Chapter 1: The Study of Change Types of Matter ● Substance ○ Substances are composed of atoms. They are not mixtures. They have a constant composition. Below are two types ofpuresubstances. ■ Element ● A substance that cannot be separated into similar substances by chemical means. ● Diatomic elements, in their natural, lowest energy state, are bonded into pairs. ○ Hydrogen, nitrogen, oxygen, fluorine, chlorine, bromine, and iodine are all diatomic elements. ■ Compound ● A substance composed of atoms of two or more elements chemically united in fixed proportions. ● Mixtures ○ Mixtures are a combination of two or more substances in which the substances retain their distinct identities. Mixtures are made up of elements and compounds, but they cannot BE just one element or just one compound. Elements and compounds cannot be mixtures, otherwise they would cease to be pure substances. ■ Homogenous mixture ● The composition of the mixture is the same throughout. ■ Heterogenous mixture ● The composition of the mixture is NOT the same throughout. ● Diatomic elements ○ A diatomic element with another element of its kind to form a diatomic molecule. Diatomic elements are:Hydrogen(H2),Nitrogen(N2) ,Oxygen(O2),Fluorine (F2),Chlorine(Cl2),Bromine(Br2),Iodine(I2).Not writing these elements as diatomic is almost always incorrect. ● States ○ Matter has three states: (s) solid, (l) liquid, and(g) gas.Bromine and mercury are liquids at room temperature. Hydrogen, nitrogen,oxygen, fluorine, chlorine, helium, neon, argon, krypton, xenon, and radonare gases at room temperature. (298 K or 25 °C) ● Properties of matter ○ Extensive(properties reliant on AMOUNT of matter) ■ Mass, energy, volume ○ Intensive(properties not reliant on amount) ■ Density, temperature ● Density = Mass/Volume Measurement ● SI Units ○ Prefixes ■ Kilo-1000 ■ Deci - 1/10 ■ Centi - 1/100 ■ Milli - 1/1000 ■ Micro (mc,µ) - 1/1000000 (one millionth, 1/10^6) ■ Nano - 1/1000000000 (one billionth, 1/10^9) ■ Pico - 1/1000000000000 (one trillionth, 1/10^12) ○ Actual units - SI Base System ■ Meter (m) - length ■ Kilogram (kg) - mass ■ Second (sec) - time ■ Ampere (amp) - electrical current (defined as one Coulomb of electrons flowing past a fixed point per second, C/s) ■ Kelvin (K) - temperature (Celsius can only be used when calculating delta T) ■ Mole (mol) - amount of substance - 6.02 x 10^23 atoms/particles ■ Candela (cd) - luminous intensity ■ Joule (J) - energy=(mv^2)/2 - the unit for kinetic energy is always J, but J does not always mean you are dealing with kinetic energy. 1 2 ● Kinetic = 2 𝑚𝑣 ● Potential = learn later ● Significant figures ○ All nonzero digits are significant, unless they have decimals after them or are in scientific notation. For instance, 100 has 1 significant digit, 1.00e2 has three, and 100.0 has four (“trailing zeroes are only significant with a decimal”). Zeroes between significant digits are significant. ○ In addition and subtraction, round the answer to the LEAST number of places in the decimal portion of any number in the problem. For example, 2.458 - 2.13 = 0.328 unrounded, but since thousandths place is uncertain, the final answer is 0.33. ○ In multiplication and division, round off to the least number of sig figs of the original numbers. Disregard ‘exact’ numbers such as the fact that there are 12 inches in a foot; those are assumed to have infinite significant digits. ● Precision and Accuracy ○ P recisionis how closely two or more measurements of the same quantity agree with one another. ○ Accuracytells us how close a measurement is to thetrue value of the quantity measured. ***In lab measurements that involve theinterpretation of scales, you always have to estimate 1 decimal place beyond what is given. ● Accuracy of Glassware ○ Volumetric Flask ■ Accurate to +0.01 or -0.01 mL ○ Graduated Cylinder ■ Accurate to +0.1 or -0.1 mL ○ Beaker ■ Accurate to +10 or -10 mL ○ Buret ■ Accurate to +0.01 or -0.01 mL ● Unit Conversionis also known asFactor Label MethodorDimensional Analysis ○ Follow the units and show work on FRQ (especially for calculations involving sig figs [nearly all of them!]) Chapter 2: Atoms, Molecules, and Ions Atoms ● John Dalton’s atomic theory ○ Elements are composed of small particles called atoms. ○ All atoms of a given element are identical. The atoms of one element are different from the atoms of all other elements. Iforgot to note: this is not always true. ○ [Law of definite proportions and Law of multiple proportions by Joseph Proust]Compounds are composed of atoms of more thanone element. In any compound the ratio of the numbers of atoms of any two of the elements present is either an integer or a simple fraction. ○ [Law of conservation of mass]A chemical reactioninvolves only the separation, combination, or rearrangement of atoms; it does not result in their creation or destruction. ● Subatomic particles ○ Electron ■ Negative charge ■ W hen emitted in large quantities (a stream), termedbeta rays; beta ray = stream of electrons ● Electrons are waves (but they are also particles); thus, they qualify as radiation, just like light. (optional: all objectswith momentum have a avelength that is inversely proportional to their momentum; because electrons have such w a small momentum, their associated wavelength is large compared to an object with a larger momentum, such as a person) ○ Proton ■ Positive charge ■ Part of the nucleus ■ Establish elemental identity (every element is defined by its number of protons, Zahl). ○ Neutron ■ Neutral charge ■ Part of the nucleus ■ Changed number of neutrons = isotopes. Ions, Isotopes, Compounds ● Some