyst in the reaction) ∆ → (triangle indicates that the reaction is being heated)

Pressure Unit Conversions

1 atm = 760 mmHg = 760 torr = 101.3 kPa

empirical formula to molecular formula

1. Calculate the empirical formula mass (EFM), which is simply the molar mass represented by the empirical formula. 2. Divide the molar mass of the compound by the empirical formula mass. The result should be a whole number or very close to a whole number. 3. Multiply all of the subscripts in the empirical formula by the whole number found in step 2. The result is the molecular formula.

process of hybridization summary

1. Draw the Lewis electron-dot structure of the molecule. 2. Use VSEPR theory to predict both the electron domain geometry and the molecular geometry of the molecule 3. Match the electron domain geometry to the appropriate hybridization for the central atom

kinetic-molecular theory (as it applies to gases)

1. Gases consist of very large numbers of tiny spherical particles that are far apart from one another compared to their size. - most of volume of gas is composed of empty space between the particles 2. Gas particles are in constant rapid motion in random directions. - fast motion of gas particles give them relatively large amount of kinetic energy (energy that an object possesses because of its motion) - particles move in straight line until collision 3. Collisions between gas particles & between particles and the container walls are elastic collisions. - elastic collision = one in which there is no overall loss of kinetic energy (maybe transfer, but no change) 4. There are no forces of attraction or repulsion between gas particles. - attractive forces are responsible for particles of a real gas condensing together to form a liquid - assumed particles of an ideal gas have no attractive forces (actual gases have London forces, but so few, it is about 0) 5. The average kinetic energy of gas particles is dependent upon the temperature of the gas. - as temperature of a gas increases, its particles move faster, increasing kinetic energies - given sample will contain particles with range of different kinetic energies - average kinetic energy of the particles in a sample is proportional to its temperature

finding density with ideal gas law

1. calculate molar mass of given gas 2. assume exactly 1 mol of given gas & calculate volume that such an amount would occupy at given temperature and pressure 3. calculate density by dividing mass of one mole of given gas by volume

calculating the theoretical yield & percent yield

1. calculate theoretical yield based on stoichiometry (know molar masses) 2. apply stoichiometry to convert from mass of reactant to mass of product (g reactant -> mol reactant -> mol product -> g product): this will get the theoretical yield of product 3. use actual yield and theoretical yield to calculate percent yield (actual yield/theoretical yield x 100%)

solving limiting reactant problems

1. convert given quantities of each reactant (typically measured by mass or by volume) to moles in order to identify the limiting reactant using atomic masses (periodic table) 2. use the balanced equation to calculate the number of moles of each substance that would be needed to react with the number of moles of the other substance present 3. compare this result to the actual number of moles of the substance present

stoichiometry with number of representative particles (ex. mass to number of particles)

1. convert mass of given to moles using periodic table 2. use mole ratio to convert moles of given to moles of unknown 3. use Avogadro's number (6.02 x 10²³ units/mol) to convert moles of unknown to number of particles of unknown

determining amount of excess reactant left over

1. convert the quantities of the substances to moles 2. find the amount of excess reactant needed using the balanced equation 3. subtract the amount (in moles) of the excess reactant that will react from amount that is originally present 4. convert moles to grams

determining quantity of product formed in a reaction

1. convert the quantities of the substances to moles 2. use the balanced equation to find the limiting and excess reactants *the limiting reactant is the reactant that will determine the amount of product that is produced 3. know the molar mass of product 4. use stoichiometry to calculate the number of moles of product produced and then convert that amount to grams (moles of substance x mole ratio from balanced equation x molar mass of product)

converting moles to mass

1. find the molar mass of substance 2. use this as conversion factor and multiply by how many moles you have 3. you should end up with mass in grams

to find how many x atoms are in a mole of y molecules

1. know the chemical formula for y 2. use conversion factors - converting from moles of particles to number of particles (6.02 x 10²³/1) - to find the number of atoms contained with each molecule (ratio of x to y)

mole-gas volume conversion

1. know the volume of the gas in L 2. use the conversion factor of molar volume (1 mole=22.4 L) to convert from L to mol

percent of water in a hydrate calculation

1. mass of water in one mole of the hydrate is the coefficient multiplied by the molar mass of H₂O 2. molar mass of the hydrate is the molar mass of the actual formula plus the mass of the associated water 3. calculate the percent by mass of water by dividing the mass of H₂O in 1 mole of the hydrate by the molar mass of the hydrate and multiplying by 100%

finding molar mass with ideal gas law

1. use ideal gas law to solve for moles of unknown gas (n) 2. divide mass of gas by the moles to get molar mass

gas stoichiometry with ideal gas law

1. write & balance chemical equation 2. calculate number of moles of given gas by stoichiometry 3. use ideal gas law to calculate volume of given gas produced *do the same for other variables

converting mass to moles

1.find the mass of the substance in grams 2. find molar mass of the substance and use this as conversion factor and multiply this by mass of substance 3. you should end up with number of moles (mol)

Geometries of Molecules & Ions in which the Central Atom has NO Lone Pairs

2 electron pairs-linear 3 electron pairs-trigonal planar 4 electron pairs-tetrahedral 5 electron pairs-trigonal bipyramidal 6 electron pairs-octahedral

Geometries of Molecules & Ions in which the Central Atom has 1+ Lone Pairs

3 electron pairs (1 lone, 2 bonding)-trigonal planar, bent 4 electron pairs (1 lone, 3 bonding) tetrahedral, trigonal pyramidal 4 electron pairs (2 lone, 2 bonding) tetrahedral, bent 5 electron pairs (1 lone, 4 bonding) trigonal bipyramidal, distorted tetrahedron/seesaw 5 electron pairs (2 lone, 3 bonding) trigonal bipyramidal, T-shaped 5 electron pairs (3 lone, 2 bonding) trigonal bipyramidal, linear 6 electron pairs (1 lone, 5 bonding) octahedral, square pyramidal 6 electron pairs (2 lone, 4 bonding) octahedral, square planar

crystal lattice

3D arrangement of a solid crystal; different arrangements of the particles within a crystal cause them to adopt several different shapes

molecular geometry

3D arrangement of atoms in a molecule; the shape of a molecule is an important factor that affects the physical & chemical properties of a compound (melting&boiling points, solubility, density, etc.); molecular geometries of molecules change when the central atom has one or more lone pairs of electrons

central atom with no lone pairs

A=central atom in a molecule; B=atoms surrounding the central atom; subscripts after the B will denote the number of B atoms that are bonded to the central A atom;

word equation

Reactants -> Products (reactants yield products) Silver + sulfur -> silver sulfide Methane + oxygen -> carbon dioxide + water

VSEPR Summary

VSEPR model can be applied to predict the molecular geometry of a given molecular compound 1. Draw the Lewis electron dot structure for the molecule 2. Count the total number of electron pairs around the central atom (Electron Domain Geometry) 3. If there are no lone pairs around the central atoms, refer to the next term 4. If there are 1+ lone pairs on the central atoms, the molecular geometry (the actual shape of the molecule) will not be the same as the electron domain geometry. Refer to the next next term. 5. In predicting bond angles, remember that a lone pair takes up more space than a bonding pair or pairs of electrons.

