Chemistry 20 Master Doc PDF

Science 10 Review: Molecular Compounds: Molecular compound – nonmetal with a nonmetal (right of staircase + hydrogen) ​ ​ ​ ( 1​ 2 3​ 4 5 6​ 7 8 ​ 9​ 1o ) ​ Uses prefixes (mono, di, tri, tetra, hecta, hexa, hepta, octa, nona, deca)...

Science 10 Review: Molecular Compounds: Molecular compound – nonmetal with a nonmetal (right of staircase + hydrogen) ​ ​ ​ ( 1​ 2 3​ 4 5 6​ 7 8 ​ 9​ 1o ) ​ Uses prefixes (mono, di, tri, tetra, hecta, hexa, hepta, octa, nona, deca) ↳if there is two or more of the first element, add a prefix. ↳the second element changes its ending to “ide” — add a prefix (including mono) ★ Exceptions – Hydrogen ↳there are a number of molecular species containing H with common names; listed in llldata booklet pages 4/5 (i.e. nitrogen trihydride – ammonia) * There are seven diatomic molecular elements: O2, Br2, Cl2, I2, H2, N2, F2 -​ P4, S8 are not included — (they change regularly) Ionic Compounds: Ionic compound – metal with a nonmetal (left of staircase + right of staircase) ​ Ions are atoms/molecules that have gained or lost electrons ↳Metals that gain or lose electrons llllbecome polyatomic ions ​ Most polyatomic ions are named for you ↳Metal ions don't change their names ↳Nonmetals change their name ( All negative, except for ammonium (NH4+) ) ​ Metals tend to lose electrons (e–) to become stable ​ Nonmetals tend to gain electrons (e–) to become stable -​ Atoms will mostly try to mimic the noble gases (group 18) by having 2, 10, or 18 e– ex. Mg = 12p+ + 12e– = 0 ​ but Mg+2 = 12p+ + 10e- =+2 N = 7p+ + 7e– = 0​ ​ but N-3 = 7p+ + 10e– = -3 ⭑ Metals keep their name when “married”, non-metal changes ( –ide suffix) ⭑ Polyatomic ions are in data booklet (above periodic table) ⭑ Multivalent ions use roman numerals (where there is more than one charge, pick first one) *capitalization kills* Ionic Formulas: ​ Balance charges to zero (when possible) ↳Most ion charges are listed ↳Reduce to lowest terms when possible ​ For multivalent ions, if you don’t know, pick the first one, (most common) [these brackets mean concentrations of something] — more on that later (these parenthesis mean multiples of ionic compounds) — as seen above Acids // Bases: -​ Acids — either starts with H ​ ​ ​ OR -​ Ends with COOH ex. Hx or XCOOH -​ Always aqueous: acids, ammonium, nitrates, most group 1 ions Except for Li, F Reaction Types: Formation – either two elements or two sides mushed together (“marriage”) *choose most likely formula* Ex. C + O2 = CO2 Decomposition – the opposite of a formation (“divorce”) Ex. CO2 – C + O2 Precipitate: two liquids forming a solid. Single/Double Replacements – Metal switches with metal, nonmetal switches with​ llllllllllllllllllllllllllllllllllllllllllllllllllllllllllllllllllllllllnonmetals + ions switch with + ions *double* two compounds make two new compounds Combustion reaction – the reaction is CO2 + HOH Balancing Chemical Equations: ​ Add coefficients to the front of each compound / element until everything balances out ​ The number of each atom must be equal on both sides ​ Reduce to simplest terms (Tip: leave elements until end) Mass Stoichiometry: Mole Conversions: ​ One mole is equal to 602 200 000 000 00 or 6.011 x 1023 ​ Any molecular / ionic compound can be measured in moles using its formula weight or relative molecular mass (Mr) or (M) Ex. C = 12.01g/mol Any molecular or ionic compound can be measured using its formula weight or relative molar mass (Mr) M = ΣAr ​ To convert mass to moles, multiply by a conversion factor ​ To convert mass to moles, multiply by a conversion factor ​ ↳Ex. 6.00 g / 1 x 1 mol / 12.01g = 0.5 moles ​ To convert moles to mass, multiply by opposite conversion factor ↳Ex. 0.75 mol / 1 x 44.01 g / 1 mol = 33.0 g Stoichiometry: “Element measured” or “calculating related amounts of compounds within a balanced chemical reaction.” ​ Every stoich question is the same: 1.​ Balance equation 2.​ Convert to moles 3.​ Mole ratio 4.​ Convert from moles ​ It is always want / have Mass of A —> (g/m) → moles of A—> (mole ratio) —> moles of B —>(Molar mass) —> mass B Gases: Introduction: ​ Gases behave differently than solids/liquids. Consider the molecular spacings between the three phases of H2O. Solid : particles in fixed position. Liquid : particles close, but not fixed.