Download AP Chemistry Study Guide I Made and more Study notes Chemistry in PDF only on Docsity!
Bookmarks
Unit 1: Atomic Structure and Properties Atoms, Molecules, and Ions Stoichiometry Atomic Structure and Periodicity Periodic Table and Trends Unit 2: Molecular and Ionic Compound Structure and Properties 2.1 Types of Chemical Bonds + Bonding Concepts 2.2 Intramolecular Force and Potential Energy 2.3 Structure of Ionic Solids (watch vid) 2.4 Structure of Metals and Alloys 2.5 Lewis Diagrams 2.6 Resonance and Formal Charge 2.7 VSEPR and Hybridization Unit 3: Intermolecular Forces and Properties 3.1 Intermolecular Properties 3.2 Properties of Solids 3.3 Solids, Liquids, and Gas 3.4 Ideal Gas Laws 3.5 The Kinetic Molecular Theory of Gases 3.6 Deviation from Ideal Gas Behavior 3.7 Solutions and Mixtures 3.8 Representation of Solution 3.9 Separation of Solubility and Mixtures Chromatography 3.10 Solubility 3.11 Spectroscopy and the Electromagnetic Spectrum 3.12 Photoelectric Effect 3.13 Beer-Lambert Law Unit 4 Chemical Reactions 4.2 Net-Ionic Equations 4.3 Representations of Reactions 4.4 Physical and Chemical Changes 4.5 Stoichiometry 4.6 Introduction to Titrations 4.7 Types of Chemical Reactions 4.8 Introduction to Acid-Base Reactions 4.9 Oxidation-Reduction Rates Unit 5 Kinetics 5.1 Reaction Rates 5.2 Rate Laws: An Introduction 5.3 Concentration Changes Over Time 5.4 Elementary Reactions 5.5 Collision Model 5.6 Reaction Profiles
5.8 Reaction Mechanism and Rate Law 5.9 Steady State Approximation 5.10 Multistep Reaction Energy Profiles 5.11 Catalysis Unit 6 Thermodynamics 6.1 Endothermic and Exothermic Processes 6.2 Energy Diagrams 6.3 Heat Transfer and Thermal Equilibrium 6.4 Heat Capacity and Calorimetry 6.5 Energy of Phase Changes 6.6 Introduction to Enthalpy of Reaction 6.7 Bond Enthalpies (Energies) 6.8 Enthalpies of Formation 6.9 Hess’s Law Unit 7: Chemical Equilibrium 7.1 Introduction to Equilibrium 7.2 Direction of Reversible Reactions 7.3 Reaction Quotient and Equilibrium Constant 7.4 Calculating the Equilibrium Constant 7.5 Magnitude of the Equilibrium Constant 7.6 Properties of the Equilibrium Constant 7.7 Calculating Equilibrium Concentrations 7.8 Representations of Equilibrium 7.9 Introduction to Le Chatelier’s Principle 7.10 Reaction Quotient and Le Chatelier's Principle 7.11 Introduction to Solubility Equilibria 7.12 Common-Ion Effect 7.13 pH and Solubility Unit 8 Acids and Bases 8.1 Introduction to Acids and Bases 8.2 pH and pOH of Strong Acids and Bases 8.3 Weak Acid and Base Equilibria 8.4 Acid-Base Reactions and Buffers 8.5 Acid-Base Titrations 8.6 Molecular Structure of Acids and Bases 8.7 pH and pKa 8.8 Properties of Buffers 8.9 Henderson-Hasselbalch Equation 8.10 Buffer Capacity Unit 9: Applications of Thermodynamics 9.1 Introduction to Entropy 9.2 Absolute Entropy and Entropy Changes 9.3 Gibbs Free Energy and Thermodynamic Favorability 9.4 Thermodynamic and Kinetic Control 9.5 Free Energy and Equilibrium 9.6 Coupled Reactions 9.7 Galvanic (Voltaic) Cells 9.8 Cell Potential and Free Energy
General Notes ● Qualitative Observations: descriptive ● Quantitative Observations: are numerical ○ A measuring system and units must be used ● A law summarizes what happens ● A theory (model) is an attempt to explain why it happens. Uncertainty in Precision ● Report a measurement by recording all the certain digits plus the first uncertain digit. ○ These numbers are called the significant figures of a measurement Exact vs Inexact Numbers ● Exact: value with no uncertainty ○ Definitions, counting, whole numbers, and simple fractions ● Inexact: value with uncertainty ○ Any measurement Measurement of Volume using a Buret ● The volume is read at the bottom of the liquid curve (meniscus). ● Meniscus of the liquid occurs at about 20.15 mL. Precision and Accuracy ● Precision: degree of agreement among several measurements of the same quantity ○ Reflects the reproducibility of a given type of measurement. ● Accuracy: the agreement of a particular value with the true value Significant Figures and Calculations ● Certain Digits: Numbers that stay the same no matter who measures them ● Uncertain digits: digits that must be estimated and therefore vary Rules for Counting Significant Figures
- Nonzero integers always count as significant figures.
