Download NYS Regents Chemistry

Study Guide for
NYS Regents Chemistry
Midterm Examination
Table of Contents:
Topic
Page
1. Scientific Method ………………………………………………………………..
1
2. Significant Figures ………………………………………………………………
2
3. Scientific Notation ………………………………………………………………
4
4. Accuracy and Precision …………………………………………………………
4
5. Graphs …………………………………………………………………………...
4
6. Density …………………………………………………………………………..
4
7. Error Calculations ……………………………………………………………….
5
8. Unit Changes …………………………………………………………………….
5
9. Standard Units ……………………………………………………………………
5
10. Atomic Structure ………………………………………………………………...
6
11. Periodic Table …………………………………………………………………...
9
12. Chemical Formulas ……………………………………………………………...
11
13. Mole Information ………………………………………………………………..
14
14. Chemical Equations ……………………………………………………………..
16
15. Matter ……………………………………………………………………………
18
16. Energy and Heat …………………………………………………………………
19
17. Gas Laws …………………………………………………………………………
21
MHS Regents Chemistry
Midterm Review 3rd Edition
1) Scientific Method – a logical, systematic approach to the solution of a scientific problem.
a) Make observations. An observation can lead to a question.
b) Hypothesis – An educated guess based on observed facts. A hypothesis can be revised
based upon experimental data.
c) Controlled Experiments – All factors or variables are held constant while only one
variable is changed at a time in order to see the effect of that variable on the experiment.
d) Data – The results of an experiment, which often include a collection of measurements
e) Theory – Provides a general explanation for the observations made of many scientists
working in different areas of science over a long period of time. Answers the “why” and
“how” questions.
f) Modify Theories if necessary.
g) Scientific Laws. A Scientific Law simply states a relationship between observed facts. It
describes natural events but does not explain why or how they occur.
Questions:
- Why should a hypothesis be developed before experiments take place?
-
What is the difference between a theory and a scientific law?
-
What happens to a theory when a new observation is made that the theory should be
able to explain but cannot?
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Midterm Review 3rd Edition
Page 1 of 23
2) Significant figures – The digits in a number that represent a quantity actually measured.
Significant figures (sig.figs.) indicate the “accuracy” in a number.
a) Obtaining laboratory measurements with correct number of significant figures
(accuracy): use the maximum number of units available – read carefully
50
40
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b) Rules for deciding the number of significant figures in a measured quantity:
(1) All nonzero digits are significant:
1.234 g has 4 significant figures,
1.2 g has 2 significant figures.
(2) Zeroes between nonzero digits are significant:
1002 kg has 4 significant figures,
3.07 mL has 3 significant figures.
(3) Leading zeros to the left of the first nonzero digits are not significant; such zeroes merely
indicate the position of the decimal point:
0.001 oC has only 1 significant figure,
0.012 g has 2 significant figures.
(4) Trailing zeroes that are also to the right of a decimal point in a number are significant:
0.0230 mL has 3 significant figures,
0.20 g has 2 significant figures.
(5) When a number ends in zeroes that are not to the right of a decimal point they are not
significant:
190 miles has 2 significant figures,
50,600 calories has significant figures.
c) Multiplication/Division Rule using Sig.Figs.
i) The number of SF in the answer cannot be more than the least number of Sig. Figs.
used in the calculation e.g. 3.25 x 0.025 = 0.081 (having 2 Sig. Figs. in the answer)
d) Addition/Subtraction Rule using Sig.Figs.
i) Line up the decimal places of each number, calculate the answer and round the
answer to the smallest decimal place in any of the numbers used.
Question
2.54 x 15 x 352 =
(50500)(2.5)/(15.0) =
2.135 + 53.24 + 873.2 =
85.967 – 0.12 =
MHS Regents Chemistry
Answer
Question
203.0 / 0.20 =
(456.03)/((2.67)(301)) =
485.23 – 32.1 =
0.0508 – 0.004 =
Midterm Review 3rd Edition
Answer
Page 3 of 23
3) Scientific Notation (SN)
a. Change the decimal to SN
i. 0.00013 to 1.3 x 10-4 : move the decimal point between the first and
second significant figure and the power of ten is the number of places
moved (right is negative and to left is positive); number smaller than
1.0000 has a negative power of ten.
b. Change SN into a decimal
i. 5.4 x 103 to 5400 positive power of ten – move decimal point to the right
the number of spaces equal to the power of ten; negative power of ten –
move decimal point to the left the number of spaces equal to the power of
ten
c. Preserve the number of significant figures when changing between decimal and
SN!
