FINAL REVIEW Chapters 1-4
Chapter 1 and 2:
1. Chemistry is:
a. The study of the composition, structure, and properties of matter
b. The study of the changes matter undergoes in a chemical reaction
c. Chemistry examines how the microscopic structure of materials affects the macroscopic
function, form, and behavior of the material in the universe
2. Physical change: characteristic of a substance that can change without the substance becoming a
different substance (odor, color, volume, state, density, melting point, boiling point)
Chemical change: substance changes into new substance with new set of properties
3. 5 evidences of a chemical change are:
a. gas given off (bubbles, fizzes, etc.)
b. precipitate forms (solid substance forms in solution)
c. color change
d. heat absorbed (endothermic) or heat given off (exothermic)
e. evolution of light or energy
4. Law of Conservation of Matter: matter cannot be created nor destroyed
5.
6. One way a mixture can be separated into its elements: filtration or distillation
7. One way a compound can be separated into its elements: chemical means
8. - Hypothesis: one or more assumptions put forth to explain observed phenomena
- Scientific Theory: an explanation of behavior; theories attempt to explain why things happen
- Scientific Method: a process used by scientists to help us solve all types of problems that we
confront in our everyday lives
9. Intensive properties: do not depend on the amount of matter that is present (boiling point, density)
Extensive properties: do depend on the amount of matter that is present (volume, mass, size)
10. Label the following as either physical or chemical changes:
a. Slicing an apple: physical
b. Boiling water: chemical
c. Two solutions mix and form a solid: chemical
d. Cooking an egg: chemical
e. Breaking glass: physical
11. Label the following as either a heterogenous or homogenous mixture, element or compound:
a. Peanut butter and jelly sandwich: heterogenous mixture
b. Water (H2O): compound
c. Copper: element
d. Completely dissolved saltwater: homogenous mixture
12. The mass of the reactants is always equal to the mass of the products.
Chapter 19.1:
1. Radioactivity: the spontaneous decomposition of a nucleus to form another nucleus; atomic
number and mass number are conserved
Alpha particles tend to be radioactive
2. Fill in the following table:
Particle Type
Description
Isotope
Notatio
n
Penetrating
ability
alpha
helium nucleus
produced in
radioactive
decay (heavy
radioactive
nuclide)
the biggest
and are least
able to
penetrate
a material
beta
radioactive
nuclide decay
that produces no
change in mass
number,
decreases atomic
number by 1
stopped by a
few
millimetres of
aluminium
but some beta
particles will
penetrate thin
aluminium
foil or paper
gamma
high-energy
photon of light
most able to
penetrate
Example equation
the same
3. Half-life: time required for half of a sample to decay from its original amount
4. A given isotope has a half-life of 5.0 minutes. If the initial mass is 280 grams, how many grams
will be left after 15 minutes? How many half-lives is this?
0 ––> 5 minutes ––> 10 minutes ––> 15 minutes
280 x (.5) ––> 140 x (.5) –––> 70 x (.5) ––> 35 g and it took 3 half-lives
5. Write a balanced nuclear decay equation for each of the following:
a. Electron capture
b. Beta decay
c. Alpha decay
d. Positron Emitter
6.
A substance has a mass of 2.50g after one half-life has occurred. What was the original mass?
2.5g / (.5) = 5 g
7. Isotopes of the same element have the same number of protons and different number of neutrons
Chapter 3:
1. Compare the parts of an atom based on location, charge, and mass:
a. Proton: a positively charged subatomic particle located in the atomic nucleus
b. Neutron: a subatomic particle in the atomic nucleus with no charge
c. Electron: a negatively charged subatomic particle
2. Define:
a. Isotope: atoms with the same number of protons, but different numbers of neutrons
b. Ion: an atom or a group of atoms that has a positive or negative charge
c. Atomic number: the number of protons in the nucleus of a given atom
d. Mass number: the total number of protons and neutrons in the nucleus of a given atom
e. Atomic mass unit (amu): a small unit of mass equal to 1.66 ‘ 10-24 grams
3. U-238 has 146 neutrons
4.