hexagonal

a = b ≠ c α = β = 90°, γ = 120° ex: beryl

pi bond (π bond)

a bond formed by the overlap of orbitals in a side-by-side fashion, with the electron density concentrated above and below the plane of the nuclei of the bonding atoms; form when the unhybridized orbital, which is oriented perpendicular to that plane, overlap with another unhybridized orbital side by side, not directly on the line between the two bonded atoms); pi bonds are perpendicular to one another

sigma bond (σ bond)

a bond formed by the overlap of orbitals in an end-to-end fashion, with the electron density concentrated between the nuclei of the bonding atoms; form when the hybrid orbitals overlap identical ones on the other atom or overlap with the 1s orbital of a hydrogen atom

covalent bond

a bond in which two atoms share one or more pairs of electrons; the point at which the potential energy reaches its minimum represents the ideal distance between hydrogen atoms for a stable chemical bond to occur; the single electrons from each of the two hydrogen atoms are shared when the atoms come together to form a hydrogen molecule (H₂); covalent bond forms when two singly occupied orbitals overlap with each other; shared electrons must have opposite spins; unpaired electrons form a covalent bond; forms when the electron clouds of two atoms overlap with each other

balanced equation

a chemical equation in which mass is conserved & there are equal numbers of atoms of each element on both sides of the equation 1. Determine the correct chemical formulas for each reactant & product. 2. Write the skeleton equation with reactants on left side of → and products on right side. 3. Count the number of atoms of each element that appears as a reactant and as a product. If a polyatomic ion is unchanged on both sides of the equation, count it as a unit. 4. Balance each element one at a time by placing coefficients in front of the formulas. No coefficient is written for a 1. It is best to begin by balancing elements that only appear in one formula on each side of the equation. Only balance equations by using coefficients; NEVER change the subscripts in a chemical formula that you already know is correct. 5. Check each atom or polyatomic ion to be sure that they are equal on both sides of the equation. 6. Make sure that all coefficients are in the lowest possible ratio. If necessary, reduce to the lowest ratio. *remember HOFBrINCl is diatomic

boiling points & bonding types

a comparison of boiling points is essentially equivalent to comparing the strengths of the attractive intermolecular forces exhibited by the individual molecules (for a substance to enter the gas phase, its particles must completely overcome the intermolecular forces holding them together); for small molecular compounds, London dispersion forces are the weakest intermolecular forces, dipole-dipole forces are somewhat stronger, and hydrogen bonding is a particularly strong form of dipole-dipole interaction; however, when the mass of a nonpolar molecule is sufficiently large, its dispersion forces can be stronger than the dipole-dipole forces in a lighter polar molecule (nonpolar Cl₂ has a higher boiling point than polar HCl)

covalent network solid

a compound in which all of the atoms are connected to one another by covalent bonds (ex. diamond is composed entirely of carbon atoms bonded in a tetrahedral geometry; does not melt at all and instead vaporizes to a gas at temperatures above 3500°C); because this type of solid is not composed of discrete molecules, melting a covalent network solid cannot be accomplished by overcoming relatively weak intermolecular forces (instead, covalent bonds must be broken, a process which requires extremely high temperatures); this covalent substance behaves quite differently than most molecular substances

hydrate

a compound that has one or more water molecules bound to each formula unit; many ionic compounds naturally contain water as part of the crystal lattice structure (those that contain a transition metal are highly colored); common for the hydrated form of a compound to be a different color than the anhydrous form (no water in its structure); hydrate can be converted to its anhydrous form by heating; formulas for hydrates have the formula for water set apart at the end with a dot, preceded by a coefficient that represents the number of water molecules per formula unit (ex: CoCl₂•6 H₂O - cobalt (II) chloride hexahydrate)

mole ratio

a conversion factor that relates the amounts in moles of any two substances in a chemical reaction; numbers in the conversion factor come from the coefficients of the balanced chemical equation in mole ratio problem, the given amount, expressed in moles, is written first; then the appropriate conversion factor is chosen in order to convert from moles of the given substance to moles of the unknown substance Given substance (mol) x Mole ratio = Unknown substance (mol) *a mole ratio represents exact quantities, since the coefficients in a balanced equation are considered to have an infinite number of sig figs *because of Avogadro's hypothesis, mole ratios between substances in a gas-phase reaction are also volume ratios mole ratio tells the quantitative relationship between reactants and products under ideal conditions, in which all reactants are completely converted into products (however, in reality, reactions may not 100% be completed or may result in side reaction, leading to different products too); theoretical stoichiometric calculations allow chemists to know maximum possible amount of product that can be generated by a reaction from a given amount of each reactant

double covalent bond

a covalent bond formed by atoms that share two pairs of electrons; (ex: occurs between the two carbon atoms in ethene C₂H₄); typically stronger and shorter than single covalent bond

coordinate covalent bonds

a covalent bond in which one of the atoms contributes both of the electrons in a shared pair; (ex: carbon monoxide CO - double bond forms between the two atoms with each atom providing one of the electrons to each bond {but this makes the carbon atom only have 6 electrons which is unstable}; so the oxygen atom contributes one of its lone pairs in order to make a third bond with the carbon atom; correctly represented by a triple covalent bond between the carbon and oxygen atoms); once formed, it is the same as any "conventional" bond

instantaneous dipole-induced dipole

a dipole with spontaneous uneven electron distribution (instantenous); this weak and temporary dipole can subsequently influence neighboring atoms through electrostatic attraction and repulsion, forming a induced dipole; the instantaneous and induced dipoles are weakly attracted to one another; the strength of dispersion forces increases as the total number of electrons in the atoms or nonpolar molecules increases

real gas

a gas that does not behave according to the assumptions of the kinetic-molecular theory; a real gas deviates most from an ideal gas at low temperatures & high pressures (high pressure forces gas molecules closer together, diminishing the empty space between particles & making the assumption that volume of particles is negligible less valid) (cooler temperature decreases kinetic energy causing them to slow down & make the attractive forces between them more prominent; continued cooling will eventually turn the gas into a liquid); gases are most ideal at high temperatures and low pressure; gases with weak attractive forces are more ideal than those with strong attractive forces

vapor pressure curve

a graph of vapor pressure as a function of temperature; boiling points of various liquids can be illustrated by these; to find the normal boiling point of a liquid, a horizontal line is drawn at standard pressure

phase diagram

a graph showing the conditions of temperature and pressure under which a substance exists in the solid, liquid, & gas phases; relationships among solid, liquid, & vapor (gas) states of a substance can be shown as a function of temperature and pressure in a single diagram; in each of the three colored regions of the diagram, the substance is in a single state/phase; dark lines that act as the boundary between those regions represent the conditions under which the two phases are in equilibrium find the X on the pressure axis & presume that the value of X is the standard pressure of 1 atm; as one moves left to right across the red line, the temperature of the solid substance is being increased while the pressure remains constant; when point A is reached, substance melts; temperature B on the horizontal axis represents the normal melting point of the substance; moving farther to the right, the substance boils at point Y, so point C on the horizontal axis represents the normal boiling point of the substance; as the temperature increases at a constant pressure, the substance changes from solid to liquid to gas starting right above point B on the temperature axis & following the red line vertically; at very low pressure, the particles of the substance are far apart from one another & the substance is in the gas state; as pressure is increased, the particles of the substance are forced closer & closer together; eventually, the particles are pushed so close together that attractive forces cause the substance to condense into the liquid state; continually increasing the pressure on the liquid will eventually cause the substance to solidify; for majority of substances, the solid state is denser than liquid state, so putting a liquid under great pressure will cause it to turn into a solid line segment R-S represents the process of sublimation, where substance changes directly from a solid to a gas (at a sufficiently low pressure, the liquid phase does not exist)

compressibility

a measure of how much a given volume of matter decreases when placed under pressure; gases can be compressed so that a relatively large amount can be forced into a small container; gases are compressible because most of the volume of a gas is composed of the large amounts of empty space between the gas particles (compressing forces the gas particles closer together)

surface tension

a measure of the elastic force in a liquid's surface; molecules within a liquid are pulled equally in all directions by intermolecular forces; molecules at the surface are pulled downward & sideways by other liquid molecules (not upward away from the surface); effect is that the surface molecules are pulled into the liquid, creating a surface that is tightened like a film; liquids that have strong intermolecular forces (like the hydrogen bonding in water) exhibit the greatest surface tension; surface tension allows objects denser than water to float on its surface; also responsible for the beading up of water droplets on wax (no attractions between the polar water molecules & nonpolar wax)

polar molecule

a molecule in which one end of the molecule is slightly positive, while the other end is slightly negative; diatomic molecules that consist of a polar covalent bond (like HF) is a polar molecule; the two electrically charged regions on either end of the molecule are called poles; when placed between oppositely charged plates, polar molecules orient themselves so that their positive ends are closer to the negative plate and their negative ends are closer to the positive plate (the molecules orient themselves to maximize the attraction between opposite charges); experimental techniques involving electric fields can be used to determine if a certain substance is composed of polar molecules & to measure the degree of polarity; for molecules with more than two atoms, the molecular geometry must also be taken into account when determining if the molecule is polar or nonpolar (ex: CO₂-linear molecule, 2 individual dipoles pointing outward from the C atom to each O atom, but since the dipoles are of equal strength and oriented linearly, they cancel each other out, and the overall molecular polarity of CO₂ is zero) (ex: H₂O-bent molecule, 2 individual dipoles point from the H atoms toward the O atom, because of the shape, the dipoles do not cancel each other out, and the water molecule is polar)