Decimal
# Sig. Figs.
Sci. Notation
Sci. Notation # Sig. Figs.
Decimal
0.0020300
5
2.0300 x 10-3
2.05 x 105
3
205,000
d. Enter SN, 1.3 x 10-4, into your NON-TI83 calculator
i. Casio calculators usually enter the number given above as “1 . 3 EXP – 4”
ii. TI calculators usually enter the number given above as “1 . 3 (2nd for
TI83) EE – 4”
4) Accuracy compared to Precision
e. Precision – The reproducibility of a series of measurements
f. Accuracy – How close a measurement is to the true or accepted value.
5) Graphs
g. Horizontal axis (abscissa) – the independent variable
i. The variable you control
h. Vertical axis (ordinate) – the dependent variable
i. The variable you measure as a result of your experiment
Example: You add water to a graduated cylinder in 5.0 mL increments. You measure
the mass of the water at each 5.0 mL increment.
The volume of water would be the independent variable and go on the x-axis
The mass of the water increments would be the dependent variable and go on the y-axis
6) Density: Formula can be found on Table t
i. D = mass / volume
i. Volume can be calculated by direct measurement: (length) x (width) x
(height)
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Midterm Review 3rd Edition
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ii. Volume can be calculated indirectly by volume displacement: (volume of
liquid with submerged substance) – (volume of liquid without submerged
substance)
j. D = (gram formula mass)/(22.4L) for gases at STP
Questions:
What is the mass of 78.2 cm3 of nickel at STP?
What is the density of SO2 gas at STP?
7) Error calculations:
k. Relative (percent) error: (Be careful to divide by the accepted value and not the
measured one.)
measured value – accepted value
Percent error = --------------------------------------- x 100
accepted value
8) Unit Changes Table C of Reference Tables
l. Using dimensional analysis (Factor Label Method)
i. Starting unit in numerator (over 1), then multiply by ratio of units such
that starting unit is in denominator and new unit is in the numerator.
Continue until desired unit is reached.
m. K H D (m) D C M
i. Kids Hate Doing Math During Certain Months
ii. Kilo Hecto Deca (unit) Deci Centi Milli
iii. k
h
da
d
c
m
iv. change units by moving from starting unit to final unit and move decimal
point in the same direction by the same number of moves.
9) Standard Units Table D of Reference Tables
n. Standard International Units (SI units)
i. Time: second
ii. Temperature: K (Kelvin)
iii. Mass: kg (kilogram)
iv. Length: m (meter)
v. Amount of substance: mole
MHS Regents Chemistry
Midterm Review 3rd Edition
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10) Atomic Structure
o. Models of the atom developed over a long time.
i. Dalton: solid atom as a ball and is the smallest indivisible unit of matter
ii. Thomson: discovered the electron; used cathode ray tube, small negative
particles (electrons) came out of all matter, Plum Pudding Model
iii. Rutherford: discovered the nucleus; Gold foil experiment shot alpha
particles at a gold foil and almost all particles went straight through only a
few deflected away; Conclusions: atoms are made up of mostly empty
space with a small, dense, positively charged center. The electrons were
somewhere outside the nucleus in the “mostly empty space.”
iv. Bohr: electrons must be in energy levels at specific distances from the
nucleus and never between these levels; concluded this from examining
the bright line spectrum of hydrogen atoms. Each energy level has a
specific energy. The further the level is away from the nucleus the greater
the energy of the electrons in it.
1. Bright line spectrum: When an electron in an atom gains just the
right amount of energy, from an outside source, electron can shift
to a higher energy state (excited state). However, the excited state
atom is energetically unstable and the electrons will return to the
ground state – giving off the absorbed energy in the form of light.
The amount of energy in the released light (wavelength) depends
on how many energy levels the electron jumps back, how many
other electrons are around and also the charge of the nucleus.
Every element has a different number of electrons and protons and
therefore will produce different light energies. The pattern of light
colors produced by an element is its bright line spectrum and is
unique to that element.
2. Although atomic line spectra were known before Bohr proposed
his model of the hydrogen atom, Bohr was able to apply
mathematics to his model and was able to account for each line in
the visible spectrum of hydrogen. The Bohr model failed to
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explain the energies absorbed or emitted by atoms with more than
one electron (hydrogen)
v. Electron Cloud (Quantum Model)(Modern Model)(Wave-mechanical
Model): electrons arranged around atomic nucleus in bands or regions
likely to contain electrons
1. Orbitals are regions around the nucleus that electrons are most
likely to be found.
p. Protons: particle in nucleus; +1 charge; mass = 1 amu (atomic mass unit)
q. Neutron: particle in nucleus; no charge; mass = 1 amu
r. Electron: particle located outside the nucleus; -1 charge; negligible mass
( 1 amu)
1837
s. 1 amu (atomic mass unit) = 1 mass of one 12C atom.