​ ​ Gas : particles far, not in fixed position. ​ Due to the large spaces between molecules in gases, they have special properties ​ ​ ​ ​ ​ ​ ​ ​ ↳can be compressed ​ Generally, particles in gases move quickly // sporadically ​ Fast moving particles (high kinetic energy) = warmer gases / vice versa ​ Temperature is the avg. kinetic energy of all particles (i.e. not all particles travel at same speed) ​ Particles collide with zero loss of energy (i.e. perfectly elastic collisions) ★ When particles collide with the sides of their container, they exert a force called pressure. ​ The above set of properties is referred to as the kinetic molecular theory of gases (KMT) and is widely accepted as good behaviour. Measuring Temperature: Temperature — average KE (speed) of the particles -​ No Celsius scale (C) -​ Kelvin is used instead (K) To convert Celsius to Kelvin, +273.15 To convert Kelvin to Celsius, – 273.15 -​ A particle with zero KE (kinetic energy) should have zero temperature -​ When particles have zero movement or temperature, this is known as Absolute Zero ​ ​ ​ ​ ​ (0°K is equal to -273.15°C)↲ -​ The magnitude of both 1°C and 1°K are the same — the scales have the same increments, but start in different places Measuring Pressure: -​ A gas generates pressure when it collides with the walls off its container -​ Pressure has many units, but we strictly use KiloPascals (kPa) -​ Other units include atmospheric (atm) and millimetres mercury (mmHg) -​ Converting between units works similarly to using mole ratios ​ ​ ​ ​ ​ (i.e. want / have)↲ ​ ​ ​ 1.00 atm = 101.3 kPa = 760 mmHg To convert pressure units: Boyle's Law -​ When you increase pressure, volume decreases by the same measure -​ When you increase volume, pressure decreases by the same measure. Basically — what you do to one side, you do the opposite to the other. (i.e. doubling P would cut T in half; multiplying P by 3 would reduce by one third, etc.) Boyle’s Law = P1V1 = P2V2 (initial)↲ ↳(final) Charles’ Law -​ Volume and Temperature are directly proportional (just like Boyle’s Law) -​ What happens to one side happens to the other Charles Law = V1 / T1 = V2 / T2 -​ You MUST use the absolute temperature in K for every calculation, if T=0K, there is no volume—but if you used a negative T, that would imply negative volume Combined Gas Law: -​ Combination of Boyle's Law and Charles Law into one equation P1V1 Combined gas Law = / T1 = P2V2 / T2 -​ NOTE: if one variable is constant (i.e. P or T) then you can omit them and get another gas law. To keep everything organized, slide everything diagonally -​ Units can be what you like, T is always in Kelvin -​ If you move something, it has to “switch floors” Ideal Gas Law: -​ If we assume gases behave ideally, then we can predict properties -​ Ideal refers to the tenets of KMT (i.e. gas particles are quickly moving in random direction with no loss of energy during collision. Ideal gas law = PV = nRT -​ R = gas constant, it dictates the units of the other variables in the equation ​ ​ P = kPa ​ ​ V = Litres ​ ​ N = Moles ​ ​ T = Temperature in Kelvin -​ STP = 100 kPa, 273 K​ ​ ​ ​ (surrounding) -​ SATP = 100kPa, 298 K — (the A stands for ambient) Avogadro’s Hypothesis: “The spaces between gaseous molecules is so vast that molecular size is irrelevant. Thus, equal volumes of any two gases at the same T and P will contain the same numbers of moles of that gas.” -​ This means that the volume of a gas is directly proportional to moles, thus, stoichiometry. -​ Essentially, this means that the volumes of gases will mix in stoichiometric quantities—it is referred to as the Law of Combining Gas Volumes Solution Stoich: What is it? ​ A solution is a homogeneous mixture, comprised of a solute dissolved in a solvent ↳Homogenous: a solution that looks, acts, etc. like one type of compound (all the same) Solute: typically a solid that separates into individual / invisible components when dissolved Solvent: typically water; the solvent is what “bonds” with solute particles and tears them into m k m​ individual entities Atypical solutions — air (N2, O2, Ar) ​ ​ – ethanol + water ​ ​ – steel / alloys (fe + c) (cu + sn), etc. ​ This unit is almost entirely aqueous solutions — (dissolved in water) ​ ​ ​ ​ ​ ​ ​ - i.e. salt water (NaCl) ​ ​ ​ ​ ​ ​ ​ - sugar water (C12H22O11) ​ Many ionic compounds + some molecular compounds dissolve