dipole

a molecule with two poles; (ex: HF); this results from the unequal distribution of electron density throughout a molecule

lone pair

a pair of electrons in a Lewis electron-dot structure that is not shared between atoms; (ex: diatomic fluorine molecule F₂ contains a single shared pair of electrons and three pairs of electrons that are not shared with the other atom); a nonbonding pair

chemical equation

a representation of a chemical reaction that displays the reactants and products with chemical formulas (CH₄ +O₂ -> CO₂ +H₂O) chemical equations must be balanced in order to satisfy the law of conservation of mass

coefficient

a small whole number placed in front of a formula in an equation in order to balance it

crystal

a substance in which the particles are arranged in an orderly, repeating, 3D pattern; majority of solids are crystalline in nature; particles of a solid crystal may be ions, atoms, or molecules, depending on type of substance; different arrangements of the particles within a crystal cause them to adopt several different shapes crystals are classified into general categories based on their shapes; crystal is defined by its faces, which intersect with one another at specific angles, which are characteristic of the given substance; edge lengths of a crystal are represented by the letters a, b, and c; angles at which the faces intersect are represented by the Greek letters α, β, and γ; each of the seven crystal systems differs in terms of the angles between the faces and in the number of edges of equal length on each face

fluid

a substance that is capable of flowing from one place to another & takes the shape of its container

triclinic

a ≠ b ≠ c α ≠ β ≠ γ ≠ 90° ex: microcline

cubic

a=b=c α = β = γ = 90° ex: pyrite

rhombohedral

a=b=c α = β = γ ≠ 90° ex: calcite

tetragonal

a=b≠c α = β = γ = 90° ex: wulfenite

Volume (V)

affects gas pressure; decreased volume = increased pressure; increased volume = decreased pressue

Temperature (T)

affects pressure; added kinetic energy by heat causes gas molecules to move faster and increase of pressure; increased temperature = increased pressure; decreased temperature = decreased pressure

(AB₆) octahedral

all of the atoms surrounding the central atom are equivalent in distance from the center; a surface covering the molecule would have 8 sides; all of those angles are 90° in an octahedral molecule, with the exception of the atoms that are directly opposite one another

ionic compounds

also crystalline in nature but with alternating cations and anions held together by attractive electrostatic forces (ionic bonds)

(AB₅E) square pyramidal

any one of the peripheral atoms in the octahedral domain geometry can be replace by the lone pair since they are all equivalent, resulting in a surface covering a square pyramidal molecule (four sided pyramid on a flat base); (ex: BrF₅)

excess reactant/excess reagent

any reactant that cannot be completely consumed by the full reaction of the limiting reactant; there is always excess reactant left over after the reaction is complete

single bonds

are always sigma bonds

double bonds

are comprised of one sigma and one pi bond

triple bonds

are comprised of one sigma bond and two pi bonds

average kinetic energy

at any given temperature, particles display wide range of kinetic energies (most near the middle of the range, but some a great deal lower/higher than average); as temperature increases, range of kinetic energies increases & distribution curve "flattens out;" at a given temperature, particles of any substance have same average kinetic energy, no matter what form (liquid, solid, gas); as sample of matter is continually cooled, average kinetic energy of its particles decreases (eventually, one would expect particles stop moving completely); Kelvin temperature of substance is directly proportional to average kinetic energy of particles of substance

hybrid orbitals

atomic orbitals obtained when two or more nonequivalent orbitals from the same atom combine in preparation for bond formation; the electron domain geometry predicted by VSEPR leads directly to the type of hybrid orbitals that must be formed to accommodate that geometry

expanded octets

atoms of elements in the 2nd period cannot have more than 8 valence electrons around the central atom; but elements of the 3rd period & beyond are capable of exceeding the octet rule (starting with the 3rd period, the d sublevel becomes available, so it is possible to use these orbitals in bonding, resulting in more than 8 electrons around the central atom); (ex: phosphorus & sulfur react with halogens to make stable compound sin which the central atom has an expanded octet; phosphorus pentachloride {central phosphorus atom has 10 valence electrons}; sulfur hexafluoride {central sulfur atom has 12 valence electrons})

intermolecular forces

attractions that occur between molecules; weaker than either ionic or covalent bonds; the varying strengths of different types of intermolecular forces are responsible for physical properties of molecular compounds such as melting and boiling points

orthorhombic

a≠b≠c α = β = γ = 90° aragonite

monoclinic

a≠b≠c α ≠ 90° = β = γ ex: azurite

heating curve

change of state behavior for any substance can be represented with a heating curve; melting and boiling points of substance can be determined by horizontal plateaus on the curve (other substances would have melting & boiling points different from those of water); exception to this exact form for a heating curve would be for a substance such as carbon dioxide, which sublimes rather than melts at standard pressure (it would only have one plateau in its heating curve, at its sublimation temperature) experiment: block of ice at -30°C, well below its melting point in closed container; as heat is steadily added to the ice block, the water molecules will begin to vibrate faster as they absorb kinetic energy; eventually ice warms to 0°C, the added energy will start to break apart the hydrogen bonding that keeps the water molecules in place in solid form; as ice melts, its temperature does not rise (all energy being put into ice goes into melting process & not into any increase in temperature); during melting, the two states - solid & liquid - are in equilibrium with one another; if system was isolated at this point & no energy was allowed to enter/leave, the ice-water mixture at 0°C would remain (temperature is always constant during a change of state) continued heating of water after ice has completely melted will increase kinetic energy of liquid molecules, causing temperature to once again begin to rise; assuming that atmospheric pressure is standard, the vaporization will rise steadily until reaching 100°C (at this point, the added energy from the heat will cause the liquid to begin to vaporize); temperature will remain steady at 100°C while intermolecular hydrogen bonds are being broken & water molecules pass from liquid to gas state (once all liquid has boiled, continued heating of the steam will increase its temperature above 100°C steam above 100°C could be steadily cooled down to 100°C, at which point it would condense to liquid water; water then could be cooled to 0°C, at which point continued cooling would freeze the water to ice; ice could then be cooled to some point before 0°C (would be diagrammed in cooling curve-reverse of heating curve)

condensation

change of state from a gas to a liquid; if water is kept in a closed container, the water vapor molecules cannot escape & the water level does not change (as some water molecules become vapor, an equal amount of water vapor molecules condense back into the liquid state)

deposition

change of state from a gas to a solid

sublimation

change of state from a solid to a gas without passing through the liquid state; solids also have a vapor pressure, though generally much less than that of a liquid; iodine & carbon dioxide sublime

5 types of chemical reactions

combination decomposition single-replacement double-replacement combustion *some reactions will fit into more than one category

condensation vs. vaporization

condensation is opposite of vaporization both represent equilibrium between liquid & gas states

solid

condensed state with particles that are far closer together than those of a gas (similar to liquid); not fluid (particles packed tightly together in an orderly arrangement); motion of individual atoms, ions, or molecules in a solid is restricted to vibrational motion about a fixed point; almost completely incompressible; densest of the three states of matter

covalent network crystals

consists of atoms at the lattice points of the crystal, with each atom being covalently bonded to its nearest neighbor atoms; 3D & contains a very large number of atoms; network solids include diamond, quartz, many metalloids, & oxides of transition metals & metalloids; network solids are hard & brittle; extremely high melting & boiling points; being composed of atoms, not ions, they do not conduct electricity well in any state

triple covalent bond

covalent bond formed by atoms that share three pairs of electrons; (ex: diatomic nitrogen gas N₂ - the three single unpaired electrons on each atom form three shared pairs of electrons; each nitrogen atom has one lone pair of electrons and six electrons that are shared with the other atom, so each atom obeys the octet rule); typically stronger and shorter than both single and double covalent bonds

classes of crystalline solids

crystalline substances can be described by the types of particles in them & the types of chemical bonding that take place between the particles 4 types: (1) ionic, (2) metallic, (3) covalent network, (4) molecule

