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AP Chemistry Exam Review
Big Idea #1
Properties of Matter
Ratio of Masses in a Pure Sample
All elements and molecules are made up of atoms
Substances with the same atomic makeup will have same average masses
The ratio of masses of the same substance is independent of size of the substance
Molecules with the same atomic makeup (ex: H2O) will have the same ratio of average atomic masses
H2O2 ratio would be different than H2O due to the different chemical makeup
Click reveals answer and explanation.
LO 1.1: Justify the observation that the ratio of the masses of the constituent elements in any pure sample of that compound is always identical on the basis of the atomic molecular theory.
Composition of Pure Substances and/or Mixtures
Percent mass can be used to determine the composition of a substance
% mass can also be used to find the empirical formula
The empirical formula is the simplest formula of a substance
It is a ratio between the moles of each element in the substance
Quick steps to solve!
% to mass, mass to moles, divide by the smallest and multiply ‘til whole!)
The molecular formula is the actual formula of a substance
It is a whole number multiple of the empirical formula
LO 1.2: Select and apply mathematical routines to mass data to identify or infer the composition of pure substances and/or mixtures.
Identifying Purity of a Substance
Impurities in a substance can change the percent composition by mass
If more of a certain element is added from an impurity, then the percent mass of that element will increase and vice versa
When heating a hydrate, the substance is heated several times to ensure the water is driven off
Then you are simply left with the pure substance and no excess water
The mass percent of oxygen in pure glucose, C6H12O6 is 53.3 percent. A chemist analyzes a sample of glucose that contains impurities and determines that the mass percent of oxygen is 49.7 percent. Which of the follow impurities could account for the low mass percent of oxygen in the sample?
a. has the lowest percent by mass of O
LO 1.3: The student is able to select and apply mathematical relationships to mass data in order to justify a claim regarding the identity and/or estimated purity of a substance.
1 mole = 6.02 x 1023 representative particles
1 mole = molar mass of a substance
1 mole = 22.4 L of a gas at STP
LO 1.4: The student is able to connect the number of particles, moles, mass and volume of substances to one another, both qualitatively and quantitatively.
Electronic Structure of the Atom: Electron Configurations
Electrons in occupy orbitals whose energy level depends on the nuclear charge and average distance to the nucleus
Electron configurations & orbital diagrams indicate the arrangement of electrons with the lowest energy (most stable):
Electrons occupy lowest available energy levels
A maximum of two electrons may occupy an energy level
Each must have opposite spin (±½)
In orbitals of equal energy, electrons maximize parallel unpaired spins
LO 1.5: The student is able to explain the distribution of electrons in an atom or ion based upon data.
Electronic Structure of the Atom: 1st Ionization Energy
1st Ionization Energy Energy (IE) indicates the strength of the coulombic attraction of the outermost, easiest to remove, electron to the nucleus:
X(g) + IE X+(g) + e–
1st IE generally increases across a period and decreases down a group
IE generally increases as #protons increases in same energy level
IE decreases as e– in higher energy level: increased shielding, e– farther from nucleus
LO 1.6: The student is able to analyze data relating to electron energies for patterns and relationships.
Electronic Structure of the Atom: Photoelectron Spectroscopy (PES)
PES uses high-energy (X-ray) photon to excite random e– from atom
KE of ejected electron indicates binding energy (coulombic attraction) to nucleus:
BE = hvphoton – KE
Direct measurement of energy and number of each electron
Lower energy levels have higher BE
Signal size proportional to number of e– in energy level
Elements with more protons have stronger coulombic attraction, higher BE at each energy level
Image from Chemistry: A Guided Inquiry by R. S. Moog and J. J. Farrell, Wiley Publishing
LO 1.7: The student is able to describe the electronic structure of the atom, using PES data, ionization energy data, and/or Coulomb’s law to construct explanations of how the energies of electrons within shells in atoms vary.
Electronic Structure of the Atom: Higher Ionization Energies
2nd & subsequent IE’s increase as coulombic attraction of remaining e–’s to nucleus increases
X+ + IE X2+ + e–
X2+ + IEX3+ + e–
Large jump in IE when removing less-shielded core electrons
LO 1.8: The student is able to explain the distribution of electrons using Coulomb’s law to analyze measured energies.