- Zeros. There are three classes of zeros: a) Leading zeros are zeros that precede (to the left) all the nonzero digits. These do not count as significant figures ● 0.032 has 2 significant figures b) Captive zeros are zeros between nonzero digits. These always count as significant figures ● 19.04 has 4 significant figures c) Trailing zeros are zeros at the right end of the number. They are significant only if the number contains a decimal point. ● 6.200 has 4 significant figures; 500. has 3 significant figures
● 6200 has 2 significant figures Exponential Notations ● Rules for Rounding
- For multiplication or division, round the answer to the least number of significant figures a) Ex:
- For addition or subtraction, round the answer to the least number of decimal places a) Ex:
- In a series of calculations, carry the extra digits through to the final result, then round. Temperature Density Mixtures ● Mixture: has a variable composition (more than one substance) ● Mixtures can be classified as homogeneous (having visibly indistinguishable parts) or heterogeneous (having visibly distinguishable parts) ○ A homogeneous mixture is called a solution. ● Heterogeneous mixtures usually can be separated into two or more homogeneous mixtures or pure substances ● Pure substance: has constant composition ○ Are either compounds or free elements Composition of Pure Substances ● All pure substances have a fixed composition : the elements present and the ratio of those element’s atoms is the same for every sample of the compound → ○ Ex: every sample NaCl has a 1:1 ratio of sodium atoms and chlorine atoms ● Fixed ratio of atoms of each element of a compound means there is a constant mass ratio of elements in every compound ○ Ex: every sample NaCl has same 40% Na and 60.66% Cl by mass
● The atoms themselves are not changed in a chemical reaction. Characterizing the Atoms Structure ● Nucleus: protons and neutrons ○ Nucleus is very small but accounts for almost all of an atom's mass ○ Protons are positive; neutrons are neutral ● Protons (+) = Electrons (-) ○ So that atom is stable and electrically neutral The Electron ● Mostly energy, negligible mass ○ Electrons have potential energy that increases the further away they are from the nucleus → The energy an electron has depends on its distance from the nucleus ● Valence electrons in outermost s and p shell is where the chemistry of an atom takes place ○ Determines how atoms react with other atoms ● Electrons (-) orbit the nucleus ● Adding energy to element excites electrons ○ Electron moves to higher energy lvl ○ To fall back electron releases energy as light, heat, sound etc Notes ● The number of protons determines the type of element an atom is and the number of electrons determines how the atom will react ○ Gaining/losing protons changes element ● Atomic Number: number of protons ● Atomic Mass/Mass Number: total number of protons and neutrons and average of isotopes, expressed in atomic mass units (u) ○ Also called mass-to-charge ratio ● Molar Mass: atomic mass of an atom expressed in grams, is equal to one mole (units: g/mol) ○ The molar mass of diff elements are different bcuz constituent particles are different but all equal one mole ● Mole: 6.022× 10^23 of some chemical unit ○ Can be atoms, particles, people etc ● Electrons are repelled by other electrons, an electron between a valence and nucleus causes the valence to be weaker (called shielding) ● Octet Rule: in order to be stable an atom must have 8 electrons in its outermost s and p shells ● Stable: unreactive, lowest energy lvl → everything in nature wants to be stable J.J Thompson ● His model of the atom had a spherical cloud of positive charge with negative electrons randomly embedded in it
One Model of The Nuclear Atom ● Results could be explained only in terms of a nuclear atom—an atom with a dense center of positive charge (the nucleus) with electrons moving around the nucleus at a distance that is large relative to the nuclear radius. Mass Spectroscopy Three Types of Questions on the AP exam
1. Calculate avg atomic mass from mass spectrum (might be fictitious element) ● avg AM = (abundance of 1st isotope × its atomic mass) + (abundance of 2nd isotope × its atomic mass) / 100 2. Identify element from mass spectrum ● Identify isotopes masses and approximate abundances from the graph → estimate the avg atomic mass → compare estimate to elements on periodic table 3. Identify isotope from mass spectrum ● Determine the element represented by the graph using technique above → determine the number of neutrons by subtracting atomic number (protons) from mass number (protons + neutrons) Graph Analysis ● Each bar represents a different isotope ● The height of the bars represents the relative abundance ● X-axis may be labeled mass, m/z, mass charge, or atomic mass
