Final Study Guide: Part 1
Chem.
Mr. Malveaux
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.
(* The study of matter and the changes it undergoes)
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?
Classify the following as physical or chemical changes:
a. spoiling of milk ___________________
b. bending wire _____________________
c. cutting paper _____________________
d. rusting of a nail ___________________
 Put the following into a graphic organizer/flowchart and define each:
 Matter
 Pure substances – elements and compounds
 Mixtures – homogeneous (solutions) and heterogeneous
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
 Calculations using the Law of Conservation of Mass for Reactions
4g H2 + ?? g O2 → 36g H2O
Atom
 For this Carbon–14 isotope, 146 C
 Atomic number = _____, Mass number = _____,
Gas

# 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%.
 Atomic Models:
Philosophers: Democritus (believed in atoms) and Aristotle (didn’t believe in atoms)
Scientists: What was the contribution of each one’s atomic model? Draw a model of each.
 John Dalton
List the four postulates of Dalton’s Atomic Theory:
 J.J. Thompson

Earnest Rutherford

Niels Bohr

Quantum mechanical model (Werner Heisenberg):
 Energy levels (n=1, 2, 3, 4,…) – represented by periods on the periodic table
 Sublevels: (s, p, d, f) – represented by blocks on the periodic table
 Orbitals – region of space where up to 2 electrons may be found
 Electron Configurations. What element has the 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 electron distribution according to Hund’s rule. The 1s2 sublevel is done for
you.
1s2_↑ ↓___
2s2 ____
2p6 ____ ____ ____
3s2 ____
3p6 ____ ____ ____
 Emission (or bright-line) Spectrums
 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 (radio – gamma) allow one to do?
 Characteristics of subatomic particles
Particle
Proton
Neutron
Electron
Periodic trends
Mass
Charge
Location in atom
 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
 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
Luster?
Malleable vs. Brittle
Conducts electricity
& heat?
Typical state(s) at
room temperature
Nonmetals
Metalloids
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
He
Electronegativity Type of Bond
2.1
Difference (x)
Li Be
B
C
N
O
F
Ne
Non–polar
0.0  x  0.4
1.0 1.5 2.0 2.5 3.0 3.5 4.0
covalent
Na Mg Al Si
P
S
Cl Ar
0.4 < x < 2.0
Polar covalent
0.9 1.2 1.5 1.8 2.1 2.5 3.0
K Ca Ga Ge As Se Br Kr
Ionic
2.0  x
0.8 1.0 1.6 1.8 2.0 2.4 2.8
N2
Lewis
Structures &
Total # of
Valence
Electrons
Structural
Formula
Shape of
Molecule
Bond Angle
Hybridization
HCl
H2S
CH2Cl2
O3
Properties of Ionic and Molecular Compounds
Molecular Compounds
Ionic Compounds
Combination of elements involved
(metals? nonmetals?)
How is bond formed?
Typical state(s) at room
temperature
Melting and boiling points
(relatively high or low?)
Conduct electricity if dissolved in
water?
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____________