Electronic Structure of the Atom: 1st Ionization Energy Irregularities
1st Ionization Energy Energy (IE) decreases from Be to B and Mg to Al
Electron in 2p or 3p shielded by 2s2 or 3s2 electrons, decreasing coulombic attraction despite additional proton in nucleus.
Same effect seen in 3d10-4p, 4d10-5p and 5d10-6p
1st Ionization Energy decreases from N to O and P to S
np4 contains first paired p electrons, e–-e– repulsion decreases coulombic attraction despite additional proton
The following explains these trends:
Electrons attracted to the protons in the nucleus of an atom
So the closer an electron is to a nucleus, the more strongly it is attracted (Coulomb’s law)
The more protons in a nucleus (effective nuclear force), the more strongly it attracts electrons
Electrons are repelled by other electrons in an atom. If valence electrons are shielded from nucleus by other electrons, you will have less attraction of the nucleus (again Coulomb’s law-greater the atomic radius, the greater the distance)
L.O. 1.9 The student is able to predict and/or justify trends in atomic properties based on location on periodic table and/or the shell model
Predictions with Periodic Trends
Nonmetals have higher electronegativities than metals --> causes the formation of ionic solids
Compounds formed between nonmetals are molecular
Usually gases, liquids, or volatile solids at room temperature
Elements in the 3rd period and below can accommodate a larger number of bonds
The first element in a group (upper most element of a group) forms pi bonds more easily (most significant in 2nd row, non-metals)
Accounts for stronger bonds in molecules containing these elements
Major factor in determining the structures of compounds formed from these elements
Elements in periods 3-6 tend to form only single bonds
Reactivity tends to increase as you go down a group for metals and up a group for non-metals.
46) Of the elements below, __________ is the most chemically reactive.
L.O. 1.10: Students can justify with evidence the arrangement of the periodic table and can apply periodic properties to chemical reactivity
Chemical Properties within a Group and across a Period
Group 1 metals more
reactive than group 2
Reactivity increases as you go down a group
Metals on left form basic oxides
Ex. Na2O + H2O → 2 NaOH
Nonmetals on right form form acidic oxides
Ex. SO3 + H2O → H2SO4
Elements in the middle, like Al, Ga, etc can behave amphoterically
If SiO2 can be a ceramic then SnO2 may be as well since both in the same group
LO 1.11: Analyze data, based on periodicity & properties of binary compounds, to identify patterns & generate hypotheses related to molecular design of compounds
Classic Shell Model of Atom vs
Quantum Mechanical Model
Developed by Schrodinger and the position of an electron is now
represented by a wave equation
Most probable place of finding an electron is called an ORBITAL (90% probability)
Each orbital can only hold 2 electrons with opposing spins (S, P, D & F orbitals)
Evidence for this theory:
Work of DeBroglie and PLanck that electron had wavelike characteristics
Heisenberg Uncertainty Principle - impossible to predict exact location of electron- contradicted Bohr
This new evidence caused the Shell Theory to be replaced by the Quantum Mechanical Model of the atom
Shell Model - Bohr
LO 1.12: Explain why data suggests (or not) the need to refine a model from a classical shell model with the quantum mechanical model
Shell Model is consistent with Ionization Energy Data
The patterns shown by
the IE graph can be
explained by Coulomb’s law
As atomic number increases, would expect the ionization energy to constantly increase
Graph shows that this is NOT observed. WHY NOT?
The data implies that a shell becomes full at the end of each period
Therefore the next electron added must be in a new shell farther away from the nucleus.
