Final Review
C H E M I S T RY 1 4 0 5
Method of Course Evaluation
• Lab Report Average
• Quiz Average
• Lab Practical
50%
20%
30%
• Lab is 20% of your final CHEM 1405 grade.
• Each student’s lowest lab report grade will be
dropped.
Final Practical
• All personal belongings (e.g. backpacks, purses) will
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be placed on the printer table or lab bench
Final exam will be require only written responses
Hands-on lab exercises will not be required
A periodic table will be provided
A calculator may be used
Smartphones, tablets, or computers are not
permitted
Long pants and close-toed shoes are needed to
comply with lab policies
Final Practical
• Format may contain any of the following:
• Fill in the blank
• Multiple choice
• Math problems
• Approximately 60 questions
• Average time for completion is 1 hour
SDS
• Safety Data Sheets
• SDS may be found in each lab or on the internet.
• They contain valuable information about chemicals
you will use in your experiments.
Density
• Density is a derived quantity.
𝑑𝑒𝑛𝑠𝑖𝑡𝑦 = 𝑚𝑎𝑠𝑠/𝑣𝑜𝑙𝑢𝑚𝑒
Reading a Meniscus
Physical Change
• A change in matter
that does NOT change
its identity.
• Example: Ice (solid)
melts into water
(liquid). The
compound remains
H 2 O despite the
change in physical
state.
Chemical Change
• Changes that produce a NEW kind of matter with
DIFFERENT properties.
• Example: Adding baking soda to vinegar results in
products of a salt, water, and carbon dioxide gas.
Evidence of Chemical Reaction
• Energy change (e.g. heat, light, electricity)
• Gas produced (e.g. evolution of carbon dioxide)
• Precipitate formed (e.g. sodium chloride solution +
silver nitrate solution forms insoluble silver chloride
precipitate)
• Color change
Element
• Simplest form of matter.
• Cannot be decomposed by chemical means.
• Each element can be found on the periodic table.
• Example: K – potassium, Fe – iron, S – sulfur, C -
carbon
Compound
• Can be decomposed into simpler substances by
chemical means.
• Compounds are composed of 2 or more elements
found on the periodic table.
• Example: H2O – water, CO2 – carbon dioxide,
NaHCO3 – baking soda (sodium bicarbonate)
Graphing
• All graphs need:
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Title
Labeled independent x-axis with units
Labeled dependent y-axis with units
Electromagnetic Spectrum
Spectral Lines
• When atoms are excited by an energy source (e.g.
electric current, heat), their electrons absorb energy.
• This causes the electrons to “jump” from their ground
state to an excited state (higher energy).
• When electrons fall back (relax) to their ground state
they emit energy in the form of a quantum of light.
• The light quantum has a characteristic frequency and
wavelength which produces a characteristic spectral
line.
Hydrogen Bright Line Spectrum
Spectral Lines
• Each element produces its own unique line
spectrum.
• Line spectra are often used to identify an unknown.
Nomenclature
• Is the chemical ionic?
• Is the chemical covalent?
• Is the chemical an acid?
• Memorize polyatomic names and charges
• Review handout
Balancing Equations
• Review lab handout
• Knowledge of molar mass calculations
• Mastery of mole concept
Symbols Used in Equations
• Reactants → Products
• → Yields or produces
• (s) solid
• (l) liquid
• (g) gas
• (aq) aqueous: a substance dissolved in water
• ↔ A reversible reaction
• ∆ Heat
Classifying Chemical Reactions
• Combination reaction (or synthesis)
• Decomposition
• Single replacement reaction
• Double replacement reaction
Combination Reactions
• Two elements or compounds → One compound
CO2 + H2O → H2CO3
MgO + H2O → Mg(OH)2
3 H2 + N2 → 2 NH3
Decomposition Reactions
• One compound → Two or more elements or
compounds
• Usually requires energy – heat, electricity
2 NaCl → 2 Na + Cl2
2 Ag2O → 4 Ag + O2
2 KClO3 → 2 KCl + 3 O2
Single Replacement Reactions
• One element + compound → Different element +
compound
• An element that is higher in the activity series will
replace one that is lower.
Mg + CuSO4 → MgSO4 + Cu
Cu + 2 AgNO3 → Cu(NO3)2 + 2 Ag
2 KBr + F2 → 2 KF + Br2
Double Replacement Reactions
• The cation (+) in one compound reacts with the
anion (-) in the other compound.
• The ions switch partners.
K2CO3 + CaCl2 → CaCO3 + 2 KCl
2 Li3PO4 + 3 BaCl2 → 6 LiCl + Ba3(PO4)2
K2SO4 + Sr(NO3)2 → 2 KNO3 + SrSO4
Diatomic Elements
• Hydrogen – H2
• Nitrogen – N2
• Oxygen – O2
• Fluorine – F2
• Chlorine – Cl2
• Bromine – Br2
• Iodine – I2
Molecular Geometry
Linear
Trigonal planar
Trigonal bypiramidal
Tetrahedral
Octahedral
Predicting Molecular Geometry
A = Central atom
X = Substituent
E = Electron Pair
Linear
Examples:
CS2; HCN; BeF2
Trigonal planar
Examples:
Examples:
SO3, BF3, NO3-, CO32-
SO2, O3, PbCl2, SnBr2
Tetrahedral
Examples:
CH4, SiCl4,
SO42-, ClO4-
Examples:
NH3, PF3
ClO3, H3O+
Examples:
H2O, SCl2
Acids and Bases
2 H2O ↔ H3O+ + OHpH = -log [H3O+]
pH range is 1 - 14
• Acid – A chemical species with a relatively high
amount of [H3O+] hydronium ion. pH range of 1 – 6.
• Base – A chemical species with a relatively high
amount of [OH-] hydroxide ion. pH range of 8 – 14.
• Neutral – pH of 7.
Buffers and Equilibrium
• Buffer – Solutions that resist pH changes when
small amounts of strong acid or strong base are
added to it.
• Buffer – A weak acid and its conjugate base OR a
weak base and its conjugate acid.
• HA ↔ H+ + A• Since this weak acid is only partially dissociated,
small amounts of acid or base will shift the
chemical reaction only slightly due to Le Chatelier’s
Principle.