Chemistry Post-Test
1.
Study Guide
Fall 2014
Atomic Structure
 Structure of the Atom
o There are 3 Subatomic Particles
Protons (+ charge) and Neutrons (neutral or no charge) make up the nucleus.
Electrons (- charge) orbit the nucleus in energy levels.
Protons and neutrons are about equal in mass. Electrons are so small that we do
not count their mass.
 Counting Subatomic Particles
o Atomic Number = # of protons
o # of protons = # of electrons in a neutral atom
o To find the number of neutrons:
Atomic Mass = # of protons + # of neutrons
# of neutrons = atomic mass - # of protons
(Hint: Mass number is always the larger number)
 Isotopes and Average Atomic Mass
o Isotopes of the same element have the same number of protons (iso = same)
but a different # of neutrons.
o To find the average atomic mass:
a.
b.
c.
Multiply the mass # x the percent. (Change % to decimal)
Repeat for all other isotopes.
Add the answers to get the average atomic mass.
 Energy and Light
o Light is a wave and a particle.
 Einstein proved, with Planck’s equation and the Photoelectric Effect, that
light is a particle.
 Double Slit experiment proves light is a wave.
o Frequency and wavelength have an inverse relationship
 Example: if frequency doubles, wavelength is cut in half
o Frequency and energy have a direct relationship
 Example: if frequency triples, energy triples
 Example: colors with high frequency also have high energy
o Photoelectric Effect – Electrons bounce off when light strikes a metal.
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2.
Electron Configurations and the Periodic Table
 Rutherford’s Model of the Atom
o protons (+) and neutrons (neutral or no charge) make up the nucleus
o electrons (-) surround the nucleus in a vast empty space
 Bohr’s Model of the Atom
o electrons orbit around the nucleus in energy levels (n)
o When all electrons are in the lowest possible energy level
it is called the ground state
o When an electron gains energy and moves to a higher
than normal energy level it is called an excited state
 Emission Spectra/Color Line Spectrum
o unique for each element, like a fingerprint
o electrons can move from one energy level to another
 electrons gain(absorb) energy to “jump” away from the nucleus
 electrons lose(emit) energy to “fall” back towards the nucleus
o the farther an electron is from the nucleus, the more energy is given off when it falls
back to the ground state
o colors are seen when an electron falls from an excited state (high to low)
o the # of electrons corresponds to the visible wavelength emission spectra
 (the more electrons present, the more colors you can see)
 de Broglie’s Hypothesis
o electrons act like waves and therefore don’t fall into the nucleus
 Quantum Mechanics
o Explains the energy states in complex atoms and molecules
o Explains the relative brightness of spectral lines
o Electrons (-) act as both a wave and a particle
o Explains why electrons (-) don’t collapse into the nucleus
 Quantum Numbers
o State of an electron (-) (ground or excited)
o Probably location of an electron
 Principle Quantum Number (n)
o Determines the energy level
o The larger the value of n, the farther from the nucleus the electron should be
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 Valence Electrons
o Are in the outermost orbital (highest n, or energy level)
o Are in the reactive orbitals
o Atoms want a noble gas electron arrangement in its highest energy level for
stability
o Atoms will gain or lose electrons to achieve noble gas electron configuration, by
forming covalent bonds or becoming ions
 Electrons in orbitals (or energy levels) below valence are already stable and therefore
nonreactive
 Electron Configuration
Energy
Level
n=1
n=2
n=3
n=4
1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
Subshell # of orbitals
(Each
orbital hold
2 e- )
s
1
p
3
d
5
f
7
Max # of
Electrons
2
6
10
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Electrons fill the energy levels/subshells according to 3 principles. These are
Aufbau Principle - electrons fill orbitals starting at the lowest available energy levels before
filling higher levels (e.g. 1s before 2s). See reference chart.
Pauli Exclusion Principle – an orbital can hold a maximum
of 2 electrons. To occupy the same orbital, two electrons
must spin in opposite directions.
Hund’s Rule
1. Every orbital in a subshell is singly occupied with one
electron before any one orbital is doubly occupied.
