Review Sheet for Chemistry First Semester Final
Refer to your class notes, worksheets, and the textbook to complete this review sheet. Study
early so that you will have time to ask questions about what you don’t understand. For the exam,
you will need to bring 2 number 2 pencils, a scientific calculator (there will be NO sharing and
NO cell phones for calculations) and you may bring a 4x6 card with formulas, equations etc that
you feel you need. This will be collected with your test and must have your name on it. (MUST
be a 4x6 CARD, no paper cut to 4x6- MUST be HAND written)
Things you need to know: In addition to the sections below:
SI Base Units- Pg 73
Metric Prefixes- Page 74
Know how to convert from one unit to another (ie: kilometers to millimeters)
KNOW YOUR ELEMENTS – The ones you were tested on! (name, symbol and spelling
counts!)
KNOW the Polyatomic Ions on Pg 257 (charge, name, formula)
KNOW the Common Metal Ions with more than one charge (table on page 255)
KNOW the Prefixes for Covalent molecule names (page 269)
BE ABLE TO USE the Significant Figures rules
Ch. 1: Introduction to Chemistry
manipulated (independent) variable
responding (dependent) variable
observation
theory
steps in scientific method
problem solving in chemistry
Ch. 2: Matter and Change
chemical change
chemical property
chemical reaction
compound
element
filtration
gas
heterogeneous mixture
homogeneous mixture
conservation of mass
liquid
mass
mixture
phase
physical change
physical property
product
reactant
Properties of Matter
Mixtures
Elements and Compounds
Chemical Reactions
Matter: Anything that takes up space and has mass
Physical Changes and Chemical Changes
Define each. How can you tell the difference between the two?
solid
solution
substance
vapor
volume
Classify the following as physical or chemical changes:
a. spoiling of milk ___________________
b. bending wire _____________________
c. cutting paper _____________________
d. rusting of a nail ___________________
Identify the following as pure substances, homogeneous mixtures or heterogeneous mixtures:
a. copper ______________________
b. sweetened tea ________________
c. sand and water _______________
d. calcium carbonate (CaCO3) ________________________
Sketch particles in the three states of matter. How close are the particles and how much do
they move?
Solid
Liquid
Gas
Calculations using the Law of Conservation of Mass for Reactions
4g H2 + ?? g O2 → 36g H2O
Ch. 3: Scientific Measurement
absolute zero
accepted value
Celsius scale
density
energy
error
experimental value
gram
International System
of Units (SI)
joule (J)
Kelvin scale (K)
kilogram (kg)
liter (L)
meter (m)
percent error
scientific notation
significant figures and rules
weight vs. mass
Be able to calculate percent error, density, and convert between Kelvin and Celsius
Be able to convert between scientific notation and decimal or integer numbers
Identify the number of sig figs; addition, subtraction, multiplication and division with
sig figs.
Ch. 4: Atomic Structure
atom
atomic mass
atomic mass unit
electron
group in Periodic Table
isotopes
mass number
neutron
proton
period in Periodic Table
Be able to find the atomic mass and atomic number on the Periodic Table.
Be able to calculate the number of neutrons.
Know how to identify the number of electrons.
Know how to identify isotopes and write their formulas.
Atom
For this Carbon–14 isotope, 146 C
• Atomic number = _____, Mass number = _____,
• # of protons = _____, # of electrons = _____, # of neutrons = _____.
Atomic Masses: What is the difference between the mass number for Carbon–14 and
carbon’s atomic mass of 12.011 amu?
Calculate the atomic mass of lithium is one isotope has a mass of 6.0151 amu and a percent
abundance of 7.59% and a second isotope has a mass of 7.0600 amu and a percent
abundance of 92.41%.
Ch. 5: Models of the Atom
atomic orbitals
aufbau principle
electromagnetic radiation
electron configurations
energy levels
ground state of electron
sublevels
types of sublevels (s, p, d, f)
Hund’s Rule
photons
quantum
Pauli exclusion principle
spectrum
principal energy level
Heisenberg uncertainty principle (Honors)
Be able to identify the electron configuration of elements.
Electron Configurations. What element has the electron configuration [Ne]3s23p1? _____
• What does the 3 mean in 3s2 ?
•
What does the s mean?
•
What does the 2 mean?
•
How many valence electrons will an atom of this element have?
•
How many electrons will an atom of this element lose to form an ion? Why?
•
Write out the orbital configuration according to the Aufbau principle, Hund’s rule and the
pauli exclusion principle.
Emission (or bright-line) Spectrums (Honors only)
• What is needed for an electron to “jump” to a higher energy level?
• What happens when an “excited” electron falls back to its ground state?
• What does an emission spectrum allow one to do?
Characteristics of subatomic particles
Particle
Mass
Charge
Location in atom
Proton
Neutron
Electron
Ch. 6: The Periodic Table
alkali metals
alkaline earth metals
anion
atomic radius
cation
electronegativity
halogens
inner transition metals
ionization energy
metalloids
metals
noble gases
nonmetals
transition metals
Know the four periodic trends of ionization energy, electronegativity, ionic size and
atomic size.
