Chemistry Semester 1 Review Outline (3x5 Note-Card)
I. Introduction to Chemistry (Unit 0: Ch. 1-4)
A. Define Chemistry: “The study of matter and the changes it undergoes.”
B. Identify the composition of the earth’s air. (78% Nitrogen (N2), 21% Oxygen (O2), etc.)
C. Identify the gases that cause acid rain.
(CO2 makes Carbonic acid, NO2 makes Nitric acid, and SO2 makes Sulfuric acid)
D. Recognize the ozone molecule (O3) and describe the reason and consequence of ozone depletion.
E. Identify the photosynthesis equation and its significance for renewing our air.
F. Identify the difference between the greenhouse effect and ozone depletion.
II. Matter and Change
1. Matter is defined as anything that has mass and occupies space.
2. Mass, the most reliable quantity of measurement, describes the amount of matter in an object.
3. The Law of Conservation of mass (Lavoisier) states that matter is neither created nor destroyed by chemical reactions.
4. Vocabulary Words Related to matter and change:
· Homogeneous: Having the same consistency throughout the substance.
· Heterogeneous: Having a different consistency throughout.
· Substance: (A pure homogeneous substance: Element or compound)
· Identify the difference between an element, compound, and mixture
· Be able to use a solubility table to predict the formation of precipitates.
· Identify the difference between a homogenous mixture and a heterogeneous mixture.
· Understand the difference between a gas and a vapor. (A vapor is the term used for the gaseous phase of a substance that is normally a liquid or solid at room temperature.)
· Symbols of phases in equations, (s) = solid, (l) = liquid, (g) = gas, and (aq) = aqueous or dissolved in water.
· Physical vs. Chemical Change: (Chemical changes describe how a new substance if formed)
· Describe the difference between reactants and products.
III. Process of Science/Scientific Method
A. Process of Science/Scientific Method
Using your senses to gather information.
A scientific guess that explains observed facts. It must be tested further with experiments.
Used to test hypotheses and theories.
4. Repeat steps 1-3
A statement that provides a general explanation for observations made over a long period of time. Theories are not facts and can be disproven.
Other Vocabulary that is related to process of science:
A statement that describes a natural event. This is usually a summary of the results of many observations and experiments. Scientific laws are different than theories. Scientific laws cannot be broken.
Identify the difference between Pure Science (study science to advance knowledge) and Technology (apply science concept to make a product that benefits society). All technology advances have both benefits and risks associated with them.
Qualitative Observations: Any observation that identifies an object without using numerical values.
Quantitative Observations: Any observation that identifies a substance by using numerical data.
A. Significant Figures and Measurement
1. Uncertainty is found in all measurements: Follow these rules with measured quantities.
· *Keep all certain digits plus one uncertain digit.
· *When mult. and dividing, round to the measured quantity with the lowest number of significant digits
· *When adding and subtracting, round to the measured quantity to the least precise number.
2. Use scientific notation in calculations.
3. Be able to use unit-conversion method to perform calculations.
4. Convert Celcius temp. to Kelvin and Kelvin to Celcius. °C = K -273 and K = °C + 273
5. Calculate the percentage error in a measurement given the equation
% error = (Accepted value - measured value)
B. Vocabulary Related to Measurement
1. Identify the difference between accuracy and precision.
2. Recognize examples of measurements that are:
· Both accurate and precise. (Three measure values for density were 2.68, 2.69, and 2.69 g/ml. The accepted value of density for this element is 2.70 g/ml)
· Precise but not accurate. (Three measurements for density were 2.32, 2.33, and 2.32 g/ml. The accepted value for the density of this element is 2.70 g/ml)
V. Elements and the Atom
1. Given a periodic table, identify an element’s
a. Chemical Symbol (Used to represent a chemical element). Only the first letter is capitalized.
b. atomic number (number of protons). In a neutral element it is also the number of electrons.
c. mass number (protons + neutrons)
d. Atomic Mass (the average of all the naturally occurring isotopes). This is different than mass number.
e. The number of electrons in a neutral element is equal to the number of protons.