terminology ○ Atomic number,Z (Zahl) ■ Number of protons in the nucleus. Also the number of electrons (for neutral atoms). ○ Mass number,A(exact) ■ Protons + neutrons (the mass of electrons is too small to be significant) ○ Atomic mass (measured with the exception of carbon-12) ■ The average of the mass numbers of all the isotopes of a given element in amu. (if a certain isotope is moreabundantthanthe other isotopes, then its mass is given more “weight” in the calculation of the atomic mass) ○ Isotopes ■ Atoms that have the same atomic number but different mass numbers, usually due to change in the number of neutrons. Hydrogen, for instance, has one proton and no neutrons; deuterium has one proton and one neutron, and tritium has one proton and two neutrons. ○ Period ■ Horizontal rows on the periodic table ○ Family/Group ■ Vertical rows on the periodic table ○ Metal ■ Good conductor of heat and electricity. Usually malleable and ductile. Have a metallic luster. ○ Nonmetal(opposite of metal) ■ Poor conductor. Not malleable nor ductile. Do not have a metallic luster. ○ Metalloid ■ P roperties intermediate between metal and nonmetal. Their conductivity increases with increasing temperature. ○ Molecule ■ An aggregate of at least two atoms in a definite arrangement held together by chemical bonds. ■ Molecular formula ● Shows the exact number of atoms of each element in the smallest unit of a substance. ● Allotrope ○ One of two or more distinct forms of an element. Example: graphite and diamond ● Empirical formula ○ Tells us which elements are present and the simplest whole-number ratio of their atoms. For instance, the empirical formula of glucose is CH2O, but the actual formula is C6H12O6. ○ Ionic compounds are always in this form ○ Ion ■ An atom or molecule with an electric charge due to losing or gaining an electron. ■ Polyatomic ion ● A charged molecule. ■ Monatomic ion ● A charged atom. Naming U ● se theStock systemof naming. ● Organic compoundscontain carbon andinorganic compoundsdo not. ● Ion naming practices ○ The first compound named is the metal cation, followed by the nonmetallic anion with the suffix -ide. ○ Naming oxyanions is different: ■ Themost common permutationhas the suffix-ate. ● Nitrate(NO3)(-), Sulfate(SO4)(-), Chlorate(ClO3)(-),Acetate (C2H3O2)(-), Carbonate(CO3)(-2), Phosphate(PO4)(-3), Chromate(CrO4)(-2), Thiocyanate(SCN)(-), Oxalate(C2O4)(-2) ■ The oxyanions withone less oxygenthan the most commonpermutation have the suffix-ite. ● Chlorite(ClO2)(-2), Nitrite(NO2)(-), Phosphite(PO3)(-3),Sulfite (SO3)(-2) ■ T he oxyanions withtwo less oxygenthan the most common permutation have the prefixhypo-and the suffix-ite. ● Hypochlorite(ClO)(-) ■ The oxyanions withone more oxygenthan the most common permutation have the prefixper-and the suffix-ate. ● Perchlorate(ClO4)(-), Permanganate(MnO4)(-) ● Molecular naming practices ○ The name of the first element, then the second element with-ideadded onto it. For instance, sodium chloride is NaCl. As a note, prefixes are not applied to the first element. Acid naming practices ● ○ ○ Screenshotted from the WebAssign for this week -Atoms,Molecules, and Ions. ○ Acids areproton donors. ● Base naming practices ○ Bases areproton acceptors. Ammonia isNH3, ammoniumion isNH4+, amide isNH2-, hydroxide isOH-, and hydride isH-. Hydrate naming practices ● ○ Name of the compound, then number + hydrate. So BaCl2•2H2O is barium chloride dihydrate. hapter 3: Mass Relationships in C Chemical Reactions Mass ● Atomic mass ○ Atomic mass is the mass of the atom inatomic massunits (amu) ○ The atomic mass on the periodic table is technically average atomic mass. See Chapter 2 outline for elaboration. ● Molar mass ○ A mole is the amount of substance that contains as many elementary entities as there are atoms in 12 grams of carbon. This is 6.022 x 1023 (atoms)/mole, also known as Avogadro’s number. ○ The molar mass of an element is the number of grams of that element for every mole of the element. Memorize these atomic masses & charges ● Charge: + ○ Hydrogen: 1.01 ○ Sodium: 23.0 ○ Potassium: 39.1 ● Charge: +2 ○ Magnesium: 24.3 ○ Calcium: 40.1 ● Charge: +3 ○ Aluminum: 27.0 ● Charge:+4 ○ Carbon: 12.0 ● Charge: -3 ○ Nitrogen: 14.0 ● Charge: -2 ○ Oxygen: 16.0 ○ Sulfur: 32.5 ● Charge: -1 ○ Chlorine: 35.45 Percent Composition ● P ercent composition is calculated by dividing the molar mass of the element by the molar mass of the compound and multiplying by 100. Experimentally, it can be deduced by figuring out the number of moles of each element and comparing the numbers. x molar mass of elementx 100 = % composition ofan element by mass n molar mass of compound U ● sing the molar masses, the formulas can be deduced. ● However, every formula that is deduced in this way is called “empirical.” For instance, the empirical formula of glucose is CH2O. The actual molecular formula is C6H12O6. To find the molecular formula, one must know the approximate molar mass of the entire compound. Balancing Equations ● B alanced equations have the same number of atoms of each element on each side. Every atom must be accounted for. ● The easiest way to balance an equation is to start with the most complex molecule and use that to extrapolate the coefficients of the less complex molecules on the other side. Stoichiometry T ● he word strikes fear into many of your hearts, but it is truly quite simple. ● Stoichiometry is the quantitative study of reactants and products in a chemical reaction. ● Using a balanced equation and the number of grams of one element involved in a reaction, one can find the number of grams of every other element