Name
Symbol
Atomic #
Mass #
# protons
# neutrons
# electrons
Sodium
Na
11
22.99 g
11
11
11
Silver
Ag
47
107 g
47
47
47
Copper II
(cation)
Cu
29
63.55 g
29
31
27
2+
Chloride
(anion)
Cl-
17
35.45 g
17
18
16
Uranium
238
(isotope)
U-238
92
238 g
92
146
92
5. The atomic mass of carbon as displayed on the periodic table is 12.011 amu. However, no single
carbon atom in nature has this mass. Explain.
No single atom of carbon has the mass 12.01 amu because carbon is a diatomic molecule so it
comes in pairs.
6. If element Z (fictitious) has two isotopes: Z-20 (20.00 amu) with 91.2% abundance, and Z-21
(21.00 amu) with 8.8% abundance. If element Z were an actual element, what mass would be
displayed on the periodic table?
0.912 x 20 amu = 18.24 and 0.088 x 21 amu = 1.84 ––> 18.24 + 1.84 = 20.08 amu
7.
8.
4 properties of metals: lustrous, ductile, malleable, conductive
4 properties of non-metals: dull, brittle, non-conductive, can be gas at room temperature
9. Aluminum touches the metalloid staircase but it is a metal
10. Where are elements with similar properties found on the periodic table (in horizontal rows, or in
vertical columns?)
11. Phosphorous:
a. period = 5A
b. group = Nitrogen group
c. atomic number = 15
d. atomic mass = 31
e. number of protons = 15
f. number of neutrons = 15
g. number of electrons = 15
h. Draw the Bohr diagram
i. Metal, or nonmetal? metal
j. What charge ion would phosphorus form? -3
12. Elements lose or gain electrons when ions are formed
13. The periodic table is organized in rows and columns by order of increasing atomic number,
which equals the number of protons in the atomic nucleus of each element.
14. Elements in the same column tend to have similar properties because they have the same number
of valence electrons.
Chapter 4:
1. a. NaBr sodium bromide
c. Fe2(SO4)3 iron sulfate
e. KBr potassium bromide
g. FeCl3 iron (III) chloride
i. aluminum sulfate Al2(SO4)3
k. ammonium carbonate (NH4)2 CO3
m. aluminum fluoride AlF3
b. FeO iron oxide
d. Mg(NO3)2 magnesium nitrate
f. CaF2 calcium fluoride
h. magnesium chloride MgCl2
j. tin (II) chloride SnCl2
l. sodium oxide Na2O
n. copper (II) fluoride CuF2
2. All of the above are a single type of compound; ionic
3. a. carbon tetrachloride CCl4
c. dichlorine heptaoxide Cl2O7
f. SO3 sulfur trioxide
b. boron trichloride BCl3
e. Cl2O dichlorine monoxide
g. P2O5 diphosphorus pentaoxide
4. All of the above are a single type of compound; covalent/molecular
5. Metal and a nonmetal will form a ionic bond, where electrons are exchanged between atoms
6.
Nonmetal and a nonmetal will form a covalent bond, where electrons are shared between atoms
7. - ionic compound metal and non-metal
- covalent compound non-metal and non-metal
- acids H atom bonded with polyatomic ion
- hydrate H2O molecule
8. 8. For the following Compounds, write the chemical formula and circle either (I) for ionic or (M)
for molecular or (A) for acid.
a) Sodium phosphide Na3P I M A
b) Aluminum oxide AlO3 I M A
c) carbon tetrachloride CCl4 I M A
d) postassium hydroxide KOH I M A
e) Ammonium chloride NH4Cl I M A
f) calcium carbonate CaCO3 I M A
g) carbonic acid H2CO3 I M A
h) phosphorous acid H3PO3 I M A
i) hydrochloric acid HCl I M A
9. a. NaClO3 sodium chlorate I M A
b. BaS barium sulfide I M A
c. Cl2O dichlorine monoxide I M A
d. SO3 sulfur trioxide I M A
e. HNO2 nitrous acid I M A
f. HF hydrofluoric acid I M A
g. SnCl2 tin (II) chloride I M A
h. Fe2(CO3)3 iron (III) carbonate I M A
i. H2SO4 sulfuric acid I M A
10. a. Dot diagram for ionic compound magnesium phosphide:
b. Ionic compound magnesium phosphide: Mg3P2
Chapter 5:
1.
