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CHEMISTRY LEVEL 2 MIDTERM REVIEW 2006

The following review is intended solely to present possible information that may appear on the midterm. Some material similar to the material included on this midterm review may appear on the midterm, some may not. Furthermore, some information not included on this review may appear on the midterm.

Chapter 1:

    • Major Topics:

      • What is chemistry and why is it considered a central science?

      • The Scientific Method- the process of scientific exploration

      • Parts of an experimental lab report.

      • Basic understanding of safety in the chemistry laboratory.

      • The Metric system-

        • Basic units in the metric system

        • Uncertainty in measurements (why and how is it written)

        • Significant figures in measurements

        • Accuracy vs. Precision

        • Scientific notation vs. extended notation

        • Percent error

        • Density and density calculations

        • Unit analysis (dimensional analysis/ factor label method)

    • Sample Calculations:

      • What is the density of an object with a mass of 4.3 kg that when dropped in a beaker of water changed the volume from 19.4 mL to 56.7 mL.

      • In an experiment, the scientists had hoped to have a yield of 24.3 grams from a single replacement experiment involving lead (II) and mercury (II) chloride. The actual yield from the experiment was only 20.0 g. What is the percent error in the experiment.

      • The density of copper is 8.96 g/cm3. If a rectangular sheet of copper is 10.3 cm wide, 46.1 cm long, and .014 cm thick, what is the mass of the copper?

      • What is the volume of an object that had a thickness of 13 cm and an area of 45 cm2

      • Convert the following using dimensional analysis.

        • 54 mL to L

        • 75 mm to dm

        • 1 mile to cm

      • Uncertainty and significant digits of the following numbers

        • .90 cm

        • 0.00120 L

        • 1000 g

      • Perform the following problems

        • 4.5 g / 2.344 L

        • 2 g + 842.2 g

      • Convert from scientific to ordinary notation or vice versa

        • 4000 g

        • 2.34 x 10-3

        • .003290

        • 3.2 x 105

Chapter 2

    • Major Topics

      • The major forms of energy- what are they and an example of each

      • The law of conservation of energy – who and when

      • The units of energy (joule vs calorie)

      • The temperature systems (Celsius, Fahrenheit, Kelvin)

        • Conversions between the systems

        • Development of the systems – who, why, and when

      • Matter

        • States of matter (5 states of matter)

        • Physical vs. Chemical Properties

        • Physical vs. chemical changes

        • The law of conservation of matter (who, what, and when)

      • Differences between Element and Compound

        • Definition of atom, molecule, and formula unit

        • Definition of ionic and covalent bonds

        • Definition of homogeneous and heterogeneous mixtures

        • Separation of mixture techniques

      • Specific Heat

        • Calculations involving specific heat

        • Know the unit of specific heat

      • Density as a unit factor

    • Sample Calculations

      • Which of the following is not a pure substance

        • Milk, hydrogen, water, oxygen

      • The SI unit of temperature is?

      • Convert the following between Fahrenheit and Kelvin

        • 45 F

        • - 23 F

      • Convert the following between Kelvin and Fahrenheit

        • 345 K

        • 832 K

      • The density of osmium is 22.5 g/cm3, what is the volume of a 34 g sample of osmium?

      • Given 25 g of Lead (II) Phosphate, how much Oxygen is needed to make that quantity of Lead (II) Phosphate

      • What is the percent composition of Iron (IV) Chloride

      • How many atoms are in Iron (IV) Chloride

      • How many calories of heat are released from the cooing of 125 g of water from 100 C to 15.5 C?

Chapter 3

    • Major Topics:

      • Early models of the atom (who, what, and when)

        • Daltons model (postulates)

        • Democritus

        • Law of Definite Composition and Law of Conservation of Matter

      • Atomic Structure

        • Relationship to electricity

        • Cathode ray experiments (who, what, and when)

        • Alpha scattering experiments (who, what, and when)

        • Significance of radioactivity

        • Oil drop experiment (who, what, and when)

      • Modern atomic theory

        • Plum Pudding Model (who, what, and when)

        • The three subatomic particles (what are they and where found)

        • Determine the number of protons, neutrons, and electrons in an atom

        • Differentiate between mass number, atomic number, and atomic mass

        • Define isotope and ion

      • Changes in the nucleus

        • Nuclear reactions

        • Define radioactivity

        • Nuclear stability and number of neutrons

        • Strong nuclear force vs. relative nuclear force

    • Sample Calculations:

      • Write the nuclear equation for the beta decay of titanium-50

      • What is the atom with 84 protons

      • What ion is formed by Barium

      • What is the number of protons, electrons, and neutrons in the following

        • 12C

        • 37Cl1-

      • Calculate the average atomic mass for iron given the following data for it’s isotopes.

 

53Fe 53.940 amu 5.82%

54Fe 55.935 amu 91.66%

55Fe 56.935 amu 2.19%

58Fe 57.933 amu .33%

      • Sketch a model of Boron

Chapter 5:

    • Sample Calculations

      • How many elements are in the 2nd period?

      • Where are the non-metals located?

      • Why do elements in group 1A from 1+ ions?

      • Why does atomic radius decrease as you move from left to right across a period?

      • Which element has a larger atomic radius- boron or magnesium

      • Which is bigger the Chlorine atom or the Chlorine ion

      • Why is a molecular radius smaller than the combined values of the individual atoms involved in the bond?