Periodic trends
Locate or define parts of the periodic table:
• Groups
•
Periods
•
Transition metals (d & f blocks) vs. Representative Elements (s & p blocks)
•
Alkali metals, Alkaline Earth metals, Halogens, Noble Gases
Periodic Trends: Increasing or Decreasing from top to bottom or left to right?
Top to Bottom in a Group
Left to Right across a Period
electronegativity
ionization energy
atomic size
Ionic size
Elements in the same ___________ have similar physical and chemical characteristics
because the
(group, period)
they have the same number of _____________________.
(atoms, protons, neutrons, electrons, valence electrons)
Draw a electron dot diagram (or Lewis Dot structure) for Be and for N
correct number of valence electrons
showing the
From their positions on the periodic table, what charges would the ions of Be and N have?
Gains or loses electrons?
Be
Symbol for ion
Gains or loses electrons?
Symbol for ion
N
Properties of Metals vs. Nonmetals vs. Metalloids
Metals
Nonmetals
Luster?
Malleable vs. Brittle
Conducts electricity &
heat?
Typical state(s) at
room temperature
Metalloids
Ch. 7: Ionic and Metallic Bonding
chemical formula
electron dot structure for ions
halide ion
ionic bonds
ionic compounds
metallic bonds
octet rule
valence electrons
p. 192 Common Anions
Key concepts: Ions
Ionic Bonds and Ionic Compounds
Bonding in Metals
Ch. 8: Covalent Bonding
covalent bond
coordinate covalent bond
diatomic molecule
double covalent bond
hydrogen bonds
molecular compound
Key concepts:
molecule
nonpolar covalent bond
polar bond
polar covalent bond
polyatomic ion
electron dot structures
8.1 Molecular Compounds
8.2 Nature of Covalent Bonding
8.4 Polar Bonds and Molecules
triple covalent bond
unshared electron pair
structural bond
molecular formula
single covalent bond
VSPER
(8.3 Honors)
Ionic vs. Molecular Compounds:
Ionic bonds are formed when a ____________ and a _________________ combine.
Metals lose electrons and form _____________ while nonmetals gain and electrons form
__________.
Molecular compounds form when a ______________ and a _______________ combine as they
share electrons.
Identify the following pairs of atoms as potentially forming an ionic or molecular compound:
Mg and Cl ____________
I and F _________________ P and Cl _______________
Ag and S _____________
K and Br ________________ Sn and O _______________
Covalent Bonding in Molecules
Draw Lewis Structures (dot diagrams) for HCl, H2S, CH2Cl2, and O3.
Use Lewis Structures to predict molecular shapes and polarity of molecules
• Identify shapes of the molecules as: Linear, Bent, Pyramidal, Trigonal Planar, or
Tetrahedral.
• Use electronegativity values to determine if the individual bonds in the molecules above
are polar.
• Look at the polarity of the bonds and the symmetry of the molecules above to determine
if the molecules are polar (if one side of the molecule will be more negative than
another).
H
2.1
Li
1.0
Na
0.9
K
0.8
He
Be
1.5
Mg
1.2
Ca
1.0
Lewis
Structures
& Total #
of
Valence
Electrons
Structural
Formula
B
2.0
Al
1.5
Ga
1.6
C
2.5
Si
1.8
Ge
1.8
N
3.0
P
2.1
As
2.0
O
3.5
S
2.5
Se
2.4
N2
F
4.0
Cl
3.0
Br
2.8
Ne
Ar
Kr
HCl
Electronegativity
Difference (x)
Type of Bond
0.0 ≤ x ≤ 0.4
Non–polar
covalent
0.4 < x < 2.0
Polar covalent
2.0 ≤ x
Ionic
H2 S
CH2Cl2
How do you know if a Bond is polar or nonpolar?
How do you know if a molecule is polar or nonpolar?
Properties of Ionic and Molecular Compounds
Molecular Compounds
Combination of elements involved
(metals? nonmetals?)
How is bond formed?
Ionic Compounds
O3
Typical state(s) at room temperature
Melting and boiling points
(relatively high or low?)
Conduct electricity if dissolved in
water?
Ch. 9: Chemical Names and Formulas
Know how to name ionic compounds, molecular compounds, and polyatomic ions.
p. 257
Naming Molecular and Ionic Compounds
Naming molecular compounds
• Name: N2O: ___________________________ and NO2
____________________________
•
Naming Ionic Compounds
Name: Li2O ___________________________ and (NH4)2SO4
__________________________
•
Name: FeO __________________________ and Sn3(PO4)4
___________________________
•
Name: NaHCO3 ______________________ and CuCl2
_______________________________
Formulas of Molecular and Ionic Compounds
Write formulas for the following molecular compounds:
Water _____________________________________ silicon dioxide
_________________________
Phosphorous trihydride _______________________ dioxygen difluoride
_____________________
Lead (II) hydroxide __________________________ chromium (III) sulfate
___________________
Write formulas for:
_________________
Ba2+ with OH– ___________ iron (III) sulfide
Na+ with OH– ___________ NH4+ with PO43– ______________ magnesium
oxide____________