Chemistry Common Exam Review

Chemistry Common Exam Review
These are the main topics you will need to know for the final exam with practice problems

Chm.1.1.1 Analyze the structure of atoms, isotopes, and ions.
Characterize protons, neutrons, electrons by location, relative charge, relative mass (p=1, n=1, e=1/2000). • Use symbols: A= mass number, Z=atomic number • Use notation for writing isotope symbols:or U-235 • Identify isotope using mass number and atomic number and relate to number of protons, neutrons and electrons. • Differentiate average atomic mass of an element from the actual isotopic mass and mass number of specific isotopes. (Use example calculations to determine average atomic mass of atoms from relative abundance and actual isotopic mass to develop understanding).

Questions 1) How many neutrons are in boron.
5 6 11 2) Select the correct statement about subatomic particles. Electrons are negatively charged and are the heaviest subatomic particle. Protons are positively charged and the lightest subatomic particle.. Neutrons have no charge and are the lightest subatomic particle. The mass of a neutron nearly equals the mass of a proton. Electrons, protons, and neutrons all have the same mass.

Questions What is the atomic mass of an element with the
following isotopes: Isotope Abundance Magnesium % Magnesium % Magnesium %

Questions What is the atomic mass of an element with the
following isotopes: amu Isotope Abundance Magnesium % Magnesium % Magnesium %

Questions 3)The nucleus of an atom ____.
is composed of protons and neutrons is composed of protons and electrons occupies a large part of the atom is the lightest part of the atom has no charge 4)The nucleus of an atom is ____. positively charged and has a high density positively charged and has a low density negatively charged and has a high density negatively charged and has a low density

Questions 5)The number of neutrons in the nucleus of an atom can be calculated by ____. adding together the number of electrons and protons subtracting the number of electrons from the number of protons subtracting the atomic number from the mass number adding the mass number to the number of electrons

Questions 6) Isotopes of the same element have different _____.
Positions on the periodic table Atomic numbers due to a different number of electrons Atomic numbers due to a different number of protons Mass numbers due to a different number of neutrons Charges 7) In which of the following is the number of neutrons correctly represented? 199F has 0 neutrons 7533As has 108 neutrons 2412Mg has 24 neutrons 19779Au has 79 neutrons 23892U has 146 neutrons

Chm.1.1.2 Analyze an atom in terms of the location of electrons.
• Analyze diagrams related to the Bohr model of the hydrogen atom in terms of allowed, discrete energy levels in the emission spectrum. • Describe the electron cloud of the atom in terms of a probability model. • Relate the electron configurations of atoms to the Bohr and electron cloud models.

Questions Draw the Bohr Diagram using the electron configuration
1s22s22p63s1

Questions From the following pictures determine which diagram is the electron cloud model. Describe everything you know about electrons based off the electron cloud model.

Questions Which of the following electromagnetic wave have the highest frequencies? Ultraviolet waves X-rays microwaves gamma rays infrared light waves

Questions 1) Determine the type of energy released when the electron drops from n= 4 to n=1. 2) What color of light is emitted when an electron falls to the 2nd energy level from the excited state of the 3rd energy level?

Questions 1) Determine the type of energy released when the electron drops from n= 4 to n=1. 97 nm 2) What color of light is emitted when an electron falls to the 2nd energy level from the excited state of the 3rd energy level? red

Chm 1.1.3 Understand that energy exists in discrete units called quanta. Describe the concepts of excited and ground state of electrons in the atom: 1. When an electron gains an amount of energy equivalent to the energy difference, it moves from its ground state to a higher energy level. 2. When the electron moves to a lower energy level, it releases an amount of energy equal to the energy difference in these levels as electromagnetic radiation (emissions spectrum). • Articulate that this electromagnetic radiation is given off as photons. • Understand the inverse relationship between wavelength and frequency, and the direct relationship between energy and frequency.

Chm 1.13 continued • Use the “Bohr Model for Hydrogen Atom” and “Electromagnetic Spectrum” diagrams from the Reference Tables to relate color, frequency, and wavelength of the light emitted to the energy of the photon. • Explain that Niels Bohr produced a model of the hydrogen atom based on experimental observations. This model indicated that: 1. an electron circles the nucleus only in fixed energy ranges called orbits; 2. an electron can neither gain or lose energy inside this orbit, but could move up or down to another orbit; 3. that the lowest energy orbit is closest to the nucleus. • Describe the wave/particle duality of electrons.

