Chemistry 1 Study Guides
Study guide for Intro to Matter
1. USE YOUR REFERENCE TABLES – Table D (units), Table C (metric prefixes), Table T
(formulas for calculations)
2. For significant figures, all non-zero numbers (1-9) are significant.
a. Any 0’s in between non-zero numbers are significant. E.g. 2007
b. In numbers bigger than 1 if there is no decimal point, the 0s are NOT significant. E.g.
45700 has 3 sig figs
c. In numbers smaller than 1, the 0s closest to the decimal point are NOT significant
d. When doing calculations, the answer should be written to the fewest significant figures.
E.g. 24.6 x 2.9836 (answer should have 3)
3. Density is found by Density=Mass/Volume
a. V= M/D
b. M = DV
4. Percent error
5. Matter is made up of substances (fixed ratio/ formula) or a mixture of substances (variable
ratios).
6. Pure substances are either elements or compounds and are all homogenous.
a. Each element has a unique symbol. Each symbol’s 1st letter is capital, all others are
lower case. Elements CANNOT be broken down/ decomposed by chemical change.
b. Compounds are made of more than element so have more than one symbol in the
formula.
7. Mixtures do not have uniform proportions so have no molecular formula.
a. They can be homogenous (all samples the same all the way through)
i. Solutions are homogenous mixtures
b. Or Heterogeneous (samples not necessarily the same all the way through)
c. They can be separated by physical means including distillation (based on boiling point),
filtration (based on solubility)
8. There are 3 states of matter
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Solids
Liquids
Gases
Particles in fixed
positions (crystals)
Particles vibrate
Particles change positions
relatively easy
Particles flow over and around
each other
Definite volume
Takes shape of container
Particles are very random
Definite volume
Definite shape
Particles move very quickly and
bump into each other often
Takes volume of container
Takes shape of container
9. Physical properties of matter include color, luster, odor, melting point, boiling point, density
and solubility (ability to dissolve).
10. Chemical properties of matter indicate how they react or combine with other substances
including corrosion, decomposition.
11. Physical changes are easily reversed and the compound/ element does not change so its
symbol, formula and properties do NOT change.
a. Phase changes are physical changes including sublimation (from solid straight to gas),
evaporation, melting, freezing, boiling, condensation, crystallization
12. Chemical changes are NOT easily reversed and create a new compound with different
formulas and properties.
13. Temperature conversion among Kelvin, degrees Celsius and degrees Fahrenheit
Conversion
Formula
Example
Celsius to Kelvin
K = C + 273
21oC = 294 K
Kelvin to Celsius
C = K - 273
313 K = 40 oC
Fahrenheit to Celsius
C = (F - 32) x 5/9
89 oF = 31.7 oC
Celsius to Fahrenheit
F = (C x 9/5) + 32
50 oC = 122 oF
14. Summary of Particle Diagrams
Type of Diagram
a. Element
b. Compound
c. Mixture
d. Diatomic
How do we know
Only one type of shape
More than one shape/color with atoms touching
More than one shape/ color with different elements/ compounds NOT touching
each other. Each pattern NOT the same
2 of the same type of shape/color, touching
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Study Guide for Atomic concepts
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6.
7.
Most info is found on the periodic table. USE THE PERIODIC TABLE
There are 3 subatomic particles
a. Protons – positively charged, found in nucleus, mass of 1 amu. Makes the nucleus positively
charged
b. Neutrons – no charge, found in nucleus, approx. mass of 1 amu
c. Electrons – negatively charged, located in orbitals outside nucleus,
mass is 1/2000 amu
Electrons were discovered first using cathode ray tubes.
Rutherford’s Gold foil experiment led to discovery that atoms is mostly empty
space and the alpha particles were deflected because of the positively charged
nucleus.
Mass number (atomic mass) is found in the upper left corner and can be calculated by
adding the number of protons and neutrons.
a. Number of protons can be found by subtracting the number of neutrons from the mass number
b. Number of neutrons can be found by subtracting the number of protons from the mass number
Atomic number is the number of protons.
Sequence of historical development is hard-sphere model, mostly empty
space, electron-shell model (electrons exist in orbitals) then the wavemechanical model (electron cloud)
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The orbital in the electron-cloud model is the region where electrons are
most likely to be found.
9. A sample composed of atoms having the same atomic number is classified
as an element.
10. Nuclear charge increases as the atomic number (number of protons)
increase.
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11. Atoms of different isotopes of the same element differ in their total number of neutrons.
12. To calculate the atomic mass:
a. Find the mass with its relative concentration.
b. Convert it to a decimal
c. Multiply the decimal by the isotope mass
d. Add the results from step c.
13. Atomic masses are a weighted average of the naturally occurring isotopes of an element.
14. Electron configuration in the ground state can be found in the lower left corner on the periodic table.
a. The first shell is full with 2 (Formula: 2n2)
b. The second shell is full with 8
c. The third shell is full with 18
d. The fourth shell is full with 32
15. Valence electrons are the outermost electrons and are the last number in the
electron configuration.
