Download 1st semester Final Exam review material. Unit 1: Measurement and

1st semester Final Exam review material.
Unit 1: Measurement and Calculations
What is Chemistry?
 Matter is anything that has mass and occupies space.
 You don’t have to be able to see something for it to qualify as matter
 Chemistry is the study of the composition of matter and the changes that matter undergoes.
 Because living and nonliving things are mad of matter, chemistry affects all aspects of life and most natural
events.
The scientific Method
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Scientific method is a logical, systematic approach to the solution of a scientific problem
Steps in the scientific method include making observations, testing hypotheses, and developing theories.
1. Making Observations
o When you use your senses to obtain information you make an observation.
2. Ask a question The observation can lead to a question
o Example: observation = flashlight won’t come on, question = what’s wrong with the flash light?
3. Developing a hypothesis
o A hypothesis is a proposed explanation for an observation.
o Example: guessing the batteries are dead would be a hypothesis
4. Testing Hypotheses
o Experiment is a procedure that is used to test a hypothesis
o independent variable is the variable that you change during an experiment
o dependant variable is the variable that is observed during the experiment.
o For the results of an experiment to be accepted they MUST be able to be reproduced.
5. Analyze Data
o Look at the information from your test to see if your hypothesis is correct
o If correct: reinforce hypothesis with additional experimentation
o If incorrect: reevaluate the hypothesis and create a new experiment
6. Developing Theories
o a hypothesis becomes a theory when it meets the test of repeated experimentation.
o Theory is a well-tested explanation for a broad set of observations.
o Theories can never be proven; they may be changed at some point in the future to explain new
observations or experimental results.
. Scientific Laws
o Scientific law is a concise statement that summarized the results of many observations and experiments
o A scientific law does not attempt to explain why something occurs
Using and Expressing Measurements
 Scientific notation, a short way of writing large or small numbers, a given number is written as the product of two
numbers: a coefficient and 10 raised to a power.
 Example: 602,000,000,000,000,000,000,000 would be written as 6.02 x 1023
Accuracy, Precision and Error
 Accuracy is a measure of how close a measurement comes to the actual or true value of whatever is measured.
 Precision is a measure of how close a series of measurements are to one another
 To evaluate the accuracy of a measurement, the measured value must be compared to the correct value. To
evaluate the precision of a measurement, you must compare the values of two or more repeated
measurements.
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The center is the true value, A has both accuracy (near center) and precision (darts close to one another. B has only
precision because darts are close to another but not the center. C has no accuracy or precision.
Accepted values is the correct value based on reliable references
Experimental value is the value measured in the lab.
Error is the difference between the experimental value and the accepted value
Error can be positive or negative
Percent error is the absolute value of the error divided by the accepted value, multiplied by 100%
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Significant Figures in Measurements
 Significant figures include all of the digits that are known, plies a last digit that is estimated
 Measurements must always be reported to the correct number of significant figures because calculated
answers often depend on the number of significant figures in the values used in the calculations.
Rules for determining whether a digit in a measured value is significant
1. Nonzero digits are significant. 5.23 has 3 significant figures
2. Zeros between nonzero digits are significant. 5001 has 4 significant figures
3. Zeros in front of nonzero digits are not significant, they are only place holders. 0.000099 has 2
significant figures
4. Zeros at the end of a number and to the right of a decimal place are significant. 1.0100 has 5
significant figures
5. Zeros to the left of an understood decimal point are not significant, they are only place holders.
55000 has 2 significant figures
6. Defined quantities and counted quantities have unlimited number of significant figures.
Significant figures in Calculations
 In general, a calculated answer cannot be more precise than the least precise measurement from which it
was calculated.
 When rounding first decide how many significant figures the answer should have.
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Next round to that number of digits, counting from the left.
If the number to right of the last significant digit is 4 or less round down, if it is 5 or up round up.
Example: each rounded to 3 significant figures 5.236 = 5.24, 8.023 = 8.02
With addition or subtraction the calculation should be rounded to the same number of decimal places (NOT digits)
as the measurement with the lease number of decimal places
Example:
[2.01 has the lease number of decimal places]
With multiplication and division the calculation should be rounded to the same number of significant figures as the
measurement with the lease number of significant figures.
Example:
[12 has only 2 significant figures]
Units and Quantities
 Meter is the SI basic unit of length
 Common metric units of length include the centimeter, meter, and kilometer
 Scientists commonly use two equivalent units of temperature, the degrees Celsius and the Kelvin.
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Conversion Factors
 Many quantities can usually be expressed different several different ways.
 Example: 1 dollar = 4 quarters = 10 dimes = 100 pennies
 Whenever two measurements are equivalent, a ratio of the their measurement will equal 1
 Conversion factor is a ratio of equivalent measurements.