Questions When an electron moves from a lower to a higher energy level, the electron _____. always doubles its energy absorbs a continuously variable amount of energy absorbs a quantum of energy moves closer to the nucleus An electron that has absorbed energy is said to be in its ? Ground state Excited State Electromagnetic State Electrostatic State None of the above

Questions What is the maximum number of electrons in the second principal energy level? 2 8 18 32 The formula 2n2 represents _____. The number of orbitals in a sublevel The maximum number of electrons that can copy an energy level The number of sublevels in any energy level None of the above

1.1.4 Use the symbols for and distinguish between alpha ( He), and beta ( e-) nuclear particles, and gamma (γ) radiation include relative mass). • Use shorthand notation of particles involved in nuclear equations to balance and solve for unknowns. • Compare the penetrating ability of alpha, beta, and gamma radiation. • Conceptually describe nuclear decay, including: 1. Decay as a random event, independent of other energy influences 2. Using symbols to represent simple balanced decay equations 3. Half-life (including simple calculations) • Compare radioactive decay with fission and fusion.

Questions Uranium-235 atoms have how many protons (atomic number 92)?
143 235 327 What is the charge on an alpha particle? neutral positive negative

Questions What is the change in atomic number when an atom emits a beta particle? decrease by 2 decrease by 1 there is no change increase by 1 increase by 2 What is the change in atomic mass when an atom emits gamma radiation? decreases by 2 decreases by 1 remains the same increases by 1 increases by 2

Questions Which of the following particles is needed to complete this nuclear equation? Mn + 21H ---> _______ n 5627 Co 2725 Mn 5526 Fe 5824 Cr To what element does polonium–214 decay to when it loses an alpha particle? (atomic number is 84) 21082 Pb 21082Po 21482 Pb 21482 Rn

Questions If the half-life of a radioactive material is 8 years, how many years will it take for one half of the original amount of material to decay? 2 years 4 years 8 years 16 years 32 years If the half-life of a sodium-24 is 15 hours, how much remains from an 8.0 gram sample after 60 hours? 0.5 grams 1.0 grams 2.0 grams 4.0 grams

1.2.1 Describe metallic bonds: “metal ions plus ‘sea’ of mobile electrons”. • Describe how ions are formed and which arrangements are stable (filled d-level, or half-filled d-level). • Appropriately use the term cation as a positively charged ion and anion as negatively charged ion. • Predict ionic charges for representative elements based on valence electrons. • Apply the concept that sharing electrons form a covalent compound that is a stable (inert gas) arrangement. • Draw Lewis (dot diagram) structures for simple compounds and diatomic elements indicating single, double or triple bonds.

Questions In which of the following is the symbol for the ion and the number of electrons it contains given correctly? S2- has 2 electrons Br – has 34 electrons Al3+ has 16 electrons Ca2+ has 18 electrons H+ has one electron Which element when combined with fluorine would most likely form an ionic compound? lithium carbon phosphorus chlorine

Questions Compounds that are composed of ions ____.
are molecular compounds have relatively high melting and boiling points are for the most part composed of two or more metallic elements fit all of the above descriptions

1.2.2 Determine that a bond is predominately ionic by the location of the atoms on the Periodic Table (metals combined with nonmetals) or when ΔEN > 1.7. • Determine that a bond is predominately covalent by the location of the atoms on the Periodic Table (nonmetals combined with nonmetals) or when ΔEN < 1.7. • Predict chemical formulas of compounds using Lewis structures

Questions Which of the following occurs in an ionic bond?
Oppositely-charged ions attract Two atoms share two electrons Two atoms share more than two electrons Like-charged ions attract A compound held together by ionic bonds is ____. Formed by sharing electrons Formed by transferring electrons Held together by Electrostatic forces Answer a and c Answer b and c

Questions Molecular compounds are usually ____.
composed of two or more transition elements composed of positive and negative ions composed of two or more nonmetallic elements exceptions to the law of definite proportions solids at room temperature

1.2.3 Explain why intermolecular forces are weaker than ionic, covalent or metallic bonds • Explain why hydrogen bonds are stronger than dipole-dipole forces which are stronger than dispersion forces • Apply the relationship between bond energy and length of single, double, and triple bonds (conceptual, no numbers).

1.2.3 continued • Describe intermolecular forces for molecular compounds.  H-bond as attraction between molecules when H is bonded to O, N, or F. Dipole-dipole attractions between polar molecules.  London dispersion forces (electrons of one molecule attracted to nucleus of another molecule) – i.e. liquefied inert gases.  Relative strengths (H>dipole>London/van der Waals).