= 5.10 cm
= 40.35 mL
2. How many significant digits does each of the following have?
a. 2300 m 2
b. 20040 m 4 c. 260.00 m 5
d. 0.00205 m 3
e. 4.65 x 10-4 m 3
3. Answer the following with the correct number of significant digits.
a. 4.535 m + 0.0251 m 4.560 m
b. 274 m - 254 m 20 m
c. 6.54 m / 3.4215 m 1.91 m
d. 30.67 m x 23 m 710 m
4. How many digits should be estimated in a measurement with a correct number of significant
figures? the lowest number
5. Temperature: measure of the random motions (Avg. KE) of the components of a substance
Heat: flow of energy due to a temperature difference
6.
Label the temp.
Kelvin
Celsius
Fahrenheit
H2O freezing point
273
0
32
H2O boiling point
373
100
212
absolute zero
0
-273
-459
Chapter 6:
1. - empirical formula the formula of a compound expressing the smallest whole-number ratio of
atoms in a compound
- molecular formula the actual formula of a compound, giving the types of atoms and the number
of each type of atom
- structural formula the representation of a molecule in which the relative positions of the atoms
are shown and the bonds are indicated by lines
- condensed structural formula the condensed formula of a molecule where symbols of atoms are
listed as they appear in the molecule's structure with bond dashes omitted or limited.
2. A compound is 35.0% nitrogen, 5.0% hydrogen, and 60.0% oxygen. What is the empirical
formula of the compound?
35 g N (1 mol N / 14 g N) = 2.5 / (2.5) = 1 x 2 = 2
N2H4O3
5 g H (1 mol H / 1 g H) = 5 / (2.5) = 2 x 2 = 4
60 g O (1 mol O / 16 g O) = 3.75 / (2.5) = 1.5 x 2 = 3
3. Mole: a unit of measurement in chemistry to represent:
1. (Quantity) Avogadro's Number of particles = 6.022 x 10ˆ23
2. (Mass) Equivalent to 12.0 grams of carbon-12
4. How many atoms are in a mole of calcium atoms?
1 mol Ca x 6.02x1023 = 6.02x10^23 Ca atoms
5. What is the mass of a mole of calcium atoms?
1 mol Ca = 40.08 g Ca / mol
6. What is the mass of a mole of Mg(OH)2?
Mg(OH)2 = (24.31 g Mg) + 2(16+1 g OH) = 58.31 g
7.
8. What is the percentage of silver in silver sulfide, Ag2S?
Ag2S = 2(108) + 32 = 248 g Ag2S
Ag2 = 216 g Ag2 / 248 g Ag2S = .87 x 100 = 87% Ag
9. How many atoms are in 10.0 grams of aluminum?
10 g Al /(27 g Al)x (6.02x10^23) = 2.23 x 10^23 atoms Al
10. How many grams is 3.4 x 10^24 carbon atoms?
3.4 x 10^24 / 6.02 x 10^23 x12 g C = 67.75 g C
11. How much copper can be purified from 750 grams of copper (I) sulfide?
750 g Cu2S / 159.1 g/mol = 4.7 mol Cu2S x (2 mol Cu/1 mol Cu2S) = 9.4 mol Cu x 63.5 g/mol =
596 g Cu
12. How many molecules of water are in 0.15 moles of H2O?
.15 moles H20 x 6.022x10^23 = 9.03 x 10^22 molecules H2O
13. How many moles of nitrogen dioxide are in a sample that contains 5 x 10^20 molecules NO2?
5 x 10^20 NO2 / 6.022 x 10^23 = 8.3 x 10^-4 moles NO2
14. Calculate the number of moles in 9.0 g of H2O.
H20 = 2(1) + 16= 18 g
9 g H20 / 18 g/mol = 0.5 moles H20
15. How many moles is 50 g of Ca3(PO4)2?
Ca3(PO4)2 = 3(40) + 2(31 + 4(16)) = 310 g/mol
50 g / 310 g = 0.161 moles
Chapter 7/8:
1. Reactants = substance on left of equation undergoing change
Products = substance on right of equation produced in chemical reaction
Subscript = tells the number of atoms
Coefficient = tells the number of molecules
Oxidation number = tells number of electrons gained or lost
2. When balancing equations, which can you change: the subscripts, or the coefficients? Why?
You can change the coefficients because changing the subscripts changes the molecule.