 Example:
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When a measurement is multiplied by a conversion factor, the numerical value is generally changed, but the
actual size of the quantity measure remains the same.
 Example: 2 hours = 120 minuets = 7200 seconds
Metric Units and Values
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Prefix
tera
giga
mega
kilo
hecta
deca
symbol
T
G
M
k
h
da
12
9
6
3
meaning
10
10
10
10
pneumonic
the
great
mad
king
2
10
Henry
base
10
1
10
died
by
0
deci
centi
milli
micro
nano
pico
d
c
m
μ
n
p
-1
-2
-3
10
10
10
drinking
chocolate
milk
-6
10
10
-9`
10-12
under
nick's
porch
Metric conversion use the following basic equation:
Prefix units have 2 or more letters: kg, pm, TJ
Base units have 1 letter: m, g, L, J, N, s
Dimensional Analysis
 Dimensional Analysis is a way to analyze and solve problems using the units of the measurements.
 Dimensional analysis provides you with an alternative approach to problem solving.
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Example: 9875 seconds equals how many hours?
 The key to dimensional analysis is to set it up so that the UNITS cancel.
Determining Density
 Density is the ratio of the mass of an object to its volume
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Unit 2: Matter and change
Identifying Substances
 Substance is matter that has a uniform and definite composition.
 Gold and copper are examples of substances and are also referred to as pure substances
 Every sample of a given substance has identical intensive properties because every sample has the same
composition
 Physical property is a quality or condition of a substance that can be observed or measured without changing the
substance’s composition.
 Physical properties include hardness, color, conductivity, malleability, melting point and boiling point.
 Physical properties can help chemists identify substances
2.1 C. States of Matter
1. Solid
 Solid is a form of matter that has DEFINITE shape and volume
 The shape of a solid doesn’t depend on the shape of a container
 The particles in a solid are packed tightly together and often in an orderly arrangement
 Solids have vibrational kinetic energy (the atoms vibrate around a fixed position)
 Solid cannot be compressed
2. Liquid
 Liquid is a form of matter that has a DEFINIE volume but NOT a definite shape
 a liquid takes the shape of it’s container
 The volume of the liquid doesn’t change as it’s shape changes, it is constant
 the particle in a liquid are in close contact with one another but not as rigid or orderly as a solid
 in a liquid the particle are free to flow past one another and have more space between the atoms
 Liquids have vibrational and rotational kinetic energy
 Liquids are slightly compressible
3. Gas
 Gas is a form of matter that takes both the shape and volume of its container
 The particles in a gas are much farther apart than the particles in a liquid.
 Gases are easily compressed into smaller volumes
 Vapor describes the gaseous state of a substance that is generally a liquid or solid at room temperature.
 Gases have vibrational, rotational and translational kinetic energy
 The atoms in gases have the most space between them
Physical Changes
 Physical change a change during which some properties of a material change, but the composition of the material
does not change
 Boil, melt, freeze, condense, breaking, splitting , grind, and cut are used to describe physical changes.
 Physical changes can be classified as reversible or irreversible.
 Melting, freezing and boiling are examples of reversible physical changes (can be changed back)
 Cutting, grinding, and breaking are example of irreversible physical changes (cannot be changed back)
A. Classifying Mixtures
 Mixture is a physical blend of two or more components (substances)
 Most samples of matter are mixtures
 Based on the distribution of their components, mixtures can be classified as heterogeneous mixtures or as
homogeneous mixtures.
1. Heterogeneous Mixture
o Heterogeneous mixture is a mixture in which the composition is NOT uniform throughout.
o Examples: salad, pizza, beach sand
2. Homogeneous mixture
o Homogeneous mixture is a mixture in which the composition IS uniform throughout.
o Example: vinegar, soda, tap water
o Solution is a homogenous mixture where solutes are dissolved in a solvent (kool-aid)
o most solutions are liquids but some can be solid (steel, bronze) and some are gases (air)
o Phase describes any part of a sample with uniform composition and properties (solid, liquid, gas)
o all homogeneous mixtures consist of a single phase
Distinguishing Elements and Compounds
 A pure substance can be either an element or a compund
 Element is the simplest form of matter that has a unique set of properties
 Element cannot be broken down into simpler substances. Examples: oxygen, nitrogen, sodium
 Compound is a substance that contains two or more elements chemically combined in a fixed proportion.
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Compounds can be chemically separated in to simpler substances. Examples: water, sugar, oil
Compounds can be broken down into simpler substances by chemical means, but elements cannot.