16. Atoms absorb energy and electrons move up to a higher electron shell.
a. The further the electron shell is from the nucleus, the more energy the
electron has.
17. When atoms in the excited state release energy and move to lower energy states, light is emitted.
a. Read a bright-line spectra like you would an electrophoresis run: matching bands mean the
same substance
18. Metals are found on the lower left of the periodic table – all are solids at room temperature except Hg
a. Nonmetals are found in the upper right – exist in all 3 states at room temperature
b. Noble gases are found in group 18.
c. The metalloids are found along the step – B, Si, Ge, As, Sb, Te, and show a mixture of metallic
and non-metallic properties.
19. Lewis electron-dot diagrams – the symbol of the element is surrounded by as many dots as valence
electrons. (in other words – 2 valence electrons means 2 dots)
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Study Guide for Periodic Table
1. Most info is found on the reference table. USE THE PERIODIC TABLE and TABLE S
2. Gold foil experiment led to discovery that atoms is mostly empty space and the alpha particles
were deflected because of the positively charged nucleus.
3. Mass number (atomic mass) is found in the upper left corner and can be calculated by adding the
number of protons and neutrons.
a.
Number of neutrons can be found by subtracting the number of protons from the mass
number
4. Atomic number is the number of protons
5. Sequence of historical development is hard-sphere model, mostly empty space, electron-shell
model (electrons exist in orbitals) then the wave-mechanical model (electron cloud)
6. The orbital in the electron-cloud model is the region where electrons are most likely to be found.
7. Nuclear charge increases as the atomic number (number of protons) increase.
8. Atomic masses are a weighted average of the naturally occurring isotopes of an element.
9. Electron configuration in the ground state can be found in the lower left corner on the periodic
table.
10. Valence electrons are the outermost electrons and are the last number in the electron
configuration.
11. Oxidation state is found on the upper right corner of the periodic table. Groups tend to have similar
ox #s, properties and reactivity
12. Atoms absorb energy and electrons move up to a higher electron shell when they become excited.
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13. Metals are found on the lower left of the periodic table – all are solids at room temperature except
Hg; form positive ions when they lose electrons; ionic radii are smaller than atomic radii
a. Nonmetals are found in the upper right – exist in all 3 states at room temperature; form
negative ions when they with larger ionic radii than atomic radii
b. Noble gases are found in group 18 and are unreactive.
c. The metalloids are found along the step – B, Si, Ge, As, Sb, Te, Po and show a mixture of
metallic and non-metallic properties.
14. Lewis electron-dot diagrams – the symbol of the element is surrounded by as many dots as valence
electrons. (in other words – 2 valence electrons means 2 dots)
15. Reactivity and electronegativity increase towards F, across the period and up the group.
16. Atomic radii increase away from F, right to left and down the group
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Study Guide for Nuclear Chemistry
1) Start with Table N, O and T to solve problems for radioactive decay
2) Start with Table E, S and periodic table to solve problems for compounds and bonding
3) Half-life is how long it takes half of the sample to decay
4) Substances will remain unchanged (stop decaying) only when it has reached a stable
isotope
5) Isotopes have the same atomic number and different mass number (atomic mass)
a. Atomic number is found on the lower left side in the periodic table
b. Mass number is found in the upper left side in the periodic table
6) To determine the original mass of a sample,
a. Identify what you have been given: time passed, fraction left/ amount left and halflife (using reference table if information is not given directly in the problem)
b. Calculate how many half-lives must have passed (time passed/ half- life)
7) Some uses for radioactive material include:
a. C-14helps us to date once living organisms
b. Co-60 and I-131 are used in medical procedures (to treat cancer and thyroid
conditions respectively)
8) Nuclear emissions are summarized in Table O with gamma rays having neither mass nor
charge as they are pure energy
9) Fission happens when heavier nuclei split while fusion happens when lighter nuclei combine
to form a heavier one (usually H combining to form He)
10) More energy is released during a fission reaction than an ordinary reaction because in a
fission reaction mass is converted to energy (gamma rays)
Study Guide for Calorimetry and Gas Laws
1. Kinetic energy (KE) is energy of motion and can be measured by the temperature i.e. increase in
temperature = increase in kinetic energy.
2. Potential energy (PE) depends on position of substance or bonds;
a. PE decreases as energy is removed during exothermic phase changes (as energy is removed);
b. PE increases during endothermic phase changes (as energy is added).
3. Endothermic reactions (require energy and so will take thermal energy from its surroundings) and
cause those surroundings to decrease in temperature.
a. Melting, evaporation and sublimation are endothermic physical changes so the molecules
have more energy at the end of these changes.
4. Exothermic reactions (release energy and so give thermal energy to its surroundings) and cause those
surroundings to increase in temperature.
a. Freezing, condensation, and deposition are exothermic physical changes so the molecules have
less energy at the end of these changes.
5. STP – standard temperature and pressure. Values found in table A of Reference table
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6. Heating Curves show us how the energy of the substance changes over time with constant heat added.
