CP Chemistry Review
Periods 2, 4, and 7
Physical Properties and Changes
● Physical properties are things you can notice about a substance
without affecting its molecular structure.
● Physical change rearranges molecules but doesn't affect their internal
structures.
o whipping egg whites, boiling water, dissolving sugar in water and
dicing potatoes
● A change that can be reversed
o ex. freezing water then melting it again
● It also does not change the substance
o ex. cutting hair
Chemical Properties and Changes
● A chemical property is something you notice about a substance only
when it undergoes a chemical change.
● Chemical change is any change that results in the formation of new
chemical substances
o milk souring, iron rusting, eggs cooking, and break rising
● These changes are Irreversible
o ex. after you bake bread you can’t get the dough back
Elements
Definition: A substance that can not be broken
up, or broken down into smaller pieces using
chemical or physical means
Compounds
Definition: a substance that is made with two or
more elements bonded together, can’t be
separated by any means.
Mixtures
Definition: A combination of substances that are not
bonded together
● Homogeneous: looks the same throughout the mixture
ex. sugar water
● Alloys: metallic material consisting of two or more elements,
cannot be separated by any physical means
● Heterogeneous: able to see different parts of the mixture are visible
ex. cereal with milk
Accuracy and Precision
Accuracy is having a true number.
(agrees with expected value)
Precision is having a reproducible number.
(if someone else measures they get same #’s and # of decimal places)
Your data can be accurate, precise,
accurate and precise, or
neither accurate nor precise.
Density
the degree of compactness of a substance
D= M/V
•Density=Mass(g)/Volume(mL or cm3)
•Your unit will be g/(mL or cm3)
Example:
A pair of sunglasses weighs 176g. They have a
volume of 5mL. What is the density of the
sunglasses.
D=M/V
D=176g/5mL
D=35.2 g/mL
Qualitative and Quantitative
Measurements that are #’s are QUANTitative.
Measurements that are words are QUALitative.
QUANTITATIVE
QUALITATIVE
0.25 L
273 K
8.83 cm
101.33 kPa
12.1 g
hot
bubbly
slow
green
solid
Scientific Notation
Scientific notation is the way that scientists easily handle very
large numbers or very small numbers.
•For example, instead of writing 0.0000000056, we write 5.6 x 10-9.
1.
6
Percent Error
Tells how wrong a measurement is.
% error = measured value - accepted value x 100
accepted value
Let’s say you measure something to be 10 cm. Let’s say the actual vaule is supposed to be 12 cm. to find your
percent error, do
10 (measured) - 12 (accepted) x 100 = -16.67% error………negative percent error means the value was lower
12 (accepted again)
than you expected it to be!
Counting Protons +,
Neutrons, and ElectronsParticles
Charge
Mass
Proton
+
1
Neutron
0
1
Electron
-
0
Example
Mass = Protons +
Neutrons
12
6
Protons =
Electrons
C
★ The atomic number is equal to
the number of protons
★ The number of electrons is
equal to the number of protons
★ The mass number is equal to
the number of protons +
neutrons
★ To find Neutrons, subtract
protons from the mass number
Radiation Particles and Equations
Alpha-stopped by clothing or skin
Beta-stopped by thin sheets of lead
Gamma-stopped by thick sheets of lead
Nuclear Reactions and Equations
Nuclear Reactions- must obey the law of
conservation of matter!
Alpha Decay195
77
He
191
75
Isotopes
-atoms of the same element that have different
numbers of neutrons.
222
98
298
98
start here
Electron Configurations
Rules:
1. Electrons fill into low energy orbitals before the high energy orbitals.
2. Electrons in the same orbital have opposite spins.
3. Electrons prefer to be alone in an orbital, rather than having a pair
Method 1:
● write out the orbitals for the element using the P, S, D and F orbitals
S Orbital: 2eP Orbitals: 6eD Orbitals: 10eF Orbitals: 14eEx.) Ar (18e-) 1s^2 2s^2 2p^6 3s^2 3p^6
Electron Configurations
Method 2:
● draw the boxes for each S, P, D, or F orbitals
● the number of boxes depends on the number of orbitals the element has
S Orbital: 1 orbital
P Orbital: 3 orbitals
D Orbital: 5 orbitals
F Orbital: 7 orbitals
Ex.) S (16e-)
Electron configurations
Method 3: Noble Gas Notation
1. Find the Noble Gas that comes before the element you are dealing with
2. Write the Noble Gas in brackets
3. Count on the diagonal rule to find the orbitals filled up by the noble gas
4. Next to brackets, start on the next orbital and fill in the remaining electrons
Ex.)