that reacted. ○ Example:N2 + 3H2→ 2NH3 ■ If we’re told 3 moles of nitrogen react, this means we have produced 3 moles of ammonia, but we used 9 moles of hydrogen because 3x3=9 ● General approach to solving stoichiometry questions: ○ 1. Write a balanced equation for the reaction ○ 2. Convert the given amount of the reactant to number of moles ○ 3. Use the mole ratio from the balanced equation to calculate the number of moles of product formed ○ 4. Convert the moles of product to grams of product Limiting reactants: In some cases, one reactant will be in excess and the other will ‘limit’ the ● amount of product that can be produced. hapter 4: Reactions in Aqueous C Solutions Of Electrolytes and Non-Electrolytes ● An electrolyte, when dissolved in water, can conduct electricity. ○ This is because the electrolyte dissociates into charged ions. When a cathode and anode are placed into the water, the ions group to the nodes, creating a closed circuit. ○ Some electrolytes are ionic compounds which dissociate fully. Ionic compounds dissociate; they do not ionize because they are already ions. ○ Some are acids or bases which ionize. For example, HCl, which is a covalently bonded molecule, ionizes in water to form H+ and Cl- ions. ○ Strong electrolytes ■ HCl, HBr, HI, HNO3, HClO4, H2SO4, NaOH (+ all otheralkali hydroxides), Ca(OH)2, Sr(OH)2, Ba(OH)2 ○ Weak electrolytes ■ CH3C OOH, HF, HNO2, NH3, NH4+, H2O (anything that isnot strong) ● A nonelectrolyte does not conduct electricity when dissolved in water. ■ (NH2) 2C 2O6, C12H22O11 O, CH3O H, C2H 5O H, C6H1 Solubility ● In a precipitation reaction, compounds containing alkali metal ions (Li+, Na+, K+, Rb+, Cs+) , ammonium ion (NH4+), nitrate (NO3-), bicarbonate (HCO3-), chlorate (ClO4-), halides (Cl-, Br-, I-), or sulfates (SO4(2-)) will usually be dissolved in solution as spectator ions. This is important so that you can figure out which ions actually reacted with each other so you can write out the net ionic equation. I actually have a mnemonic regarding solubility, but it’s quite complicated. ● The basic rule is:All sodium, potassium, ammonium,and nitrate salts are soluble in water. Metathesis ● Molecular equations, Ionic equations, and Net Ionic equations ○ Let’s write theprecipitationof lead(II) iodide in3 different ways. Why? Because it’s fun to be confusing! ○ No, because each equation offers important information impossible to place in the other equations. ○ Method 1: Molecular ■ Pb(NO3)2 + 2KI → PbI2 + 2KNO3 ■ Writing it this way shows us what happens in a molecular interpretation. However, since these are ionic compounds, this is not an accurate representation. ○ Method 2: Ionic ■ Pb(2+) + 2NO3(-) + 2K(+) + 2I(-) → PbI2 + 2K(+) + 2NO3(-) ■ This method is unnecessarily confusing and includes some repetition, which will be eliminated in the final method. ■ How did we know that PbI2 would be the only undissociated substance? Because potassium nitrate is soluble (thanks to nitrate and potassium) and will instead take the form of spectator ions. (also because PbI2 is insoluble) ○ Method 3: Net Ionic ■ Pb(2+) + 2I(-) → PbI2 ■ This method shows us the bare-bones equation, no spectator ions to be confusing. Acid-Base Neutralization Reactions F ○ un fact: aspirin is an acid and milk of magnesia is a base. ○ Acids ionize in water to produce H+ ions; bases produce OH- ions. ○ In a neutralization reaction, the acid and base react to form water and a salt, an ionic compound made up of a cation other than H+ and an anion other than OH-. The net ionic equation will sometimes be H(+) + OH(-) → H2O (for strong acid-base neutralizations, the NIE will always be H(+) + OH(-) → H2O) Chapter 5: Properties of Gases Review: Pressure = force/area - more particles bumping (and the harder they bump) against the walls of a container increases the force per unit area - decreasing surface area increases pressure What is a Gas? H ● 2, N2, O2, O3, F2, Cl2, He, Ne, Ar, Kr, Xe, Rn are gases at room temperature ● Gases are compressible; they assume the volume and shape of the containers and mix evenly. ● Density of a gas is expressed in g/L instead of g/mL (translation: gases are mostly empty space!). Gas Laws ● Boyle’s Law ○ This law relates pressure and volume. They have an inverse relationship. PV = k ■ with constant n and T ○ The specific law is P1V1 = P2V2 ● Charles's law ○ Temperature and volume have a direct relationship. V1/T1 = V2/T2 ■ with constant P and n ○ Pressure and temperature have a direct relationship. P1/T1 = P2/T2 ● Avogadro’s Law ○ Volume and number of moles have a direct relationship. V1/n1 = V2/n2 ■ with constant P and T ○ 22.4 L in every mole ONLY AT STP conditions (1 atm and 273 K (0 degrees C)) ● The Ideal Gas Law ○ PV=nRT ○ n=PV/RT ○ Gases are most likely to be ideal under high temperature(low intermolecular forces)and low pressure conditions. ● Absolute Zero ○ A graph of the volume or pressure of a gas (constant [P or V] AND n) plotted against its temperature will give a straight line, when extrapolated, intersects the x-axis at the temperature -273[.15]°C, 0K, or absolutezero. ● Stoichiometry ○ Volumes and pressures can be treated just like moles when working with gases. ● Dalton’s Law of Partial Pressures ○ Total Pressure = Partial pressure 1 + Partial pressure 2 + Partial pressure 3… ● Maxwell-Boltzmann speed distribution curve ○ Every gas molecule moves at a different speed, and the speed distribution curve shows us the range of speeds. ○ The area under the curve represents the moles of the compound. ○ Less massive molecules are moving faster than more massive molecules at a given