3. Balance the reaction: ____O2 + _____H2 _____H2O
4. Balance the reaction: _____Li + _____Ca3N2 _____Li3N + _____Ca
5. Describe what happens in each type of reaction. Give an example of each.
a. synthesis (combination): 2 reactants combine to form 1 product
A + B ---------> AB
b. decomposition: 1 reactant breaks down into 2 or more products
AB ------------> A + B
c. single replacement: 1 cation displaces another cation in aqueous solution or 1 anion displaces
another anion in aqueous solution
A + BC ------------> AC + B
d. precipitation (double replacement): Both cation and anion exchange in aqueous solution
AB + CD --------------> AD + CB
e. acid-base neutralization: A reaction between a strong acid and strong base to produce water and
an ionic salt
HX + MOH ----------> MX + OH
f. combustion of a hydrocarbon: compound combines with oxygen gas to produce carbon dioxide,
heat, and light energy
X + O2 -----------> CO2 + H2O + Heat + Light Energy
6. 4 driving forces in an aqueous solution:
1. Formation of a solid precipitate (ppt) or (s)↓
2. Formation of a gas in aqueous solution (g) ↑
3. Formation of water H2O(l)_
4. Transfer of electrons (Oxidation = process where element loses electrons, Reduction
= process where element gains electrons)
Chapter 9:
1. Mole ratio = whole-number ratio of moles of substances undergoing reaction found from the
coefficients from a balanced equation
2. Limiting reactant = the reactant that is completely consumed in a chemical reaction; the limiting
reactant determines the maximum amount of product
Excess reactant = the reactant that is not completely consumed and has left over amounts in a
chemical reaction; the excess reactant does not determine the product
Actual yield (experimental yield) = the amount produced in the chemical reaction and is
measured in the experiment (the given amount)
Theoretical yield = the maximum amount of product that is produced from the limiting reactant if
there is no percent error or loss of product
Percent yield = [Actual yield/Theoretical yield] x 100%
Percent error = /100% - Percent yield/
Chapter 10:
1. Describe an exothermic and endothermic reaction in terms of energy of bond making and bond
breaking:
2. Draw the energy diagram for an endothermic and exothermic reaction:
- Does the temperature increase or decrease in an endothermic reaction? decrease
- Does the enthalpy of the system increase or decrease in an endothermic reaction? increase
3. For which type of reaction (endothermic or exothermic) is the sign of the enthalpy change
negative? ( H < 0) exothermic
Chapter 11:
Light, Photon Energies, and Atomic Spectra
E = hv nano = 10-9 h = Planck’s constant = 6.63 x 10-34 Js micro = 10-6 c = speed of light =
3.00 x 108 m/s milli = 10-3
1. Light: a range of electromagnetic radiation that can be detected by the human eye.
2. In the modern model of light, light behaves as a _________ and as a ___________.
3. Calculate the wavelength of light that has a frequency of 3.20 x 1014 s -1
4. What is the frequency of electromagnetic radiation with a wavelength of 520 nm?
5. As the wavelength of light increases, the frequency increases.
Electromagnetic Spectrum
6. Electromagnetic radiation: radiant energy that exhibits wavelike behavior and travels through
space at the speed of light in a vacuum
How Light Interacts with Matter
7. Bohr’s Model: The atom contains a small positive nucleus with electrons orbiting around the
nucleus in circular orbits, similar to the planets orbiting the sun
Rutherford’s Model: Consists of a tiny, dense nucleus at the center of the atom with electrons that
occupy most of the volume of the atom
8. What is the v..difference between a continuous and a linespectrum?
9. Describe how energy is converted into light of specific wavelengths in the Bohr model of the
hydrogen atom.