Periodic Table Set up/ Features
●
●
●
●
Families in the periodic table are the vertical columns and are also referred to as groups and
rows.
In the periodic table of the elements, elements are arranged in a series of rows (or periods) so
that those with similar properties appear in a column.
Valence Electrons: Counting the families except transition metals you go up by one.
Charges: counting the families but skipping transition metals by adding one all the way to four
then subtracting by one.
ion charges:
Valence Electrons:
Electronegativity, Ionization Energy, and Atomic Radius
•Ionization energy describes the minimum amount
of energy required to remove an electron from the
atom or molecule in the gaseous state.
Ions
● How are ions formed?
o
Positive ions have lost electrons (cations)
o
Negative ions have gained electrons (anions)
● All atoms become stable when they have an octet of valence
electrons.
o
How many valence electrons does sodium have?
o
1s²2s²2p63s1
o
1 valence electron so no octet and not stable
Ionic vs Covalent Properties
Ionic
● made of ions→ one cation
(metal), one (non-metal)
● trade e● crystal lattice
● bonds are strong
● melting points are high
● hard, brittle, rigid→ (not
bending)
● excellent conductors
● ionic substance dissociate in
Covalent
● made of atoms → 2 non-metals,
no charges
● share e● molecules
● bond strength varies
● melting points are low →
average
● softer, more flexible
● poor conductors
● covalents do not dissociate in
water (stay together)
Naming Ionic Compounds
1. Write the metal down (cation)
2. Then, write the name of the non-metal (anion)
3. After the non-metal, put a “-ide” ending
Naming Ionic Compounds
with
Transition
Metals
1. Write the formula of the ionic compound
2. Find the charge of the metal, make sure to write the
charge with a roman numerals in between the two
names.
3. Write the name of the metal
4. Put the first and second name together
Naming Ionic Compounds
with Polyatomics
1. Write the formula for the polyatomic ion.
2. Find the charge of the metal.
3. Write the name of the metal,then write the name of the
nonmetal ions
4. Combine the two names
How to know if you are looking at a covalent
molecule or an ionic molecule
● Ionic molecules gain or lose electrons while covalent
molecules share electrons.
● H20 is a covalent bond because it shares electrons
(both substances are non-metals).
● FeO is an ionic molecule because it contains a metal,
which is a positive ion, and a non-metal negative ion,
which trade electrons.
Naming Covalent Molecules
1st Step: Name the 1st Element Listed
● Put a prefix in front to tell if there is more than
one of that atom.
● Di-2 tri-3 Tetra-4 Penta-5 Hexa-6 Hepta-7
Octa-8 Nona-9 Deca-10
2nd Step: Name the Second element
● Change to -ide ending
● Use a prefix to tell how many.
● If only one of the atom, use mono-1
Examples
● PS - Phosphorus monosulfide
● PS2 - Phosphorus disulfide
● P2S - Diphosphorus monosulfide
Naming Acids
● Look for ending
o “--ide”
o “--ate”
o “--ite”
Hydro___ic acid
___ic acid
___ous acid
Lewis Dot Diagrams
-Method for drawing diagrams that show electrons in covalent molecules
1. Add valence electrons for all atoms in the molecule
2. Arrange atoms with the least amount of electronegativity and the most
amount of symmetry atom in the center
3. Draw your shared pairs between each atom. Subtract bonded electrons
for total
4. Draw unshared pairs around each atom
Single Bond ----> 1 shared pair
Double Bond ---> 2 shared pairs
Triple Bond ------> 3 shared pairs
*Hydrogen only needs 2 electrons to make it stable*
Resonance happens when you can draw more than the picture showing
double/triple bonds
Lewis Dot Diagram Example
Center
Atom
Double bond or 2 shared pairs
Single bond or one shared pair
Unshared Pairs
are the dots.