temperature. ● Deviations ○ The assumption that molecular volume in a gas is negligible does not work at high pressures and concentrations. ○ The assumption that gas molecules do not attract each other does not work at lower temperatures. ○ Differing measurements from what was predicted by the Ideal Gas Laws and why: ○ Lower pressures than predicted = intermolecular attractions ○ H igher pressures than predicted = higher than normal molecular volume compared to the volume of empty space . Chapter 6: Thermochemistry Types of energy ● Kinetic energy ○ Energy produced by a moving object ● Potential energy ○ Energy available by an object’s position (stored energy) ● Chemical energy (a type of potential energy; also nuclear energy) ○ Energy stored within chemical substances ● Thermal energy ○ Energy from random movement of atoms and molecules ● Radiant energy ○ Energy from the sun (also called solar energy) Enthalpy (H) ● What is enthalpy? ○ Enthalpy (H) or chemical energy, is stored potential energy in substances. It is almostimpossible to measure, so the∆H (change inenthalpy per mole of reaction as written) is measured instead. ○ While calculating the change in enthalpy, the pressure of the reaction must remain constant. ● Types of reactions ○ Exothermic reactions(withdrawals of H) ■ Stored potential energy is lost and is gained as kinetic energy by its surroundings ○ Endothermic reactions(deposits of H) ■ Stored potential energy is gained by taking kinetic energy from its surroundings ● Different ways to find∆H ○ ∆H = H(products)-H(reactants) (by definition) ○ ∆H = kJ/molrxn ■ ∆H>0 = endothermic reaction (e.g. 100 kJ/molrxn) ■ ∆H<0 = exothermic reaction (e.g. -100 kJ/molrxn) ● Hess’ Law ○ Finding change in enthalpy requires the use of this equation: ΔHrxn = ΔHformation products - ΔHformation reactants ○ ΔHf products and ΔHf reactants can be found usingthe Heat of Formation tables. ● Calculating heat ○ Using the equation q=mcΔT, the heat required to change temperature can be found. ○ q = heat in any units (J, c/Calorie, etc.) ○ m=mass ○ c=specific heat capacity ○ ΔT = change in temperature Introduction to thermodynamics ● H, enthalpy ○ ∆H energy changes in reactions only involving heat (constant pressure) ■ AP chemistry only deals with ∆H ● First law of thermodynamics (law of conservation of energy) ○ Energy can be converted from one form to another, but cannot be created or destroyed. ○ The application of this law in AP chemistry ■ An increase in potential energy is accompanied by a decrease in kinetic energy. ∆Esystem = -(∆Esurroundings) ■ An increase in kinetic energy is accompanied by a decrease in potential energy. ● Rules for writing thermochemical equations ○ Always specify physical states (e.g. solid, liquid, gas, etc.) ○ If both sides of the equation are multiplied by a factor,∆H must also change by the same factor, because molrxn “as written” changes. ○ When the reaction is reversed, the sign of ∆H changes. ■ H2O (s) ⟶ H2O (l) ∆H = 6.01 kJ/mol ■ H2O (l) ⟶ H2O (s) ∆H = -6.01 kJ/mol Chapter 7: Electronic Structure of Atom Two important equations: ○ c=𝛌v ■ c = speed of light:2.998e8 m/sor 2.998e17 nm/s ■ 𝛌 = wavelength (generally provided in nm) ■ v = frequency (given in 1/s or Hz) ○ E=hv ■ E = energy, in joules ■ h = Planck’s constant: 6.626e-34 J s ■ v =frequency (given in 1/s or Hz) ○ What is wavelength? What is frequency? ■ Wavelength is the distance between successive crests of a wave. ● Visible light has a wavelength of between 400-700 nm. Violet color has lower wavelength, and red color has higher wavelength. ● Decreasing wavelength and increasing frequency ○ Radio waves ○ Microwaves ○ Infrared ○ ROYGBIV ○ Ultraviolet ○ X-rays ○ Gamma rays ■ Frequency is the number of waves that pass a fixed point per second. ● Wavelength and frequency have an inverse relationship. Higher wavelengths correspond to lower frequency, and vice versa. c=𝛌v Electron configuration ○ Principle energy level ■ These levels are numbers: 1, 2, 3, 4… ■ They correspond to period on the periodic table (horizontal rows). ○ Subshell Subshells are mini-categories within the principle energy level. ■ ■ They are indicated by the letters s, p, d, f… ○ Orbital ■ Smaller still are the orbitals within the subshells which can hold up to two electrons each. ■ The s subshell has but one orbital; p has three; d has five, and f has seven. ■ Within the orbitals, the electrons have different spins. One electron has the up spin and the other has the down spin. ■ According the Hund’s Rule, the electrons will fill one orbital at a time, then they will come back to double up. Just like passengers in a bus, they don’t want to start sitting with other passengers til all the seats are occupied. Example:Helium has an electron configuration of 1s2. ○ ■ More illustratively, magnesium has 1s2 2s2 2p6 3s2. ■ Do not write aufbau - put all the numbers together. ○ Abbreviate using the closest noble gas. ■ Magnesium has an abbreviated electron configuration of [Ne] 3s2. hapter 8: Periodic Relationships C Between Elements Valence electrons ○ T hese guys are in the top energy level and have the most impact on atomic behavior, as they are generally the most exposed. Atomic size ○ A toms increase drastically in size at the beginning of each period because a new principle energy level has been added. ○ From left to right, atoms generally decrease in size because more protons exert a higher coulombic attraction. ■ Exception: halogens (O, S, Se, Te, Po) are larger than their previous atoms because the first paired electrons’ repulsions are more significant than the added proton’s attraction. Ionization