in water (see solubility chart in data booklet) ​ Solvent molecules (H2O) are attracted to the solvent particles (and vice versa) ​ That attraction separates the solute into their individual components ​ ​ ​ ​ ​ ​ ​ ↳ions for ionic compounds ​ ​ ​ ​ ​ ​ ↳individual molecules for molecular lllllllllllllllllllllllllllllllllllllllllllllllllllllllllllllllllllllllllllllllllllllllllllllllllllllllllllllllllllllcompounds ​ When a solute dissolves, there are two types of bonds involved ↳the bonds holding the solute llllparticles together are being lll ​ Breaking bonds releases energybroken ​ Making bonds absorbs energy↳New bonds between the solute lllland solvent are beingformedBreaking bonds absorbs energy -​ Making bonds releases energy -​ Energy = heat -​ When heat is absorbed = endothermic -​ When heat is released = exothermic -​ Dissolving a solid in a solvent will be partially endothermic (breaking solute-solute bonds) + partially exothermic (making solvent-solute bonds)  -​ Some solute x2 bonds are very hard to break, and will absorb more energy that they release when they dissolve = cold -​ Some solute x2 bonds are very easy to break. And will release more energy than they consume when they dissolve = hot -​ There is no general rule for which solutions absorb or release energy when they dissolve, just know that some heats of solution are endothermic while others are exothermic. Solubility: clours : one two three four five six -​ Solution = homogeneous mixture composed of a solute and a solvent. Ions = for ionic compounds Molecules = for molecular compounds -​ Solubility describes the maximum amount of solute that can dissolve in a given solvent (units are g/100mL or g/L) -​ - since solubility is temperature dependent, values are reported at specific temperatures. -​ The solubility of NaCl in water at 25C is 36g/100mL or 360g/L -​ The solubility of sucrose in water at 25C is 200g/mL; at 50C it jumps to 260g/L Not all substances increase their solubility with increasing temp. -​ It is also possible to super saturate a solution if you heat it until all the solid dissolves, if you let it cool slowly, the extra solute won’t precipitate and thus — supersaturation -​ Supersaturated solutions are very unstable, a bump/disturbance causes the extra solute to precipitate immediately – usually releasing heat. Factors affecting solubility: -​ Higher temperatures = higher solubility (i.e. a warmer solvent can dissolve more solute) — however, that is not true for all substance Okay okay okay, gotta lock in. Bonding: Bond character - Bonds are not ionic or covalent, they are a mix of both - some bonds are more covalent (they are less ionic) and vice versa Its all decided by how much electronegativity (eneg.) they posses, their character is decided by their electronegativity ​ ​ ​ Fluorine is the most electronegative ​ ​ ​ Francium is the least A chemical bond is essentially a fight over electron, Non polar covalence = to equal sharing between the electrons Polar covalence = unequal sharing Ionic = stealing Ionic bonds Ionic bonds form between metals and nonmetals ​ ​ ​ ​ ​ ​ Technically they form between anions and cations Anion – negatively charged ion​ ​ They form when an atom GAINS an electron Cation – positively charged ion​ ​ They form when an atom LOSES an electron Usually, metals form anions, nonmetals form cations Polyatomic ions are molecules that gain/lose electrons Ionic solids exist in lattice cubes, not as individuals ​ ​ ​ Lattice is made up in a ratio of cations:anions When ionic solids dissolve or melt, you end up with free moving ions Ionic solids do not conduct electricity, ionic solutions and molten liquids do. Covalent bonds Equal Sharing — (i.e. F-F) Non-polar covalent Unequal Sharing — (i.e. Ag-Cl) Polar covalent No sharing (stealing) — (i.e. Ionic When calculating, find higher #, and find difference. ​ a covalent bond is an electrostatic bond attraction between a shared pair of electrons and their nuclei. Only the valence e- participates. ​ Lewis diagrams (or electron dot diagrams) show how e- are arranged around an atom -​ Pretend theres a square around the element, electrons are added without pairing at first to each side of the squar. Once each side has an e- , then the pairing starts -​ A single e- is called a bonding electron and the pairs are called lone pairs