Download 1 km = 1 000 m 1 m = 100 cm 1 cm = 10 mm 1 m = 1 000 mm

Lyon Midterm Review Packet
Things we haven’t done yet will not be on the midterm:
Measurement
Quantity
Being
Measured
Units of Measurement
Unit Symbols
mass
kilogram, gram, centigram, milligram
kg, g, cg, mg
volume
kiloliter, liter, milliliter, cubic meter,
cubic decimeter, cubic centimeter, cubic
millimeter
kL, L, mL, m3,
dm3, cm3, mm3,
length
kilometer, meter, centimeter, millimeter
km, m, cm, mm
energy
kilojoules, joules;
chemist unit for heat: calorie
kJ, J;
cal
pressure
millimeters of mercury, Pascals,
kilopascals, atmospheres
mmHg, Pa,
kPa, atm
density
temperature
mass__
volume
degree Celsius;
g__;
mL
Kelvin
o
C;
g__
cm3
K
1 km = 1 000 m
1 m = 100 cm
1 cm = 10 mm
1 m = 1 000 mm
1 kg = 1 000 g
1 g = 100 cg
1 cg = 10 mg
1 g = 1 000 mg
1 kL = 1 000 L
1 L = 1 000 mL
1 dm3 = 1 L
1 mL = 1 cm3
Dimension Analysis
Dimensional Analysis (also called Factor-Label Method or the Unit Factor Method) is a
problem-solving method that uses the fact that any number or expression can be multiplied by
one without changing its value. It is a useful technique. The only danger is that you may end up
thinking that chemistry is simply a math problem - which it definitely is not.
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Unit factors may be made from any two terms that describe the same or equivalent "amounts" of
what we are interested in. For example, we know that
1 inch = 2.54 centimeters
Note: Unlike most English-Metric conversions, this one is exact. There are exactly
2.540000000... centimeters in 1 inch.
We can make two unit factors from this information:
Now, we can solve some problems. Set up each problem by writing down what you need to find
with a question mark. Then set it equal to the information that you are given. The problem is
solved by multiplying the given data and its units by the appropriate unit factors so that only the
desired units are present at the end.
(1) How many centimeters are in 6.00 inches?
(2) Express 24.0 cm in inches.
You can also string many unit factors together.
(3) How many seconds are in 2.0 years?
(4) Convert 50.0 mL to liters. (This is a very common conversion.)
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(5) What is the density of mercury (13.6 g/cm3) in units of kg/m3?
We also can use dimensional analysis for solving problems.
(6) How many atoms of hydrogen can be found in 45 g of ammonia, NH3?
We will need three unit factors to do this calculation, derived from the following information:
1. 1 mole of NH3 has a mass of 17 grams.
2. 1 mole of NH3 contains 6.02 x 1023 molecules of NH3.
3. 1 molecule of NH3 has 3 atoms of hydrogen in it.
There are many, many more examples in your textbook!
QUIZ:
Question
1
Question
2
Question
3
How many millimeters are present in 20.0 inches?
The volume of a wooden block is 6.30 in3. This is equivalent to how many cubic
centimeters?
A sample of calcium nitrate, Ca(NO3)2, with a formula weight of 164 g/mol, has
5.00 x 1027 atoms of oxygen. How many kilograms of Ca(NO3)2 are present?
Answers: (1) 508 mm (2) 103 cm3 (3) 227 kg
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Matter and Change
Physical means of separating a mixture include: filtration, evaporation, using known freezing
points and boiling points to separate different liquids, distillation (boiling off the liquid to leave
the solid component, and then condensing the vapor back to the liquid state).
Physical states of matter:
Solid: particles packed very tightly together, particles are “fixed” in position relative to each
other
Liquid: particles still very close together but particles can move around each other
Gas: particles very far apart from each other
Physical changes involve changes in physical state (solid ↔ liquid ↔ gas: melting, boiling,
condensation, freezing), cutting or crushing a large sample into smaller pieces, dissolving in an
appropriate solvent.
Chemical changes involve the rearrangement of the atoms of one or more substances to form
one or more new substances.
Physical properties of matter include: density, color, physical state at a given temperature,
boiling point, freezing point, odor, malleability, brittleness, hardness, solubility in a given
solvent, crystal shape.
Chemical properties involve how the substance behaves in the presence of another substance:
Does the substance give up electrons easily to another substance? Does the substance take
electrons from another substance? When the substances are mixed together does a new
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substance form? Does a new substance form when the two substances are heated together?
Does the substance combine readily with oxygen gas to form new substances?
Atomic Structure and Nuclear Chemistry
Atomic Structure
Atom: smallest particle of an element; composed of protons, neutrons, and electrons.
Atomic Particle
Location
Electrical Charge
Mass (atomic mass units)
proton
nucleus of atom
+1
1 a.m.u.
neutron
nucleus of atom
0
1 a.m.u.
electron
outside of nucleus
─1
0 a.m.u.
(mass is too small to be
significant)
In a neutral atom, the number of protons equals the number of electrons.
Number of protons + number of neutrons = mass number of the atom (atomic mass in a.m.u.)
The atomic number of an element is determined by the number of protons in every atom of that
element.
If “X” represents the symbol of the element, the mass number is written as a superscript at the
upper left of the symbol, and the atomic number is written as a subscript at the lower left of the
symbol:
mass number
X
atomic number
Isotopes: atoms of the same element that have the same atomic number (number of protons) but
have different numbers of neutrons and therefore different mass numbers.
and 157N are isotopes of nitrogen. Both have atomic number “7” but because the different
numbers of neutrons, the atoms have different mass numbers, 14 and 15 respectively.
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7N
Another way of writing isotopes is to use the element symbol followed by the mass number:
N-14, and N-15.
Percent Composition
There are two different ways to describe the composition of a compound: in terms of the number
of its constituent atoms (like C2H6) and in terms of the percentages (by mass) of its elements.
When showing the constituent atoms of a molecule, you can either show the chemical formula,
which shows the real number of atoms in the molecule, like C2H6, or show the empirical
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formula, which merely shows their relative amounts in a substance, so the above molecular
formula would be expressed as CH3.
You can describe the composition of a compound in terms of the weights of its constituent
elements by determining the percent composition of particular elements in the molecule. To
calculate percent compositions, you would find the weight of each constituent atom, then figure
out what percent of the total molecular weight it makes up. Consider ethanol, C2H5OH. Taking
subscripts into consideration, you have 2 moles of carbon, 6 moles of hydrogen (5 + 1), and 1
mol of O. Now convert moles into grams for each constituent element as well as for the entire
molecule:
Mass of C = 2
Mass of H = 6
Mass of O = 1
12.01
= 24.02 g
1.01
= 6.06 g
16.00
= 16.00 g
Mass of 1 mol of C2H5OH = 46.08 g
Now use the formula you learned above to find the percent compositions of the constituent
elements:
Mass percent of C:
100% = 52.14%
Mass percent of H:
100% = 13.15%
Mass percent of O:
100% = 34.77%
Gas Laws
Kinetic Molecular Theory