Questions How many electrons are shared in a single covalent bond?
1 2 3 4 8 Which element forms diatomic molecules held together by a triple covalent bond? hydrogen carbon oxygen fluorine nitrogen

Questions Which compound would you expect to have the higher boiling point, CH4 or NH3 ? Explain

Questions Which compound would you expect to have the higher boiling point, CH4 or NH3 ? Explain NH3 because it has hydrogen bonding which is stronger than london disperson/van der waals

Questions What element can form single, double or triple bonds with itself? hydrogen carbon oxygen fluorine nitrogen

Questions Which of the forces of molecular attraction is the strongest? dipole interaction dispersion hydrogen bond

1.2.4 Write binary compounds of two nonmetals: use Greek prefixes (di-, tri-, tetra-, …) • Write binary compounds of metal/nonmetal* • Write ternary compounds (polyatomic ions)* • Write, with charges, these polyatomic ions: nitrate, sulfate, carbonate, acetate, and ammonium. Know names and formulas for these common laboratory acids: HCl, HNO3, H2SO4, HC2H3O2, (CH3COOH)

Questions Which set of chemical name and chemical formula for the same compound is correct? ammonium sulfite, (NH4)2S iron(lll)phosphate, FePO4 potassium chloride, K2CI magnesium dichromate, MgCrO4 lithium carbonate, LiCO3 The correct formula for sulfuric acid is _____. H2SO3 H2SO4 H2SO2 H2SO H2S

Questions What is the formula unit of aluminum oxide?
AlO Al2O Al3O AlO3 Al2O3 Which of the following compounds has the formula KNO3 Potassium nitrate Potassium nitride Potassium nitrite Potassium nitrogen oxide

1.2.5 Explain how ionic bonding in compounds determines their characteristics: high MP, high BP, brittle, and high electrical conductivity either in molten state or in aqueous solution. • Explain how covalent bonding in compounds determines their characteristics: low MP, low BP, poor electrical conductivity, polar nature, etc. • Explain how metallic bonding determines the characteristics of metals: high MP, high BP, high conductivity, malleability, ductility, and luster

1.2.5 continued Apply Valence Shell Electron Pair Repulsion Theory (VSEPR) for these electron pair geometries and molecular geometries, and bond angles - Electron pair - Molecular (bond angle); Linear framework – linear; Trigonal planar framework– trigonal planar, bent; Tetrahedral framework– tetrahedral, trigonal pyramidal, bent; Bond angles (include distorting effect of lone pair electrons – no specific angles, conceptually only)

1.2.5 continued • Describe bond polarity. Polar/nonpolar molecules (relate to symmetry) ; relate polarity to solubility—“like dissolves like” Describe macromolecules and network solids: water (ice), graphite/diamond, polymers (PVC, nylon), proteins (hair, DNA) intermolecular structure as a class of molecules with unique properties.

Questions What shape and polarity does a water molecule have?
pyramidal and polar bent and polar linear and nonpolar linear and polar What is the shape and polarity of methane, CH4? tetrahedral and polar tetrahedral and nonpolar pyramidal and nonpolar

Questions Which one of the following molecules has
tetrahedral geometry? a. XeF4 b. BF3 c. AsF5 d. CF4 e. NH3

Questions According to VSEPR theory, the shape of the PH3
molecule is best described as a. linear. b. trigonal planar. c. tetrahedral. d. bent. e. trigonal pyramidal

1.3.1 Using the Periodic Table, Groups (families)
• Identify groups as vertical columns on the periodic table. • Know that main group elements in the same group have similar properties, the same number of valence electrons, and the same oxidation number. • Summarize that reactivity increases as you go down within a group for metals and decreases for nonmetals.

1.3.1 continued Periods • Identify periods as horizontal rows on the periodic table. Metals/Nonmetals/Metalloids • Identify regions of the periodic table where metals, nonmetals, and metalloids are located. Classify elements as metals/nonmetals/metalloids based on location.

1.3.1 Representative elements (main group) and transition elements
• Identify representative (main group) elements as A groups or as groups 1, 2, • Identify alkali metals, alkaline earth metals, halogens, and noble gases based on location on periodic table. • Identify transition elements as B groups or as groups 3-12.

Questions A mystery element Q is a non-lustrous solid and a poor conductor of electricity. To what category of elements does it belong? Metals Nonmetals Transition metals Semimetals metalloids A column of elements in the periodic table is known as a ____. row list group transition period

Questions Identify each of the following as a metal, nonmetal, or metalloid and as representative elements or transition elements: A. Neon b. Arsenic c. Zinc d. Magnesium

Questions Identify each of the following as a metal, nonmetal, or metalloid and as representative elements or transition elements: A.Neon b. Arsenic c. Zinc d. Magnesium Nonmetal metalloid metal metal Represent Represent Transition represent

Questions Which of the following elements is in the same period as phosphorous? carbon magnesium nitrogen oxygen

1.3.2 Using the Periodic Table, Atomic and ionic radii
• Define atomic radius and ionic radius. • Know group and period general trends for atomic radius. • Apply trends to arrange elements in order of increasing or decreasing atomic radius. Explain the reasoning behind the trends. • Compare cation and anion radius to neutral atom.

