Chemistry- CST Review



Chemistry- CST Review

Standard 1: Atomic and Molecular Structure

1. Complete this table.

Remember that atoms have same number of protons and electrons and mass number is close to the average atomic mass on the periodic table. Isotopes have a mass number not close to the average atomic mass, because they have different number of neutrons. Ions have different number of electrons. Ions are negative because they gain electrons.

|Symbol of |Atomic Number|Mass Number |Number of |Number of |Number of |Atom, Ion, or |

|Element | | |Protons |Neutrons |Electrons |Isotope |

|F-1 |9 |19 |9 |10 |10 | Ion |

|Si-64 |14 |64 |14 |15 |14 |Isotope |

|Mn |25 |55 |25 |30 |25 |Atom |

|Zn-64 |30 |64 |30 |34 |30 |Isotope |

|Sb-3 |51 |122 |51 |71 |54 |Ion |

|Hg |80 |201 |80 |121 |80 |Atom |

2. What is the size and mass of the nucleus in an atom?

The nucleus is very small and dense (heaviest) part of the atom.

3. How many electrons are available for bonding in the following atoms? For metals, you determine electrons for bonding by group number. For nonmetals, you determine electrons for bonding by 8 minus group number.

a) barium 2 b) sodium 1 c) aluminum 3

d) oxygen 2 e) germanium 4 f) chlorine 1

g) argon 0 h) bismuth 5

4. Name the following groups: Group 1A, Group 2A, Group 7A, and Group 8A. Alkali metals, Alkali earth metals, halogens, and noble gases

5. Where are the transition metals? Write a general statement identifying the locations of metals, nonmetals, and metalloids?

The metals are on the left side and the nonmetals are on the right side of the periodic table. The transition metals are in the middle. The metalloids are near the diagonal line.

6. Where on the periodic table would you find the elements with large atomic numbers and large atomic masses?

The elements are on the right side and bottom of the periodic table.

7. Indicate which element in each pair has the larger atomic radius. (Remember the size of the atom is larger moving to the left and down a group of the periodic table.)

a) sodium, lithium b) strontium, magnesium

c) carbon, oxygen d) selenium, bromine

e) Na, Na+ f) S, S-2

g) I, I- h) Al, Al+3

8. Indicate which element in each pair has the greater electronegativity/ionization energy, ability to attract electrons for bonding. (Remember the polar bears, high electronegativity, is on the top right and penguins, low electronegativity on the bottom left of the periodic table.)

a) lithium, boron b) magnesium, strontium

c) cesium, aluminum d) fluorine, chlorine

Standard 2- Chemical Bonds

1. Compare ionic and molecular (covalent) bonds. How are each formed? Ionic compounds are metals and nonmetals bonding because they lose or gain electrons. Molecular (covalent) compounds are nonmetals bonding by sharing electrons.

2. What type of compound (ionic or covalent) is each of the following?

a) Cl2O covalent b) SrSO4 ionic c) NH3 covalent

d) SnO2 ionic e) N2H4 covalent f) PI3 covalent

3. Draw Lewis dot structures for the following molecular compounds. .. .. ..

a) difluorine monosulfide :F:S:F:

۬ ۬ ۬ ۬ ۬ ۬

b) water [pic]

c) carbon tetrachloride [pic]

Standard 3- Conservation of Matter and Stoichiometry

1. Write a balanced chemical equation for each reaction below, and then identify the type of reaction: synthesis, decomposition, single replacement, double replacement, and combustion.

a) __2__ NaCl + ____ F2 ( __2__ NaF + ____ Cl2

Single Replacement reaction

b) ___2_ H2 + ____ O2 ( __2__ H2O

Synthesis (combination) reaction

c) ____ Pb(OH)2 + __2__ HCl ( __2__ H2O + ____ PbCl2

Double replacement reaction

d) ____ CH4 + ____ O2 ( ____ CO2 + __2__ H2O Combustion reaction

2. Define one mole? One mole equals 6.02x1023 atoms or the atomic mass of a substance.

3. What is the molar mass of NaCl, sodium chloride? 58.44g/mol

4. How many moles are in 15.45 g copper, Cu?

15.45g x 1 mole / 63.55 g = 0.2431 moles

4. How many molecules are in 2 moles of H2O, water?

2 moles (6.02x10 23) = 1.02x1024 molecules

5. Convert 85.0 L Cl2, chlorine, gas to grams at STP. (Remember 22.4L at STP).

86.0 L x 1mole x 70.9 g = 269 grams

22.4 L 1 mole

6. CaCO3 + 2HCl → CO2 + H2O + CaCl2

a) How many moles of CaCO3, calcium carbonate, would be needed to react completely with 3 moles of HCl, hydrochloric acid?