The gas consists of objects with a defined mass and zero volume.
The gas particles travel randomly in straight-line motion where their movement can be
described by the fundamental laws of mechanics.
All collisions involving gas particles are elastic; the kinetic energy of the system is
conserved even though the kinetic energy among the particles is redistributed.
The gas particles do not interact with each other or with the walls of any container.
The gas phase system will have an average kinetic energy that is proportional to
temperature; the kinetic energy will be distributed among the particles according to a
Boltzmann type of distribution.
Temperature is defined as the degree of hotness or coldness of a body or environment. In reality
temperature really is the measure of the average kinetic energy of the particles in a sample of
matter. In other words, temperature is a measure of activity and the frequency of collisions of
molecules. Therefore the higher the temperature the faster the molecules move.
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Pressure is defined as the amount of force per unit area. Gas molecules apply a force by
slamming into container walls. The faster they move the harder they hit the more force they
apply the greater the pressure. So temperature and pressure are directly related.
Gas Laws
Go to http://home.comcast.net/~rstine/gaslaws3.html For a complete review of the 7 gas
laws.
Lussac’s Law:
Pressure vs Temperature
Boyle Law:
Pressure vs Volume
Charles Law:
Volume vs Temperature
Combined Gas Law: Pressure, Volume and Temperature are related.
Ideal Gas Law:
PV=nRT
Dalton’s Law:
Ptot= Pa + Pb + Pc….. and Pa = XPtot Where X is the mole fraction or
percentage
Graham’s Law
Velocity of a molecule is proportional to it’s mass.
Atomic Theory
Dalton, Thompson and Rutherford and relevant experiments in determining atomic
structure
Dalton (1808)- Based on experimental results from previous century, devised scheme for each
element being distinct, indivisible ("atomos" in Greek), and with a unique set of properties,
including being very small, and combing with other atoms in distinct ways. He studied the ratios
in which elements combine in chemical reactions. He discovered that all elements are composed
of tiny invisible particles called atoms (an element is composed of only one kind of atoms, and a
compound is composed of particles that are chemical combinations of different kinds of atoms);
He came up with the first atomic theory which consisted of the following points





