Honors Chemistry
1 st Semester Final Exam Review 2014-2015
 The structure of the atom
Name : _____________________
 Subatomic particles- know location and how to calculate amounts given atomic number or atomic mass
 Atomic number = number of protons = number of electrons for a neutral atom  identifies what the element is. All atoms of the same element have the same atomic number!
 Atomic mass = protons + neutrons
 Isotopes- same number of protons and electrons, different number of neutrons; therefore mass number different, atomic number the same. Be able to distinguish how two isotopes differ. I.e. Carbon –12 and Carbon-14
 Be able to view a picture of an atom and determine its charge, atomic number, atomic mass
 Ion- charged particle resulting from the gain or loss of electrons. Results in an unequal number of electrons and protons in the atom o Positive ions = cations o Negative ions = anions o Be able to explain what happens to any element in the P-Table to form an ion
Lab Safety
 Bunsen burners- how do you safely light a burner?
 Luminous vs. Non-luminous flames- which can you read by? Which represents complete combustion of the gas? Which is the hotter flame?
Lab Skills
 Measurement- what is significant in a measurement? Measure to the readable value and estimate the last digit
 Volumetric measurement- Know how to read the meniscus
 Be able to identify basic lab equipment and know its use  which piece of glassware is most appropriate for which task?
 Significant digits- be able to determine the number of sig figs. in a number or for a final answer in a calculation
 Know your rules for sig figs and sig fig calculations!- Alt – Pac Rule . +/- determine the lowest decimal
position. X and / lowest number of significant figures.
 Appreciate that the electron can be considered to have wave like properties as well as particle like properties
 Understand the concept of electrons in regions of probability and use quantum numbers to explain the location and characteristics of electrons
Valence electrons Vs kernel electrons o Valence electrons determine the chemical reactivity of the atom
Ground state electrons Vs excited state electrons o Photons, quanta, bright line spectrum vs. continuous spectrum
Electron configurations
Be able to identify elements based on their electron configurations o o Recall and understand the rules for filling orbitals and determining electron configurations including the Aufbau principle, the o
Pauli exclusion principle and Hund’s rule.
Be able to determine the electron configuration (Noble gas configuration) of the most common ion of an atom o Be able to identify the location of the element in the periodic table based on its electron configuration
The Periodic Table o Know the 1 st periodic table was developed by Mendelev and the modern periodic table was developed by Mosley o Understand the regular, repeatable patterns occur across periods and within groups on the periodic table o o
Basic layout- periods (horizontal rows) vs. groups (vertical columns)
Know the names and basic properties of the groups in the periodic table
 1A = Alkali Metals = most active metals
 2A = Alkaline Earth Metals
 1B –10B = Transition Metals
 Metalloids = staircase  separates the metals from the nonmetals – has characteristics of each
 7A = Halogens = most active nonmetals
 8A = Noble Gases
 Lanthanide and Actinide series
Know the blocks in the P-Table = s, p, d, and f and be able to use electron configurations with the P-Table o o Know the types of ions that form in the various groups- i.e. Metals will form cations, nonmetals will form anions
Trends in the P-Table
 Recall the definition of effective nuclear charge and shielding and use these terms to describe the trends below o Electronegativity- ability of an atom to attract another atom’s valence electrons. Decreases down a group, increases across a period

o Ionization energy- energy necessary to remove an atom’s outer shell electrons. Decreases down a group, increases across a period o Atomic radius- increases down a group, decreases across a period o Know how to use the P-Table to determine the oxidation number of an element  this will help tremendously when it comes to writing formulas!
 Atoms bond to gain stability  octet rule  Nobel Gas configuration
 Ions- formed from the gain or loss of electrons
 Cations = + ions
 Anions = - ions
 Ionic Bonds- formed as a result of the electrostatic force (attraction of oppositely charged ions)
 Ionic Compounds are Crystals  organized into an array- regular geometric pattern
 Ionic compounds are NOT molecules!