Stoichiometry
Notes ● Reaction Stoichiometry: mass relationships between reactants and products in a chemical reaction ● Mole Ratio: conversion factor that relates amount in moles of two substances in chemical reaction ● Coefficients convey ratio of substances needed for the reaction to occur (in terms of moles) ○ One mole water = 18 grams; 1 molecule of water = 18 amu ○ Oxygen Ratio: 1 mole = 16 grams = 6.02 x 10^23 atoms ■ Round up to 100th place ● Ex: 4:5:6:4 can mean 4 molecules react with 5 molecules to produce 6 molecules and 4 molecules or 4 moles react Molar Mass Molar Mass of a Compound
- Determine how many of each atom
● Moles to mass (grams) → use molar mass from periodic table ● Moles to Volume (liters) → use 22.4 liters/mol but only gasses at standard temp and pressure Calculating Masses of Reactants and Products in Reactions
- Balance the equation for the reaction
- Convert the known masses of reactant or product to moles
- Use balanced equation to set up mole ratios
- Use the mole ratios to calculate the number of moles of desired reactant or product.
- If required, convert moles back to grams The Limiting Reactant ● Limiting Reactant: the one that is consumed first and thus limits the amount of product ● You know you have a limiting reactant problem anytime you are given amounts of both reactants ● To determine how much product can be produced in a reaction, we have to look for the reactant that is limiting ● Theoretical yield: The amount of a product formed when the limiting reactant is completely consumed ○ What you should have gotten if everything was perfect ● Actual yield: what you actually get from a reaction ● AP exam will never give you the theoretical yield → will always have to calculate it ○ Given mass of reactant
- Write balanced equation
- Process: mass of reactant given x (grams of product/grams of reactant from balanced equation ● Determination of Limiting Reactant U sing Reactant Quantities
- Balance the equation
- Do two stoichiometry problems
- Figure out how much product each reactant makes
- The one that makes the least is the limiting reagent (the other is the excess reagent)
- The lesser amount of product is the true amount made Atomic Structure and Periodicity Electromagnetic Radiation (EMR) ● EMR: energy that exhibits wavelike behavior and travels thru space at the speed of light in a vacuum ○ Ex: light from the sun, X-rays ○ All EMR travel at speed of light ○ Each form of EMR is only different in wavelength ● Photon: tiny particle of light that acts as a carrier of energy ● Shorter wavelength = higher frequency = more energy ○ So there is an inverse relationship between wavelength and frequency Essential Formulas
● Formula: ● 3 Characteristics of Waves Wavelength ( ) Frequency ( ) Speed of Light ( ) ● Distance between two consecutive peaks ● Unit = Meters ○ Is often given in nm but must always be converted ■ 1m = 10⁹ nm ● The number of waves (cycles) per second ● Unit = s⁻¹ (Hertz)
● 2.9979 X 10⁸
Albert Einstein ● The intensity of light is a measure of the number of photons in a beam ○ Greater intensity = more photons are available to release electrons ● → Energy has mass! ● Mass (kg) of a particle: ○ v = velocity ○ Can also use equation to calculate wavelength of a particle ● The Dual Nature of Light: EMR can show both wave properties and particulate matter properties ○ Electrons seem to move in an interference pattern (like a wave) and have the ability to carry energy and momentum when in motion (like a particle) → an electron is both a wave and particle The Bohr Model ● Predicted that electrons orbit the nucleus at fixed radii ○ Atoms absorb energy in the form of electromagnetic radiation → electrons move to higher energy lvl ● As electrons become more tightly bound, its energy becomes more negative ○ As the electron is brought closer to the nucleus, energy is released from the system ● Is wrong bcuz does not take into account for sublevels (s, p, d, f), orbitals, or electron spin → electrons aren't actually locked into orbits The Quantum Mechanical Model of the Atom ● Quantum Mechanical Model: specifies the probability of finding an electron in the 3D space