B. The Atom (The fundamental building block of nature)
1. Identify the three subatomic particles found in an atom.
Mass (u) Charge Location
a. Proton 1 u + 1 nucleus
b. Neutron 1 u none nucleus
c. electron 1/1800 u - 1 very far outside nucleus
2. What is a cathode ray tube and what subatomic particle is represented by cathode rays? electrons
3. Define an isotope: Same element with a different mass.
i.e. hydrogen has three isotopes:
· Hydrogen-1 (Protium): 1 proton, 0 neutrons, 1 electron-- Atomic # = 1 and Mass # = 1
· Hydrogen-2 (Deuterium): 1 proton, 1 neutron, and 1 electron-- Atomic # = 1 and Mass # = 2
· Hydrogen-3 (Tritium): 1 proton, 2 neutrons, and 1 electron-- Atomic # = 1 and Mass # = 3
VI. Electric Charge, Laws and Atomic Theories (Unit 1: Ch. 5)
A. Electric Charge
1. Two types of charge (positive and negative)
2. Like charges repel each other
3. Unlike charges attract each other
4. Electrically neutral objects have equal amounts of positive and negative charge.
5. Ions are particles that have excess charge.
a. Cations have extra positive charge. (These ions are usually metals that have lost one or more of their valence (outer) electrons)
b. Anions have extra negative charge. (These ions are usually nonmetals that have gained extra valence electrons)
B. Laws (Dalton’s atomic theory was based on these three laws)
1. Law of conservation of mass (Lavoisier) (see earlier definition)
2. Law of definite proportions.
a. States that the proportion by mass of elements in a given compound is always the same
3. Law of multiple proportions. (Dalton)
a. State that the mass of one element that combines with a fixed mass of the other element form simple, whole number ratios.
C. Atomic Theories
1. The Atom
a. Idea of the atom originated from the Greek “atomos” which means fundamental and indivisible. First proposed by the Greek philosopher Democritus.
b. Since Democritus had no scientific evidence to support the atom, the idea of the atom was lost for 2000 years due to a stronger following of Plato and Aristotle.
2. Dalton’s atomic model
a. Based on Greek idea of indivisible matter
· The atom is indivisible (This was later disproven!!)
· Same elements are exactly alike. (This was later disproven!!)
example: one carbon atom is exactly like another carbon in every aspect of properties and mass.
· Different elements are different.
· Compounds are made by two or more elements linking together on an atomic level.
3. Thompson’s model (Plum Pudding Model)
a. The discovery of the proton by E. Goldstein in 1886 and J.J. Thompson’s later discovery of the electron in 1897 proved that the atom consisted of smaller charged particles.
b. The discovery of the neutron by James Chadwick in 1932 also proved that atoms could have slightly different masses creating isotopes of those elements.
4. Rutherford’s atomic model (Gold Foil Experiment)
1. Be able to describe Rutherford’s Gold Foil Experiment.
1. Describe the experiment Set-up.
2. Describe the experiment results.
2. Changes Rutherford made to the atomic model
· He placed the proton and neutron in the center of the atom. (most of the alpha particles went straight through the gold foil.)
· He said the nucleus of the atom contained most of the mass and was charged positively. (Some of the + charged alpha particles were deflected (like charges repel) and some bounced straight back (indicated concentrated mass)).
· Electrons were placed a large distance from the nucleus and moved in a circular orbit around the nucleus. (Only explanation of why the electrons did not crash into the nucleus) This was not proven experimentally and would later lead to the revision of the theory.
5. Bohr’s atomic model (electrons orbit in energy levels)
1. Still have a nucleus.
2. The electrons orbit the nucleus at set energy levels. (Later disproven)
6. Current atomic model (electrons locations are predicted using probability)
1. Still have a nucleus.
2. Uses a wave equation to predict the location of electrons using probability.
3. Uses the Quantum idea. (e- jump specific energy levels. Photons are a quantum of light energy)
VII. Orbital Diagrams, Electron Configurations, and Dot Structures (Units 2-13: Ch. 6, 13, 14, and 16)
a. Principle Energy Levels (1, 2, 3, 4, 5 ...)
b. Energy Sublevels (s, p, d, f, ...)