This is supported by the fact that the ionization energy drops despite the addition positive charge in the nucleus
LO: 1.13 Given information about a particular model of the atom, the student is able to determine if the model is consistent with specified evidence
Mass Spectrometry - evidence for isotopes
Mass spectrometry showed
that elements have isotopes
This contradicted Dalton’s early model of the atom which stated that all atoms of an element are identical
3 Br2 & two Br isotopes shown in diagram
The average atomic mass of the element can be estimated from mass spectroscopy
LO 1.14: The student is able to use the data from mass spectrometry to identify the elements and the masses of individual atoms of a specific element
Using Spectroscopy to measure properties associated with vibrational or electronic motions of molecules
IR Radiation - detects different types of bonds by analyzing molecular vibrations
UV or X-Ray Radiation
Causes electron transitions
Transitions provides info on
LO: 1.15 Justify the selection of a particular type of spectroscopy to measure properties associated with vibrational or electronic motions of molecules
A = abc
A = absorbance
a = molar absorptivity (constant for material being tested)
b = path length (cuvette = 1 cm)
c = concentration
Taken at fixed wavelength
LO1.16: Design and/or interpret the results of an experiment regarding the absorption of light to determine the concentration of an absorbing species in solution
Beer-Lambert Law - used to measure the concentration of colored solutions
Law of Conservation of Mass
N2 + 3H2 → 2NH3
LO1.17: Express the law of conservation of mass quantitatively and qualitatively using symbolic representations and particulate drawings
Use Mole Ratio in balanced equation to calculate moles of unknown substance
LO1.18: Apply the conservation of atoms to the rearrangement of atoms in various processes.
Buchner Filtration Apparatus
How much lead
(Pb2+) in water?
Pb2+(aq) + 2Cl-(aq) → PbCl2 (s)
By adding excess Cl- to the sample, all of the Pb2+ will precipitate as PbCl2
Solid product is filtered using a Buchner Filter and then dried to remove all water
Mass of PbCl2is then determined
This can be used to calculate the original amount of lead in the water
LO 1.19: Design and/or interpret data from, an experiment that uses gravimetric analysis to determine the the concentration of an analyte in a solution.
Using titrations to determine concentration of an analyte
At the equivalence point, the stoichiometric molar ratio is reached
LO1.20: Design and/or interpret data from an experiment that uses titration to determine the concentration of an analyte in a solution.
Big Idea #2
Properties Based on Bonding
http://my.hrw.com/content/hmof/science/high_school_sci/na/gr9-12/hmd_chem_9780547708089_/dlo/virtuallab/c06_00vl18/index.htmlLab to explore properties based on bond type (click on perform)
Not all ionic compounds are soluble, but those containing ammonium, nitrate, alkali metals, and halogens (except bonded to Ag, Hg and Pb) are typically
LO 2.1: Students can predict properties of substances based on their chemical formulas, and provide explanations of their properties based on particle views
Binary Acid Strength
LO 2.2: student is able to explain the
relative strengths of acids and bases based on
structure, IMF’s, & equilibrium.
The increased number of oxygen atoms pulls negative charge away from the O-H bond, weakening the attraction of the proton for the electron pair and thus strengthening the acid.
The greater the size of the negative ion, the weaker its attraction for the proton, and so the stronger the acid, and the weaker the conjugate base. HI is the strongest binary acid.
Behaviors of Solids, Liquids, and Gases
LO 2.3: The student is able to use particulate models to reason about observed differences between solid and liquid phases and among solid and liquid materials.
Kinetic Molecular Theory (KMT)
IF the temperature is not changed, no matter what else is listed in the problem, the average kinetic energy of a gas does not change. That is the definition of temperature!
All gases begin to act non-ideally (aka real) when they are at low temperatures and/or high pressures because these conditions increase particle interactions
Under the same conditions, the stronger the intermolecular attractions between gas particles, the LESS ideal the behavior of the gas
LO 2.4: The student is able to use KMT and IMF’s to make predictions about the macroscopic properties of gases, including both ideal and non-ideal behaviors
Properties of a Gas - Factors
Don’t worry about individual gas law names, but do worry about the effect of changing moles, pressure and temperature on a sample of gas
LO 2.5: Refine multiple representations of a sample of matter in the gas phase to accurately represent the effect of changes in macroscopic properties on the sample
The Ideal Gas Law
LO 2.6: The student can apply mathematical relationships or estimation to determine macroscopic variables for ideal gases
LO 2.7: The student is able to explain how solutes can be separated by chromatography based on intermolecular interactions.
Dissolving/Dissociation: Solute and Solvent
When drawing solute ions:
pay attention to size (Na+ is smaller than Cl-)
Draw charges on ion, but not on water
draw at least 3 water molecules around each
the negative dipole (oxygen side) points toward cation and the postive dipoles (H side) points towards the anion
LO 2.8: The student can draw and/or interpret representations of solutions that show the interactions between the solute and solvent.