1.3.2 continued Metallic Character
• Compare the metallic character of elements. • Use electron configuration and behavior to justify metallic character. (Metals tend to lose electrons in order to achieve the stability of a filled octet.) • Relate metallic character to ionization energy, electron affinity, and electronegativity.

1.3.2 Electron configurations/valence electrons/ionization energy/electronegativity • Write electron configurations, including noble gas abbreviations (no exceptions to the general rules). Included here are extended arrangements showing electrons in orbitals. • Identify s, p, d, and f blocks on Periodic Table. • Identify an element based on its electron configuration. (Students should be able to identify elements which follow the general rules, not necessarily those which are exceptions.)

1.3.2 continued Determine the number of valence electrons from electron configurations. • Predict the number of electrons lost or gained and the oxidation number based on the electron configuration of an atom. • Define ionization energy and know group and period general trends for ionization energy. Explain the reasoning behind the trend. • Apply trends to arrange elements in order of increasing or decreasing ionization energy. Define electronegativity and know group and period general trends for electronegativity. Explain the reasoning behind the trend. • Apply trends to arrange elements in order of increasing or decreasing electronegativity.

Questions What happens to electronegativity as I go down a group.
Electronegativity decreases Electronegativity increases Electronegativity first increases then decreases Electronegativity first decreases then increases none of the above If only two electrons occupy two p orbitals, what is the direction of the spins of these two electrons? Both up and down Opposite Parallel None of these

Questions Which element in EACH pair has the larger ionization energy? Why? A. Li or N B. Kr or Ne C. Cs or Li Why do atomic radii increase as you move down a group?

Questions Which element in EACH pair has the larger ionization energy? Why? A. Li or N more valence electrons B. Kr or Ne more stable C. Cs or Li higher up on family Why do atomic radii increase as you move down a group? Because you increase the number of energy levels

Questions Which of the following particles are free to drift in metals? protons electrons neutrons pions cations What is the basis of a metallic bond? the attraction of metal ions for mobile electrons the attraction between neutral metal atoms the neutralization of protons by electrons the attraction of oppositely-charged ions the sharing of two valence electrons between two atoms

Questions How does atomic radius change from left to right across a period in the periodic table? it tends to decrease it tends to increase it does not change it first increases, then decreases it first decreases, then increases

Questions What is the electron configuration of sulfur?
1s22s22p63s23p3 1s22s22p63s23p4 1s22s22p63s23p5 1s22s22p63s23p6

Questions How many unpaired electrons are there in a sulfur atom (atomic number 32)? 1 2 3 4

Questions What is the electron configuration of the calcium ion?
1s22s22p63s23p6 1s22s22p63s23p4 4s2 1s22s22p63s23p5 4s1 1s22s22p63s2 1s22s22p63s23p3 What does nitrogen do in order to achieve a noble-gas electron configuration? gains 2 electrons gains 3 electrons loses 2 electrons loses 3 electrons none of these

Questions The electron configuration of potassium is ____.
1s22s22p23s23p24s1 1s22s22p103s23p3 1s22s23s23p63d1 1s22s22p63s23p64s1 Which of the following electron configuration is the most stable 4d55s2 4d45s1 4d35s2 4d25s4

Questions Which of the following factor contributes to the relatively greater atomic size of the higher-atomic-number elements within a particular family of the periodic table? more shielding of the outer electrons by the inner electrons larger nuclei greater number of protons smaller number of valence electrons Atomic size generally ____. increases as you move from left to right across a period decreases as you move down a group remains constant within a period decreases as you move from left to right across a period

Questions The periodic law states that there is a periodic repetition of the physical and chemical properties of elements _____. When they are arranged in order of increasing atomic mass If only metals are considered When they are arranged in order of increasing atomic radii When they are arranged in order of increasing atomic number

Questions Which of the following factors contributes to the greater ionization energy of the elements on the right side of a period in the periodic table? more shielding by inner electrons larger nuclei greater number of protons in nuclei with electrons at the same energy level smaller number of valance electrons Why is the radius of a positive ion always less than the radius of its neutral atom? The nucleus pulls the remaining electrons in closer The number of protons is increased The atomic orbitals contract all by themselves Electron speed are reduced

Questions Which of the following elements has the lowest electronegativity? lithium carbon oxygen fluorine

2.1.1 Explain physical equilibrium: liquid water-water vapor. Vapor pressure depends on temperature and concentration of particles in solution. (conceptual only – no calculations) • Explain how the energy (kinetic and potential) of the particles of a substance changes when heated, cooled, or changing phase. • Identify pressure as well as temperature as a determining factor for phase of matter. • Contrast heat and temperature, including temperature as a measure of average kinetic energy, and appropriately use the units Joule, Celsius, and Kelvin.