Note that the number of electrons is equal to the # of valence electrons. ​ Single bond e- are very reactive/unstable (like me), and so Lewis diagrams for these molecules only show lone pairs in the form of lines, dots, dashes, x’s, or any combination of those. Conditions for Lewis Diagrams: -​ The total number of valence electrons is conserved -​ The “octet rule” is conserved ​ This means that every atom has a full valence shell, for most that means 8 e- (4 pairs), but theres exceptions ​ H needs 1, Be needs 2, B needs 3 Ex. CH4 — 4 + 4 = 8e- or 4 pairs ​ For many molecules , atoms will share more than one pair of e- -​ Two pairs is a double bond -​ Three pairs is a triple bond ​ THIS IS THE HIGHEST LEVEL OF BONDS ​ Single bonds are usually long and weak ​ Triple bonds are short and strong Ex. N2 — 5 + 5 = 10e- or 5 pairs Note that the # of bonds around each atom does NOT always equal the # of bonding e- ​ Lewis diagrams for polyatomic atoms are a hair different, you simply add or subtract the amount of electrons gained/lost, then add a bracket + the charge. For cations, you subtract electrons, for anions, you add electrons Ex. OH— VSEPR: Valence Shell Electron Shell Repulsion -​ Bonds and lone pairs all occupy 3D spaces called electron domains. Since each ED is filled with a negative charge, the EDs repel each other and spread out as much as possible, creating that shape you have to remember ​ To figure out the shape, draw Lewis diagram, then count the total # of eds around the central atom -​ Classify each Ed as a lone pair or a bonding pair -​ Double and triple bonds only count as 1 pair 1 / 0 — linear​ ​ ​ 2 / 0 — bent 1 / 1 — linear​ ​ ​ 2 / 1 — bent 1 / 2 — linear​ ​ ​ 3 / 0 — trigonal planar 1 / 3 — linear​ ​ ​ 3 / 1 — trigonal pyramidal 2 / 0 — linear​ ​ ​ 4 / 0 — tetrahedral You are responsible to know this… Molecular Polarity: -​ Covalent bonds can be either polar or nonpolar based on ∆eneg. The polarities are represented with an arrow pointing to the atom with the most eneg -​ As stated as above — Equal Sharing — (i.e. F-F) Non-polar covalent No charge Unequal Sharing — (i.e. Ag-Cl) Polar covalent No sharing (stealing) — (i.e. Ionic ​ A nonpolar bond (NPC) has a uniform distribution of charge in ED ​ An ionic bond is a stolen electron, one atom has a positive (+) charge and one has a negative (–) charge. ​ A polar bond is between the two extremes, the result is a partial charge -​ These can be denoted by δ+ and δ– ​ A polar bond is thus said to have a permanent dipole ( two opposite poles) ​ A molecule that has all dipoles pointing to one end will produce a polar ​ molecule, however, if there is a break, there is not dipole molecule Ex. ​ ​ H–CΞN (dipoles add up)​ ​ ​ O = C = O (dipoles do not add up) →→→​ ​ ​ ​ ​ ←​ → ​ ​ ​ (polar)​ ​ ​ ​ ​ ​ (nonpolar) Intermolecular Forces (IFs) Some physical properties of molecules are determined by shape and polarity Melting point (mp) and boiling point (bp) and solubility/miscibility ​ Miscibility is the ability for two liquids to mix (i.e. oil and water are immiscible) Melting, boiling, and dissolving all involve the breaking of intermolecular bonds ​ If you melt, boil, or dissolve a molecule, you still have a collection of free moving molecules, not free moving atoms. ​ If you melt, boil, or dissolve an ionic compound, you have a collection of free  moving ions. ​ Covalent bonds are intr-A-molecular — meaning “within the molecules.” Intermolecular forces (IFs) are broken into 3 different categories: -​ London Dispersion Forces (LDFs) -​ Dipole-dipole interactions (d.d.) -​ Hydrogen bonding (H bond) The polarity of a molecule determines`what kind of IF is present. ​ Everything has London Dispersion Forces, but nonpolars ONLY LDs ​ Regular polars had dip-dips. ​ Very polar molecules (H-F, H-O, and H-N) Non polar molecules have an even distribution of e-, their electron clouds are symmetrical Through random motion of e-, sometimes they become temporarily polarized. These temporary dipoles induce temporary polarization of neighbouring clouds and temporary attraction occurs (like a situationship) Temporary flashes of attraction are on and off, this is known as London Dispersion Forced ​ Because LDFs are dependant on the number of electrons, molecules with larger masses have more LDFs ⇧in MASS =⇧ in LDF More attraction = mp/bp#