energy ○ T his is the energy required to remove an electron from an atom in its gaseous state. IE is usually given in kJ/molelectrons, andis always a positive number. ○ Ionization energy is smaller when the atom is farther away from a complete valence shell, and larger when the atom is very close to a complete valence shell. PES Graphs ● How to interpret ○ Each peak is the number of electrons in the shell. ○ The peaks with the higher binding energy are closer to the middle of the atom and thus subjected to higher binding energy and live at lower shells. ○ The peaks on the right, with lower binding energy, are closer to the outer shell of the atom and thus subjected to lower binding energy, living at higher shells. ○ As the number of protons increases, the peaks move leftward or toward the higher binding energies, as the nuclear attractions on the electrons increase. Chapter 9: Chemical Bonding 1 Electronegativity trends ● Electronegativity is the ability of an atom to attract toward itself the electrons in a chemical bond. ● As you move down on the periodic table, electronegativity decreases. ● As you move from left to right, electronegativity increases. Bond Energy ● Lattice energy= The energy required to separate amol of a solid ionic compound into gaseous ions. You're breaking bonds here, so it will always be endothermic. ● ● Bond (dissociation) energy = literally the energy it takes to break bonds. Again, breaking bond = endothermic. ● ● Heat of solution (ch. 6 outline page 12) ● ● ● Ionization energy = The energy required to ionize an atom aka remove an electron. Again, this is breaking a bond. So, it will always be endothermic. ● ● Breaking a bond = Endothermic, you must put energyin. The reaction converts kinetic energy taken from the environment into potential energy. ● Forming a bond = Exothermic, energy is released, usuallyin the form of heat. Potential energy is converted into kinetic energy and released to the environment. ● ● Think about the direction of the energy/ heat when looking at these kinds of questions. Octet (Lewis-dot diagrams) T ● he octet rule dictates that every atom (bonded or not) should have eight electrons. ● There are many exceptions: ○ Hydrogen and Helium are satisfied with only two electrons, because 1s only has two electrons. ○ E lectrons from third period down also have a d shell so they can accommodate more than eight electrons. ○ Elements in column 13, such as boron, aluminum, etc can also just have 6 valence electrons. Formal Charge F ● ormal charge is the relative charge of each atom in the molecule. ● The charges added up make up the formal charge of the entire molecule. ○ Formal charge is calculated by this formula: ■ Normal valence electrons of atom - (bonding electrons/2) - lone pairs Resonance ● R esonance structures are “special” bonds on molecules. When one of two adjacent bonds is a double bond, and neither structure is preferable in terms of formal charges, the pi bond(s) can ‘switch’ from atom to atom. ○ In reality, the bond formed is an intermediate bond, having a bond length between that of a single and double bond. In Chapter 10, we will learn that the electrons in these resonance structures are delocalized over both/all of the resonance sites. Chapter 10: Chemical bonding 2 Molecular Geometry ● VSEPR structures ○ Linear, Trigonal planar, Tetrahedral, Trigonal bipyramidal, Octahedral, ○ See handout for the full list. ● Having a lone pair exerts a strong repulsive force, decreasing bond angles. Dipole moments ● Dipole moments are the polarity of the molecule. Polarity is the difference between the formal charge on one atom and another in the same molecule. ● In general, an asymmetrical molecule has a dipole moment, and a symmetrical one does not. Hybridization ● In bonds, the orbitals overlap. ● Hybrid orbitals are atomic orbitals obtained when two or more nonequivalent orbitals of the same atom combine in preparation for covalent bond formation. ● Examples include sp, sp2, and sp3. hapter 11: Intermolecular Forces and C Liquids and Solids ● Intramolecular attractions ○ Are the covalent bonds between the atoms in the molecule ○ Measured in hundreds of kJ ● Intermolecular attractions ○ Are weaker thanintramolecular attractions and Coulombicattractions ○ Measured in tens of kJ ○ Accounts for liquid and solid phases of molecules Types of intermolecular forces ● London Dispersion Forces (LDF) ○ There are natural fluctuations in electron densities within every molecule (including nonpolar molecules) ○ Charges in one molecule induces charges on the surrounding molecules ○ Positive sides of the molecule face the negative side of another molecule (LDF are attractive forces between molecules) ○ The charges fluctuate, they are not permanent ○ LDF increases with surface area of electron clouds around atoms. The larger the electron cloud, the more polarizable the molecule becomes (greater LDF) ○ The larger the LDF, the higher the melting temperature. ○ The only intermolecular attractions found in symmetrical molecules, monoatomic gases, and any hydrocarbon are LDF. ○ Are considered the weakest of intermolecular attractions, but increase with molecular size and can become stronger than dipole-dipole and hydrogen bonding. ○ It is possible for a substance with only LDF to have a higher boiling temperature than a substance with other types of intermolecular attractions ○ Do not use molar mass in explaining the magnitude of the intermolecular forces. ● Dipole-Dipole Forces ○ Permanent attractive force between polar molecules ○ Permanent polar charges on molecules (dipoles) produce these intermolecular