QUESTION: Rank the six solutions above in order of increasing molarity. Pay attention to volume, and some have equal concentration
C,D, and E (tied); A and F (tied); most concentrated is B
Click reveals answer
LO 2.9: The student is able to create or interpret representations that link the concept of molarity with particle views of solutions
Molarity and Particle Views
Distillation to Separate Solutions
In the diagram above, ethanol has lower IMF’s and a resulting lower boiling point than water, so it can be heated, vaporized and condensed easily.
Ethanol hydrogen bonds as water does and is polar, but part of the ethanol has only weaker LDF’s because it’s nonpolar resulting in a lower boiling point
LO 2.10: Design/interpret the results of filtration, paper/column chromatography, or distillation in terms of the relative strength of interactions among the components.
London Dispersion Forces and Noble and Nonpolar Gases
This answer is VITAL! Remember with increased number of ELECTRONS a particle becomes more polarizable, not with increased mass!
LO 2.11: The student is able to explain the trends in properties/predict properties of samples consisting of particles with no permanent dipole on the basis of LDF’s.
Deviations from Ideal Gas Behavior
When watching the video, don’t concern yourself with Van der Walls – AP Exam focuses on LDF’s instead
LO 2.12: The student can qualitatively analyze data regarding real gases to identify deviations from ideal behavior and relate these to molecular interactions
Hydrogen bonding is seen in the following molecules: water, DNA, ammonia, HF, and alcohols. H-bonding is an attraction or force not a true intramolecular bond.
Hydrogen bonds are like a sandwich with N, O, and/or F as the bread. H will be in a intramolecular (same molecule) bond with one N, O, and/or F and have an intermolecular attraction (different molecule) with the other.
Remember this tip:
hydrogen bonds just wanna have FON
LO 2.13: The student is able to describe the relationships between the structural features of polar molecules and the forces of attraction between the particles.
Coulomb’s Law and Solubility
Ionic compounds can dissolve in polar liquids like water because the ions are attracted to either the positive or negative part of the molecule.
There is a sort of tug-of-war involved with species dissolved in water. The water pulls individual ions away from the solid. The solid is pulling individual ions back out of the water. There exists an equilibrium based on how strongly the water attracts the ions, versus how strong the ionic solid attracts the ions.
We can predict the degree of solubility in water for different ionic compounds using Coulomb's law. The smaller the ions, the closer together they are, and the harder it is for the water molecules to pull the ions away from each other. The greater the charge of the ions, the harder it is for the water to pull them away as well.
QUESTION: Predict which of the following pairs should be more soluble in water, based on Coulombic attraction.
LiF or NaF
NaF or KF
BeO or LiF
LO 2.14: Apply Coulomb’s law to describe the interactions of ions, & the attractions of ions/solvents to explain the factors that contribute to solubility of ionic compounds.
Entropy in Solutions
Do NOT say “like dissolves like.” You’ll, like... get no points.
DO refer to LDF’s, hydrogen bonding and dipole-dipole interactions
Generally speaking. There are exceptions
LO 2.15: Explain observations of the solubility of ionic solids/molecules in water and other solvents on the basis of particle views that include IMF’s and entropic effects.
Physical Properties and IMF’s
LO 2.16: Explain the properties (phase, vapor pressure, viscosity, etc.) of small and large molecular compounds in terms of the strengths and types of IMF’s.
Bonding and Electronegativity
Differences in electronegativities lead to different types of bonding*:
0.0 – 0.4: Bond is generally considered nonpolar
0.5 – 1.7: Bond is generally considered polar
> 1.7: Bond is generally considered ionic
Electronegativities are assigned values and are relative to fluorine. Electronegativity is a function of shielding / effective nuclear charge.
*Values presented are one possibility – other scales exist.
LO 2.17: The student can predict the type of bonding present between two atoms in a binary compound based on position in the periodic table and the electronegativity of the elements.
Ranking Bond Polarity
LO 2.18: The student is able to rank and justify the on the ranking of bond polarity on the basis of the locations of the bonded atoms in the periodic table.
Ionic Substances and their Properties
Ionic compounds are brittle. As the crystal structure is struck, the ions become displaced. The displaced ions will repel like charges and fracture.