Standard temperature and pressure (STP)

defined as 0°C (273.15 K) and 1 atm of pressure; at STP, one mole (6.02x10²³ representative particles) of any gas occupies a volume of 22.4 L; (this is needed because gas volume is affected by pressure {compressible as particles are forced closer together, empty space decreases and volume reduces with higher pressure} & temperature {heat makes the molecules move faster, expanding the gas}; gas volumes must be compared at the same temperature and pressure because of this variation)

pressure

defined as the force per unit area on a surface; Pressure = force/area

molar mass (g/mol)

defined as the mass of one mole of representative particles of a substance (atomic mass=the mass of one mole of that element in grams; relative scale of atomic masses in amu is also a relative scale of masses in grams); molar mass is measured in the units grams per mole (g/mol); mole is the amount of a substance that contains as many representative particles as the number of atoms is exactly 12 g or carbon-12 (exactly 12 g of carbon-12 contains one mole, or 6.02 x 10²³ atoms of carbon-12) {molar mass is the atomic mass on the periodic table FOR ELEMENTS); use molar masses rounded to the hundredths place; molar mass of an element is its atomic mass expressed in grams and is equal to the mass of one mole of atoms of that element; molar mass of a compound is the mass of one mole of representative particles of the compound; molar mass of any substance is the mass in grams of one mole of representative particles of that substance

gas density (g/L)

density of a gas at STP is dependent only on its molar mass (since gasses all occupy the same volume on a per mole basis); gas with a small molar mass will have a lower density than a gas with a large molar mass; can be calculated by combining molar mass and molar volume 1. know the molar mass 2. molar mass (g/mol) divided by molar volume (1 mol/22.4 L) yields the gas density (g/L) at STP 1. know the density 2. density (in g/L) multiplied by molar volume (22.4 L/mol) equals molar mass (g/mol)

deposition vs. sublimation

deposition is opposite of sublimation both represent equilibrium between solid & gas states

manometer/U-tube

device similar to barometer used to measure pressure of enclosed gas sample; pressure of gas in bulb determined by difference in height of mercury between two arms of U-tube

factors affecting gas pressure (PV = nRT)

elastic collisions with walls of container define gas pressure; 4 variables: PRESSURE (P), VOLUME (V), TEMPERATURE (T), AMOUNT OF GAS, as measured by the number of moles of gas particles (n)

hybridization of elements in the third period

elements in the third period and beyond are capable of expanding their octet to form molecules with either trigonal bipyramidal or octahedral electron domain geometries (previously unoccupied orbitals in the d sublevel of the central atom are involved in the hybridization process);

diatomic elements

elements whose natural form is of a diatomic molecule; seven (hydrogen, nitrogen, oxygen, fluorine, chlorine, bromine, and iodine); HOFBrINCl; by forming a diatomic molecule, both atoms in each of these molecules satisfy the octet rule, resulting in a structure that is much more stable than the isolated atoms

skeleton equation

equation that shows only the formulas of the reactants and products with nothing to indicate the relative amounts; 1st step to writing a chemical equation (making sure that the formulas of all substances involved are written correctly); also put the appropriate symbol in parentheses after each formula: (s) for solid, (l) for liquid), (g) for gas, & (aq) for aqueous solution

pure metals

exist as extended 3D structures of close packed metal cations with mobile valence electrons

sp³d hybridization

for electron domain geometries that are trigonal bipyramidal; electron promotion is from the 3s orbital to an empty 3d orbital; hybridization then occurs when a single s orbital, three p orbitals, and a single d orbital (set of five hyrbid orbitals are called sp³d hybrids); (ex: PCl₅); other molecular geometries derived from trigonal bipyramidal electron domain geometry (seesaw, T-shaped, linear) also display the same hybridization

sp³d² hybridization

for molecules with an octahedral electron domain geometry; a 3s and a 3p electron are each promoted to two empty 3d orbitals; hybridization now occurs with one s orbital, three p orbitals, and two d orbitals (resulting set of six equivalent orbitals are called sp³d² hybrids); (ex: SF₆); other molecular geometries derived from an octahedral electron domain geometry (square pyramidal, square planar) also exhibit this type of hybridization

intramolecular forces

forces that act within a molecule or crystal (i.e. covalent and ionic bonds)

anhydrous form

form of hydrate with no water in its structure; converted to this form by heating

structural formula

formula that shows the arrangement of atoms in a molecule and represents covalent bonds between atoms by dashes (ex: H-H)

freezing vs. melting

freezing is opposite of melting both represent the equilibrium between the solid & liquid states

sp hybridization

geometry of these sp hybrid orbitals is linear with the large lobes of the two orbitals pointing in opposite directions along one axis; remaining p orbitals do not hybridize and remain unoccupied; (ex: BeH₂, CO₂, C₂H₂)

sp² hybridization

geometry of these sp² hybrid orbitals is trigonal planar, with the large lobe of each orbital pointing toward one corner of an equilateral triangle; angle between these lobes is 120°; (ex: BF₃, O₃)

polyatomic ions

group of covalently bonded atoms that carries an overall electrical charge; when drawing the Lewis structure of a polyatomic ion, the charge of the ion is reflected in the total number of valence electrons in the structure; (ex: ammonium NH₄⁺ - formed when a hydrogen ion (H⁺) attaches to a lone pair of ammonia (NH₃) via a coordinate covalent bond; 1 N atom=5 valence electrons, 4 H atoms=4 valence electrons, subtract 1 electron to give ion overall charge of 1+, total of 8 valence electrons in the ion); customary to put the Lewis structure of a polyatomic ion into a large set of brackets, with the charge of the ion as a superscript outside the brackets; add/subtract electrons in the structure to satisfy the octet rule and give the ion its charge

critical temperature (Tc)

highest temperature at which the substance can possibly exist as a liquid; at 373.99°C, particles of water in the gas phase are moving very, very rapidly; at any temperature higher than that, the gas phase cannot be made to liquefy, no matter how much pressure is applied to the gas

hybridization in molecules w/ double/triple bonds

hybridization model helps explain molecules with double or triple bonds; in a double bond (one is a sigma bond, formed by a direct overlap of bonding orbitals, while the other is a pi bond, formed by a side-to-side overlap of unhybridized p orbitals); in a triple bond (one is a sigma bonds, while the other two are pi bonds)

ideal gas

imaginary gas whose behavior perfectly fits all the assumptions of the kinetic-molecular theory (in reality, gases are not ideal, but are very close to being so under most everyday conditions); ideal gas assumed to occupy insignificant volume compared to space between particles; gases are easy to study because since there is so much space between particles, intermolecular forces can largely be ignored, vastly simplifying any analysis of the motion exhibited by individual particles a gas that follows the gas laws at all conditions of temperature and pressure (occupies zero volume with no attractive forces whatsoever toward each other); not true in reality

incomplete octet

in some compounds, the number of electrons surrounding the central atom in a stable molecule is fewer than eight (ex: beryllium, which forms primarily molecular compounds when combined with many other elements; does not typically attain a full octet by sharing electrons since it only has two valence electrons)

(AB₄E₂) square planar

in the octahedral domain geometry, two of the atoms are replace by two lone pairs that occupy positions directly opposite each other; the four remaining atoms are in the same plane as the central atom and all angles are equal to 90° (ex: XeF₄)

(AB₂E₃) linear

in the trigonal bipyramidal domain geometry, all three equatorial atoms have been replaced by lone pairs (in addition to the two axial atoms); result is a linear ion (the central atom and the two axial atoms); (ex: I₃⁻)

(AB₃E₂) T-shaped

in the trigonal bipyramidal domain geometry, two of the 3 equatorial atoms is replaced by lone pairs (in addition to the 2 axial atoms); creates a T-shape (ex: ClF₃)