2.1.2 Define and use the terms and/or symbols for: specific heat capacity, heat of fusion, heat of vaporization. • Interpret the following:  heating and cooling curves (noting both significance of plateaus and the physical states of each segment  Phase diagrams for H2O and CO2, • Complete calculations of: q=mCpΔT, q = mHf , and q = mHv using heatling/cooling curve data. • Explain phase change calculations in terms of heat absorbed or released (endothermic vs. exothermic processes).

Questions What is the amount of heat required to raise the temperature of 100 g of aluminum by 10C? (specific heat of aluminum = 0.21 cal /g x C 0.21 cal 2.1 cal 210 cal 21,000 cal None of the above Two objects are sitting next to each other in direct sunlight. Object A gets hotter than object B. Object A has a higher specific heat than object B. Object A has a lower specific heat than object B. Both objects have the same specific heat.

Questions Calculate the amount of heat necessary to raise
the temperature of 12.0 g of water from 15.4°C to 93.0°C. The specific heat of water = 4.18 J/g·°C. a J b. 324 J c. 389 J d. 931 J e. 3,890 J

2.1.3 Draw phase diagrams of water and carbon dioxide (shows how sublimation occurs). Identify regions, phases and phase changes using a phase diagram. • Use phase diagrams to determine information such as (1) phase at a given temperature and pressure, (2) boiling point or melting point at a given pressure, (3) triple point of a material.

Phase Diagram

Phase Diagram Questions
Determine the NORMAL melting point Determine the NORMAL boiling point Determine the triple point conditions Determine the critical point conditions

Phase Diagram Questions
Determine the NORMAL melting point Zero Celsius Determine the NORMAL boiling point 100 celsius Determine the triple point conditions atm and 0.01 celsius Determine the critical point conditions atm and C

2.1.4 Recognize that, for a closed system, energy is neither lost nor gained only transferred between components of the system. • Complete calculations of: q=mCpΔT, q = mHf , q = mHv, and q lost=(-q gain) in water, including phase changes, using laboratory data

2.1.5 Identify characteristics of ideal gases.
• Apply general gas solubility characteristics. • Apply the following formulas and concepts of kinetic molecular theory. 1. 1 mole of any gas at STP=22.4 L 2. Ideal gas equation (PV=nRT), Combined gas law (P1V1/T1 = P2V2/T2) and applications holding one variable constant: for PV=k, P1V1 = P2V2; for V/T=k, V1/T1= V2/T2; for P/T=k, P1/T1 = P2/T2. Note: Students should be able to derive and use these gas laws, but are not necessarily expected to memorize their names.

2.1.5 continued 3. Avogadro’s law (n/V=k), n1/V1 = n1/V2
4. Dalton’s law (Pt=P1+P2+P3 …) 5. Vapor pressure of water as a function of temperature (conceptually).

Questions What happens to the pressure of a gas inside a container if the temperature of the gas is lowered? The pressure increases The pressure does not change The pressure decreases As the temperature of a fixed volume of a gas increases, the pressure will____. vary inversely decrease be unchanged increase What happens to the temperature of a gas when it is compressed The temperature increases The temperature does not change The temperature decreases

Questions A sample of gas occupies 40.0 mL at -123C. What volume does the sample occupy at 27C? 182 mL 8.80 mL 80.0 mL 20.0 mL If a balloon containing 1,000.0 L of gas at 50.0 C and kPa rises to an altitude where the pressure is 50.5 kPa and the temperature is 10.0 C. the volume of the balloon under these new conditions would be ____. 401L 569 L 1760 L

2.2.1 Explain collision theory – molecules must collide in order to react, and they must collide in the correct or appropriate orientation and with sufficient energy to equal or exceed the activation energy. • Interpret potential energy diagrams for endothermic and exothermic reactions including reactants, products, and activated complex.