12
t. Atomic Number: defines an element; number of protons in the nucleus of an
atom
e.g. 7N nitrogen has 7 protons in its nucleus; also is equal to the number
of electrons in any uncharged atom.
u. Mass Number: (number of protons) + (number of neutrons) in the nucleus of an
atom. Mass Number = number of nucleons
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Midterm Review 3rd Edition
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v. Number of neutrons in a nucleus calculated by (Mass Number)-(Atomic Number)
w. Isotope: atoms of a same element (atomic number) but have a different mass
number; different number of neutrons e.g. 12C and 14C; also indicated as C-12 and
C-14
x. Average atomic mass: the weighted average mass of all naturally occurring
isotopes, shown on periodic table as mass e.g. 14.0067N;Average atomic mass =
(decimal of percent abundance)(mass of that isotope)+(decimal of percent
abundance)(mass of that isotope)+(decimal of percent abundance)(mass of that
isotope) etc. for each isotope. (NOTE: if isotope masses are not given then use
the isotope mass numbers instead) – Use the link below for a tutorial and video:
http://www.kentchemistry.com/links/AtomicStructure/atomicmasscalc.htm
Example: The element copper has naturally occurring isotopes with mass numbers of 63
and 65. The relative abundance and atomic masses are 69.2% for mass = 62.93amu, and
30.8% for mass = 64.93 amu. Calculate the average atomic mass of copper.
(69.2 x 62.93) + (30.8 x 64.93) =
100
4354.76 + 1999.84 =
100
63.55
y. Nucleons: particles in the nucleus are protons and neutrons
MHS Regents Chemistry
Midterm Review 3rd Edition
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z. Ions
i. Negative ions (anions) are formed when atoms gain electrons; larger than
atom
Atomic radius < Ionic radius
ii. Positive ions (cations) are formed when atoms lose electrons; smaller than
atom
Atomic Radius > Ionic radius
Questions:
- Calculate the atomic mass of magnesium given the following information:
Isotope
magnesium-24
magnesium-25
magnesium-26
-
-
Relative Abundance
78.70%
10.13%
11.17%
Atomic Mass
23.985
24.986
25.983
How many protons, neutrons and electrons do each of the following have?
o 39K
o 33S-2
What is the symbol for an atom or ion with 10 neutrons, 10 electrons and 9 protons?
MHS Regents Chemistry
Midterm Review 3rd Edition
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11)
Periodic Table
aa. Elements were at first arranged according to mass
bb. Today elements are listed in order of atomic number (number of protons in the
nucleus)
cc. Elements in the same row (period) have the same number of occupied principal
energy levels (period number also equals number of outermost principal energy
level containing electrons), but have different chemical and physical properties
dd. Elements in the same group (column) have the same number of electrons in their
outermost level and therefore have similar chemical properties
i. Group 1 “Alkali Metals”; Group 2 “Alkaline Earth Metals”; Groups 3-11
“Transition Metals”; Group 17 “The Halogens”; Group 18 “Noble Gases”
ee. Valence shell: the outermost level containing electrons
i. Lewis Dot Structure: Is the atomic symbol with its valence electrons
drawn around it.
1. Atoms (example):
2.
Ions (examples):
..
Ca
and
+2
[Ca]
..
ּ O.. ּ
..
; [: O :] -2
..
ff. Kernel – part of atom NOT including valence electrons
gg. Three types of elements: metals, non-metals and metalloids
hh. Metals and non-metals separated by “staircase” beginning at Group 13
i. Metals to the left of the “staircase” (except H) (most elements are metals)
ii. Non-metals to the right of the “staircase” (including H)
ii. Properties of Metals:
i. Are mostly solids (one liquid, Hg)
ii. Lose electrons easily (low ionization energy)
iii. Don’t want to gain more electrons (low electronegativity)
iv. Good conductors of heat and electricity (due to mobile electrons)
v. Ductile – ability to be drawn into wires
vi. Malleable – ability to be rolled or hammered into thin sheets
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vii. Luster – glossy finish, reflect light
viii. Want to give away electrons to form positive ions (cations) with smaller
radii
jj. Metallic properties are more pronounced with increasing atomic size and fewer
valence electrons.
kk. Transition elements are elements in groups 3 through 11
i. These elements generally have a color when they are ions (either in a
compound or dissolved in water).
ll. Properties of Non-metals:
i. Are mostly gases (some solids, one liquid, Br2)
ii. Want to gain electrons and form negative ions (anions) with larger radii
(high electronegativity)
iii. Don’t want to give up electrons (high ionization energy)
iv. Poor conductors of heat and electricity
v. Solids tend to be brittle
vi. Solids generally have a dull finish
mm.