(AB₄E) distorted tetrahedron/seesaw

in the trigonal bipyramidal geometry, one of the 3 equatorial atoms is replaced by a lone pair because of the greater repulsion of a lone pair (in addition to the 2 axial atoms); creates a seesaw effect around the central atom (ex: SF₄)

measuring amount of matter

in three ways (by mass, by volume, by number of particles)

general phase diagram vs water phase diagram

in water's diagram, the slope of the line between the solid & liquid states is negative rather than positive (reason is that water is an unusual substance in that its solid state is less dense than the liquid state - ice floats in liquid water); a pressure change has the opposite effect on those two phases (if ice is relatively near its melting point, it can be changed into liquid water by the application of pressure - water molecules are actually closer together in the liquid phase than they are in the solid phase)

perspective drawing

indicated the 3D character of shapes of some molecules; dotted line bonds should be visualized as receding into the page, while the solid triangle bonds should be visualized as coming out of the page

barometer

instrument used to measure atmospheric pressure; traditional mercury barometer is an evacuated tube immersed in a container of mercury (air molecules from atmosphere push down on outer surface of mercury, but because tube is a vacuum, there is no corresponding downward push on mercury in tube); mercury rises inside the tube (height to which the mercury rises ids dependent on external air pressure)

hydrogen bond

intermolecular attractive force in which a hydrogen atom, that is covalently bonded to a small, highly electronegative atom, is attracted to a lone pair of electrons on an atom in a neighboring molecule (in water, occurs between hydrogen atom and the other's oxgyen's lone pair); very strong compared to other dipole-dipole interactions; about 5% as strong as a covalent bond; attractive force between water molecules is an unusually strong type of dipole-dipole interaction (water contains hydrogen atoms that are bound to a highly electronegative oxygen atom, making for very polar bonds; the partially positive hydrogen atom of one molecule is then attracted to the oxygen atom of a nearby water molecule); occurs only in molecules where hydrogen is covalently bonded to one of three elements: fluorine, oxygen, or nitrogen (these three elements are so electronegative that they withdraw the majority of the electron density from the covalent bond with hydrogen, leaving the H atom very electron-deficient; because the hydrogen atom does not have any electrons other than the ones in the covalent bond, its positively charged nucleus is almost completely exposed, allowing strong attractions to other nearby lone pairs); the hydrogen bonding in water leads to unusual, important properties (liquid due to strong hydrogen bonds even though most molecular compounds with a similar mass are gases; multiple hydrogen bonds occur simultaneously in water because of its bent shape and the presence of 2 hydrogen atoms per molecule; as a liquid, hydrogen bonds of water break and reform but when water is cooled the molecules begin to slow down until the water is frozen to ice and the hydrogen bonds become more rigid and form a network; bent shape of molecules leads to gaps in hydrogen bonding network of ice giving ice the property of being less dance than the liquid state; hydrogen bonds also play an important biological role in the physical structures of proteins & nucleic acids)

London dispersion forces

intermolecular forces that occur between all atoms and molecules due to the random motion of electrons; also considered a type of van der Waals force and are the weakest of all intermolecular forces; the dispersion forces are strongest for iodine molecules because they have the greatest number of electrons (relatively stronger forces result in melting and boiling points which are highest of the halogen group and can hold iodine molecules close together in the solid state at room temperature); dispersion forces are progressively weaker for bromine (liquid), chlorine (gas), fluorine (gas), as illustrated by their steadily lower melting and boiling points; because gaseous molecules are so far apart from one another, intermolecular forces are nearly nonexistent in the gas state, and so the dispersion forces in chlorine and fluorine only become measurable as the temperature decreases and they condense into the liquid state

critical point

intersection point of the critical temperature and the critical pressure

ionic crystals

ionic crystal structure consists of alternating positively charged cations & negatively charged anions; ions may either be monatomic or polyatomic; generally, ionic crystals form from a combination of metal cations & Group 16/17 nonmetal anions, although nonmetallic polyatomic ions are also common components of ionic crystals; hard; brittle; high melting points; do not conduct electricity as solids, but conducts when molten/dissolved

hybridization-electron domain geometry matching

linear - sp (2 hybrid orbitals) trigonal planar {bent} - sp² ( 3 hybrid orbitals) tetrahedral {trigonal pyramidal, bent} - sp³ (4 hybrid orbitals) trigonal bipyramidal {seesaw, T-shaped, linear} - sp³d (5 hyrbid orbitals) octahedral {square pyramidal, square planar} - sp³d² (6 hybrid orbitals)

activity series

list of elements in decreasing order of their reactivity; for a single-replacement reaction, given element is capable of replacing an element below it in the activity series *can be used to predict if a reaction will occur

Mass to Moles

mass of given → moles of given → moles of unknown 1. mass of given substance converted into moles by using molar mass from periodic table 2. moles of given substance converted into moles of unknown using mole ratio from balanced equation

Mass to Mass

mass of given → moles of given → moles of unknown → mass of unknown 1. mass of given substance converted into moles by using molar mass from periodic table 2. moles of given substance converted into moles of unknown using mole ratio from balanced equation 3. moles of unknown converted into mass (in grams) by using its molar mass from periodic table

Mass to Volume & Volume to Mass

mass of given → moles of given → moles of unknown → volume of unknown//volume of given → moles of given → moles of unknown → mass of unknown 1. convert mass of given to moles using periodic table 2. apply mole ratio to convert moles of given to moles of unknown 3. convert moles of unknown to liters using molar volume of a gas (22.4 L/mol) --- 1. convert volume of given to moles using molar volume of a gas 1 mol/22.4 L) 2. use mole ratio to convert from moles of given to moles of unknown 3. convert moles of unknown to grams using periodic table

vapor pressure

measure of the pressure exerted by the vapor that forms above its liquid form in a sealed container; considered as a property of the liquid & is constant for a given substance at a set temperature; vapor pressure of substance at given temperature is based on the strength of its intermolecular forces (liquid with weak forces evaporates more easily & has high vapor pressure; liquid with stronger forces does not evaporate easily & has lower vapor pressure); vapor pressure can be measured with a manometer dynamic equilibrium between the liquid & vapor phases (when partially filled container of liquid is sealed, some liquid molecules at the surface evaporate into the vapor phase; vapor molecules cannot escape from the container though; over time, some of the molecules lose energy through collisions with other molecules or with walls of container; at this point, the less energetic vapor molecules are trapped by the attractive forces of the molecules in the liquid & begin to condense back into liquid form; eventually system reaches a point where the rate of evaporation is equal to rate of condensation); because they cannot escape the container, the vapor molecules above the surface of the liquid exert a pressure on the walls of the container

Amount of Gas (n)

measured by the number of moles of gas particles; affects gas pressure; increased number of gas particles = increased pressure; decreased number of particles = decreased pressure; gas will move from area of HIGHER PRESSURE to area of LOWER PRESSURE until the pressures are equal

metallic crystals

metal cations surrounded by a "sea" of mobile valence electrons; delocalized electrons do not belong to any one atom & are capable of moving through entire crystal; good conductors of electricity; can have wide range of melting points

mmHg

millimeter of mercury; unit of gas pressure; equivalent to the torr (Torricelli)

hybridization

mixing of the atomic orbitals in an atom to produce a set of hybrid orbitals; result of the mixing of nonequivalent orbitals (s & p can hybridize, but p & p cannot hybridize); (ex: in CH₄, we promote one of the 2s electrons to the empty 2p orbital to make four bonds possible; to make the molecular geometry {tetrahedral} correct, hybridization must occur as the single 2s orbital hybridizes with the three 2p orbitals to form a set of four hybrid orbitals, called sp³ hybrids, which are all equivalent to one another); (ex: in NH₃, same thing occurs as in methane, but one of the hybrid sp³ orbitals already contains a pair of electrons, leaving only three half-filled orbitals available to form covalent bonds with three hydrogen atoms making a trigonal pyramidal)

mole road map

mole is at the center of any calculation involving the amount of a substance; outlines various possible conversions between mass, number of representative particles, and gas volume

diatomic molecule

molecule that contains exactly two atoms; example is hydrogen (H₂)

(AB₃) trigonal planar

molecule where atoms are positioned at the vertices of an equilateral triangle surrounding the central atom (ex: BF₃ with the F-B-F angle at 120° & all 4 atoms lying on the same plane)

(AB₄) tetrahedral

molecule where each of the 4 atoms lies at the corners of a tetrahedron surrounding the central atom (each face of a tetrahedron is an equilateral triangle & the atoms do not lie in the same plane); (ex: methane CH₄ with the H-C-H bond angles at 109.5°, which is larger than the 90° that they would be if the molecule was planar);

(AB₂) linear molecule

molecule where the atoms will arrange themselves on directly opposite sides of the central atom (ex: H-Be-H for BeH₂); the H-Be-H bond angle is 180° because of its linear geometry; (ex: O=C=O for CO₂)

(AB₅) trigonal bipyramidal

molecule where three of the atoms lie in a plane (w/ bond angles of 120°; trigonal planar arrangement; referred to as the equatorial atoms) & the other two atoms are oriented exactly perpendicular to the plane formed by the central atom and the equatorial atoms (called the axial atoms; form a vertical axis w/ central atom; 90° angle between axial bonds and equatorial bonds); (ex: phosphorus pentachloride PCl₅); surface covering the molecule would take the shape of two 3-sided pyramids pointing in opposite directions