Questions If the heat involved in a chemical reaction has a negative sign, ____. heat is lost to the surroundings heat is gained from the surroundings no heat is exchanged in the process The following equation shows the reaction that occurs when nitroglycerine explodes. 4C3H5O9N3  12CO2 + 6N2 + O2 + 10H2O kcal This reaction is ____. endothermic exothermic a combination reaction a combustion reaction

Questions What is the standard heat of reaction for this reaction:
Zn(s) + Cu2 + (aq)  Zn2 + (aq) +Cu(s) (Hto for Cu2 = kJ/mol; Hto for Zn2 + = kJ/mol; Hto for Zn = 0 kJ/mol; Hto for Cu = 0 kJ/mol ) kJ per mole -88.0 kJ per mole 88.0 kJ per mole 216.8 kJ per mole

2.2.2 Precipitate formation (tie to solubility rules)
• Product testing - Know the tests for some common products such as oxygen, water, hydrogen and carbon dioxide: burning splint for oxygen, hydrogen or carbon dioxide, and lime water for carbon dioxide. Include knowledge and application of appropriate safety precautions. • Color Change – Distinguish between color change as a result of chemical reaction, and a change in color intensity as a result of dilution. Temperature change – Tie to endothermic/exothermic reaction. Express ΔH as (+) for endothermic and (–) for exothermic.

Questions Predict which one of the following compounds would be insoluble in water? NaCl HCl CF4 CuSO4

2.2.3 Write and balance chemical equations predicting product(s) in a reaction using the reference tables. • Identify acid-base neutralization as double replacement. • Write and balance ionic equations. • Write and balance net ionic equations for double replacement reactions.

2.2.3 continued Recognize that hydrocarbons (C,H molecule) and other molecules containing C, H, and O burn completely in oxygen to produce CO2 and water vapor. • Use reference table rules to predict products for all types of reactions to show the conservation of mass. • Use activity series to predict whether a single replacement reaction will take place. • Use the solubility rules to determine the precipitate in a double replacement reaction if a reaction occurs.

Questions Which of the following is true for all chemical reactions?
The total mass of the reactants increases. The total mass of the products is greater than the total mass of the reactants. Water is given off The total mass of the reactants equals the total mass of the products.

Predict the products for the following
2AgNO3+H2S  2) Ni (s) + MgCl2(aq)  3) 2C4H10(l) +13 O2(g) 

Predict the products for the following
2AgNO3+H2S  Ag2S +2HNO3 2) Ni (s) + MgCl2(aq) NiCl2 + Mg 3) 2C4H10(l) +13 O2(g) 8CO2 +10H2O

Write the net, complete and spectator ions
Pb(NO3)2(aq) + H2SO4(aq)PbSO4(s) +2HNO3(aq)

Write the net, complete and spectator ions
Pb(NO3)2(aq) + H2SO4(aq)PbSO4(s) +2HNO3(aq) Complete: Pb2+(aq) + 2NO3-(aq) + 2H+(aq) +SO42-(aq)  PbSO4(s) + 2H+(aq) + 2NO3-(aq) Net: Pb2+(aq) + SO42-(aq)  PbSO4(s) Spectator: NO3- and H+

Questions The coefficients are missing from the skeleton equation below. NH3(g) + O2  N2(g) + H2O(l) The correct order of the missing coefficient is ___. 4, 3, 2, 6 2, 1, 2, 3 1, 3, 1, 3 3, 4, 6, 2

Questions Write a balanced equation for the combination reaction that takes place when iron(III) oxide is formed from it constituent elements. Fe2+O3  Fe2O3 2Fe +3O  Fe2O3 4Fe +3O2  2Fe2O3 3Fe +O  Fe3O Fe+O3  FeO

Questions Identify the indicators of chemical reactions that would help you distinguish between these two reactions. Write a balanced chemical equation for each reaction (include phases). Identify the type of reaction. 1. Sodium metal dropped into a beaker of water. 2. Silver nitrate is added to sodium chloride

Questions Identify the indicators of chemical reactions that would help you distinguish between these two reactions. Write a balanced chemical equation for each reaction (include phases). Identify the type of reaction. 1. Sodium metal dropped into a beaker of water. Na (s) + H2O (l)  NaOH (aq) +H2(g) Gas released 2. Silver nitrate is added to sodium chloride AgNO3 (aq) + NaCl (aq)  AgCl (s) + NaNO3 (aq) Formation of a precipitate

Questions In order for the reaction 2Al +6HCl  2AlCl3 +3H2 (g) to occur, which of the following must be true? Al must be above Cl on the activity series. Al must be above H on the activity series. Heat must be supplied for the reaction A gas must be formed.