attractions. ○ Unsymmetrical molecules have permanent dipole-dipole attractions ○ Molecules with permanent dipole moment align with opposite polarities ○ C an cause molecules to condense into a liquid and even lock together to make a solid. ○ Is weaker than covalent bonds and is shown as a dotted line in illustrations (covalent bonds are shown as solid lines) ● Ion-Dipole Forces ○ Attraction of an ion to one of the dipoles of a polar molecule ○ Major reason ionic compounds dissociate in water. ● Hydrogen Bonds ○ Hydrogen bonds are a special type of dipole-dipole interaction between the charged hydrogen atom and another polar molecule’s intense charge. ○ Hydrogen atoms can only attain an intense enough charge when bonded with N, O, or F, the three most electronegative elements. Crystals ● Molecular crystals ○ Connected by dispersion forces ● Covalent crystals ○ A single crystal is a giant covalent molecule. ○ Diamond and graphite are examples. ○ Silicon is a covalent network solid ■ Its electrical conductivity increases as temperature increases. ■ N-doping ● Adds electrons ● Done with phosphorus, arsenic, and antimony ■ P-doping ● Creates electron deficit ● Boron, aluminum, gallium ● Ionic crystals ○ Atoms held together by ionic bonds, three dimensional arrays of cations and anions. ○ Conducts electricity only when molten. ● Metallic crystals ○ Metallic crystals are bonded by a sea of electrons. They are good conductors of heat and electricity. Chapters 24-25: Orgo ● O rganic molecules come in many shapes and sizes. Since this is chemistry, people have categorized the molecules and named each category, and it is now our responsibility to remember their names! ○ Numbering ■ 1 - meth- (form-) ■ 2 - eth- (acet-) ■ 3 - prop- ■ 4 - but- ■ 5 - pent- ■ 6 - hex- ■ 7 - hept- ■ 8 - oct- ■ 9 - non- ■ 10 - dec- ○ Aliphatic Hydrocarbons ■ Alkanes ● Alkanes have the general formula CnH2n+2 ● They are known as saturated hydrocarbons because they contain the max number of hydrogen atoms that can possibly bond with the carbon atoms. ● Name them according to the number of carbon atoms. For instance, CH4 is methane, C2H6 is ethane, C3H8 is propane, and C4H10 is butane. ● Should one or more of the hydrogens be replaced with another functional group, the name of the compound indicates which carbon atom it’s been joined to. If the methyl group is attached to the second carbon from the left on pentane, it would be called 2-methylpentane. Use the smallest number possible ● These rules are strange, but they mean that just from looking at the name of the compound you can figure out its formula AND shape! ● Cycloalkane ○ With cycloalkanes, the carbon atoms are joined in rings. They have the general formula of CnH 2n. ○ Steroids like testosterone, estrogen, progesterone, etc, are generally cycloalkanes. ■ Alkenes ● Alkenes have a carbon-carbon double bond. They have the general formula of CnH2 n. ● A number is usually involved to indicate where the double bond is. ■ Alkynes ● Alkynes have a carbon-carbon triple bond. They have the general formula of CnH 2n-2. The simplest alkyne is acetylene. ● ○ Aromatic Hydrocarbons ■ Aromatics usually involve a carbon ring. ■ Monosubstituted benzenes are just named for a single group of atoms. ● Ethylbenzene, for instance. ■ Benzenes with more groups of atoms have to use numbers. Each of the six carbons on the ring are assigned a number 1-6. Functional Groups ○ ■ Alcohols ● Alcohols contain the hydroxyl functional group - OH ■ Ethers ● Ethers contain the R-O-R’ linkage, where R and R’ are hydrocarbon groups. Basically they have an embedded oxygen. ■ Aldehydes/Ketones ● The functional group is >C=O ● In an aldehyde, at least one hydrogen is bonded to the carbonyl group. ● In a ketone, the carbon is bonded to two hydrocarbon groups ■ Carboxylic Acids ● These compounds have the carboxyl group, COOH ■ Esters ● Esters have the general formula of R’COOR where R’ can be hydrogen, or a hydrocarbon group and R has to be a hydrocarbon. ■ Amines ● Amines have the general formula R3N Chapter 12: Solutions ● Types of solutions ○ Saturated - maximum amount of solute dissolved in solution. ○ Unsaturated - less solute dissolved in solution ○ Supersaturated - More solute dissolved than in a saturated solution ○ Crystallization is the process in which dissolved solute comes out of the solution and forms crystals. ● Solubility rules ○ All sodium, potassium, ammonium, and nitrate salts are soluble in water. ● Thermodynamics of the solution process ○ If the solute-solvent attraction is stronger than the solvent-solvent attraction and solute-solute attraction, the solution process is favorable or exothermic (deltaH of solution < 0). ○ If the solute-solvent interaction is weaker than the solvent-solvent and solute-solute interactions, then the solution process is endothermic (deltaH of solution > 0). ○ Like dissolves like, butdon’t use this as justificationon FRQs. ● Henry’s law - the solubility of a gas in a liquid is proportional to the pressure of the gas over the solution. Raoult’s law (not tested anymore) - the vapor pressure above the solution is the mole ● fraction multiplied by the vapor pressure of the pure solvent and the solute and then added together. 