Download Unit 1: Chapter 3

MHS Chemistry
Semester Exam Review Packet 2013
Unit 1: Chapter 3
Sig Figs, Scientific Notation, SI Units, & Measurement
1. Significant Figures: know the rules of significant figures for:
a. Addition/subtraction: the answer can be no more significant than the weakest number
b. Multiplication/Division: the answer should have the same number of significant digits as the
number with the smallest number of sig figs.
2. Precision vs. Accuracy: Repeated measurements that are in
agreement are precise. A measurement that matches the
actual number is accurate.
Neither
Precise
Precise & Accurate
3. Uncertainty in Measurement: When reading an instrument, you are only able to estimate ONE DIGIT
beyond the smallest increment.
For example: On the ruler below, the measurement would be 2.90 cm, where the 0 is estimated as one
digit beyond the smallest increment.
4. Percent Error: Percent Error = ((Accepted – Experimental)/Accepted) x 100%
Note: Percent Error does not have a sign, since it is based on the absolute
value of the difference.
Metric Prefixes Symbol
Tera
1012
T
5. SI Unit Conversions:
Giga
109
G
a. Know how to convert the SI Units for length, mass, & volume
Mega
106
M
based on the following prefixes:
Kilo
103
k
b. Know how to convert Celcius and Kelvin temperatures. oC = K + 273
Hecto
102
h
Deca
101
da
Deci
10-1
d
6. Density: Know how to calculate density, mass, and volume given the
Centi
10-2
c
-3
density formula: D=m/V
Milli
10
m
Micro
10-6
µ
Nano
10-9
n
Unit 1 Sample Questions:
-12
Pico
10
p
1. Calculate the volume of a mercury sample with a density of 11.0 g/mL and has a mass of 440g.
2. Know your metric units and how to convert from one to another: For example, how many
millimeters in 13.4 meters?
a. _________ kg = 546 g
b. ____________L = 345 mL
c. _________cm = 4.21 m
d. ____________mg = 0.0789 kg
3. How many significant figures in the following numbers?
a. 0.00450
c. 250
b. 2.034
d. 4.5678 X 10-9
4. Calculate the following:
a. 8.7 g + 15.43 g + 19 g =
b. 4.32 cm x 1.7 cm =
c. 8.53.2 L - 627.443 L =
d. 38.742 g ÷ 0.421 mL=
_________________
_________________
_________________
_________________
5. During September the average temperature was 21.8 degrees Celsius. What is the average
temperature in Kelvin?
The ceiling in Jan’s living room is 2.5 m high. She has a hanging lamp that
hangs down 41 cm. Her husband is exactly 2 m tall. Will he hit his head on the
hanging lamp? Why or why not?
Unit 2: Atomic Structure and Electrons Chapters 4, 5, & 6
Part One: Atomic Structure
1. Describe the various models in the historical development of modern atomic theory:
a. Democritus: The first to say that matter is composed of atom, or “atomos.”
b. Dalton: Had five basic principles in his model of the atom
c. Thomson: discovered the charge of the electron by deflecting the flow of electrons through
his Cathode Ray Tube with magnetic and electrical fields, and theorized that the atom was a
“plum pudding” of electrons and positive charges.
d. Rutherford: Used the Gold Foil experiment to prove the Plum Pudding Model. He shot Alpha
particles (positively charged He nuclei) at Gold foil only to discover that the atom has a
nucleus, filled with protons, and other significant mass. Since most of the alpha particles
passed through, he concluded that the atom is mainly empty space. The small number of
particles that bounced back proved that the nucleus had mass, and positively charged
protons.
e. Bohr: Electrons move about the nucleus in “orbits” in Quantum Energy levels. The maximum
number of electrons per energy level is 2n2, where n = the energy level number.
f. Quantum Mechanical Model: electrons spin around the nucleus in clouds of probabilities.
2. Distinguish among protons, neutrons, and electrons in terms of their relative masses, charges, and
location with respect to the nucleus.
3. Infer the number of protons, neutrons, and electrons using the atomic number and mass number of
an element from the periodic table and symbol notation i.e.
F has 9 protons, 9 electrons (both
determined by the Atomic Number), and 10 neutrons (Mass Number – Atomic Number)
4. Explain how isotopes of an element differ: Isotopes are atoms of the same element with the same
atomic number but different Mass Numbers. Thus, isotopes have the same # of protons, but different
numbers of neutrons.
5. Explain why the atomic masses of elements are not whole numbers: Atomic Masses are “Weighted
Averages” of the naturally occurring isotopes.
6. BE ABLE to calculate the average atomic mass of an element from isotope data. Predict which of the
naturally occurring isotopes is most abundant from average atomic mass.
7. Differentiate between an atom and an ion. MAKE SURE that you understand that a “+” ion has LOST
electrons, while a “-“ ion has gained electrons.