3 moles HCl x 1 mole CaCO3 = 1.5 moles CaCO3

2 moles HCl

b) How many grams of CO2, carbon dioxide, are produced when 10.0 g of CaCO3, calcium carbonate, reacts?

10.0 g x 1 mole x1 mole CO2 x 44.01 g CO2 = 4.397 g CO2

100.09g CaCO3 1 mole CaCO3 1 mole

7. 2Ca + O2 → 2CaO

a) How many moles of O2, oxygen, are needed to produce 4.50 moles of CaO, calcium oxide?

4.50 moles CaO x 1 moles O2 = 2.25 moles O2

2 moles CaO

b) How many grams of Ca, calcium, would be needed to make 14.5 g of CaO, calcium oxide?

14.5 g CaO x 1 mole CaO x 2 moles Ca x 40.08 g Ca= 10.4 g Ca

56.08 g CaO 2 moles CaO 1 mole Ca

Standard 4 – Gases and their Properties

1. What causes gas pressure in terms of kinetic theory?

Gas pressure is caused by the random motion of the gas molecules.

2. If someone sprays perfume at the front of the room, will the people in the back of the room eventually be able to smell it? Why? Explain completely. Yes, the perfume will be smelled by the people in the back of the room. The perfume gas molecules will mix with the air through random motion.

3. What values represent standard temperature and pressure (STP)? O˚C (273 K) and 1 atm

4. What is absolute zero? What happens at this temperature? Are there any temperatures below absolute zero? Absolute zero is O Kelvin and there is no temperature lower. At this temperature, gas molecules move the slowest speed.

5. Convert the following.

a) 100 °C to K 373K b) 250 K to °C -23°C

c) -35 °C to K 288 K d) 50 K to °C -223°C

e) 273 K to °C 0°C

6. How does changing the amount of gas, volume of gas, and temperature affect the gas pressure? Gas pressure increases when the amount of gas and temperature increases but the volume is decreased.

For Q’s #9-14, name the gas law and show all your work.

7. The pressure on 2.00 L of anesthetic gas changes from 100 kPa to 40 kPa. What will be the new volume if the temperature remains constant?

100kPa (2.00L) = 40kPa (V2)

V2 = 5.00L

8. If a sample of gas occupies 6.55 L at 300 °C, what will be its volume at 25 °C if the pressure does not change?

6.55 L = V2

573 298

V2 = 3.42 L

9. A gas at 790 mm Hg and 25 °C occupies a container with an initial volume of 1.20 L. By changing the volume, the pressure of the gas increases to 1500 mm Hg as the temperature is raised to 125 °C. What is the new volume?

790mm Hg (1.20L) = 1500 mmHg V2

298 398

V2 = 0.844L

10. A 500 mL air sample at a temperature of -50 °C has a pressure of 1.3 atm. What will be the new pressure if the temperature

is raised to 102 °C and the volume expands to 700 mL?

1.3 atm (500 mL) = P2 (700mL)

223 375

P2 = 1.6 atm

Standard 5- Acids and Bases

1. Classify the following properties as those belonging to an acid or base or both.

a) bitter taste base

b) sour taste acid

c) H+ ion donating acid

d) OH- ion donating base

e) pH greater than 7 base

f) pH less than 7 acid

g) H+ ion accepting base

h) strong electrolyte (hint: an electrolyte dissolves into ions in water and therefore conducts electricity) strong acid and strong base

i) weak electrolyte weak acid and weak base

j) feels slippery base

2. What is the difference between a strong acid or base and a weak acid or base? A strong acid has a low pH and a strong base has a high pH. Strong acids and bases fully dissociate and strong electrolyte (conduct electricity).

3. Which substances are hydrogen ion donating, hydrogen ion accepting, or neither.

a) HCl Hydrogen ion donating b) CO2 neither

c) KOH hydrogen ion accepting d) H2O hydrogen ion accepting and donating

Standard 6- Solutions

1. Define solute and solvent. Salt is dissolved in a glass of water. Which is the solute? Which is the solvent?

Solute is the substance being dissolved and it is present in lesser amount. The solvent is usually a liquid and present in the greater amount. Salt is a solute and water is a solvent.

2. Explain what you would do to quickly dissolve cube sugar in a cup of coffee (Like changes in temperature and surface area, breaking up the cube sugar). In order to dissolve a cube of sugar, you would increase surface area by breaking it up and increase temperature by heating the mixture.

3. What effect would increasing concentration (adding more solute) have on the dissolving process? Lowering concentration? Explain. Increasing concentration of solute will slow the dissolving process, because there would not be enough solvent molecules to disperse the solute molecules. Lowering the concentration of solute would speed up the dissolving process because less solute molecules will have to disperse into the solvent.