 Know how to draw a Lewis Dot diagram for atoms and show the transfer of electrons in an ionic bond
 Know how to name ions and ionic compounds
 Anions- replace the ending of the element with “-ide” ion
 Cations- name element and include the word ion
 Use Table F to name polyatomic ions or determine their oxidation numbers
 Binary compounds- cation followed by name of anion
Ionic Compounds-
 Metals + nonmetals or substances that contain polyatomic ions
 Transfer electrons
Naming
 Binary ionic compounds- metal first, add nonmetal + suffix “-ide”
 Polyatomic ionic compounds- metals first, add name of polyatomic (see TABLE E)
*NOTE: if the ionic compound contains a transition metal, Roman numerals are used in the name to show the metal’s charge. examples: FeCl
2
FeCl
3
– Iron (II) chloride
– Iron (III) chloride
)
3
– Gold (III) sulfate Au
2
(SO
4
Writing formulas
 Metal ion first, then nonmetal
 Look up oxidation numbers
 Crisscross method- used with binary and polyatomic ionic compounds
Mg +2 + Cl  MgCl
Li + + MnO
4
2 Li
2
2
MnO
4
Covalent Compounds-
 Nonmetal + nonmetal
 Share electrons
Naming
 Least electronegative element first, second element add suffix “-ide”
 Use prefixes to indicate number of atoms in each molecule
 SO
3
 NO
2
– Sulfur trioxide
– Nitrogen dioxide
Writing formulas
 Least electronegative element first
 Use prefixes to determine the numbers of atoms
 Carbon dioxide – CO
2
 Carbon tetrachloride – CCl
4
 Write Empirical formulas- lowest whole number ratio of atoms in a crystal
 Determine Molecular formulas for covalent compounds
 Covalent compounds- share electrons to form bonds
 Octet Rule also applies to covalent compounds
 Draw structural formulas for covalent compounds  remember double and triple bonds are possibilities in drawing structural formulas
 Covalent compounds can form multiple bonds- double and triple
 Know exceptions to Octet Rule- Boron and Sulfur
 Know how to name covalent compounds- use prefixes
Polarity
 Know how to determine bond polarity vs molecule polarity- molecules can have polar bonds, but are symmetrical in shape and therefore nonpolar molecules

 VSEPR- valence shell electron pair repulsion theory- electrons want to get as far away from each other as possible- form specific shapes with max distance between electrons
 Use Structural formulas to determine the number of electron pairs around the central atom
 Count the number of lone and bonding pair electrons to determine the VSEPR numbers and the molecule’s shape
 The shape of the molecule will determine the molecules’ polarity
 Asymmetrical molecules are polar
 Symmetrical molecules are nonpolar
Comparison of Ionic Vs Covalent Substances
Ionic Substances
 High melting points- due to strong electrical attraction of the ions for one another
 Dissolve in water
Covalent Substances
 Lower melting points than ionic substances- weaker bonds
 Some will dissolve in water, some do not (Polar Vs
Nonpolar)
 Form crystals
 Hard and brittle
 Conduct electricity in solution (dissolved in water- aqueous) and in the liquid phase
 Do not conduct electricity in solid phase because of the rigidity of the crystal
 Soft solid phase
 Some conduct electricity in aqueous phase, but not as well
 as ionic substances (Polar Vs Nonpolar)
Do not conduct electricity in liquid or solid phase
 Matter has a natural tendency to be disordered (have a great deal of entropy). If a substance is to be held in the solid or liquid phase, then some sort of force must be present to keep the individual atoms, molecules, or ions in the solid or liquid in place. In other words, these forces prevent particles from “flying around” randomly, as they do in gases. These forces are called intermolecular forces
 Metallic substances- positive ions immersed in a sea of electrons
 Mobility of electrons gives the metals their characteristics
 Atomic mass vs molecular mass vs molar mass- remember molar mass is the heaviest- g/mol
 Perform mole conversion without mole map
 Define a mole
 Determine number of atoms or moles of atoms within a molecule using molecular formula- Cu(C
2
H
3
O
2
) moles of atoms?
 Percent composition calculations from chemical formulas
 Determine empirical formulas from percent by mass
 Determine molecular formulas from empirical formula and molecular mass
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2
- how many total
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