8. Determine ion charge given proton, neutron, and electron data.
1. Give the number of protons, neutrons, and electrons for the nuclide Na-23.
2. How many electrons are in the S2- ion?
3. Calculate the number of neutrons for each of the isotopes of hydrogen, H-1, H-2, and H-3.
Part Two: Electrons and Electron Configurations
1. Explain the atomic emission spectrum of an element using Bohr’s model of the atom. Electrons go
from the ground state to the excited state when energy is absorbed. Then the electron moves from
the excited state back to the ground state, light is emitted!
2. Calculate the Energy, frequency and wavelength of electromagnetic radiation (light) using the
equations: E = hν, and c = λν. Where h = Planck’s constant (6.63x10-34 J-s) & c = 3 x 108 m/s.
3. What are the parts of the Wave? Predict the relationships between wavelength, frequency, and
energy for a wave.
4. Describe the properties of the different types of electromagnetic radiation.
5. Describe the Quantum mechanical Model of the atom.
6. Apply the Aufbau principle (filling of electron levels from low to high energy), the Pauli Exclusion
Principle (two electrons occupying the same orbital must have opposite spins), and Hund’s rule (1
electron must be given to each of equivalent orbitals; ie. 3 p and 5 d orbitals) to write electron
configurations and orbital diagrams (1s2, 2s2, 2p6, 3s2, 3p6, …) of elements. Do not forget the
exceptions to the Aufbau principle, specifically the d4 and d9 exceptions.
7. Identify the electron configuration for an Excited Atom, where an electron has moved from the
Ground State to a higher energy level. For example, the electron configuration for the ground state of
20
Ca is 1s22s22p63s23p64s2. An possible electron configuration for an excited 20Ca atom may be
1s22s22p63s23p54s25s1 .
8. Complete “Orbital Filling Diagrams” for an element’s electron configuration using the Aufbua, Pauli,
and Hund’s Rules.
9. Don’t forget the Quick equations:
a. Maximum number of electrons within an (n) energy level = 2n2
b. Maximum number of orbitals (s, p, d, f) within an (n) energy level = n2
Practice Questions:
Calculate the atomic mass of oxygen given the following data.
Isotope
Oxygen-16
Oxygen-17
Oxygen-18
Mass
16.00 amu
17.03 amu
18.34 amu
% Abundance
99.76%
0.037%
0.204%
Electron Configurations:
1) Write out the complete electron configuration for:
a.
b.
c.
2)
chlorine
manganese
barium
Write out the short-hand (noble gas) electron configuration for:
a.
b.
c.
bromine
cobalt
radium
3)
Identify the element corresponding to the following configurations:
a.
b.
c.
4)
Name the element(s) that have:
a.
b.
c.
Symbolic
Notation
59
1s22s22p4
1s22s22p63s23p64s1
1s22s22p63s23p64s23d7
Electrons in a 4s orbital but not in 3d
Electrons in a partially-filled 4d level
Only unpaired electrons in a 6p orbital
Element
Name
Atomic
Mass
Atomic
Number
#
Protons
#Electrons #Neutrons
Co
27
19
F 1-
9
23
Na 1+
11
Neon
15
18
5) Calculate the wavelength of radiation with a frequency of 1.50 x 1013 Hz
6) What is the frequency of radiation with a wavelength of 5.00 x 10-8m?
7) What type of electromagnetic radiation has the highest energy? The lowest?
Part Three: Periodic Table
1. The Periodic Table: What is the Periodic Law? The Periodic Law states that when the elements are
arranged in order of increasing atomic number, there is a periodic repetition of their chemical and
physical properties.
a. The horizontal rows are called the periods. There are seven periods. Going across a period
from left to right, elements are filling that energy level’s “s & p” orbitals, eventually getting to
a full octet at the noble gas.
b. The vertical columns are called groups or families. Elements within the same group have the
same number of valence electrons, and thus have similar properties.
2. Where are the metalloids? Where are the Transition Metals? The Lanthanides? The Actinides? The
Halogens? The Nobel Gases? The Alkali Metals? The Alkaline Earth Metals?
3. Describe the origins of the modern periodic table. Describe the organization of the periodic table
(periods, groups, periodic law, s block, p block, d block, & f block), and categorize the elements as
halogens, alkali metals, alkaline earth metals, noble gases, transition metals, inner transition metals,
and representative elements (all of the Group A elements). As a refresher, the Periodic law states that
when elements are arranged by their Atomic Number, there is periodic repetition of properties within
a group or family.
4. Contrast the physical and chemical properties of metals (LOSE ELECTRONS), nonmetals (GAIN
ELECTRONS), and metalloids (DO BOTH), and locate them on the periodic table
5. Explain the relationship between the electron configuration of an element, its position on the periodic
table, and its chemical properties. Simply put, be able to determine the electron
configuration using the Periodic Table, and predict the charge of the ion based on the Group A
number: i.e. You would expect a Group 2A metal, which has 2 valence electrons, to lose both of those
electrons and become a 2+ ion.