4. Calculate the molarity of each of the following solutions:

a) 0.60 mol of NaCl dissolved in 1.6 L of solution.

0.60mol / 1.6 L =0.375M

b) 25.2 g of potassium nitrate, KNO3, in enough water to make 150.0 mL of solution.

25.2g KNO3x 1 mole/ 101.11g = 0.2492 moles

0.2492 moles / 0.15 L = 1.66 M

5. Calculate the number of grams of solute needed to prepare each of the following solutions:

a) 4500.0 mL of a 2.0M solution of potassium hydroxide, KOH.

4.5L x 2.0M = 9 moles KOH

9 moles KOH x 56.11g/ 1 mole = 505 g KOH

b) 2.0 liters of 3.0M nitric acid, HNO3, solution.

2.0L x 3.0 M = 6.0 moles HNO3

6.0 moles HNO3 x 63.02g/ 1 mole = 378 g HNO3

Standard 7- Chemical Thermodynamics

1. Compare exothermic and endothermic processes. Exothermic process releases heat and an endothermic process absorbs heat.

2. Is heat released or absorbed during these processes?

a) melting absorbed b) freezing released

c) boiling absorbed d) condensing released

3. How much heat is required to raise the temperature of 20.55 g of mercury 50°C? The specific heat of mercury is 0.14 J/g °C.

Q = m(ΔT)C

Q = 20.55 g (50˚C) (0.14 J/g˚ C)

Q = 143.85 J

4. Calculate how much heat 35.0 g of water absorbs when it is heated from 20°C to 80°C. (Specific Heat of water = 4.180 J/gºC)

Q=35.0g (60˚C)(4.180 J/g˚C)

Q = 8778 J

5. What is the specific heat of a 15.0 g substance that absorbs 350 J of heat when the temperature is raised 40 °C?

C = Q/m (ΔT)

C= 350J/15.0g (40˚C)

C= 0.583 J/g˚C

Standard 8- Reaction Rates

1. What is the rate of a reaction? The rate of reaction is the increase in concentration (amount) of products over time.

2. What factors increase the rate (speed) of a reaction? (Like concentration, pressure, and temperature) The increase of concentration of reactants, pressure, and temperature are factors that increase reaction rate.

3. What does a catalyst do to the reaction rate? A catalyst decreases the activation energy of reaction therefore increases the reaction rate.

Standard 9- Equilibrium

1. What is dynamic equilibrium? Dynamic equilibrium is when a reversible reaction in balance with the forward and reverse reactions occurring at the same rate (speed).

2. Given the following system at dynamic equilibrium:

2SO2 (g) + O2 (g) ↔ 2SO3 (g) + heat

Determine the effect of each of the following changes on the equilibrium position (shifts left or right) and on the amount of O2 that would result (increase or decrease).

| |Reaction Shift (left or |Amount of O2 (increase or |

| |right) |decrease) |

|Increasing temperature |Left |Increase |

|Decreasing pressure |Left |Increase |

|Adding SO2 |Right |Decrease |

|Removing SO3 |Right |Decrease |

|Increasing pressure |Right |Decrease |

|Adding SO3 |Left |Increase |

|Decreasing temperature |Right |Decrease |

|Removing SO2 |Left |Increase |

Standard 10- Organic Chemistry (Smells Unit)

1. What following compounds contain simple repeating units?

a) nucleic acid b) proteins

c) lipid d) starch

e) water f) salt

2. Amino acids are building blocks for ___proteins___.

3. How many bonds does the carbon atom form? Carbon can form single, double, and triple bonds.

4. What atoms does carbon commonly form bonds with? Hydrogen, nitrogen, oxygen, and another carbon commonly form bonds with carbon.

Standard 11- Nuclear Processes

1. What elements have radioactive isotopes? Elements with atomic number 84 and above are radioisotopes. There are more like carbon which has a radioisotope of carbon-14.

2. What is the difference between a chemical and nuclear reaction? In a nuclear reaction matter can be destroyed and created but not in a chemical reaction. Nuclear reactions created great amounts of energy too.

3. What is nuclear fission and nuclear fusion? Nuclear fusion is when two nuclei are fused together. Nuclear fission is when a radioactive element breaks up.

4. Explain the three different types of nuclear decay: alpha decay, beta decay, and gamma.

Alpha decay is when a helium atom (42He) is released.

Beta decay is when an electron is released (-10e), because a neutron split into a proton and an electron.

Gamma decay released no matter but energy.

5. How danger is the radiation from alpha, beta, and gamma particles?

An alpha particle can be stopped by paper.

A beta particle can be stopped by wood.

A gamma ray can only be stopped by lead. It is most dangerous.

................
................

In order to avoid copyright disputes, this page is only a partial summary.

Google Online Preview   Download