ü Locations of sublevels (s, p, d, f)
ü Shapes of orbitals (s, p)
c. Orbital (holds a maximum of 2 electrons)
d. Electron Configurations
i. Stable Noble Gas configuration s2p6 (closed shell)
1. Be able to recognize the electron configuration for an element.
Example: Titanium Z=22 1s22s22p63s23p64s23d2 or shorthand version [Ar] 4s23d3
2. Be able to identify the Lewis dot structure for an element and a compound.
Examples: (See Unit 3 practice test for more)
3. In molecular compounds, atoms share valence electrons to form covalent bonds to attain the noble gas configuration. (Octet Rule)
4. Double and triple bonds form in order to satisfy the octet rule.
a. Sigma bond are the 1st covalent bond between atoms. Pi bonds are any additional bonds formed in double or triple bonds.
b. Double bonds have a sigma and a pi bond.
c. Triple bonds have a sigma and 2 pi bonds.
5. Resonance structures happen when there are several equally valid ways of drawing the molecule.
ii. Orbital Diagrams: Be able to use an Aufbau diagram or periodic table to draw an orbital diagram. (See unit 2 notes)
1. Aufbau Principle (electrons fill lowest energy level 1st)
2. Hund’s Rule (When filling a sublevel with multiple orbitals the electrons fill empty orbitals within that sublevel before doubling up.)
3. Pauli Exclusion Principle (only 2 electrons can occupy one orbital and they must have opposite spin)
iii. Heisenberg Uncertainty Principle (You cannot be certain of the exact location or momentum of an electron at the same time)
e. Energy associated with electrons
i. Greater changes in quantum levels produced greater energy in the form of light color.
1. Visible light ends in a transition of n=2. (5→2 is greater energy than 3→2)
2. Electromagnetic spectrum includes a variety of radiation (radio to gamma)
3. Visible light represents only a small portion.
a. Violet is the highest energy for visible light (wavelength l = 400 nm)
b. Red is the lowest energy for visible light. (wavelength l = 700 nm)
VIII. Periodic Table:
a. Describe the origin of the periodic table. (Mendeleev and Moseley)
b. Identify the position of groups, periods, and the transition metals in the periodic table.
i. Families of elements
1. Alkali Metals (Most reactive group of metals)
2. Alkaline Earth Metals
3. Carbon family
4. Boron family
5. Nitrogen family
6. Oxygen family
7. Halogens (Most reactive group of nonmetals)
8. Noble gases (Least reactive group of elements)
ii. Location of elements
1. Transition metals
2. Inner transition metals
3. Location of metals
4. Location of nonmetals
5. Location of metalloids
c. Other related vocabulary words.
i. Periodic: a repeating pattern.
ii. Periodic table: The arrangement of elements based on similarities in their properties.
iii. Periodic Law: When elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties.
iv. Periods: Horizontal row of elements that correspond to principle energy level.
v. Groups: Vertical group of elements with similar electron configurations.
vi. Valence electrons: electrons found in the outermost energy level.
X. Chemical Formulas (Unit 4: Part 1) Ch. 6
A. Formulas are used to represent both qualitative and quantitative information.
1. Qualitative relates to which elements are in the formula.
2. Quantitative relates to the number of atoms of each element in a molecule.
3. Ionic compounds contain a metal combined with a nonmetal. (formula units)
4. Ionic compounds are always represented by empirical formulas (The simplest ratio of ions)
5. Molecular compounds contain two nonmetals. (molecules)
B. Rules for writing correct chemical formulas. (use ion sheet)
1. Diatomic elements , like chlorine (Cl2) are written with the subscript 2 to show that they are found in pairs. There are 7 diatomic elements. ( N2, O2 , H2 , F2 , Cl2 , Br2 , I2 )
2. Cation (+) is written first and anion (-) is written second.
3. Make sure that the charges balance out so the compound is neutral.
example: The ions Mg2+ and Cl- will make the formula MgCl2
4. If you need more than one polyatomic ion use parentheses ( ) and subscripts. Mg2+ and NO3- will make the formula Mg(NO3)2
C. Naming chemical compounds. (Use ion sheet to help)
1. Stock System
a. Roman Numerals are only used to identify the charge on the cation if it is possible for the cation to have more than one charge.
examples: Ionic compound
CuCl2 is named copper (II) chloride
b. Roman numerals are not used if there can be only one charge.
example: AlCl3 is named aluminum chloride
2. Traditional system
a. Ionic compounds (a metal with a nonmetal)
· metals with the lower charge have an ous ending like cuprous