𝑃𝑇 = 𝑋𝐴 𝑃𝑝𝑢𝑟𝑒𝐴 + 𝑋𝐵 𝑃𝑝𝑢𝑟𝑒𝐵 +... ● Colloids ○ Colloids are dispersions of particles (size usu. 1-100 nm) of one substance throughout a dispersing medium made of another substance. ○ An aerosol consists of liquid droplets dispersed in a gas. ○ Mayonnaise is an emulsion, liquid droplets dispersed in another liquid. ○ Milk of magnesia is a sol, a suspension of solid particles in a liquid. Chapter 13: Kinetics Reaction Rate ● Chemical kinetics is the area of chemistry concerned with the speeds or rates at which a chemical reaction occurs. ● Rate law: ○ rate=k[A]x[ B]y ○ k = rate constant ○ A = one element, B = the other ○ X = the order of A, Y = the order of B ● Reaction order ○ Reaction order is the sum of the orders of the reactants ○ A first order reaction is a reaction whose rate depends on the reactant concentrations raised to the first power. ■ No transcendental function - zero order ■ ln(x) - first order ■ 1/x - second order ● Half-lives ○ The rate constant of a reaction (k) is equal to the negative slope of the graph of ln[M]t versus time. ○ Half life can be calculated using this equation: (half life) = 0.693/(k) ■ 0.693 is the natural log of 2, by the way. ● Reaction Mechanism ○ You need to understand some basics of thermochem: ■ Exothermic reactions = release heat = negative dH ■ Endothermic reactions = absorb heat = positive dH ○ Rate laws always include the reactants. ■ A → products ● Rate = k[A] ■ A + B → products ● Rate = k[A][B] ■ 2A → products ● Rate = k[A]2 ○ Activation energy ■ Catalysts lower the activation energy for a reaction to get started ● There are many types of catalysts ○ Heterogenous catalyst - the reactants and catalyst are in different phases, e.g. catalytic converter ○ Homogenous catalyst - same phase ○ Enzyme catalyst - the catalyst is an enzyme, or a protein ■ The stage of the reaction with the lower activation energy is the “fast step.” Chapter 14: Chemical Equilibrium quilibrium is a state in which there are no observable changes as time goes by. It is rare for a E chemical reaction to proceed in only one direction; (dynamic) equilibrium is when the forward and backward rates are equal. Equilibrium Constant ● Kc, or equilibrium concentration constant. ○ At equilibrium, the concentration of the products over the concentration of the reactants raised to their stoichiometric coefficients. ● Kp, or equilibrium pressure constant ○ At equilibrium, the pressure of the products over the pressure of the reactants raised to their stoichiometric coefficients ● The rate constant of the forward reaction divided by the rate constant of the reverse reaction is the Kc. Reaction Quotient ● R eaction Quotient, or Q, is the concentration of the products divided by the concentration of the reactants. ● It’s a test that allows you to plug in the existing conditions to assess whether the reaction is at equilibrium. ● Follow the alligator - if the Q<K, reaction proceeds toward the products, vice versa, if Q>K the reaction proceeds toward the reactants. Le Chatelier’s Principle ● A n important concept in equilibria is Le Chatelier's Principle, which states thatif an external stress is applied to a system at equilibrium, the system adjusts in such a way that the stress ispartiallyoffset as the systemreaches a new equilibrium position. ● LeChat cannot be used on FRQs - instead, one should explain using Q and K. hapter 15 & 16: Acids, Bases, Buffers C & Titration Acids ● ● ● ● ● ronsted acids are substances capable of donating protons. B Strong acids dissociate readily (100%). Water can be an acid or a base. Acids, when dissolved in water, create a solution with pH lower than 7. Strong acids: ○ HI, HBr, HCl, HNO3, HClO4, H2SO4 Bases B ● ronsted bases are substances capable of accepting protons. ● Strong bases dissociate readily (100%). This usually happens with a soluble ionic bond such as that found in NaOH. ● Bases have pH higher than 7. ● Strong bases: ○ OH ion (NaOH, KOH, etc). pH and pOH ● ● ● ● H = -log[H+] p pOH = -log[OH-] pH + pOH = 14.0 @ 25°C [OH-][H+] = 1.0e-14 @ 25°C Ka and Kb ● Ka = [H+][A-]/[HA] ○ Using Ka, you can find the concentration of any of the reactants/products. ● Kb = [HB+][OH-]/[B] K ● w = Ka x Kb = 1e-14 @25ºC ● Percent ionization is the ionized acid, or proton concentration at equilibrium divided by the initial concentration of the acid, multiplied by 100%. ● pKa = -logKa ● pKa + pKb = 14 @ 25°C Buffers ● B uffers are weak acids in which the conjugate base is present in significant concentrations. This keeps the pH ‘locked’ because if H+ ions from a strong acid are added, they will be neutralized by OH- ions/the conjugate base; if OH- ions from a strong base are added, they will be neutralized by H+ ions. ● Buffers are important enough to warrant their own equation, the Henderson Hasselbalch equation. ○ pH = pKa + log [base]/[acid] ○ [H+]=Ka x [acid]/[base] ● Perfect buffer: pH = pKa [HA] = [A-] [H+] = Ka Titration ● T itration begins with an analyte in the flask. You add a titrant to it until the analyte is neutralized. The purpose of titration is usually to find the concentration of the analyte. ● Half-equivalence point, [HA] = [A-] ○ Once you’re halfway to the equivalence point, this is called the half-equivalence point. On a graph of pH versus volume of titrant added, it occurs at half the equivalence volume. ○ Properties of the solution at this point include that the pH will be equal to the pKa (Henderson-Hasselbalch equation). ● Equivalence point, [A-] = [HA]initial or [HB+] = [B]initial ○ The equivalence point is reached once the analyte has been neutralized. Also known as the equimolar point. ○ The moles of titrant = moles of analyte ● [Acid] x Volume Acid = [Base] x Volume Base. hapter 18: Entropy, Free Energy, and C Equilibrium Three Laws of Thermodynamics . Energy can never be created or destroyed 1 2. The entropy of any isolated system always increases in a spontaneous reaction. (the entropy of the system plus surroundings always increases) 3. The entropy of a system