hold a total of 8 electrons (s=2 & p=6)→ octet rule orbitals label. Nodes ● Point where the probability of finding an electron is 0 ● The number of nodes is always one less than the principal quantum number ○ # number of nodes increases with n ● Standing wave can have a different number of nodes→ allows patterns to repeat themselves when there are more electrons around the nucleus Quantum Numbers ● Each of these orbitals is characterized by a series of numbers called quantum numbers Name Symbol Allowed Values Notes Principle quantum Number n 1, 2, 3, 4, … Describes the size and energy level of an orbital
- relative distance from nucleus Is equal to the number of sublevels ● Ex: n = 2 (2nd electron shell) → 2s & 2p orbitals n increases → orbital becomes larger and the electron spends more time farther from the nucleus + higher energy bcuz the electron is less tightly bound to the nucleus, and the energy is less negative. Angular momentum/azimuthal quantum number l 0 ≤ L ≤ n- ● Only (+) values Describes the shape of an orbital L = 0 → s orbital L = 1 → p orbital L = 2 → d orbital L = 3 → f orbital L ≤ n- ● Ex: n = 2 ○ L = 0 →
s sublvl ○ L = 1 → p sublvl Magnetic quantum number m L ≤ m ≤ L Describes the orientation of the orbital Ex: ● S sublvl has 1 orbital and l = o → m = 0 ● P sublvl has 3 orbitals & l = 1→ -1≤ m ≤ 1 Each orbital corresponds to one of 3 orientations Spin quantum number m +½. -½ Describes the spin of an e-→ an e- can only spin in a clockwise or counter-clockwise direction
- ½ ( up arrow) or -½ (down arrow) Example: 2p^5 → find n, l , m , m ● N = ● L = ● m → draw our number of orbitals → place electrons → 0 ● Ms → -½ Important Principles ● Pauli Exclusion Principle: two electrons which share an orbital cannot have the same spin → have different values of m ● Aufbau principle: electrons are placed in orbitals, shells, and subshells of increasing energy ● Hund’s rule : Every orbital in a sublevel is singly occupied before any orbital is doubly occupied + All of the electrons in singly occupied orbitals have the same spin ○ Because of the symmetrical distribution of electrons, orbitals in which the subshell is exactly half-filled or filled are more stable → so removing an electron from these atoms requires more energy Radial Probability, Penetration, and Electron Repulsion ● Electrons are attracted to the nucleus at the same time as electrons repel each other. ● Penetration: The ability of an electron to get close to the nucleus ○ Penetration depends on the shell ( n ) and subshell ( ml)
Periodic Table and Trends An Introduction to the Periodic Table Column/Families/Groups ● Tell us how many valence electrons are in the outermost s and p shells ● Elements with similar properties are found in the same group Rows/Periods ● Tell us which shell the valence electrons are found; each row has valence electrons in same energy lvl Groups of Periodic Table Group 1: Alkali Metals ● Low density, soft, silver, very reactive, all salts, form strong bonds with water Group 1: Alkali Earth Metals ● Stronger, denser, less reactive, one more valence electron Group 3-12: Transition metals ● D-block elements Diagonal row: Metalloids ● Poor conductors, both properties Next: Nonmetals ● Nonconductors, dull, brittle Group 7: Halogens ● Most reactive cuz have 7 ve- Group 8: Noble Gasses ● Non reactive cuz have 8 valence electrons and are stable ○ 2 in the s and 6 in the p ● Most gasses at room temperature ● When solid take crystal form
● Can run electricity thru noble gasses but won’t react Polyelectronic Atoms ● Polyelectronic atoms: atoms with more than one electron ● Three energy contributions in the description of an atom ○ The kinetic energy of the electrons as they move around the nucleus ○ Effective Nuclear Charge ○ The potential energy of repulsion between electrons. Effective Nuclear Charge ● Effective nuclear charge (Zeff): attraction between the nucleus and the valence electrons ○ Note: nuclear charge is the total (+) charge of the protons and attraction to all the e- ○ 2nd Note: Nuclear charge increases both down a group and across a period; Zeff weakens down a group and increases across a period ■ Zeff depends more on distance than # of protons ● When justifying trends talk about Zeff and why Weakens moving down a group ● Attraction between valence electrons and nucleus weakens bcuz of increasing number of filled energy lvls and thus distance between them ○ Valence electrons are shielded more from the nucleus ○ Coulomb's law: attraction between 2 charged particles is inversely related to the distance Strengthens from left to right across a period ● Electrons are being added to the same principal energy lvl & are more