Atoms of the same element are identical;
The atoms of any one element are different from those of any other element;
Atoms of different elements can physically mix together or can chemically combine to
one another in simple whole number rations to form compounds; Law of Definite
Proportions
Atoms can combine in different ratios to form different compounds. Law of Multiple
Proportions
When reactions occur atoms rearrange, they are not created nor destroyed. Law of
Conservation of Mass
Thomson, Joseph John (J.J.)(1890's)- He performed experiments that involved passing
electronic current through gases at low pressure. He knew that opposite charges attract and like
charges repel, so proposed that a cathode ray is a stream of tiny, negatively charged particles
moving at high speeds..
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Cathode Ray Tube Experiment
Since these particles were smaller than atoms, but seemed to come from them, they must be subatomic parts. He concluded that electrons must be parts of atoms of all elements So, he
discovered the first subatomic particle (electron), the atom is no longer indivisible, and
developed the "plum pudding" (or "chocolate chip cookie") model.
Plum Pudding Model of the Atom
Rutherford, Ernest(1900's)- Discovered atoms are not homogeneous, very small part (the
nucleus) is hard and heavy, electrons on the outside are very light and occupying the vast
majority of the space of the atom. All from the gold foil experiment in which he directed a
narrow beam of alpha particles at a very thin sheet of gold foil and noticed that the alpha
particles should have passed easily through the gold, with only a slight deflection due to the
positive charge thought to be spread out in the gold atoms.
Gold Foil Experiment
He concluded that most of the alpha particles passing through pass through the gold foil
because the atom is mostly empty space. The mass and positive charge is concentrated in a
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small region of the atom. Rutherford called this region the nucleus. Particles that approach
the nucleus are greatly deflected.
Solar System Model of the Atom
Electrons in Atoms
Soon after Rutherford disclosed his model of the atom it was discovered that electrons may
“jump” from one orbit to a higher one. Later it was discovered that this jump happens ONLY IF
the electron absorbs ENOUGH energy to make the jump. This is known as quantized energy.
The Energy released when the electron fell back to the ground state was just as specific as the
energy absorbed by the atom. In order to truly understand what is going on in the atom you must
first understand electromagnetic radiation.
Electro Magnetic Radiation
Electromagnetic Radiation is energy in the form of waves. The electromagnetic spectrum is the
range of all possible electromagnetic radiation. Also, the "electromagnetic spectrum" (usually
just spectrum) of an object is the frequency range of electromagnetic radiation that it emits,
reflects, or transmits. The electromagnetic spectrum, shown in the chart, extends from just below
the frequencies used for modern radio (at the long-wavelength end) to gamma radiation (at the
short-wavelength end), covering wavelengths from thousands of kilometers down to fractions of
the size of an atom. Electromagnetic energy at a particular wavelength λ (in vacuum) has an
associated frequency  and photon energy E. Thus, the electromagnetic spectrum may be
expressed equally well in terms of any of these three quantities. They are related according to the
equations:
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wave speed (c) = frequency x wavelength
c = 
and
E = h
or
where:

c is the speed of light, 299792458 m/s

h is Planck's constant,
.
.
So, high-frequency electromagnetic waves have a short wavelength and high energy; lowfrequency waves have a long wavelength and low energy.
Bohr Atom
Niels Bohr quickly seized upon this used it to propose a quantized description of the atom.
1. Bohr proposed that while circling the nucleus of the atom, electrons could only occupy
certain discrete orbits, that is to say energy levels. Bohr used Max Planck's equations
describing quanta of radiation to determine what these discrete orbits would have to be.
As long as electrons stay in these energy levels, they are stable.
2. Further, Bohr said electrons give or take energy only when they change their energy
levels. If they move up, they take energy (say from light), and if they move down, they
release energy. This energy itself is released in discrete packets called photons.
3. Furthermore, Bohr also said that an electron which is not in its native energy level (in
other words, which has been excited to a higher energy level) always has to fall back to
its original, stable level. Electrons may be excited by heat, light, electricity, or any other
form of energy. The photons released by excited electrons returning to their normal
energy levels accounts for the colors we see in flame tests, fireworks, any fire such as that
in a fireplace or a lit match, and in the colors of our clothes (the electrons of the atoms in
dye molecules are excited by light energy).
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Bohr Model of the Atom
Bohr interpreted the lines in the spectra of gases as formed by the transitions of electrons to and
from various energy-levels.
Hydrogen Spectrum Using the Bohr Atom
Bohr also assumed that the electron can change from one allowed orbit to another
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The Periodic Table
Facts about the Periodic table