approaches a constant value as time goes on. Entropy ● A measure of disorder, or freedom. The more energetically-equivalent states exist for a system, the more entropy that system has. ● The entropy of a perfect pure crystal at 0 kelvin is zero. (i.e. zero movement = zero microstates) ● A decrease in entropy is thermodynamically unfavored; an increase in entropy is favored. However, entropy alone cannot predict whether areactionis thermodynamically favored (spontaneous). ● The unit of entropy is J/K mol. Enthalpy T ● he heat of reaction ● The enthalpy of an element in its standard state is zero. ● Negative enthalpy indicates that the reaction is exothermic and therefore thermodynamically favored (not really); positive enthalpy means that it is endothermic and thermodynamically unfavored (not really). As with entropy, enthalpy alone cannot predict whether areactionis thermodynamically favored(spontaneous). ● The unit of enthalpy is kJ/mol ̀[ Gibbs Free Energy ( change in) Gibbs free energy is ΔH - TΔS. ● ● If Gibbs free energy is negative, the reaction is spontaneous (thermodynamically favored) in the forward direction. ● If Gibbs free energy is positive, the reaction is nonspontaneous and favored in the opposite direction. ● If Gibbs free energy is zero, the system is at equilibrium. ● There is a relationship between Keq and ΔG. ○ ΔG °= -RTlnK(notice negative sign) ○ Keq = e-ΔG ° /RT = e-Δ G °/2.5 kJ/mol (at room temperature) ● IfΔG is positive, the reaction is endergonic, meaningthat the reverse reaction is favored nd requires coupling to move forward. a ● If ΔG is negative, the reaction is exergonic, meaning that the forward reaction will be spontaneous. With +25 kJ/mol or -25 kJ/mol for ΔG, you get extreme values for equilibrium - the ● reaction will go to completion in the forward or backward reaction. ○ Example (room temperature) with ΔG°= -25 kJ/mol:Ke q = e25 kJ/mol/2.5 kJ/mol = e10. = 2.2e4 (essentially all products) Chapter 19: Electrochemistry Oxidation Number Rules ● E lements containing only one element have an oxidation number of zero. O2, F2, H2, etc. ● Monatomic ions have an oxidation number equal to the charge on the ion. Cl-, Mg2+ ● In compounds ○ Alkali metals always have an oxidation number of +1 ○ Alkaline earth metals have an oxidation number of +2 ○ Aluminum is always +3 ○ Fluorine always has an oxidation number of -1 (will bully other elements into submission to this oxidation number, like a Tyrannosaurus Regina). ○ Oxygen is almost always -2 ■ Unless it’s in a peroxide, which is -1. ■ Or in a compound with fluorine, which will make it positive. ○ Hydrogen has a number of +1 (but can be -1 in hydrides) ● In a neutral compound, oxidation numbers add up to zero. ● In a polyatomic ion, the sum of oxidation numbers is equal to the charge of the ion. Red Ox Reactions ● O xidation - Electrons are lost (ie oxidation number becomes less negative) (Loss of Electrons is Oxidation) LEO ● Reduction - Electrons are gained (ie oxidation number becomes more negative) (Gain of Electrons is Reduction) GER ● LEO the lion goes GER ● The oxidized substance loses electrons. It is also the reducing agent, because it gives the electrons to the substance that is subsequently reduced. ● The reduced substance gains electrons. It is also the oxidizing agent, because it gets the electrons from the substance that is oxidized. ● OA2+ + RA → OA + RA2+ in this example, the oxidizingagent is reduced, and the reducing agent is oxidized. Galvanic Cell ● G alvanic cells are thermodynamically favored reactions which have been harnessed to create an electric current. ∆G = - ● Reduction potential ○ Reactions are ranked for their ability to reduce, or gain electrons. ○ The opposite of reduction potential is oxidation potential. Multiply every number by -1 to get the oxidation potentials. ● The voltage of the cell is obtained by subtracting the reduction potential of the anode from the reduction potential of the cathode. Eoc ell = Eoc athode - Eoa node ● The cathode, the site of reduction, normally has thehigherreduction potential ● The cathode is where reduction occurs, while the anode is where oxidation occurs. ○ Red Cat ○ An Ox ● The wire between the two half-cells is where the electrons move, always towards the cathode (as the anode produces electrons through oxidation). ● ΔG °=-nFEcell ● Nernst equation ○ Ecell = E°cell - (0.0592/n)logQ ○ Products = the cathode = the reduced ○ Reactants = the anode = the oxidized ○ If the reactants are larger, then the voltage will increase ■ Q < 1 → logQ will be negative, resulting in a positive term ○ If the products are larger, then the voltage will decrease ■ Q > 1 → logQ will be positive, resulting in a negative term Electrolytic Cell ● E lectrolytic cells have a negative voltage and are not thermodynamically favorable. ∆G= + , as opposed to Galvanic cells, in which ∆G = - . They require a charge to be run through them to work. (Think back to coupling reactions. The outside current is actually a favorable reaction.) ● In contrast to Galvanic cells, the cathode, the site of reduction, normally has thelower reduction potential ● In general, these cells are used to break apart compounds. The amount of products produced is stoichiometrically proportional to the number of electrons, measured in coulombs, run through the cell. ● Use the equation I = q/t to figure out how many amps need to be run in order to produce the products. ● I is expressed in amps, A, which is actually C/s, coulombs per second ● F = Faraday’s number = 96485 coulombs per mole of electrons
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