Metalloids: 6 elements on the “staircase” except Al, At, and Po which are
metals; have properties of both metals and non-metals (semimetals or
semiconductors)
nn. All elements with atomic numbers greater than ‘83’ have no known stable
isotopes. They are all radioactive.
MHS Regents Chemistry
Midterm Review 3rd Edition
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oo. Diatomic elements: Br I N Cl H O F – Bromine, Iodine, Nitrogen, Chlorine,
Hydrogen, Oxygen, and Fluorine.
pp. Gas, Liquid, and Solid Elements
i. 11 Gases at STP: Fluorine, oxygen, hydrogen, chlorine, nitrogen (last 5 of
Br I N Cl H O F) plus all 6 of the Noble Gases
ii. 2 Liquid elements at STP: Bromine and Mercury (Hg)
iii. ALL OTHER ELEMENTS ARE SOLIDS at STP
qq. Atomic radius
i. Increase size going down any group (new electron shells and increased
nuclear shielding)
ii. Decrease size going left to right across any Period (row) (increased
nuclear charge)
rr. Electronegativity: the “pull” atoms of an element have to gain another electron
ss. First Ionization Energy: the energy necessary to remove the first valence electron
from an atom in the gaseous state.
tt. Fluorine is the most electronegative element; most active non-metal
uu. Francium is the least electronegative element; most active metal
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vv. Highly reactive elements – found only as compounds in nature
i. Group 1: Alkali Metals
ii. Group 2: Alkaline Earth Metals
iii. Group 17: Halogens (has elements in all three phases)
iv. The most reactive of these can only be purified by electrolysis
ww.
Group 18: Noble Gases (Monatomic Molecules)
i. Generally not reactive – already have the stable octet valence electron
configuration
xx. Allotropes: one or more molecular forms of the same element in the same state.
Examples: carbon as diamond, graphite and coal; oxygen as O2 or ozone, O3
yy. Ground state electron configuration: when electrons of an atom are occupying the
lowest possible energy levels, which is the electron configurations given on the
Periodic Table. For example: Mg is 2-8-2
zz. Excited state electron configurations: when an electron(s) are in a higher energy
level than the lowest possible energy levels. Example: excited Mg 2-7-3
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12) Chemical Formulas
aaa.
Binary compounds: contain atoms of only 2 different elements; name
ends with “-ide”
bbb.
Ternary compounds contain more than 2 elements: will contain a
polyatomic ion if carbon is not one of the elements (Table E)
ccc.
First element (ion) has a positive oxidation number (charge if ion); the last
element (ion) has a negative oxidation number (charge if ion)
ddd.
Writing empirical formula from the name
i. Use symbols for elements (Table S) or formula given for polyatomic ion
(Table E)
ii. Determine oxidation number of each element or ion in the compound
1. Oxidation number for metal either given in the name or is the only
one possible and shown on the Periodic Table
2. Oxidation number for non-metal atoms are usually the first one
listed on the Periodic Table for the element
3. Work with a polyatomic ion as a whole using its charge
iii. Criss-cross oxidation number (charge if ion) to the subscripts and reduce
by the greatest common factor. (remember to use parentheses around
polyatomic ions if there are more than one in a compound)
1. For ionic compounds subscripts in its chemical formula MUST be
reduced to their lowest whole number ratio. E.g. Ti2O4 needs to
be reduced to TiO2
eee.
Writing the name from the formula
i. Compounds with metal and nonmetals OR ionic compounds (Stock
System Method)
1. Determine the oxidation numbers of all individual element in
compound (or charge if ion - from Table E)
a. Group 1 metals always have a +1 oxidation state
b. Group 2 metals always have a +2 oxidation state
c. Fluorine always has a –1 oxidation state
d. All uncombined elements have an oxidation state = 0
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e. Hydrogen has a +1 oxidation state except when combined
with metals has a –1 oxidation state
f. The total of all oxidation states must = the charge of the
substance
2. Write name of metal then oxidation state of metal as roman
numeral in parentheses if more than one is possible – not necessary
if only one choice possible (according to Periodic Table)
3. Write name of second atom or polyatomic ion
a. If a single element then use the element name with an “ide” ending, e.g. FeCl3 is Iron(III)chloride
b. If a polyatomic ion – just use the name of the polyatomic
ion, e.g. Fe(NO3)2 is Iron(II)nitrate
ii. Compounds containing only NONMETAL elements (Traditional Method)
(On the Regents Exam Stock System Method is generally used.)