The periodic table was discovered in 1869 by Dimitri Mendeleev.

The periodic table can be divided into groups and periods. The vertical columns are
called GROUPS or “families” and are numbered left to right as 1-18. The horizontal
rows are called PERIODS and are numbered top to bottom as 1-7 on the current periodic
table.

Elements in the same group or family have similar chemical and physical properties.

Group 1 is the ALKALI METAL group.

Group 2 is the ALKALINE EARTH METAL group.

Groups 3-12 are the TRANSITION METAL groups. These metals often form
compounds that are colored (red, green, yellow, etc.)

Group 13 is known as the BORON group.

Group 14 is called the CARBON group.

Group 15 is the NITROGEN group.

Group 16 is the OXYGEN group.

Group 17 is known as the HALOGEN group.

Group 18 is the NOBLE GAS group.
Electron Configuration
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
How many electrons can fit into the n=4 principle energy level?

Write the orbital notation, electron configuration notation, and Noble Gas notation for
electrons in different elements

Determine an element based on its electron configuration or Noble Gas notation

Explain the 3 rules that govern electron configuration
Chemistry
Midterm Exam Review 2012
Name
_____
Date _____/______/______
Unit 1. Matter and Change
1. Identify each of the following as an element, a mixture, or as a compound:
a) iced tea
b) ice
c) table sugar
d) silver
2. Classify each of the following as either a physical change or as a chemical change:
a) bending a piece of glass
b) melting an ice cube
c) cooking a steak
d) cutting grass
e) burning wood
f) sugar dissolving in water
g) boiling water
3. Give the correct symbols for each of the following elements:
a) sodium
e) copper
b) aluminum
f) magnesium
c) chlorine
g) iron
d) sulfur
h) nitrogen
4. List the three common phases or states of matter: _________________________,
___________________________, and _________________________.
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5. Which of the following is NOT a physical property of matter? _______________________
density
texture
color
flammability
odor
malleability
melting point
luster
boiling point
Unit 2. Measurement
1. Rank these measurements from the smallest to the largest: (hint: change all numbers to
scientific notation)
a) 5.3 x 104 m
d) 0.005 7 m
b) 7.7 x 103 m
e) 5.1 x 10-3 m
c) 4.9 x 10-2 m
f) 0.072 m
Correct order:
2. For each of the following pairs of units, which is the larger unit?
a) centigram or milligram
b) liter or centiliter
c) calorie or kilocalorie
d) millisecond or centisecond
e) milliliter or kiloliter
f) mm3 or m3
Show all work.. units on every number, cancel units where possible.
3. An object has a volume of 3.5 cm3 and a mass of 27.2 g. What is the density of the
object?
4. What is the volume of an object having a density of 1.05 g/mL and a mass of 4.85 mL?
Show all work. units on every number, cancel units where possible, pay attention to sig. figs
and units in answer.
5. Write the following measurements in scientific notation:
a) 572.5 km
b) 0.005 725 m
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6. Write the following measurements in standard notation:
a) 4.45 x 10-3 g
b) 4.45 x 107 mg
Unit 3. Gas Laws
Abbreviations
atm - atmosphere
mm Hg - millimeters of mercury
torr - another name for mm Hg
Pa - Pascal (kPa = kilo Pascal)
K - Kelvin
°C - degrees Celsius
Conversions
1 cm3 (cubic centimeter) = 1 mL (milliliter)
1 dm3 (cubic decimeter) = 1 L (liter) = 1000 mL
Standard Conditions
0.00 °C = 273 K
1.00 atm = 760.0 mm Hg = 101.325 kPa = 101,325 Pa
1. A balloon has a volume of 2500.0 mL on a day when the temperature is 60.0 °C. If the
temperature at night falls to 20.0 °C, what will be the volume of the balloon if the pressure
remains constant?
2. 300.0 mL of a gas are under a pressure of 600.0 torr. What would the volume of the gas
be at a pressure of 1000.0 torr?
3. If a gas in a closed container is pressurized from 12.0 atmospheres to 17.0 atmospheres
and its original temperature was 25.0 °C, what would the final temperature of the gas be?
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4. A container with two gases, helium and argon, is 60.0% helium. Calculate the partial
pressure of helium and argon if the total pressure inside the container is 4.00 atm.