6. State the trends of properties of elements within periods and groups of the Periodic Table. Interpret
the trend shown by atomic radii, ionic radius, electronegativities, and 1st, 2nd, & 3rd ionization
energies within the period.
Decreasing ionization energy and
electronegativity
Increasing ionization energy and electronegativity
i
Decreasing atomic & ionic radius
7. Determine the number of valence electrons &/or predict the stable ion formed by a representative
element, using the periodic table. Remember, the representative elements are the GROUP A
elements.
8. Predict Element location based on period and valence electrons. For example, a valence configuration
of 4s24p5 tells me that the element is in the 4th period, in Group 7A. The element is therefore Br
(Atomic No. 35).
Practice Questions:
1. Write the group names of the following:
a. IA
b. IIA
c. VIIA
d. VIIIA
2. Write the short hand electron configuration for each of the following and give the number of
valence electrons
a. Cl
b. Kr
c. Hg
3. Put the following in order of increasing ionization energy: P, S, As, Se
4. Put the following in order of decreasing radius: F, Cl, Br, I
5. Two characteristics of metals are:
6. Two characteristics of nonmetals are:
Unit 3: Bonding - Chapters 7 & 8
Part One: Ionic Bonding
1. Predict the formation of cations & anions from a representative atom, based on the location on
the Periodic Table; more specifically, based on the atom’s Group and the atom’s desire to obtain
an OCTET of electrons (or a noble-gas electron configuration). Be able to draw a Lewis Dot
Structure for the atom and ion.
 Group 1A: loses s1 to become a 1+ cation
 Group 2A: loses s2 to become a 2+ cation
 Group 3A: loses s2p1 to become a 3+ cation
 Group 5A: gains 3 electrons to become s2p6, and a 3- anion
 Group 6A: gains 2 electrons to become s2p6, and a 2- anion
 Group 7A: gains 1 electron to become s2p6, and a 1- anion
2. Predict the formula unit of an ionic compound based on the expected charge of the ions, and
write the Lewis Dot Structure for the Ionic compound.
3. What are the characteristics/properties of an ionic bond? Ionic compounds occur when very
reactive metals transfer 1 or more electrons to reactive nonmetals. The resulting positive metal
ion and nonmetal negative ions are highly attracted, and strongly bonded together.
4. Ionic compounds exist as solids in a crystalline structure at room temperature, are dense, and
have high & specific melting points. Ionic compounds are conductive only when melted or
dissolved in water.
Part Two: Covalent Bonding
1. Describe the formation of a covalent bond between two nonmetallic elements. These bonds
care either nonpolar covalent (for diatomic molecules; i.e. H2, F2, Cl2), or polar covalent (all
others that are not ionic, or nonpolar).
2. Identify bonds as ionic, polar covalent, non- polar covalent using electronegativity values If
asked to judge the polarity of a covalent bond based on the difference in electronegativity, use
the following chart as a guide:
3. Describe double and triple covalent bonds and draw Lewis structures to represent covalent
bond structures containing single, double, triple bonds, including exceptions to the octet rule.
 Hydrogen (1 bond, 2 valence electrons only)
 Boron (3 bonds, 6 valence electrons only)
 Sulfur (can exceed the octet rule, and can have 6 bonding atoms, or 12 valence
electrons)
 Phosphorus (can exceed the octet rule, and can have 5 bonding atoms, or 10 valence
electrons)
4. Explain the modern interpretation of resonance bonding. Be able to draw/predict resonance
structures from the molecular formula.
5. Be able to determine the shape of a molecule using the VESPR theory. See the chart below:
Electron
Domains
2
In molecules where
the outside
molecules are
different, shapes
that tend to be
nonpolar usually
become polar.
Linear Triatomic, Usually nonpolar
CO2, HCN
3
Trigonal Planar: BF3, SO32-, NO3120˚
Bent, 12O˚ Usually polar
Remember to count
the number of
“clouds” of electrons,
not the actual
number of electrons.
A double or triple
bond counts as one
effective pair.
Also: If there ever
is a two molecule
atom (diatomic)
that molecule’s
polarity depends
upon the
electronegativity
difference of the
atoms
NO24
Tetrahedral; 109˚: Usually
nonpolar CH4
CF4
Trigonal Pyrimidal: 107˚
Usually polar: NH3, PCl3
Bent: 104.5˚
Usually polar: H2O, OF2
1. Draw the Lewis dot structures for the following molecular compounds:
a. CCl4
b. SF6
c. O2
d. CNe. SO3
2. Decide if the bonds are polar or nonpolar:
a. H-Br
b. N-N
c. P-Cl