(AB₂E) bent molecular geometry

molecules with three electron pairs have a domain geometry that is trigonal planar, but the lone pair on the central atom repels the electrons in the two bonds, causing the atom to adopt a bent molecular geometry; (ex: ozone O₃ w/ O-O-O angle slightly less than 120°); within the context of the VSEPR model, lone pairs of electrons can be considered to be slightly more repulsive than bonding pairs of electrons, due to their closer proximity to the central atom (lone pairs take up more space)

Moles to Mass

moles of given → moles of unknown → mass of unknown 1. moles of given converted to moles of unknown using mole ratio from balanced equation 2. moles of unknown converted into mass (in grams) by using its molar mass from periodic table

noble gases

monatomic and exist as individual atoms

aneroid barometer

more convenient & measures pressure by expansion & contraction of small spring within an evacuated metal capsule

energy in bond formation

nature favors chemical bonding because most atoms attain a lower potential energy when they are bonded to other atoms than when they are isolated; two hydrogen atoms separated by a distance large enough to prevent interaction between them has the potential energy equal to zero; as the atoms approach one another, their electron clouds gradually overlap and interactions begin (single electrons possessed by each hydrogen atom begin to repel each other, causing an increase in potential energy; however, attractive forces begin to develop between each electron and the positively charged nucleus of the other atom, causing a decrease in potential energy); at first, the attractive force is stronger than the repulsive force, so potential energy decreases and system becomes more stable (as the two atoms move closer, the potential energy continues to decrease until a position is reached where the potential energy is at its lowest possible point); if the hydrogen atoms move any closer together, the repulsive force between the two positively charged nuclei (3rd interaction) begins to dominate (resulting repulsive force is very strong & sharp rise in potential energy); the point at which the potential energy reaches its minimum represents the ideal distance between hydrogen atoms for a stable covalent bond to occur

Avogadro's number

number of representative particles in a mole; 6.02 x 10²³; an experimentally determined number; a mole of any substance contains Avogadro's number (6.02 x 10²³) of representative elements; mole is the amount of a substance that contains as many representative particles as the number of atoms is exactly 12 g or carbon-12 (exactly 12 g of carbon-12 contains one mole, or 6.02 x 10²³ atoms of carbon-12)

triple point

only temperature/pressure pairing at which the solid, liquid, & vapor states of a substance can all coexist at equilibrium; point TP

valence shell

outermost occupied shell of electrons in an atom; this shell holds the valence electrons, which are the electrons that are involved in bonding and shown in a Lewis structure

liquids

particles of a liquid much closer together than of a gas (very little empty space between them); essentially not compressible; far denser; intermolecular attractive forces are the only thing that keeps the particles close together (in contrast, according to the kinetic-molecular theory for gases, any attractive forces between the particles of a gas are so minor that they can usually be ignored); liquids & solids are condensed states of matter; liquids & gases are both fluids

percent composition

percent by mass of each element in a compound % by mass = (mass of element/mass of compound) × 100% percent composition also used to determine the mass of a certain element that is contained in a sample whose total mass is known (by using conversion factors based on the percent by mass of each element)

molecular vs ionic compounds

physical state and properties of a particular compound depend largely on the type of chemical bonding it displays; molecular compounds display a wide range of physical properties due to the different types of intermolecular attractions; melting and boiling points of molecular compounds are generally quite low compared to those of ionic compounds (because the energy required to disrupt the intermolecular forces between molecules is far less than the energy required to break ionic bonds in crystalline ionic compounds); since molecular compounds are composed of neutral molecules, their electrical conductivity is generally quite poor, whether in solid or liquid state; ionic compounds do not conduct electricity well in the solid state because of their rigid structure, but conduct well when molten or dissolved in water (liquid state); water solubility of molecular compounds is variable and depends primarily on the type of intermolecular forces involved (substances that exhibit hydrogen bonding or dipole-dipole forces are generally water soluble, whereas those that exhibit only London dispersion forces are generally insoluble); most, but not all, ionic compounds are quite soluble in water; molecular compounds can be gases, liquids, or solids at room temperature while ionic compounds are solids at room temperature; the representative unit of a molecular compound is a molecule while the ionic compound is a formula unit; molecular compounds are made up of nonmetals only with a covalent (sharing of pairs of electrons between atoms) bond while ionic compounds are made up of metals and nonmetals with an ionic (transfer of electrons between atoms) bond

atmospheric pressure

pressure exerted by gas particles in Earth's atmosphere as those particles collide with objects; at sea level, atmospheric pressure reported as 760 mmHg; at higher altitudes, atmospheric pressure is lower & column of mercury will not rise as high; atmospheric pressure also slightly dependent on weather conditions

critical pressure (Pc)

pressure that must be applied to the gas the critical temperature in order to turn it into a liquid; for water, the critical pressure is very high - 217.75 atm

gas pressure

pressure that results from collisions of gas particles with an object; ex: outward pressure exerted by gas in balloon keeps it inflated

elemental analysis: determining the empirical formula of a compound

procedure where an unknown compound can be analyzed in the laboratory to determine the percentages of each element contained within it; these values can be used to find the molar ratios of the elements, which gives us the empirical formula 1. Assume a 100 g sample of the compound so that the given percentages can be directly converted into grams. 2. Use each element's molar mass to convert the grams of each element to moles. 3. In order to find the whole-number ratio, divide the moles of each element by the smallest value obtained in step 2. 4. If all the values at this point are whole numbers (or very close), each number is equal to the subscript of the corresponding element in the empirical formula. 5. In some cases, one or more of the values calculated in step 3 will not be whole numbers. Multiply each of them by the smallest number that will convert all values into whole numbers. (All values must be multiplied by the same number so that the relative ratios are not changed.) These values can then be used to write the empirical formula.

percent yield

ratio of the actual yield to the theoretical yield, expressed as a percentage; measure of the efficiency of the reaction; this measurement is used by chemists to indicate how successful a reaction has been (chemical reactions in experiments usually have many different things, like experimental errors, incomplete reactions, and undesirable side reactions, that contribute to the formation of less product than would be predicted) Percent Yield=Actual Yield/Theoretical Yield x 100% percent yield is very important in the manufacture of products (one step with a low percent yield can cause a large waste of reactants and money); percent yields greater than 100% are possible if the measured product of the reaction contains impurities that add to the mass of the pure product (when a chemist synthesizes a desired chemical, he/she must always be careful to purify the products of the reaction)

limiting reactant/limiting reagent

reactant that determines the amount of product that can be formed in a chemical reaction; reaction proceeds until the limiting reactant is completely used up

combination (synthesis) reaction

reaction in which 2+ substances combine to form a single new substance (A+B→AB); ex: two elements combining to form a compound (2Na+Cl₂→2NaCl); ex: metal/nonmetal reacting with oxygen to form an oxide (2Mg+O₂→2MgO); ex: nonmetals often can combine in different ratios to produce different products/molecular compounds (S+O₂→SO₂, 2S+3O₂→2SO₃); ex: transition metals are capable of adopting multiple positive charges within their ionic compounds and can form different products in combination reactions (2Fe+O₂→2FeO, 4Fe+3O₂→2Fe₂O₃); ex: element reacts with compound to form new compound composed of larger number of atoms (2CO+O₂→2CO₂); ex: two compounds may also react to form a more complex compound, such as oxides with water (CaO+H₂O→Ca(OH)₂, SO₃+H₂O→H₂SO₄)

decomposition reaction

reaction in which a compound breaks down into 2+ simpler substances (AB→A+B); requires an input of energy in form of heat, light, or electricity; ex: binary compound decomposes into its elements (2HgO→2Hg+O₂); ex: metal carbonate decomposes into metal oxide & carbon dioxide gas (CaCO₃→CaO+CO₂); ex: metal hydroxide decompose to yield metal oxides and water (2NaOH→Na₂O+H₂O); ex: unstable acid decomposes to produce nonmetal oxides and water (H₂CO₃→CO₂+H₂O) *reaction still a decomposition one even when 1+ products is still a compound

combustion reaction

reaction in which a substance reacts with oxygen gas, releasing energy in form of light & heat *must involve O₂ as one reactant (ex: combustion of hydrogen gas producing water vapor 2H₂+O₂→2H₂O) *many combustion reactions occur with a hydrocarbon, a compound made up solely of carbon & hydrogen, producing carbon dioxide & water (ex: C₃H₈+5O₂→3CO₂+4H₂O) *products of combustion of an alcohol are carbon dioxide & water *water produced in combustion reaction is in gaseous form

