Download General Chemistry First Semester Review General

General Chemistry Semester One Final Review
The following review is a summary of the important topics learned during first semester. Strictly studying the questions at the
end of the review will not guarantee a good grade. Correctly completing each question should give you a good idea of what
your strengths and weaknesses are. Then work on improving your weaknesses by working on the necessary unit (perhaps
redoing a unit exam review). You need to know all the topics covered on the final review! ALL UNIT EXAM REVIEWS ARE
POSTED ON MY WEBSITE.
What should you expect to be on the chemistry final? The chemistry final is a cumulative exam testing you on topics from each
of the units given during 1st semester. It will be somewhat longer than a regular unit exam. However, each student should
have more than enough time to finish it during the allotted class time. Chemistry is a sequential topic. For example, in unit one
we learned about the symbols of elements and their atomic structure. In unit two we used the known structure of atoms to
explain why they “need” to bond and how they bond together. Finally, we did a short topic on dimensional analysis (with
significant figures) and the “mole”. Relevant sections from the book have been selected and listed below. Read these sections
to fill in any areas you’re unsure of or just to get the big picture.
Unit One: Elements, Atoms, & the Periodic Table
Reading: 3.1-3.5 4.1-4.11
Concepts:
Qualitative and quantitative data
Metric measurements using units of mass, volume, and distance
Numerical value of prefixes (milli-, centi-, kilo- etc.)
Uncertainty of measurement
- accuracy and precision
- significant figures (sigfigs) in measurements and calculations
Scientific notation
- writing measurements in “standard” scientific notation
- rewriting scientific notation into decimal (long) form
- add/subtract/multiply/divide numbers in scientific notation
Physical properties of matter
Specific physical properties of metals, nonmetals, and metalloids
Elements, Compounds and Mixtures (heterogeneous and homogeneous)
Density (including calculating density from measurements)
Physical change and chemical change
Atomic structure
- What did Dalton, Thomson, Rutherford, Bohr, and Chadwick have to say about atoms?
- Charges, locations, and relative sizes of subatomic particles (protons, neutrons, electrons)
Periodic table
- Organization by atomic number (modern)
- Periodic trends across the table (periods) and down the table (groups)
How atoms change
- Ions (adding or subtracting electrons)
- Element identity (changing number of protons)
Average atomic mass
Unit Two: Naming and Bonding
Reading: 5.1-5.5, 5.7
Concepts:
Naming type I compounds (metal/nonmetal pairs) [with our without polyatomic ions]
Naming type II compounds (metal/nonmetal pairs) [with or without polyatomic ions]
- Understand that the difference between type I and II compounds is the metal part. Type II compounds
possess a metal that has two or more charge choices. Include a Roman numeral in the name to identify which
charge the metal possesses.
Naming type III compounds (nonmetal/nonmetal pairs)
- prefixes identify the number of atoms in the compound.
Writing chemical formulas from the name (when given)
Bonding
a) Ionic bond b) Covalent bonds c) polar covalent bonds
a) Valence electrons b) Lewis structures
Unit Three: Chemical Reactions
Reading: 6.1-6.3 7.1-7.3, 7.7
Concepts:
Chemical reactions neither create nor destroy matter (elements, atoms)
Balanced chemical equations
Types of chemical reactions – a)Synthesis b) Decomposition c) Combustion d) Single replacement e) Double replacement
Identify factors that would change the reaction rate - a) surface area b) temperature c) concentration d) catalyst
Phase notations (solid, liquids, gases, aqueous)
- aqueous (aq) is written if a solution is used
- pure liquids (not a mixture of something) use (l)
- solid: This could refer to a multitude of different substances: metals, flakes, crystals, and precipitates; use (s)
- gas: Use (g), these are usually diatomic molecules such as O2, H2, Cl2, etc.
- read given information carefully. Water vapor is noted with a (g) because it is a gas, not a liquid.
Double replacement reactions are between two aqueous solutions that produce an insoluble precipitate. Solubility rules can
predict which substance is the solid. The solubility table is on p. 178 in the textbook. You do not need to memorize it, just be
able to interpret this table.
Reaction rates: the time it takes a chemical reaction to run to completion can be affected by several different variables 1)
temperature 2) concentration 3) catalysts 4) surface area 5) stirring. This results in more reactions. Decreasing the
temperature lowers the speed of the reactant particles, which collide less frequently for a given amount of time. This results in
fewer reactions. Reactions rates are all determined by the number and frequency of collisions (for instance, when
temperature increases, particles collide faster and more requently).
Unit 4: Dimensional Analysis and the Mole
Reading: 1.1-1.5 2.1-2.6, 2.8
Concepts:
Dimensional Analysis - this is a required skill! Use the Given X conversion factor format
- changing metric units to metric units, changing between metric and English units
- changing between units of measure in a scenario or story problem
Mole conversion problems
We are just going to get started on Moles so I will let you know specifically what you need to know for the Semester Final.
Question Sampler
Know all of the topics listed above. If a particular question does not appear in this sampler, it is not mean that it will
not appear on the Final.
1. Write the following numbers in standard scientific notation:
(a) 7770 (b) 0.0075 (c) 0.125
2. Express the following in decimal form (ordinary number):
(a) 2.5 X 10-3 (b) 3.20 X 102 (c) 6.7 X 105
3. Identify the number of significant figures in each quantity:
(a) 150 m (b) 0.008 m (c) 5.01 X 10-7 km (d) 30.0 seconds
4. Solve showing dimensional analysis:
(a) 32,500 g = ? kg (b) 100 g = ? mg (c) 2.1 inches = ? Km (1” = 2.54 cm)
5. Sketch and label the location and charges of the subatomic particles in an atom of oxygen-15.
6. What physical properties distinguish metals from nonmetals?
7. Elemental oxygen forms diatomic molecules (O2). Draw a Lewis structure for an oxygen molecule (that’s showing the total
valence electrons and how they are arranged to satisfy the “octet rule”).
8. What kind of bonds (ionic or covalent) are most likely holding a particle of magnesium chloride? How do you know?
9. How many valence electrons does an atom of silicon possess?
10. Are metals more likely to accept or donate electrons?
11. What electrical charge does a Br atom form when it becomes an ion?
12. What happens to the number of valence electrons as you go down a group of elements? What happens to the number of
valence electrons as you go from left to right across the periodic table?
13. How is the modern periodic table organized?
14. Draw a Lewis structure of CO2.
15. Balance the following equation: Al2O3(s) + C(s)  Al(s) + CO2(g)
16. Magnesium metal when mixed with silver nitrate solution produces magnesium nitrate and silver metal. Write a balanced
chemical equation for the reaction. Include phase notations.
17. Write the correct systematic name for: (a) Al2(SO4)3 (b) NBr3 (c) NaNO3 (d) FeCO3
18. Write the correct chemical formula for: (a) potassium nitrate (b) diphosphorus pentoxide (c) lead(II)carbonate
19. When solutions of sodium iodide and lead(IV)chloride mixed and brilliant yellow precipitate immediately forms.
(a) Write the balanced chemical equation for the reaction. b) Identify the precipitate and include other phase notations.