1. Use the names of the elements (Table S)
2. Use prefixes: 1-“mono”, 2-“di”, 3-“tri”, 4-“tetra”, 5-“penta”, 6“hexa”, 7- “hepta”
a. Except never use “mono” for first element in the name
b. When the prefix ends in a or o and the name of the anion
begins with a vowel, the a or o is often dropped.
3. Examples:
a. CO2: carbon dioxide
b. CO: carbon monoxide
c. SO3: sulfur trioxide
d. N2O5: dinitrogen pentoxide
e. CO2 (using Stock System): Carbon (IV) oxide
fff. Gram formula mass (GFM): the mass of 1 mole of compound; calculated by
adding the atomic masses of all atoms (including multiples) in the compound.
ggg.
Empirical formulas are formulas with the lowest possible ratio of elements
(subscripts); all ionic compounds (metal and a nonmetal) must be written as
empirical formula
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hhh.
Molecular formulas are multiples of empirical formulas; have different
physical and chemical properties from each other; determined by the ratio of gram
molecular masses.
Example:
If a compound has an empirical formula of CH2 and a molecular mass of 84g/mol, what is the
compound’s molecular formula? Answer 84/14 = 6. 14 is the mass of CH2
Therefore: 6(CH2) = C6H12
iii. Hydrates – a crystal consisting of a solid substance combined chemically with
water in a definite ratio. Example CuSO4 • 5H2O
jjj. Anhydrous – without water; usually applied to the product obtained after
removing the water from a hydrate.
kkk.
Percent mass composition of compounds for each element may be
calculated from their formula – the total atomic mass of the element divided by
the gram formula mass and then multiplied by 100% (Table T)
i. When calculating % water in a hydrate remember to include the mass of
the water in the total GFM
Example: What is the percent of water in CuSO4 • 5H2O?
Name of element
Copper
Sulfur
First oxygen
Hydrogen
Second oxygen
g/mol of element
64
x
32
x
16
x
1
x
16
x
Number of atoms
1
1
4
10
5
Total
=
=
=
=
=
Total
64
32
64
10
80
250 g/mol
Next – you must divide the part by the whole.
mass of 5 H2O
X 100%  ? %
Example (for 1mole of compound)
mass of CuSO 4  5H2O
90 g
x 100%  36%
250 g
Example: Using the experimental findings below, calculate the % mass water in the
hydrate tested.
Mass of crucible and cover = 79.463 g
Mass of crucible, cover, and hydrated salt = 94.362 g
Constant mass of crucible, cover, and anhydrous salt = 89.852 g
MHS Regents Chemistry
Answer:
Midterm Review 3rd Edition
Mas of the hydrate = (94.362) – (79.463) = 14.899 g
Mass of water in hydrate = (94.362) – (89.852) = 4.510 g
Page 16 of 23
lll. Empirical formulas may be determined from % composition of compounds by
i. Assuming a 100 gram sample – therefore all % compositions become
grams
ii. Change grams of each element into moles of each element (Table T)
iii. Normalize the number of moles by dividing each number of moles by the
smallest number of moles – these final numbers become the subscripts in
the empirical formula
iv. If a result in step iii is between 0.4-0.6 then multiply each answer by 2
Example:
What is the empirical formula of a compound made up of 75% carbon and 25% hydrogen
by mass? Treat % as grams.
Carbon 75g C
1 mol C
x ------------ = 6.25 mol C
12.0 g C
6.25 mol
----------6.25 mol
25 mol
----------6.25 mol
Hydrogen 25g H
1 mol H
x ---------- = 25 mol H
1.0 g H
divide by the smallest number
Answer is CH4
13) Mole Information
mmm.
1 mole = 1 gram formula mass (GFM) (gram molecular mass as well)
nnn.
1 mole = 6.02 x 1023 molecules of a compound, or atoms of an element;
this number can be thought of the same way we think of a dozen – there can be a
dozen of anything – it is just a special group size.
ooo.
1 mole = 22.4 L of a gas at STP (Table A)
ppp.
The number of moles in a sample may be determined by dividing the
given mass of the sample by the compound’s GFM (Table T)
qqq.
To convert between units of matter (above) the Factor Label Method may
be used
Example: What would be the volume (L) of 52 grams of N2 gas at STP?
