Download 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
Element
Atomic
Number
Mass
Number
Number
of
Protons
Number
of
Neutrons
Number
of
Electrons
Atom,
Ion, or
Isotope
F-1
Si-64
Mn
Zn-64
Sb-3
Hg
9
14
25
30
51
80
19
64
55
64
122
201
9
14
25
30
51
80
10
15
30
34
71
121
10
14
25
30
54
80
Ion
Isotope
Atom
Isotope
Ion
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, Ih) 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
c) cesium, aluminum
b) magnesium, strontium
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. What is the molar mass of NaCl, sodium chloride? 58.44g/mol
3. Draw Lewis dot structures for the following molecular
compounds.
.. .. ..
a) difluorine monosulfide
:F:S:F:
۬۬ ۬۬ ۬۬
4. How many molecules are in 2 moles of H2O, water?
2 moles (6.02x10 23) = 1.02x1024 molecules
b) water
4. How many moles are in 15.45 g copper, Cu?
15.45g x 1 mole / 63.55 g = 0.2431 moles
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.
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
c) carbon tetrachloride
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.
CaCO3 + 2HCl → CO2 + H2O + CaCl2
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 right)
Increasing
temperature
Decreasing
pressure
Adding SO2
Removing SO3
Increasing
pressure
Adding SO3
Decreasing
temperature
Removing SO2
Left
Amount of O2
(increase or
decrease)
Increase
Left
Increase
Right
Right
Right
Decrease
Decrease
Decrease
Left
Right
Increase
Decrease
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.