2. All electrons in singly occupied orbitals have the same spin.
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3.
The Periodic Table and Periodic Trends
 Ions
o When a neutral atom loses electrons it becomes a cation (positive ion). A cation is
always smaller than its parent atom as it has lost electrons which means less same
charge repulsion which causes shrinking of the ion.
o When a neutral atom gains electrons, it becomes an anion (negative ion). An anion
is always larger than its parent atom as it has gained electrons increasing the same
charge repulsion which causes expansion of the ion.
 Isoelectric Series - ions with the same number of electrons
 Electronegativity - measure of the ability of the nucleus to attract electrons
o INCREASES as you go ACROSS (L to R) a period (smaller size = higher
electronegativity)
o DECREASES as you go DOWN a group (larger size = lower electronegativity)
4.
Ionic Bonding and Ionic Compounds
 Chemical Bonds
o Ionic bonds – the electrostatic attraction between ions
 Opposite charges
 Metal + nonmetal
 Octet Rule
o Atoms want to have a complete outer shell (8 valence electrons aka noble gas
configuration)
 Ions – atoms gaining or losing electrons to form an octet
o Cations – positive, left side, metals (lose electrons)
o Anions – negative, right side, nonmetals (gain electrons)
 Writing Chemical Formulas
1. Write the symbols for the elements (positive charge goes first)
2. Write the charge after the symbol
3. “Criss cross” the charges and drop the (+) and (-)
Example: Al (+3) S (-2) = Al2S3
 Naming Binary Ionic Compounds
o write the name of the cation (+)
o change the ending of the anion (-) to –ide
Example: NaCl = sodium chloride
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 Cations formed by Transition Elements
o most transition metals have more than one ionic charge (See Chart)
 Tin (Sn) and Lead (Pb) act like transition metals
 Writing Formulas with Transition Metals
o Write ion with charge.
Example: tin (IV) oxide = Sn4+ O2o Criss – Cross charges.
Sn2O4
o Reduce if necessary.
SnO2
 Naming Formulas with Transition Metals
o Find the charge of the cation
o Compounds must be neutral (equal to 0)
(charge of cation)(# of cations) + (charge of anion)(# of anions) = 0
*Example: Find the charge on the cation in Fe2O3
(charge of Fe = x)(# of Fe atoms = 2) + (charge of O = -2)(# of O atoms =3) = 0
(x)(2) + (-2)(3) = 0
2x + (-6) = 0
x = +3
Since x is the charge of the cation, Fe must have a charge of 3+ or Fe3+
5.
Covalent Bonding and Molecular Compounds
 Chemical Bonds – hold atoms together to create chemical compounds
o Covalent bonds
 nonmetal + nonmetal
Prefix
 share electrons to become an octet
1 – mono
 electronegativity is < 1.7
2 – di
3 – tri
o Ionic Bonds
4 – tetra
 metal + nonmetal
5 – penta
 gain or lose electrons to become an octet
6 – hexa
 electronegativity is > 1.7
7 – hepta
8 – octa
 Naming Binary Molecular (Covalent) Compounds
9 – nona
10 - deca
1. Use prefixes to show # of atoms
(*if there is only one of the first atom, the mono- is left off)
2. Second element ends in –ide
**Example: CO = carbon monoxide
CO2= carbon dioxide
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 Lewis Structures – diagrams that show valence electrons as dots
1. Find the total number of valence electrons (outer shell)
(label columns #1-8 on periodic table to show valence)
2. Central atom is the least electronegative
(H can never be the central atom)
3. Connect the other atoms to it by single bonds
4. Count each single bond as a pair of electrons
5. Add electrons to the outer atoms to give each one 8 (octet rule)
6. Add electrons to the central atom to give it 8
7. Check to make sure all valence electrons are used
 Lewis Structure for Ions
o Negative ions – ADD the charge to the total of valence electrons
o Positive ions – SUBTRACT the charge to the total of valence electrons
 Double and Triple Covalent Bonds
o Atoms that share two or three pairs on electrons to form an octet
1. Determine # of valence electrons
2. Form single bonds
3. Place lone pairs on electrons on atoms to get 8
4. Check to make sure all valence electrons are used
5. If there are too many electrons, form double bonds
Covalent Bonds
Type Electrons Strength Length
Shared
2
weak
long
4
6
medium medium
strong
short
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6.