single-replacement (single-displacement) reaction

reaction in which one element replaces a similar element in a compound (metals (A,B): A+BC→AC+B) (nonmetals (Y,Z): Y+XZ→XY+Z) *metals must replace metals & nonmetals must replace nonmetals *only occurs when element doing the replacing is more reactive than the element that is being replaced 1. metal replacement: more reactive metal replaces less reactive metal (ex: Mg+Cu(NO₃)₂→Mg(NO₃)₂+Cu) 2. hydrogen replacement: hydrogen in acid replaced by active metal; metals react easily with acids producing hydrogen gas and compound (ex: Zn+2HCl→ZnCl₂+H₂); some metals (group 1) are so reactive that they are capable of replacing hydrogen in water, producing metal hydroxide and hydrogen gas (ex: 2Na+2H₂O→2NaOH+H₂) 3. halogen replacement: more reactive halogen replaces less reactive halogen (ex: Cl₂+2NaBr→2NaCl+Br₂)

double-replacement (double-displacement) reaction

reaction in which the positive and negative ions of two ionic compounds exchange places to form two new compounds (AB+CD→AD+CB, where A & C are positive cations and B & D are negative anions) *generally occur between substances in aqueous solution *one of products is usually a solid precipitate, a gas, or a molecular compound like water 1. formation of precipitate: precipitate forms in double-replacement reaction when cations from one of the reactants combine with the anions from the other reactant to form an insoluble ionic compound (ex: 2Kl+Pb(NO₃)₂→2KNO₃+PbI₂; very strong attractive forces between Pb²⁺ & I⁻ ions result in precipitate) 2. formation of gas: some double-replacement reactions produce a gaseous product which bubbles out of the solution, escaping into the air (ex: Na₂S+2HCl→2NaCl+H₂S) 3. formation of molecular compound: double-replacement reaction that produces a molecular compound (usually water) as one of its products (ex: HCl+NaOH→NaCl+H₂O) *occasionally, reaction will produce both a gas & molecular compound (ex: Na₂CO₃+2HCl→2NaCl+CO₂+H₂O)

empirical formulas

show the lowest whole-number ratio of the elements in ionic and molecular compounds; subscripts in a formula represent the molar ratio of the elements in that compound

amorphous solid

solid that lacks an ordered internal structure, unlike a crystalline solid; some examples include rubber, plastic, & gels; glass is a very important amorphous solid that is made by cooling a mixture of materials in such a way that it does not crystallize (supercooled liquid rather than a solid); amorphous solids do not have a distinct melting point like crystalline solids do (as heated, slowly softens and can be shaped); glass object shatters in very irregular way (crystalline solids on other hand breaks along specific planes as dictated by its crystal system)

multiple covalent bond

some molecules are not able to satisfy the octet rule by making only single covalent bonds between the atoms; (ex: ethene C₂H₄ - carbon atoms are bonded together & each carbon is also bonded to two hydrogen atoms; if the LEwis electron dot structure were drawn with a single bond between the carbon atoms and lone pairs added until satisfying the octet rule, it would be incorrect {containing 14 valence electrons & the atoms in this molecule only have a total of 12 valence electrons available}; can be corrected by eliminating one lone pair and moving another lone pair to a bonding position {two carbon atoms share two pairs of electrons instead of just one}

atmosphere (atm)

standard atmospheric pressure is 1 atm; equal to 760 mmHg & 101.3 kPa

pascal (Pa)

standard unit of pressure; one pascal is a very small amount of pressure; more useful unit of everyday gas pressures is kilopascal (kPa)

VSEPR (valence-shell electron pair repulsion model)

states that a molecule will adjust its shape so that the valence electron pairs stay as far apart from each other as possible; makes sense based on the fact that negatively charged electrons repel one another; for the purposes of this model, a double or triple bond is no different in terms of repulsion that a single bond; within the context of the VSEPR model, lone pairs of electrons can be considered to be slightly more repulsive than bonding pairs of electrons, due to their closer proximity to the central atom (lone pairs take up more space)

valence bond theory

states that covalent bonds are formed by the overlap of partially filled atomic orbitals; overlapping orbitals do not have to be of the same type (ex. HF: 1s orbital of hydrogen atom overlaps with 2p orbital of fluorine atom); covalent bond forms when the electron clouds of two atoms overlap with each other (the electron in one atom becomes attracted to the nucleus of the other in the molecule as the atoms come closer until reaching an optimum distance between the two nuclei, equal to the bond length, and the potential energy reaches a minimum forming a stable covalent bond); # of bonds = # of unpaired electrons; however, in some molecules, one must promote one of the electrons from a lower orbital to a higher one to have enough unpaired electrons to form the number of bonds required (this "costs" a small amount of energy, but the process of bond formation is accompanied by a decrease in energy; thus, there is a lower overall energy and greater stability for the molecule)

Avogadro's hypothesis

states that equal volumes of all gases at the same temperature and pressure contain equal numbers of particles (since the total volume that a gas occupies is primarily composed of the empty space between the particles, the actual size of the particles themselves is nearly negligible); given volume of a gas with small light particles (H₂) contains same number of particles as the same volume of a heavy gas with larger particles (SF₆);

Lewis electron-dot structures

structures of molecules that are held together by covalent bonds can be diagrammed by these; (ex: hydrogen molecule - the shared pair of electrons is shown as two dots in between the two H symbols {H:H}; only VALENCE electrons are portrayed

molecular compounds

take the form of individual molecules; generally comprised of 2+ nonmetal atoms; molecular formula shows the quantity of each atom that occurs in a single molecule of that compound

molecular formulas

tell how many atoms of each element are present in one molecule of a molecular compound (many times, molecular formula is same as empirical formula); sometimes molecular formula is a simple whole-number multiple of the empirical formula; to determine a compound's molecular formula, it is necessary to know the molar mass of the compound (by mass spectrometer)

melting point

temperature at which a solid changes into a liquid; as a solid is heated, the avg kinetic energy of its particles still increases, but due to their relatively fixed positions, this manifests itself as stronger and more rapid vibrations; eventually, organization of the particles within the solid structure begins to break down and the solid starts to melt; at melting point, the disruptive vibrations of the particles in solid overcome the attractive forces operating within the solid; melting point dependent on strength of attractive forces; melting point of solid is same as freezing point of corresponding liquid (at this temperature, the solid & liquid states of the substance are in equilibrium) -- for water, the equilibrium occurs at 0°C

absolute zero

temperature at which motion of particles particles theoretically ceases; absolute zero = 0 K (Kelvin temperature scale based on this theoretical limit); Kelvin temperature of substance is directly proportional to average kinetic energy of particles of substance

normal boiling point

temperature at which the vapor pressure of the liquid is equal to the standard atmospheric pressure of 760 mmHg; because atmospheric pressure can change based on location, the boiling point of a liquid changes with the external pressure; normal boiling point is constant because it is defined relative to a standard pressure (760 mmHg, 1 atm, or 101.3 kPa)

mole (mol)

the amount of a substance that contains 6.02 x 10²³ representative particles of a substance; the SI unit for amount of a substance; use conversion factors to convert between number of particles & moles of those particles (use sig figs)

actual yield

the amount of product that is actually formed when the reaction is carried out in the laboratory

dipole-dipole forces

the attractive forces that occur between polar molecules; result from the attraction between the positive end of one dipole and the negative end of a neighboring dipole; involve only partial charges (much weaker than ionic bonds, although they are similar)

stoichiometry

the calculation of amounts of substances in a chemical reaction from the balanced equation stoichiometry problems characterized by (1) the information given & (2) the information to be solved for, the unknown; amounts of the substances can be expressed in moles or in laboratory situation, in grams *mass and the total number of each atom must be conserved in any chemical reaction; the number of molecules is not necessarily conserved

evaporation

the conversion of a liquid to its vapor form below the boiling temperature of the liquid;for a molecule to escape into the gas state, it must have enough kinetic energy to overcome the intermolecular attractive force in the liquid; given liquid sample will have molecules with a wide range of kinetic energies (liquid molecules with a kinetic energy that is above a certain threshold are able to escape the surface and become vapor); only the highest energy molecules are leaving the liquid state, so the collection of molecules that remain in the liquid now have a lower average kinetic energy; as evaporation occurs, the temperature of the remaining liquid decreases; given liquid will evaporate more quickly when heated (because heating process results in a greater fraction of the liquid's molecules having the necessary kinetic energy to escape the surface of the liquid); number of molecules that have required kinetic energy to evaporate increases with higher temperature liquid than lower

