Regents Chemistry Content Syllabus
Overview of units:
1.
2.
3.
4.
5.
6.
Scientific measurement
Atomic structure
Periodic table
Bonding
Matter and energy
Math and solution calc.
7. Kinetics and equilibrium
8. Acids and bases
9. Redox and electrochemistry
10. Organics
11. Nuclear chemistry
12. Review
Unit 1 – The Science of Chemistry
I. Brief history of measurement
II. International system of units
A. Scientific notation
B. Significant figures
1. Rounding
III. Temperature, Volume, mass and Density measurements
IV. Classification of Matter and Energy
A. Matter
B. Substances
1. Elements
2. Compounds
3. Mixtures
C. Energy
D. Forms of energy
Unit 2 – Atomic structure
I.
II.
8/06
Matter (defined)
A. Solids, Liquids and gases
1. Physical and chemical properties
B. Elements
C. Mixtures
1. Compounds
2. Calculating percentages
Atoms
A. Introduction
B. Subatomic particles
1. Electrons
2. Nucleons (particles found in the nucleus)
a. protons (+)
b. neutrons (think Switzerland)
C. Atom structure
1. “empty space”
2. nucleus
a.
b.
c.
d.
atomic number
isotopes
mass number
atomic mass (weight)
D. Atomic structure models
1. Brief summary of the history
2. Principal Energy Levels
3. Quanta
4. Spectral lines
E. Orbital model
1. Energy levels
a. Principle quantum numbers
b. Sublevels
c. Orbital
2. electron configurations
3. valence electrons
F. Introduction to the Mole
Unit 3 – Periodic table
I.
II.
III.
IV.
V.
VI.
Development of
Element properties – Physical
a. Covalent atomic radius
b. Ionic radius
c. Metals
d. Non-metals
e. Metalloids or semi-metals
Element properties – Chemical
a. Valence electrons lost
b. Filling the outer ring or shell
Group chemistry
a. Groups 1 and 2 (most metallic element)
b. Groups 15 and 16
c. Group 17 (Most non-metallic element)
d. Group 18
e. Transition elements
Period chemistry
Naming the elements
Unit 4 – Bonding
I.
Nature of bonding – Why do chemical bonds form?
E. Chemical energy
F. Energy changes
G. Bonding and stability
H. Electronegativity differences
II.
Bonding between atoms
a. Ionic
i. Names
ii. Formulas
b. Covalent
i. Names
ii. Formulas (Molecular and Empirical)
iii. Non-polar
iv. Polar
v. Coordinate covalent
c. Metallic
i. Elements that contain metallic bonds
III.
Molecular attraction
a. Dipole-dipole
b. Vanderwaals
c. Hydrogen forces (not really a bond!)
IV.
Polarity and symmetry of covalent bonds
a. Electron dot diagrams
Types of Chemical reactions
a. Parts of the chemical formula
V.
Unit 5 –Kinetic Molecular Theory (KMT) and Energy
I.
II.
III.
Matter (A review)
Energy
a. Forms of energy
b. Energy and chemical changes
i. Exothermic reactions
ii. Endothermic reactions
c. Measurement of energy
i. Joules
ii. Thermometry
Phases of matter
a. Gases
i. Boyle’s law
ii. Charles law
iii. Combined gas law (equation)
iv. Standard temperature and pressure (STP)
v. Kinetic theory
vi. Gas law deviations
vii. Avogadro’s hypothesis
b. Liquids
i. Vapor pressure
ii. Boiling point
iii. Heat of vaporization
c. Solids
i. Heat of fusion
ii. Sublimation
Unit 6 – Mathematics of chemistry
I.
II.
III.
Mole interpretations
Using moles
a. Gram atomic mass
b. Gram molecular mass
c. Molar volume of a gas
d. Mole road map
Stoichiometry
a. Types of reactions (a review)
b. Balancing chemical equations
c. Problems involving formulas
i. Percent composition
d. Problems involving equations
i. Mole ratios
ii. Mass-mass problems
iii. Mass-volume problems
iv. Volume-volume problems
IV.