Polarity
Polar Molecules do NOT share electrons
Non-Polar molecules share evenly
Electronegativity tells how much an atom wants an electron
For a bond: Compare E.N of atoms if same or really close its non-polar. If
there is a big difference then it IS polar
For a Molecule: First is there a lone pair on the center atom?
YES
POLAR
NO
YES
NON-POLAR
are the atoms attached the
same
NO
POLAR
VSEPR Theory
This theory states that electron pairs repel each other whether or not they
are in bond pairs or in lone pairs. Thus, electron pairs will spread
themselves as far from each other as possible to minimize repulsion.
Molecular shape is determined by the center atom, not the whole
molecule.
Look at the center atom
of a lewis dot diagram.
● How many lone
pairs are on it?
● How many
bonded atom are
on it (constituent
groups)?
Predicting Products of Reactions
Once you identify the type of reaction, you can
predict the products of the reaction.
For DR, SR, and Synthesis:
Check charges of the substances you pair up!
For Combustion:
ALWAYS write carbon dioxide and water.
For Decomposition:
Break it up into the elements.
Empirical Formulas
The empirical formula is the simplest
formula for a compound. The
empirical formula is the simplest
formula for a compound. An empirical
formula is often calculated from
percent composition data. The weight
percentage of each of the elements
present in the compound is given by
this percent composition
Molecular Formulas
A molecular formula is the same as or
a multiple of the empirical formula, and
is based on the actual number of
atoms of each type in the compound..
For example, if the empirical formula
of a compound is C3H8 , its molecular
formula may be C3H8 , C6H16 , etc.
The ratio of C:H:O has been found to be 1:2:1, thus the empirical formula is: CH2O. Suppose we know that the molecular
weight of this compound is 180 g/mol. The formula weight of the empirical formula is 30 g/mol. Divide the molecular weight
by the empirical formula weight to find a multiple:
The molecular formula is a multiple of 6 times the empirical formula:
C(1 x 6) H(2 x 6) O(1 x 6) which becomes C6H12O6
Percent Composition
Finding any percent requires a format of
% = (part / whole)x 100
For percent of an element in a molecule, find the molar mass,
then use the mass of the element and the molecule mass to find
percent.
Remember, if it says “find the percent composition” and does not
specify a specific element, you must find the percent of every
element in the molecule!
Mole Ratio
A mole ratio is ratio between the amounts in moles of any two compounds
involved in a chemical reaction.
Mole ratios are the central step in performing stoichiometry because they allow us to convert
moles of one substance to moles of another substance
To determine the mole ratio between two substances all you
need to do is look at balanced equation for the coefficients in
front of the substances you are interested in.
Made by: Samantha Yaccarino & Dan Yu
When to Use Stoichiometry
● Stoichiometry should be used when trying to find the
amount of molecules or moles for a DIFFERENT substance
than you have information about.
● To identify a Stoichiometry problem, look for questions with
a similar format to: “How many _____are in a sample of X
that contains this much Y.”
Tips for solving stoichiometry problems:
1. Convert from grams, liters, or particles to
moles.
1. Use a mole ratio.
Always be
careful to read
the problem! If
moles are given,
don’t do step 1.
1. Convert back to grams, liters or particles!
Limiting Reactants (LR)
● The substance that is
used completely in the
reaction
Excess reactants (XS)
● Left over at the end of
the reaction
How to select the limiting and excess reactants
1) Compare the amounts of each reactant in
units of moles (this is how much you have)
2) Use a mole ratio on only one of the
reactants to compare same amounts of
moles (this is how much you need)
3) Compare the HAVE vs. the NEED values
(a)
(b)
have more than you need (XS) excess
need more than you have (LR) limiting
Percent Yield
If you have done stoichiometry to find the mass of a product, you have
calculated a THEORETICAL YIELD.