Questions A double replacement reaction takes place when aqueous cobalt(III) chloride reacts with double aqueous lithium hydroxide. One of the products of this reaction would be____. Co(OH) 3 Co(OH) 2 LiCo3 LiCl3 Cl3OH

2.2.4 Interpret coefficients of a balanced equation as mole ratios.
• Use mole ratios from the balanced equation to calculate the quantity of one substance in a reaction given the quantity of another substance in the reaction. (given moles, particles, mass, or volume and ending with moles, particles, mass, or volume of the desired substance)

Questions How many moles of tungsten atoms are there in 4.8 x 1025 atoms of tungsten? a) 1.3 x 10-2 b) 8.0 x 101 c) 1.5 x 104 d) 2.6 x 1023 e) 2.9 x 1049 How many atoms are there in 5.20 mol of hafnium? a) 8.64 x 10-24 b) 2.91 x 10-2 c) 9.31 x 102 d) 1.16 x 1023 e) x 1024

Questions How many moles of aluminum are needed to react completely with 1.8 mol of FeO? 2Al(s) + 3FeO(s)  3Fe(s) + Al2O3(s) 0.9 mol 1.2 mol 1.8 mol 2.7 mol 3.6 mol

Questions What is the mass in grams of 5.90 mol C8H18? a) g b) 19.4 g c) 389 g d) 673 g e) 3.55 x 1024 g What is the volume, in liters, of mol of C3H8 gas at STP? L 16.8 L 29.9 L 739 L 1310 L

2.2.5 Calculate empirical formula from mass or percent using experimental data. • Calculate molecular formula from empirical formula using molecular weight. • Determine percentage composition by mass of a given compound. • Perform calculations based on percent composition. • Determine the composition of hydrates using experimental data.

Questions What is the gram formula mass of chromic sulfate, Cr2(SO4)3?
a) 148 g b) 200 g c) 288 g d) 344 g e) 392 g

Questions When glucose is consumed it reacts with oxygen in the body to produce carbon dioxide, water, and energy. How many grams of carbon dioxide would be produced if 45 g of C6H12O6 completely reacted with oxygen? C6H12O6 + 6 O2  6 CO2 + 6 H2O 1.5 g 1.8 g 11 g 66 g 12,000 g

Questions If 60.2 grams of Hg combines completely with 24.0 grams of Br to form a compound, what is the percent composition of Hg in the compound? 28. 5% 39.9% 71.5% 60.1% 251% All of the following are empirical formulas EXCEPT ____. Na2SO4 C6H5Cl N2O4 Sn3(PO4)4

Questions What is the empirical formula of a compound that is 50.7% antimony and 49.3% selenium by weight? SbSe SbSe2 Sb2Se Sb2Se3 Sb3Se2

3.1.1 Understand qualitatively that reaction rate is proportional to number of effective collisions. • Explain that nature of reactants can refer to their complexity and the number of bonds that must be broken and reformed in the course of reaction. • Explain how temperature (kinetic energy), concentration, and/or pressure affects the number of collisions.

3.1.1 continued Explain how increased surface area increases number of collisions. • Explain how a catalyst lowers the activation energy, so that at a given temperature, more molecules will have energy equal to or greater than the activation energy.

Questions Why does a catalyst cause a reaction to proceed faster?
There are more collisions per second The collisions occur with greater energy The activation energy is lowered There are more collisions per second and the collisions are of greater energy

3.1.2 Define chemical equilibrium for reversible reactions.
• Distinguish between equal rates and equal concentrations. • Explain equilibrium expressions for a given reaction. • Evaluate equilibrium constants as a measure of the extent that the reaction proceeds to completion.

3.1.3 Determine the effects of stresses on systems at equilibrium. (Adding/ removing a reactant or product; adding/removing heat; increasing/decreasing pressure) • Relate the shift that occurs in terms of the order/disorder of the system.

Questions What happens to a reaction at equilibrium when more reactants are added to the system? The reaction makes more products The reaction makes more reactants The reaction is unchanged

Questions In an exothermic reaction at equilibrium, what is the effect of lowering the temperature? The reaction makes more products The reaction makes more reactants The reaction is unchanged What is the effect of adding more water to the following equilibrium reaction. 4 C3H5O9N3 <---> 12CO2 + 6N2 + O2 + 10H2O kcal More C3H5O9N3 is made CO2 concentration increases The equilibrium is pushed in the direction of the products Nothing

3.2.1 Distinguish between acids and bases based on formula and chemical properties. • Differentiate between concentration (molarity) and strength (degree of dissociation). No calculation involved. • Use pH scale to identify acids and bases.

3.2.1 continued Interpret pH scale in terms of the exponential nature of pH values in terms of concentrations. • Relate the color of indicator to pH using pH ranges provided in a table. • Compute pH, pOH, [H+], and [OH-]. Distinguish properties of acids and bases related to taste, touch, reaction with metals, electrical conductivity, and identification with indicators such as litmus paper and phenolphthalein.