Solution: (remember the GFM for N2 is 28 g/mole)
52 g N2
MHS Regents Chemistry
X
22.4 L
28 g
=
41.6 L N2
Midterm Review 3rd Edition
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Homework Assignment
Moles and Formulas
Regents Chemistry Review
1. Name the following compounds using the Stock System
a. BaF2 ______________________
b. Pb(NO3)2 _______________________
c. ZnCO3 _______________________
d. PbO2 ______________________
2. Name the following compounds using the Traditional Method
a. NO2 ___________________________
b. PCl5 __________________________
c. S2O3 __________________________
3. Write the formula for each of the following chemical names.
a. Aluminum hydroxide ______________________
b. Chromium (III) nitrate ______________________
c. Phosphorous trihydride ______________________
d. Manganese (IV) oxide ______________________
4. Using the compound, BaCO3 • 4H2O, answer the following questions.
a. What is the GFM of the compound?
b. How many moles of oxygen atoms are there in each mole of compound?
c. What is the percent mass composition of oxygen in the compound?
d. What is the percent water in the compound?
5. What is the empirical formula for a compound that contains 30.4% Nitrogen and 69.6%
Oxygen?
6. If the gram molecular mass of the molecule found in question 5 above was 92 g/mole,
what is its molecular formula?
7. What is the mass of 5.0 x 1024 molecules of O2?
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Midterm Review 3rd Edition
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14) Equations
rrr. Reactants are substances used in a chemical reaction and are on the left side of the
arrow
sss. Products are substances formed in a chemical reaction and are on the right side of
the arrow
ttt. Balancing a chemical equation: (Law of Conservation of Matter) the total number
of each type of atom must be the same on each side of the arrow (equation);
coefficients are placed in front of whole compounds in order to achieve
i. Choose to balance an element that needs a coefficient to be added first
ii. Balance polyatomic ions separately as whole ions if the polyatomic ions
did not change during the reaction… do not balance them as individual
atoms.
iii. Balance the element that occurs as a separate element alone in the
equation last.
iv. Always use only whole numbers to balance equations; if fraction is needed
– then when done multiply every coefficient by the denominator to
eliminate the fraction
Exercise: Balance the following equation.
___ Na3PO4 + ___ CaCl2 → ___ NaCl + ___ Ca3(PO4)2
uuu.
Coefficients in front of reactants and products in a balanced equation
represent moles of that substance used relative to the others in the reaction
vvv.
Types of reactions
i. Synthesis: N2 + 3H2  2NH3 (A + B  AB)
ii. Decomposition: 2KClO3  2KCl + 3O2 (AB  A + B)
iii. Single replacement: Cu + 2AgCl  CuCl2 + Ag (A + BC  AC + B)
1. Spontaneous if single element reactant is higher up on Table J than
the element being replaced. E.g. Sr can replace Ti in compounds;
Al can replace Cr in compounds; Fe CANNOT replace Mn in
Exercise: Complete/balance the following single replacement reactions that will happen.
AlCl3 + Zn →
Cu + HCl →
MHS Regents Chemistry
NaI + Cl2 →
Midterm Review 3rd Edition
Mg + CaCl2 →
Page 20 of 23
compounds.
iv. Double replacement: NaCl(aq) + AgNO3(aq)  AgCl(s) + NaNO3(aq)
(AB+CDAD+CB)
1. Spontaneous if a solid (precipitate), liquid compound, or gas is
formed as a product; use Table F to determine if a precipitate is
formed
2. All acid base neutralization reactions are double replacement
reactions. (Produces salt and liquid water as products.)
Exercise:
Are the following soluble in water? (Table F)
LiOH
MgCO3
NH4OH
AgI
BaSO4
Which of the following reactions will occur?
NaCl(aq) + KBr(aq) →
Pb(NO3)2(aq) + ZnCl2(aq) →
NaOH(aq) + H3PO4(aq) →
www.
Equation Stoichiometry: predicting amount of reactant needed or product
produced according to a balanced chemical equation.
Example:
How many moles of NH3 will be produced according to the following reaction if 3.75
moles of hydrogen are consumed?
N2 + 3 H2  2 NH3
a) Place 3.75 under the coefficient in front of hydrogen in the balanced equation
b) Place a “X” under the coefficient in front of NH3 in the balanced equation
1 N2 + 3 H2  2 NH3
3.75
X
c) Solve for “X” in the proportion: X= 2.5 moles
Exercise: How many moles of nitrogen gas are need to completely react with 3.75 moles of
hydrogen?