You can use the theoretical yield to compare how much product you actually
collect in a lab to see how effective you were. The measured amount of
product is called the EXPERIMENTAL YIELD.
Percent Yield =
THEORETICAL YIELD * 100
EXPERIMENTAL YIELD
Parts Of Solutions
● Two Parts of Every Solution
o
o
Solute- things that dissolve (disappears!)
Solvent- does the dissolving (surrounds the solute)
● Concentrated solutions contain a LOT of
solute
● Diluted Solutions have less solute per
volume
Molarity
● To quantify concentration, use molarity (M)
Intermolecular Forces
What is a Dipole?
● Most often found in polar molecules, in which the
electrons are unevenly shared
● This uneven sharing gives one side of the molecule a
partially positive charge (δ+) and the other side a
partially negative charge (δ-)
What is Hydrogen Bonding?
● Hydrogen bonding is an especially strong
form of dipole-dipole interaction
● When a hydrogen atom is covalently bonded
to nitrogen, oxygen, or fluorine, a very strong
dipole known as Hydrogen bonding is
formed
What is Dispersion?
● Dispersion forces, also known as van der
Waals forces, help neutral atoms and
nonpolar molecules attract each other
● The dipole-dipole interactions among
momentary dipoles are known as dispersion
forces
Phase Diagrams
Triple point
Critical Point
Melting/Freezing Points
Boiling Points
Sublimation Points
Temperature and Pressure
Temperatures can be converted: Kelvin = oC+ 273
A temperature of 0 Kelvin (-273oC) is called absolute zero.
Standard temperature and pressure is 273 Kelvin and 1 atm.
Pressure can be converted using conversion factors. These
factors will be given on the final exam reference page.
Boyle’s Law (Pressure and Volume)
Boyle’s Law states that the volume of a given amount of gas held at constant temperature varies inversely with the applied
pressure when the temperature and mass are constant.
●
As pressure decreases, volume increases
●
Equation: P1V1=P2V2
Ex.) A balloon contains 30.0L of helium gas at 103kPa. What is the volume of the helium when the balloon rises to an altitude
where the pressure is only 25.0kPa?
V1 = 30L
P1 = 103kPa
V2 = ?
P2 = 25kPa
(103kPa)(30L) = (25kPa)(V2)
3090 = 25(V2)
V2 = 123.6L
Charles's Law (Temperature and Volume)
This law states that the volume of a given amount of gas held at constant pressure is directly proportional to the Kelvin
temperature.
● AS volume goes up, Temperature goes up
● Equation: V1/T1=V2/T2
Ex.) If a sample of gas occupies 6.80L at 325C , what will its volume be at 25C (3.39L)
V1 = 6.80L V2 = ?
T1 = 325C T2 = 25C
6.80L/325C = V2/25C
170 = 325(V2)
V2 = .52L
Gay-Lussac’s Law (Pressure and Temperature)
●
●
●
Gives the relationship between pressure and temperature when volume and amount are held constant.
As temperature increases, pressure increases
Equation: P1/T1 = P2/T2
Ex.) The gas in a used aerosol can is at a pressure of 103kPa at 25C. If the can is thrown into a fire, what will the
pressure be when the temperature reaches 928C?
P1 = 103kPa P2 = ?
T1 = 25C T2 = 928C
103kPa/25C = (P2)/928C
95584 = 25(P2)
P2 = 3823kPa
Collision Theory
Occurs when molecules crash into with powerful Kinetic
energy. This theory is only dealing with gas-phase
chemical reactions. we’re assuming:
1. All molecules are going through space in a straight line
2. All molecules are rigid spheres.
3. The reactions are between only between two molecules
4. The molecules need to collide with each other
Effusion and Diffusion
Diffusion: The process of a gas substance spreading out
evenly to fill its environment.
Effusion: The process of gas molecules escaping from a
small hole in the environment.
There is a Graham's Law of Effusion which states that the
more massive the gas, the slower it effused. For example,
Helium would release form the hole faster than Oxygen
would release out the same size hole and environment.