strongly attracted to the nucleus due to additional protons ○ Valence electrons are less shielded from the nucleus (due to added e- pair repulsions) ○ Coulomb's law: attraction between 2 charged particles in directly related to the magnitude of their charges Atomic Radius ● Size of an atom, distance from the nucleus to the outermost electrons (half the distance between two nuclei of a molecule) ● Is determined by how much the electrons are attracted to the (+) nucleus ○ Number of protons and electrons determines the size of atoms and ions Atomic Radius Trends ● Group Trend: increases as you go down ○ Bcuz Zeff decreases (less attraction = larger size) ● Periodic Trend: decreases as you go across ○ Bcuz Zeff increases → (more (+) nuclear charge = more attraction= smaller size) Ionic Radius ● Size of an atom when it is an ion Ionic Radius Trend ● Metals lose e-, so more p+ than e- (Stronger Zeff so…) ○ Ionic radius<atomic radius ● Nonmetals gain → more e- that p+ (weaker Zeff bcuz of stronger electron-electron repulsions so…) ○ Ionic radius>atomic radius ● Group Trend: increases as you go down ● Period Trend: decreases as you go across
○ Because Zeff is weakening ● Periodic Trend: increases across a period ○ Bcuz Zeff strengthens Electronegativity ● Measure of how much an atom in a molecule attracts shared electrons to itself; affected by atomic radius and atomic number ○ Associated with the production of a negative ion Electronegativity Trend ● Group Trend: decreases as you go down ○ Zeff decreases ● Period Trend: increases across ○ Zeff increases; atom less likely to give up electron Electron Affinity ● The energy change associated with adding an electron to a gaseous atom ○ ○ The more negative the energy, the more energy is released (exothermic) Metals ● Easier to lose electrons because the nucleus doesn't have a strong attraction to valence electrons ● Metals tend to have positive electron affinities while nonmetals tend to have negative electron affinities Reactivity ● “Ability of an atom or compound to undergo a chemical reaction with another atom” ● Metals lose e- when react so reactivity based on low ionization energy ○ Low I.E = high reactivity ● Nonmetals gain e- when react so reactivity based on high electronegativity ○ High electronegativity = high reactivity Properties of Nonmetals ● The ability to gain one or more electrons to form an anion when reacting with a metal. Thus nonmetals are elements with large ionization energies and the most negative electron affinities. ● Non-conductors (no flow of electrons) ● Dull, brittle ● Most gasses at room temperature and form crystals when solid Metalloids ● exhibit both metallic and nonmetallic properties under certain circumstances The Properties of a Group: The Alkali Metals Alkali Metals ● Most reactive of metals, low density, soft, silver, all salts, form strong bonds with water
● Note: even though hydrogen is in Group 1A of the periodic table, it behaves as a nonmetal, bcuz of its very small size → electron in the small 1s orbital is bound tightly to the nucleus
Unit 2: Molecular and Ionic Compound Structure and Properties
2.1 Types of Chemical Bonds + Bonding Concepts (watch video) Ionic Bonding Ionic Bonds ● Electrons are transferred from one atom to another creating ions ● Metal + nonmetal ○ Metals form cations (lose e- bcuz low IE) and nonmetals form anions (gain e- bcuz high IE) ○ Cations are attracted to anions; (+) & (-) attract ○ Nonmetal achieves electron configuration of next noble gas and valence orbitals of metal are emptied ● Ions are usually more stable than atoms but still unstable because are electrically charged ● Before losing/gaining electrons, atoms are neutral (no charge) ● All elements with more than one charge are metals and give away e- (+) Types of Ions ● Ions: electrically charged particle Monatomic = one type of atom (same element) ○ H+, Ca^2+, N^3- Polyatomic = many types of atoms (different elements) with a charge; small charged molecules ● Held by covalent bonds and net charge is not zero ● Elements are in imperfect bonding → slightly more stable than by itself ○ So they are ready to react with a better bond Lewis Dot Structures ● Valence electrons represented by dots, no more than two per side ● Can show rearrangement of electrons during chemical reactions ● Note: → has lost an e- so C will have only 6 ve- Binary Ionic Compounds ● Contains ions of only two elements ● Formula: cation first, then anion ○ Charges of atoms written as superscript (on top) ○ Number of atoms written as subscripts (on bottom)