5. 2.00 liters of hydrogen, originally at 25.0 °C and 650.0 mm of mercury, are heated until a
volume of 10.0 liters and a pressure of 2.50 atmospheres is reached. What is the new
temperature?
6. A 10.0 g sample of gas occupies 16.2 L at STP. What is the molecular weight of this gas?
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7. What is the formula weight of a gaseous compound if at room temperature it effuses
though a pinhole 2.86 times as rapidly as xenon, (Xe)? Which compound is it, CH4, C2H4,
C3H4, C3H6, C3H8
8. What is temperature?
9. What does a decrease in the temperature of a substance mean with respect to the average KE
of the particles of the substance?
10. What does the term absolute zero on the Kelvin scale mean?
Unit 4. Atomic Structure and Nuclear Chemistry
1. Complete the following table – assume neutral atoms. Use your periodic table as a reference.
Element
K
# protons
Mass #
# electrons
Atomic #
# neutrons
31
20
12
17
2. For each of the following atoms, give the atomic number and mass number:
a)
36
18Ar
atomic #:
mass #:
b)
79
35Br
atomic #:
mass #:
Complete the following sentences:
3. The atomic number is determined by the number of
4. In a neutral atom, the number of
in an atom.
equals the number of
5. The mass number is determined by the total number of
and
.
.
6. The difference in mass of isotopes of the same element is due to different numbers of
________________________ in the nucleus of the atoms.
_____ 7. The major weakness in Rutherford’s model of the atom was that
A) negatively charged electrons orbited the positively charged nucleus
B) there was no explanation of why there were no neutrons in the nucleus
C) there was no explanation of why the negatively charged electrons didn’t fall into the
positively charged nucleus
D) there was a dense positively charged nucleus
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Match the atomic model with the appropriate scientist or theory name
_____ 8. John Dalton
A) dense positively charged nucleus with negatively
charged electrons moving around the nucleus
_____ 9. Ernest Rutherford
B) positively charged sphere in which negative charges
are uniformly distributed
_____ 10. J.J. Thomson
C) tiny, indivisible bits of matter; different kinds of
matter made of different kinds of atoms
_____ 11. Niels Bohr
D) tiny invisible spherical bits of matter having a
uniform density and which combine in definite
proportions to form compounds
_____ 12. ancient Greeks
E) dense positively charged nucleus with negatively
charged electrons moving at fixed distances from the
nucleus; electrons could move from one level to a
higher level if sufficient energy is absorbed
_____ 13. modern atomic theory
F) dense positively charged nucleus around which
electrons move in areas of highest probability
Unit 7 Periodic Trends
1. In the periodic table, a column of elements is known as a
2. The alkali metals are located in Group
.
3. The alkaline earth metals are located in Group
4. The halogens are located in Group
5. The noble gases are located in Group
.
.
.
.
6. The element located in Group 15, Period 4 is
.
Two characteristics of metals are
7)
8)
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Two characteristics of nonmetals are
9)
10)
11.
If an atom or group of atoms gains or loses one or more electrons, an
is formed.
Unit 6. Electrons in Atoms
_____ 1. The light energy emitted by an “excited” electron is
A) lost as the electron returns to a higher energy level
B) lost as the electron returns to a lower energy level
C) gained as the electron returns to a lower energy level
D) gained as the electron returns to a higher energy level
_____ 2. When an element is heated sufficiently to excite its electrons and the emitted light is
passed through a diffraction grating or a prism, you would expect to observe
A) a continuous spectrum
C) a single color
B) a bright-line spectrum
D) white light
_____ 3. A single burst of light is released from an atom. Which statement BEST explains
what happened in the atom?
A) An electron pulled a neutron out of the nucleus.
B) An electron moved from a higher energy level to a lower energy level.
C) An electron changed from a particle to a wave.
D) An electron pulled a proton out of the nucleus.
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