MHS Regents Chemistry
Midterm Review 3rd Edition
Page 21 of 23
MHS Regents Chemistry
Midterm Review 3rd Edition
Page 22 of 23
15) Matter
xxx.
Element: can’t be decomposed by chemical change. First letter is always
capitalized and if there is a second letter it is always lower case.
yyy.
Compound: 2 or more elements chemically combined that can be
decomposed by chemical change (not physical change). A chemical change
produces matter with a different composition than the original matter.
zzz.
A chemical property is the ability of a substance to undergo a specific
chemical change. A physical property can be observed or measured without
changing the substance’s composition.
aaaa.
A substance is either an element or compound and is the same throughout
the sample (homogeneous)
bbbb.
Mixtures are combinations of elements and/or compounds (mixtures of
substances)
i. May be homogeneous; samples taken anywhere in the mixture will contain
the same ratio of components (the same everywhere); all solutions (aq) are
homogeneous mixtures
ii. May be heterogeneous; samples differ when taken from different locations
in the mixture; e.g. vinegar and oil
iii. Mixtures may be separated by physical means using techniques such as
centrifugation, distillation (different b.p.), filtration (example solid/liquid
mixtures), magnets, etc.
cccc.
Extensive properties: Ex. mass, volume, shape (dependent on amount of
substance)
dddd.
Intensive properties: Ex. melting/boiling point, vapor pressure, density,
etc. (specific to substance and may be used to identify substance)
eeee.
Solids (s) have a definite shape and volume.
ffff.
Liquids (l) have definite volume but no definite shape
gggg.
Gases (g) have no definite shape and no definite volume
hhhh.
Aqueous (aq) means dissolved completely in water (no solid remains) (a
homogeneous mixture)
Exercise: Which
sample represents a
substance?
MHS Regents Chemistry
Midterm Review 3rd Edition
Page 23 of 23
16) Energy/Heat
iiii. Exothermic processes release heat as a product (A + B  C + 35kJ)
jjjj. Endothermic processes absorb heat just like a reactant (A + B + 35kJ  C)
kkkk.
All steps on a heating curve are endothermic; potential energy increases
during phase changes
llll. All steps on a cooling curve are exothermic; potential energy decreases during
phase changes
mmmm. Melting/freezing and vaporization/condensation are steps involving
Potential Energy changes (no Kinetic Energy changes);
i. Melting/freezing use: q = mHf
ii. Vaporization/condensation use: q = mHv respectively
iii. During a phase change both phases are in equilibrium with each other.
Ex. At 0oC water is freezing and ice is melting at the same rate.
nnnn.
During temperature change heating (or cooling) water – Average Kinetic
Energy (Temperature) is changing; use equation for heat exchange: q = mCT
oooo.
Use constants for water on Table B
pppp.
Joules is the unit of measurement for heat
qqqq.
One degree Celsius is equal to one Kelvin; the two scales just start at
different reference points and are 273 degrees shifted (Table T)
rrrr.
Sublimation: changing directly from a solid to a gas; substances that
sublime generally have very weak intermolecular forces between their molecules
and therefore very high vapor pressures.
i. CO2 and I2 are both solids that sublime at STP
ii. Sublimation is an endothermic process
ssss.
Deposition: changing directly from a gas to a solid
i. Deposition is an exothermic process
tttt. Heat always moves from HOT to COLD!
MHS Regents Chemistry
Midterm Review 3rd Edition
Page 24 of 23
Homework Assignment
Equations, Matter and Energy
1. Complete and balance the following equations:
a) ____ Al(NO3)3 + ___ Na2CO3 
b) ____ Li
+ ____ Al2(SO4)3 
c) ____ KClO3 
____ KCl + ____ O2
d) ____ N2 + ____ H2 + ____ Cl2  _____ NH2Cl
2. What type of equation is each of the above equations?
3. How can you determine if some chemical in a beaker (liquid) is a substance or a mixture?
4. If 15.0 grams of water are heated to 55oC when 1215 Joules of heat are added, what was the
initial temperature of the water?
5. How much energy is required to melt 25.0 grams of ice at 0oC?
6. When ammonium chloride crystals are dissolved in water, the temperature of the water
decreases. Does the temperature change indicate that the dissolving of ammonium chloride in
water is endothermic or exothermic?
MHS Regents Chemistry
Midterm Review 3rd Edition
Page 25 of 23
17) Gas Laws
uuuu.
When using temperature must always be in Kelvin!
vvvv.
Avogadro’s Hypothesis: Equal volumes of gases under the same
conditions of temperature and pressure contain the same number of molecules.
wwww. Boyle’s Law; PV=k (constant); P1V1 =P2V2 ; is non-linear, inverse
relationship
V
Pressure
xxxx.