10. Explain why we see colors when an electron transitions from a higher energy state to a lower
energy state. Include the terms, photon, excited state, and quantified energy.
11. Why do different elements have different line spectra?
12. How does an excited state electron configuration differ from a ground state electron
configuration?
Quantum Mechanics
13. Heisenberg uncertainty principle: It is impossible to determine both the speed and location of an
electron with certainty because the size of the electron is so small and its speed is so fast.
14. The Quantum-Mechanical Model of the Atom:
- Assumes electrons have both particle and wave properties
- Electrons occupy "orbitals", the electrons do not "orbit" the nucleus as in Bohr's model Orbitals = areas of high probability of finding an electron
- The size of an atom is described as 90% probability of finding an electron at a given point in
space
- most similar to describing the region of space as a "cloud" around the nucleus
15. How many orbitals are in the following sublevels?
3p = 3 orbitals 2s = 3 orbitals 4p = 3 orbitals
16. Arrange the following sublevels in order of decreasing energy: 2p, 4s, 3s, 3d, 3p
17. Identify the element that corresponds to the following electron configuration: 1s2 2s2 2p5
fluorine
18. Write the ground state electron configurations for
F 1s2 2s2 2p5
Mg 1s2 2s2 2p6 3s2 or Ne 3s2
Fe [Ar] 3d6 4s2
Pb [Xe] 4f14 5d10 6s2 6p2
O 2- 1s2 2s2 2p4
Ca2+ 1s2 2s2 2p6 3s2 3p6
Fe2+ [Ar]3d6
19. Draw an Aufbau diagram (orbital box diagram) for oxygen.
Periodic Trends
20. Atomic radius: measures the average distance between the nuclei of 2 atoms
-the atomic radius increases going DOWN a group on the periodic table
- the atomic radius decreases going ACROSS a row on the periodic table
21. Ionization energy: the amount energy required to remove a valence electron from the outer shell
of an atom in the gas phase
- ionization energy decreases going DOWN a group on the periodic table
- ionization energy increases going ACROSS a row on the periodic table
22. Electronegativity: the amount of attraction an atom has for a valence shell electron of another
atom; also related to the electron affinity = the amount of energy to add an electron to the valence
shell of an atom (opposite to the ionization energy)
- electronegativity decreases going DOWN a group on the periodic table
- electronegativity increases going ACROSS a row on the periodic table
23. For the following elements – a. Ba, Ca, Ra
b. P,Si,Al
- List the 3 elements in order from the largest atomic radius to the smallest atomic radius.
a. Ra, Ba, Ca
b. Al, Si, P
- List the 3 elements in order from the highest ionization energy to the lowest ionization energy.
a. Ca, Ba, Ra
b. P, Si, Al
Chapter 12:
1. - ionic bond: electrostatic attraction between closely packed, oppositely charged ions
- covalent bond: an equal sharing of electrons by the nuclei of two non-metal atoms
- non-polar covalent bond: two atoms share a pair of electrons with each other
- polar covalent bond: an unequal sharing of electrons by the nuclei of two atoms
2. Dipole: δ = separation of positive and negative charges
3. How does the electronegativity value determine the polarity of a bond?
The polarity of the bond increases as the difference in electronegativity values increases
12.2 - Characteristics of Ions and Ionic Compounds
4. Which is generally larger than their parent atom - cations or anions? Why?
Anions are larger because they gain electrons
5. Polyatomic ion: an ion consisting of two or more atoms bound together
12.3, 12.4 - Lewis Structures & VSEPR Theory
6. Octet Rule: Elements will typically fill their valence orbital with 8 electrons maximum
7.
8. VSPER: Valence Shell Electron Pair Repulsion model is useful to help predict the molecular
geometry and overall shape that the molecule will exist in 3-dimensional space
*Electrons Repel* - electrons will position themselves as far apart as possible in 3-dimensional
space to minimize the repulsion between electron pairs.
9.