Questions If the hydrogen ion concentration is 10-1 M, what is the nature of the solution? acidic alkaline neutral If the [H+] in a solution is 1 x 10-1 mol/L, then the [OH-] is _____. 1 x 10-1 mol/L 1 x mol/L 1 x mol/L cannot be determined

Which of these solutions is the most basic?
[H+] = 1 x 10-2 M [OH-] = 1 x 10-4 M [H+] = 1 x M [OH-] = 1 x 10-13 If the hydroxide ion concentration s M, what is the pH of the solution? 1 4 7 10 14 The pH pf a solution in which [OH-] = 1 x 10-4 M is ______. 4.0 10.0 –4.0 –10.0

Questions Which of these solutions is the most basic?
[H+] = 1 x 10-2 M [OH-] = 1 x 10-4 M [H+] = 1 x M [OH-] = 1 x 10-13 If the hydroxide ion concentration s M, what is the pH of the solution? 1 4 7 10 14 The pH pf a solution in which [OH-] = 1 x 10-4 M is ______. 4.0 10.0 –4.0 –10.0

Questions When an acid reacts with a base what compound(s) is/are formed? A salt only Water only Metal oxides only A salt and water What is the name of H2SO4? Hyposulfuric acid Hydrosulfuric acid Sulfuric acid Sulfurous acid Hydrosulfite acid

3.2.3 Compute concentration (molarity) of solutions in moles per liter. • Calculate molarity given mass of solute and volume of solution. • Calculate mass of solute needed to create a solution of a given molarity and volume. • Solve dilution problems: M1V1 = M2V2.

3.2.3 Perform 1:1 titration calculations: MAVA = MBVB
• Determine the concentration of an acid or base using titration. Interpret titration curve for strong acid/strong base

Questions Explain what is happening in the graph below. What type of acid and base are present? How can you tell?

Questions Explain what is happening in the graph below. What type of acid and base are present? How can you tell? As you add more base, the pH increases. The equivalence point is where the reaction has completely neutralized. This is a strong base with a strong acid because the equivalence point is right at a pH of 7

questions What is the number of moles of solute in 250 mL of a 0.4M solution? 0.1 mol 0.16 mol 0.62 mol 1.6 mol 1 mol

Questions How many mL of 3M HCl is needed to make 300 mL of 0.1M HCl?
What mass of sucrose, C12H22O11, is needed to make 500 mL of 0.2 M solution? 34.2 g 100 g 17.1 g 68.4 g

3.2.4 Identify types of solutions (solid, liquid, gaseous, aqueous).
• Define solutions as homogeneous mixtures in a single phase. • Distinguish between electrolytic and nonelectrolytic solutions. • Summarize colligative properties (vapor pressure reduction, boiling point elevation, freezing point depression, and osmotic pressure).

Questions CIRCLE the condition that causes sugar to dissolve FASTER in water. a) As a whole cube OR in granulated form b) At higher temperature OR at a lower temperature c) Allowed to stand OR when stirred

Choose the electrolytes from the following list
NaCl(aq) NH3 (aq) KBr(aq) Fe(OH)3 (aq) Ca(OH)2 (aq) HCl (aq) HC2H3O2 (aq)

Choose the electrolytes from the following list
Choose the electrolytes from the following list. Also tell WHY they are good electrolytes NaCl(aq) NH3 (aq) KBr(aq) Fe(OH)3 (aq) Ca(OH)2 (aq) HCl (aq) HC2H3O2 (aq)

3.2.5 Use graph of solubility vs. temperature to identify a substance based on solubility at a particular temperature. • Use graph to relate the degree of saturation of solutions to temperature.

Questions How much Sodium Nitrate can be dissolved in 50 grams of water at 80 °C What type of solution is created if 25 grams of NaCl are dissolved in 100 grams of water at 40°C

Questions How much Sodium Nitrate can be dissolved in 50 grams of water at 80 °C ~63 grams What type of solution is created if 25 grams of NaCl are dissolved in 100 grams of water at 40°C unsaturated

Questions Which of the following substances is less soluble in hot water that cold water? CO2 (g) NaCl (s) NaNO3 (s) KBr (s)

3.2.6 Develop a conceptual model for the solution process with a cause and effect relationship involving forces of attraction between solute and solvent particles. A material is insoluble due to a lack of attraction between particles. • Describe the energetics of the solution process as it occurs and the overall process as exothermic or endothermic. • Explain solubility in terms of the nature of solute-solvent attraction, temperature and pressure (for gases).

Questions What is primarily responsible for the water’s ability to dissolve substances? Diffusion Water’s polarity Ionic attractions between the ions Hybridization

Chemistry Common Exam Review

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