Math involving solutions - concentration
a. Methods of expressing concentration
i. Molarity (M)
ii. % by volume
iii. Parts per million (ppm)
b. Effects of solute on solvent (colligative properties)
i. Crystals
ii. Melting and freezing point (colligative properties)
iii. Behavior of electrolytes
Unit 7 – Kinetics and equilibrium
I.
II.
III.
Kinetics
a. Role of energy
i. Activation energy
ii. Heats of reaction
iii. PE diagrams
b. Factors affecting reaction rates
i. Nature of reactants (ionic vs. covalent)
ii. Concentration
iii. Temperature
iv. Surface Area
v. Catalysts
Equilibrium
a. Phase-change equilibrium (One substance)
b. Solution Equilibrium
i. Gases in liquids
ii. Solids in liquids
iii. Solubility
c. Chemical equilibrium
i. LeChatelier’s Principle
1. effect of concentration
2. effect of pressure
3. Effect of temperature
4. effect of a catalyst
ii. Keq – equilibrium expressions
Spontaneous reactions
a. Energy changes
b. Entropy changes
Unit 8 – Acids and bases
I.
II.
III.
IV.
Electrolytes
Acids and bases
a. Arrhenius theory of
b. Bronsted Lowry
Acid-base reactions
a. Neutralization reaction
i. Acid-base titrations
ii. Salts
Ionizations
a. pH and concentrations of ions
Unit 9 – Redox and electrochemistry
I.
II.
Redox
a. Oxidation
b. Reduction
c. Oxidation numbers
d. Redox ½ reactions and balancing
Electrochemistry
a. Half-reactions
b. Half-cells
c. Chemical cells
i. Electrodes
1. cathode
2. anode
d. Electrolytic cells
i. electroplating
Unit 10 – Organics
I.
II.
III.
IV.
Definitions
Compounds and characteristics
a. Bonding
b. Structural formulas
c. Isomers
d. Saturated, Unsaturated
Homologous series
a. Alkanes
b. Alkenes
c. Alkynes
Other compounds
a. Alcohols
i. Monohydroxy alcohols
V.
1. primary
2. secondary
3. tertiary
ii. dihydroxy alcohols
iii. trihydroxy alcohols
b. Organic acids
c. Aldehydes
d. Ketones
e. Ethers
f. Amides
g. Polymers and polymerization
i. Condensation polymerization
ii. Addition polymerization
h. Esters
Organic reactions
a. Substitutions
b. Addition
c. Fermentation
d. Esterfication
e. Saponification
f. Combustion
g. Others
Unit 11 – Nuclear chemistry
I.
II.
III.
Natural radioactivity
a. Emanations
i. Alpha decay
ii. Beta decay
iii. Gamma radiation
b. separating emanations
c. Detecting radioactivity
d. Half-lives
Artificial radioactivity
Nuclear energy
a. Fission
b. Fusion
c. Uses of radioisotopes
Tentative lab schedule: You are only allowed to miss 3 labs and still be eligible for the NYS Regents Exam.
Lab #
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
Lab Name – concept
Safety – Reference table preview
Graphing and math practice
Lab Techniques
Density and error calculation
Mendeleev and Indirect observations – development of the periodic table
Flame tests – ground state vs. excited state
Beanium lab – Calculating average atomic mass
Periodic trends
Puzzle pieces lab – Making and naming ionic compounds
A chemical reaction: metal and S Move to lab # 17:Hydrate mass – calculating
% by mass
Molecular symmetry
Conductivity – electrolytes and solution conductivity (write up #1)
Separation of mixtures and solutions
Heat transfer – Q=mCΔT
Vapor pressure – Using reference Table H
Physical and Chemical Changes / Types of chemical reactions
Determining chemical compound composition - Coefficients
Predicting mass of the product
Mass of the solute in solution
Colligative properties
Solubility of curves – Using reference table G
Cooling curves – moved to end
Reaction rates
Properties of Acids and Bases
Titration – (Write up #2)
Redox – electrochemistry
Organic models
Saponification - Esterfication
Making Ice Cream
Reference table review