7. While phase changes are happening equilibrium (balance) exists between at least 2 phases.
8. The energy required to change a unit mass of a solid to a liquid at constant temperature is called the
heat of fusion.
9. The heat of vaporization is the energy required to change a unit mass of a liquid to a gas at a constant
temperature.
a. The heat of vaporization and the heat of fusion and can be seen in the horizontal segments on a
heating curve.
Heat, when given mass, specific heat capacity and change in temperature
Mass, when given Heat, specific heat capacity and change in temperature
Specific heat capacity, when given mass, Heat and change in temperature
Change in temperature, when given mass, specific heat capacity and Heat
Heat, when given mass and heat of fusion
Mass, when given heat and heat of fusion
Heat of fusion, when given mass and heat
Heat, when given mass and heat of vaporization
Mass, when given heat and heat of vaporization
Heat of vaporization, when given heat and mass
q=m.c.ΔT
m= q/c.ΔT
c=q/m.ΔT
ΔT = q/m.c
q=m.Hf
m=q/ Hf
Hf = q/m
q= m.Hv
m= q/Hv
Hv= q/ m
10. To find (use Table B when possible)
11. Kinetic molecular theory describes the behavior of an ideal gas:
b. Gases are made molecules that are extremely tiny and far apart from each other. This is why
gases can be compressed.
c. Gas particles move in a constant straight-line motion.
d. Any collisions a gas particle makes will be elastic that is, there no loss of energy. Each time a gas
molecule collides against an obstacle, it bounces off with the same speed that it hit with.
e. There are no intermolecular attractive force between gas particles
f. The average speed of the particles is directly proportional to the kelvin (absolute) temperature.
(the gas molecules move faster when it’s hotter)
12. Ideal gas behavior: Under conditions of HIGH temperature and LOW pressure, molecules behave more
like an "ideal" gas.
g. The smaller the gas molecule is, the more ideal it will behave, therefore,
h. Hydrogen (H) and helium (He) are the gases that behave the most ideal.
i. Ideal gas law: PV = nRT
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13. The strongest intermolecular forces of attraction exist in liquids with low vapor pressure (so have
higher heats of vaporization and boiling points).
14. ‘normal’ boiling point is @101.3kPa on Table H
15. As atmospheric pressure increases, temperature also increases but volume decreases.
j. As temperature goes up volume goes up (molecules are able to move farther apart)
k. As temperature goes up pressure goes up (more collisions with container)
16. Avogadro’s constant: equal numbers of particles of a gas will occupy equal volume under the same
conditions of temperature and pressure. (when V1=V2 and P1=P2 and T1 =T2)
17. To perform Combined Gas Law calculations:
a. Read through the problem to determine what they are asking for.
b. Identify what you have. Write down all the variables and fill them in as you find them.
c. Write the gas law equation down. Actually write it right there on the page. In the blank spot
(variable).
d. Rearrange the equation to solve for the missing value.
e. Substitute the numbers into the equation.
f. Cancel as many units as possible.
g. Plug the numbers into the calculator. Put all the numerators in (in parentheses) then divide by
denominator (in parentheses).
h. Round your answer off. Since this is all multiplication and division, round your answer off to the
same number of sig figs as the data you put in that had the fewest number of sig figs!
18. Combined Gas Law Equation P1V1/T1 = P2V2/ T2
P1= P2V2T1/ V1T2
P2 = P1V1T2/ V2T1
V1= P2V2T1/ P1T2
V2 =P1V1T2/ P2T1
T1= P1V1T2/ P2V2
T2 =P2V2T1/ P1V1
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Study guide for 6 (Bonding)
1) Formula mass is the total of all the atomic masses in the formula
2) Co-efficient are the numbers at the front of a formula, while subscripts are found to the
lower right of the element within the formula
3) The number of atoms in the formula is determined by multiplying the subscripts of each
element then adding each. If there is no subscript there is only one atom of that element.
4) Both mass and charge are conserved during a chemical reaction
5) To balance an equation, you add co-effecients. DO NOT CHANGE THE SUBSRIPTS!!!!
6) There are 3 types of bonds/ compounds as electrons are either gained, lost or shared:
a. Metallic bonding occurs between 2 metals (even the same element if it is bonding
with itself)
b. Covalent bonding where 2 nonmetals are sharing valence electrons
c. Ionic bonding where the nonmetal is taking valence electrons from the metal
7) To determine which type of bonding is happening, look at the elements:
a. If they are both metals, metallic bonding is happening
b. If they are both nonmetals, they are sharing electrons in a covalent bond
c. It one is a metal and the other is a nonmetal there is ionic bonding (the nonmetal has
taken electrons creating charged particles or ions)
8) Metals lose electrons so their cations (positive ions) have fewer electrons and smaller atomic
radii compared with the atoms
9) Nonmetals gain electrons so their anions (negative ions) have more electrons and larger
atomic radii compared with their atoms.
10) Hydrates have water added. To calculate the percent composition of water in a hydrate:
a. Determine the mass of the original sample
b. Find out how much water is in that sample (original mass – salt)
c.
Divide the mass of the water by the original sample then multiply result by 100
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