LO 2.19: The student can create visual representations of ionic substances that connect the microscopic structure to macroscopic properties and/or use representations to connect microscopic structure to macroscopic properties (e.g., boiling point, solubility, hardness, brittleness, low volatility, lack of malleability, ductility, or conductivity).
Metallic Properties – Sea of Electrons
“The metallic bond is not the easiest type of bond to understand, so an analogy may help. Imagine filling your bathtub with golf balls. Fill it right up to the top. The golf balls will arrange themselves in an orderly fashion as they fill the space in the tub. Do you see any spaces between the balls? If you turn on the faucet and plug the drain, the water will fill up those spaces. What you now have is something like metallic bonding. The golf balls are the metal kernals, and the water represents the valence electrons shared by all of the atoms.”
LO 2.20: The student is able to explain how a bonding model involving delocalized electrons is consistent with macroscopic properties of metals (e.g., conductivity, malleability, ductility, and low volatility) and the shell model of the atom.
Lewis Diagrams / VSEPR
LO 2.21: The student is able to use Lewis diagrams and VSEPR to predict the geometry of molecules, identify hybridization, and make predictions about polarity.
Ionic or Covalent? Bonding Tests
Low except for some giant covalent molecules
Conduct electricity in molten and in aqueous solution
Does not conduct electricity in any state when pure, may conduct in aqueous solution (i.e., acids)
Solubility in water and organic solvents
Soluble in water
Insoluble in organic solvent
Insoluble in water, except for some simple molecule
Soluble in organic solvent
As the type of particles and forced of attraction in ionic and covalent compounds differ, their properties also differ!
Use properties of compounds to differentiate them from one another. Other tests may be performed to positively identify the compound, but are not necessary to observe types of bonds present.
http://www.mrpalermo.com/virtual-lab-conductivity.htmlhere to do a virtual lab on bonding type (chart pictured below)
LO 2.22: The student is able to design or evaluate a plan to collect and/or interpret data needed to deduce the type of bonding in a sample of a solid.
Great Lab Example
Crystal Structure of Ionic Compounds
LO 2.23: The student can create a representation of an ionic solid that shows essential characteristics of the structure and interactions present in the substance.
The +2 and -2 ions attract each other more strongly than +1 attracts -1.
The ions Mg+2 and O-2 are smaller than Na+1 and Cl-1, therefore the ions can get closer together, increasing their electrostatic attractions.
LO 2.24: The student is able to explain a representation that connects properties of an ionic solid to its structural attributes and to the interactions present at the atomic level.
LO 2.25: The student is able to compare the properties of metal alloys with their constituent elements to determine if an alloy has formed, identify the type of alloy formed, and explain the differences in properties using particulate level reasoning.
Alloys and their Properties
LO 2.26: Students can use the electron sea model of metallic bonding to predict or make claims about macroscopic properties of metals or alloys.
Metallic Solids - Characteristics
LO 2.27: The student can create a representation of a metallic solid that shows essential characteristics of the structure and interactions present in the substance.
Properties of Metallic Solids
LO 2.28: The student is able to explain a representation that connects properties of a metallic solid to its structural attributes and to the interactions present at the atomic level.
Covalent Compounds - Interactions
Graphite are sheets of carbon atoms bonded together and stacked on top of one another. The interactions between sheets is weak, much like the substance itself.
Diamond’s carbon atoms are more connected in a three dimensional structure, adding strength to the network.
LO 2.29: The student can create a representation of a covalent solid that shows essential characteristics of the structure and interactions present in the substance.
LO 2.30: The student is able to explain a representation that connects properties of a covalent solid to its structural attributes and to the interactions present at the atomic level.
Polar Covalent compounds align according to dipole-dipole interactions.
Non-Polar Covalent compounds align according to LDF’s as a solid.
LO 2.31: The student can create a representation of a molecular solid that shows essential characteristics of the structure and interactions in the substance.
Molecular Compounds - Interactions
Molecular Compound Interactions
a. Covalent bonds
b. Hydrogen bonds
c. Dipole-dipole interactions
d. London Dispersion Forces
LO 2.32: The student is able to explain a representation that connects properties of a molecular solid to its structural attributes and to the interactions present at the atomic level.
LO 2.26: Stud