Charles’ Law; V/T = k (constant); V1/T1 = V2/T2; linear, direct
relationship
V
Temp, K
yyyy.
Gay Lussac’s Law; P/T = k (constant); P1/T1 = P2/T2; linear, direct
relationship
P
Temp, K
zzzz.
Combined gas law P1V1/T1 = P2V2/T2 ; turns into one of the above laws
when either temperature, pressure or volume is held constant
Example:
If the temperature of a 2.0L sample of helium gas at STP is increased to 27C and
the pressure was decreased to 80.kPa, then the new volume of the helium gas would
be?
(101.3kPa) x (2.0L)
(80.kPa) x (V2)
------------------------ = --------------------273K
300K
Answer is 2.8L
Remember to remove any variable from both sides of the equation , that according
to the question, is held constant (or can be assumed that it is held constant).
MHS Regents Chemistry
Midterm Review 3rd Edition
Page 26 of 23
aaaaa.
Kinetic Molecular Theory(KMT) for an Ideal Gas – used to explain the
behavior of gases. It describes the relationships of pressure, volume, and
temperature with velocity, frequency and forces of gas molecule collisions.
i. Gas molecules are separated by great distances relative to their size,
therefore the volume of the gas molecules themselves is negligible (next to
nothing)
ii. Gas molecules move in constant, random, straight-line motion
iii. Gas molecules have no attraction for each other
iv. Gas molecules collide and may lead to the transfer of energy between gas
molecules, but the collisions are referred to as “perfectly elastic” meaning no loss of energy.
v. In order for a gas to be truly “ideal” it must adhere to all of the above
requirements. Therefore there are no truly ideal gases. All real gases
deviate from these ideal requirements. Gas molecules have attractive
forces between them and do have some volume. Real gases act more like
ideal gases under certain conditions. Those conditions are high
temperature (kinetic energy), low pressure (concentration), and small
molecular size (volume).
bbbbb.
Ideal gases
i. Take up no volume
ii. Have perfectly elastic collisions, travel in constant, random, straight lines,
and have no attraction between gas molecules
iii. High temperature and low pressure best approximate the conditions 1 & 2
above for any real gases; smaller molecules will also approximate “ideal
conditions” better.
ccccc.
Density (grams/Liter) of an ideal gas at STP = GFM/22.4L
ddddd.
Vapor Pressure of Liquids (Table H)
i. Surface liquid molecules evaporate at all liquid temperatures
ii. Increased temperature means liquid molecules are moving faster and more
molecules will be able to evaporate; vapor pressure increases
iii. Vapor pressure: the pressure of the liquid vapor molecules above a liquid
at a specific temperature in a closed container.
iv. Liquids with strong intermolecular forces, high Hv, will evaporate more
slowly and have a lower vapor pressure at all temperatures compared to
liquids with weaker intermolecular forces, lower Hv
MHS Regents Chemistry
Midterm Review 3rd Edition
Page 27 of 23
eeeee. Boiling Point: The temperature at which the vapor pressure of a liquid is equal to the
external pressure and liquid vaporizes under its surface.
fffff.
Normal Boiling Point: The boiling point of a liquid substance at normal atmospheric
pressure (standard pressure), which is 101.3 kPa, 1 atm, 760 mm Hg, or 760 torr.
Exercises:
Answer questions 1 and 2 using the information given below.
Four identical balloons contain equal volumes of gas at STP.
Balloon #1 contains H2 gas
Balloon #2 contains He gas
Balloon #3 contains O2 gas
Balloon #4 contains N2 gas
1. Which balloon, if any, would weigh the most? Explain your answer by describing what is
equal and/or different in each balloon.
2. According the theory of gas molecule movement, Kinetic Molecular Theory, why would
the balloons expand if they were heated?
3. A gas has a volume of 1,400 milliliters at a temperature of 293K and a pressure of
1.0atm. What will be the new volume when the temperature is changed to 323K and the
pressure is changed to 0.80 atm? (Show all work)
4. If the pressure on 36.0 milliliters of a gas at STP is changed to a pressure of 35.3 kPa at
constant temperature, what will the new volume be?
5. What must the atmospheric pressure be if water is boiling at 55oC?
MHS Regents Chemistry
Midterm Review 3rd Edition
Page 28 of 23
6. What is the vapor pressure of ethanol at the normal boiling point of propanone?
MHS Regents Chemistry
Midterm Review 3rd Edition
Page 29 of 23