bond energy

the energy required to break a covalent bond between two atoms; breaking a chemical bond requires an input of energy (because the formation of a chemical bond results in a decrease in potential energy); high bond energy means a strong bond & molecules with these are likely to be more stable and less reactive; reactive compounds often contain at least one bond with low bond energy; halogen elements exist naturally as diatomic molecules (relatively small amounts of bond energy, very reactive); triple bond>double bond>single bond

theoretical yield

the maximum amount of product that could be formed from the given amounts of reactants

sp³ hybridization

the mixing of an s orbital with a set of three p orbitals to form a set of four sp³ hybrid orbitals; each large lobe of the hybrid orbitals points to one corner of a tetrahedron (the electron domain geometry that the hybrid orbitals must match); if there are lone pairs (trigonal pyramidal, bent), some of the hybrid orbitals are filled; angle between these lobes is 109.5°; sp³ hybrids are all equal to each other; one of the p orbitals is left unhybridized; (ex: methane, ammonia, water)

vaporization

the process in which a liquid is converted to a gas

unit cell

the smallest portion of a crystal lattice that shows the 3D pattern of the entire crystal; crystal can be thought of as the same unit cell repeated over and over in 3D; unit cells occur in different varieties (cubic crystal system is composed of 3 different types of unit cells: simple cubic, face-centered cubic, body-centered cubic

representative particle

the smallest unit in which a substance naturally exists (atom, ion, molecule, or formula unit); for majority of pure elements, representative particle is the atom; representative particles of HOFBrINCl is the diatomic molecule; molecule is representative particle of all molecular compounds; for ionic compounds, the representative particle is the formula unit; a mole of any substance contains Avogadro's number (6.02 x 10²³) of representative elements

formula mass

the sum of the masses of all the atoms represented in a chemical formula (applicable to molecular compounds, ionic compounds, or ions)

boiling point

the temperature at which the vapor pressure of a liquid is equal to the external pressure; as a liquid is heated, the average kinetic energy of its particles increases; the rate of evaporation increases as more & more molecules are able to escape the liquid's surface into the vapor phase; eventually point is reached when molecules all throughout liquid have enough kinetic energy to vaporize, beginning to boil boiling point of liquid also correlates to strength of its intermolecular forces (weak forces means it does not require a large input of energy to make the liquid boil.. vice versa for strong forces which require more energy) boiling points are affected by external pressure; at higher altitudes, the atmospheric pressure is lower & with less pressure pushing down on the surface of the liquid, it boils at a lower temperature

electron domain geometry

the total number of electron pairs, both bonding and lone pairs, lead to this; (linear, trigonal planar, tetrahedral, trigonal bipyramidal, or octahedral); when one or more of the bonding pairs of electrons is replaced with a lone pair, the actual shape of the molecule is altered

resonance

the use of two or more Lewis structures to represent covalent bonding in a molecule (ex: ozone O₃ - by distributing the total of 18 valence electrons, two structures can be drawn; the structures can be converted by shifting electrons without altering the position of the atoms; consists of one single bond and one double bond that are actually identical; the properties of each bond are in between those expected for a single bond and a double bond between two oxygen atoms {stronger and shorter than a typical O-O single bond but longer and weaker than an O=O double bond}); the true structure of a molecule that displays resonance is an average or a hybrid of all the resonance structures (ex: ozone O₃ - each of the covalent bonds is thought of as "one and a half" bonds); "half-bond" can be shown as a dotted line; many polyatomic ions also display resonance (ex: nitrate NO₃⁻ has a true structure that is an average of three valid resonance structures; the bond lengths between the central N atom and each O atom are identical & each bond is approximated as a "one and one-third" bond)

Van der Waals forces

the weakest intermolecular force and consist of dipole-dipole forces and dispersion forces

kinetic-molecular theory

theory that explains the states of matter & is based on the idea that matter is composed of tiny particles that are always in motion; theory helps explain observable properties & behaviors of solids, liquids, & gases; theory most easily understood as it applies to gases; theory applies specifically to the ideal gas

odd-electron molecules

there are a number of molecules whose total number of valence electrons is an odd number; not possible for all of the atoms in such a molecule to satisfy the octet rule (ex: nitrogen dioxide NO₂ - each oxygen atom contributes 6 valence electrons and the nitrogen atom contributes 5, for a total of 17

bond polarity

true nature of a chemical bond often falls somewhere in between the two extremes of a covalent bond and an ionic bond (bonding between atoms of different elements is rarely purely ionic or purely covalent); degree to which a given bond is ionic or covalent is determined by calculating the difference in electronegativity between the two atoms involved in the bond; a difference in electronegativity (∆EN) greater than 1.7 results in a bond that is mostly ionic in character [∆EN>1.7 mostly ionic, ∆EN 0.4-1.7 polar covalent, ∆EN<0.4 mostly nonpolar covalent]

molecular crystals

typically consist of molecules at the lattice points of the crystal, held together by relatively weak intermolecular forces (dispersion forces in case of nonpolar substances or dipole-dipole forces in case of polar substances); some, like ice, have molecules held together by hydrogen bonds; when one of the noble gases is cooled & solidified, the lattice points are individual atoms rather than molecules (because atoms are held together dispersion forces and not by covalent or metallic bonds, the properties of such a crystal are most similar to crystals of molecular substances); in all cases, the intermolecular forces holding the particles together are far weaker than either ionic or covalent bonds; melting & boiling points of molecular crystals are much lower; lacking ions or free electrons, they are poor electrical conductors

number of particles, moles, mass interrelationship

use two consecutive conversion factors converting grams of substance to moles and then moles to number of molecules

vapor pressure & temperature

vapor pressure is dependent upon temperature; when the liquid in a closed container is heated, more molecules escape the liquid phase & evaporate; the greater number of vapor molecules strike the container walls more frequently, resulting in an increase in pressure; temperature dependence of vapor pressure is not linear

ideal gas constant

variable R in the ideal gas law equation; value of this variable depends on the units chosen for pressure, temperature, and volume in the ideal gas equation; numerical value of R depends on the units in which pressure is expressed; (8.314 J/K•mol or L•kPa/K•mol) because kPa x L = J; (0.08206 L•atm/K•mol); (62.36 L•mmHg/K•mol)

molar volume of a gas

volume of one mole of the gas at a given temperature and pressure; at STP, one mole (6.02x10²³ representative particles) of any gas occupies a volume of 22.4 L

volume to volume

volume ratios can be easily used when the volume of one gas is known and the volume of another gas is unknown; pressure and temperature must be held constant over the course of the reaction, but the reaction does not need to be run at STP (Avogadro's hypothesis is true regardless of the pressure and temperature being used) 1. use volume ratios to convert volume of given gas to volume of unknown gas

octet rule

when ions form, they conform to this rule by either losing or gaining electrons in order to achieve the electron configuration of the nearest noble gas; similarly nonmetal atoms share electrons by forming covalent bonds in such a way that each of the atoms involved in the bond can attain a noble-gas electron configuration (the shared electrons are "counted" for each of the atoms involved in the sharing); ex: for hydrogen (H₂), the shared pair of electrons means that each of the atoms is able to attain the electron configuration of the noble gas helium (which has 2 electrons); for atoms other than hydrogen, the sharing of electrons will usually provide each of the atoms with 8 valence electrons

central atoms with lone pairs

when one or more of the bonding pairs of electrons is replaced with a lone pair, the molecular geometry (actual shape of the molecule) is altered; E=lone pair on the central atom (A); subscript will be used when there is more than one lone pair; lone pairs on the surrounding atoms (B) do not affect the geometry; within the context of the VSEPR model, lone pairs of electrons can be considered to be slightly more repulsive than bonding pairs of electrons, due to their closer proximity to the central atom (lone pairs take up more space)

single covalent bond

when two atoms are joined by the sharing of one pair of electrons; can be shown by a dash in between the two symbols (H-H); covalent bond forms when two singly occupied orbitals overlap with each other; shared electrons must have opposite spins; unpaired electrons form a covalent bond