Chemical Reactions
 Chemical Equations
o Reactant + Reactant
Product + Product
o “Skeleton” equations are unbalanced
 Symbols Used in Chemical Equations
Symbol Meaning
+
separates two reactants or two products
yields
reversible reactions
(s)
solid
(l)
liquid
(g)
gas
(aq)
dissolves in water
heat is used in the reaction
Pt
catalyst (speeds chemical reactions)
 Law of Conservation of Mass
o Mass is neither created or destroyed
o The atoms (or mass) are rearranged in a chemical reaction.
 Balancing Equations
1. Write a skeleton equation
2. Count up the # of each type of element on each side (make a T chart)
3. Pick an element that is not balanced and find the side that needs more of it
4. Add a coefficient (a number before the element that multiplies) to that side
5. Reevaluate the # of each element from step 2
6. Continue with these steps until all elements are balanced
 Types of chemical reactions
o Acid/Base: reaction between acids and bases; Acid – donates H+; base accepts H+
Example: HCO3-   +   HF    F- +   H2CO3
o Precipitation: formation of insoluble product out of aqueous solutions.
Refer to the back of the green periodic table for a solubility chart.
Precipitation reactions are often called "double replacement" or "double displacement" reactions
because it appears as if the ions switch places with each other.
Example: NaCl(aq) + AgNO3(aq)
AgCl(s) + NaNO3(aq)
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o Oxidation/Reduction: involve a transfer of electrons between atoms/ions.
Aka Redox reactions. See more below.
 Types of Redox Reactions
a. Synthesis or Combination - two or more substances combine
Example: 2Mg(s) + O2(g)  2MgO(s)
b. Decomposition - a substance breaks into two or more compounds
Example: 2LiClO3(s)  2LiCl(s) + 3O2(g)
c. Disproportionation - the same material is both oxidized and reduced
Example: Hg2Cl2 → Hg + HgCl2
d. Combustion - a substance reacts with oxygen in a rapid reaction that produces a flame.
Most often involve hydrocarbons reacting with oxygen in the air. Other elements may also
react with O2. O2 is always one of the reactants. Combustion may be complete or incomplete
depending upon the amount of oxygen that is available.
o Complete combustion – products are CO2 and water.
o Incomplete Combustion – products are CO and water.
Examples:
C(s) + O2(g)  CO2(g)
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
C3H8 (g) + 5O2 (g) 
3 CO2 (g) + 4 H20 (g)
Combustion reactions may also be classified as combination/synthesis reactions.
Examples:
Mg (s) + O2 (g)  MgO (s)
N2 (g) + 2 O2 (g)  2 NO2 (g)
4Al(s) + 3O2(g)

2Al2O3
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 More on Redox Reactions
4Al(s) + 3O2(g)

2Al2O3
Each Al loses 3 electrons
Each O gains 2 electrons
Oxidation
Reduction
element loses electrons
element gains electrons
Oxidation States = charge on the atom or ion
Which species was oxidized/reduced?
Oxidation = Loss of electrons
Na --> Na+ + e
0 --> +1
Notice that an element becomes more positive when it gets oxidized.
Reduction = Gain of electrons
Mg 2+ + 2e - --> Mg
+2
--> 0
Notice that an element becomes less positive (more negative) when it gets reduced.
Which element was oxidized and which one was reduced in the following reaction?
Find oxidation states (charges).
2KClO3
2KCl + 3O2
+1 +5 -2
+1 -1
0
Chlorine went from +5 --> -1, becoming less + so it must have  gained electrons or reduced.
Oxygen went from -2 --> 0, becoming more + so it must have lost electrons or oxidized.
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