Chapter 7/ 8:

    • Major Topics:

      • Ionic Bonding

        • Describe the characteristics of an ionic bond

        • The octet rule

        • The duet rule

        • Anions vs cations

      • Covalent Bonding

        • Describe the characteristics of a covalent bond

        • Polar vs non-polar covalent bonds

      • Naming compounds

        • Naming ionic compounds (rules)

          • Binary ionic

          • Ternary ionic

        • Naming covalent compounds (rules)

        • Naming aqueous acids

          • Binary acids

          • Ternary acids

        • The polyatomic ions

      • Empirical vs Molecular Formulas for Ionic and Covalent Compounds

Chapter 9/ 11

    • Major Topics:

      • The Nature of Chemical Reactions

        • The characteristics of a chemical reaction

        • Reactants vs products

        • The importance of catalysts

        • The law of conservation of matter

      • Chemical equations

        • Balancing a chemical equation

        • Coefficients

        • Word equations vs formula equations

      • Classification of reaction types

        • Single replacement, double replacement, combustion, neutralization, synthesis (direct combination) and decomposition.

        • Solubility rules for double replacement reactions

        • Reactvity series for single replacement reactions

        • Special cases of synthesis and prediction of products

          • Metal and oxygen

          • Non-metal and oxygen

          • Metal and non-metal

        • Special cases of decomposition

          • Metal hydrogen carbonate decomposition

          • Metal carbonate decomposisiont

        • The combustion reaction

          • Hydrocarbon and oxygen reacting to produce carbon dioxide and water

        • The neutralization reaction

          • Acid and base reacting to produce a salt and water

        • Special cases of Single replacement reactions

          • Active metal and water

          • Metal and an aqueous acid

          • Metal and aqueous solution

      • Stoichiometry

        • Definition and significance

        • Stoichiometry and the balanced chemical reaction

      • Solving stoichiometry problems

        • Mass mass problems

        • Mass volume problems

        • Volume volume problems

      • Limiting reactants and percent yield

        • Define and determine limiting reactant

        • Calculation of product formed from a given reaction based on limiting reactants

        • Calculations of percent yield based on actual and expected yields

    • Sample Calculations:

      • When octane (C8H18) is burned in oxygen, carbon dioxide and water are produced. If 320 g of octane are burned and 392 g of water are recovered, what is the percent yield of the experiment.

      • In the reaction between copper (II) sulfate and aluminum, if 2 moles of aluminum enter the reaction with unlimited copper (II) sulfate, how much copper can be produced.

      • What mass of barium chloride is needed to react completely with 46.8 g of sodium phosphate.

      • To produce 196.5 g of aluminum sulfate, what mass of aluminum must react with sulfuric acid?

      • What volume of oxygen at STP is necessary to burn 500. g f glucose?

      • A student places an iron nail with a mass of 2.32 g into a flask of copper (II) sulfate. The nail reacts completely, leaving a quantity of copper metal in the bottom of the flask. The student finds the mass of the recovered copper to be 2.51 g. What is the expected and percent yield?

      • 549.6 g of potassium carbonate react with 429.2 g of Magnesium nitrate, how many grams of potassium nitrate can be produced?

Chapter 10:

    • Major Topics:

      • Chemical Measurements

        • The mole and its importance

        • Avagadro’s number and its use

        • Molar mass and atomic mass and formula mass

      • Mole conversions

        • Convert between number of particles, number of moles, and mass of a substance

        • Molar volume, STP, and its use

      • Empirical and Molecular Formulae

        • Percent Composition

        • Percent Composition to find a formula from an experiment

        • Calculations of empirical and molecular formula

    • Sample Calculations

      • Decomposition of an unknown compound gave the following data, what is the empirical formula: 9.00 g carbon, 16.8 L hydrogen, and 2.80 L oxygen. If the molar mass of the compound is 116 g/ mol. What is the molecular formuls?

      • What is the percent composition of sucrose (C12H22O11)?

      • Phenylalanine is an essential amino acid whose chemical composition is 65.5% carbon, 6.67% hydrogen, 8.48% nitrogen and 19.4 % oxygen. What is the empirical formula for pheynalaline?

Chapter 12

    • Major Topics

      • Chemical Reactions that Involve Heat

        • What is the difference between heat and energy

        • What is the difference between surroundings and system

        • What is the difference between exothermic and endothermic

        • How is heat transferred in terms of kinetic energy

        • Where is kinetic energy stored in chemicals

        • When is heat written as a reactant and when is it written as a product

      • Heat Changes and Enthalpy

        • What is enthalpy

        • How is it represented

        • What is the significance of changes in enthalpy

        • What is delta H, how is it calculated

        • What does the sign of delta H tell you about the type of reaction

        • Which has more energy in which type of reaction (products vs reactants)

      • Hess’s law

        • What is hess’s law

        • How and why is it used

      • Calorimetry

        • What is calorimetry

        • What is the difference between calorie, specific heat, and heat capacity

        • If the heat flow is known for the calorimetry solution, how can it be figured out for the substance in the calorimeter.

        • Foods as fuels- which ones supply/ store the most energy

      • What is Heat

        • Know the scientists critical for the development of the kinetic theory of heat

        • What was the caloric theory- why was it disproved

    • Sample Calculations

      • The specific heat of silver is .24 J/ g°C. How much heat must be added to a silver block of mass 86 grams to raise its temperature 9.0 C

      • Sunita performed a calorimeter experiment to determine the enthalpy change for the dissolving of Potassium permanganate. Her calorimeter contained 50.0 g of water at an initial temperature of 17.4°C. After 1.10g of potassium permanganate, the temperature fell to 16.0°C. What value of enthalpy can Sunita obtain from her data