1. Breaking Down Compounds
o Chemical change is a change that produces matter with a different composition than the original matter
o Heating a substance can be used to break down compound into simpler substances.
o Example: heating sugar with give you carbon and water
2. Properties of Compounds
o The properties of compounds are usually quite different from those of their component elements.
o Carbon: black tasteless solid, sugar: white sweet solid
o Hydrogen: flammable gas, Oxygen: color gas that supports burning, water: liquid that can stop
materials from burning
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Symbols and Formulas
 Chemists use chemical symbols to represent elements, and chemical formulas to represent compounds
 Chemical symbol is the one or two letter abbreviation for an element
 The first letter of a chemical symbols is ALWAYS capitalized and the second letter is always lowercase.
 Chemical symbols provide a short hand way to write the chemical formulas of compounds
 Subscripts are used to show the number of atoms of a given element in a compound’s chemical formulas.
 Example: C12H22O11 is the formula for table sugar, it has 12 Carbon atom, 22 Hydrogen atoms, and 11 Oxygen
atoms.
Chemical Changes
 Chemical property is the ability of a substance to undergo a specific chemical change
 Words that signify a chemical change: burn, rot, rust, decompose, ferment, explode, and corrode
 Chemical properties can be used to identify a substance
 During a chemical change, the composition of matter always change
Recognizing chemical changes
 Possible clues to chemical changes include a transfer of energy, a change in color, the production of a gas, or
the formation of a precipitate
 Precipitate is a solid that forms and settles out of a liquid mixture
Conservation of Mass
 During any chemical reaction, the mass of the products is always equal to the mass of the reactants.
 Mass is also held constant during a physical change
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The law of conservation of mass states that in ANY physical change or chemical change the mass is conserved
(stays the same).
 Mass is neither created nor destroyed
Phase Change and Temperature
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During a phase change the temperate of a substance remains the same. C= temperature that the substance is
melting, G = temperature that the substance is boiling.
Chapter 4: Atomic Structure
Subatomic Particles
 Three kinds of subatomic particles are electrons, protons, and neutrons.
1. Electrons
o 1897 J.J. Thomson discover electrons using the cathode ray tube
o Electrons are negatively charged subatomic particles
2. Protons and Neutrons
o Atoms have not net electric charge, they are electrically neutral
o Electric charges are carried by particles of mater
o Electric charges always exist in whole-number multiples of a single basic units
o When a given number of negatively charged particles combines with an equal number of positively
charged particles
o Protons are positively charged subatomic particles
o Neutrons are subatomic particles with not charge but with a mass nearly equal to that of a proton
The Atomic Nucleus
 When subatomic particles were discovered, scientists wondered how these particle were put together in an atom
 J.J. Thompson’s model was known as the ‘plum-pudding model”
1. Rutherford’s gold-foil Experiment
o 1911 Rutherford decided to test the current theory of atomic structure
o He shot alpha particles (positive particle) at thin sheet of gold
o The majority of the particles passed straight through
o A small fraction bounced off the gold foil at very large angels
2. The Rutherford Atomic Model
o Based on the experimental results Rutherford suggested a new theory of the atom
o He suggested that the atom is mostly empty space (explaining the lack of deflection of most alpha
particles)
o Also that all the positive charge and mass is concentrated in a small region
o Nucleus is the tiny central core of an atom and is composed of protons and neutrons.
o In the nuclear atom, the protons and neutrons are located in the nucleus. The electrons are
distributed around the nucleus and occupy almost all the volume of the atom
Atomic Number
 Elements are different because they contain different numbers of protons
 Atomic number of an element is the number of protons in the nucleus of an atoms in their element
 Remember that atoms are electrically neutral
 Because of that the number of electrons must equal the number of protons
Mass Number
 Mass number is the total number of protons and neutrons in an atom.
 If you know the mass number and the atomic number of any atom you can determine the number of neutrons in the
atom
 The number of neutrons in an atoms is the difference between the mass number and the atomic number
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Short hand notation:
,
or gold-197
Isotopes
 Isotopes are atoms that have the same number of protons but different number of neutrons
 Because isotopes of an element have different numbers of neutrons, they also have different mass numbers.
 Isotopes are chemically alike because they have identical numbers of protons and electrons (which are responsible
for chemical behavior)
Unit 4: Modern atomic theory and Periodic Trends
Atomic Orbitals
 An atomic orbital is often thought of as a region of space in which there is a high probability of finding an electron
 Each energy sublevel corresponds to an orbital of a different shape, which describes where the electron is
likely to be found.
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s orbitals are spherical, the probability of finding an electron at a given distance from the nucleus in an s orbital
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does not depend on direction
p orbitals are dumbbell-shaped. The three kinds of p orbital’s have different orientations in space.
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Four of the five kinds of d orbitals have clover leaf shapes. The shapes of f orbitals are more complicated than for
d orbitals.
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The numbers and kinds of atomic orbitals depend on the energy sublevel.
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The lowest principal energy level (n = 1) has only one sublevel, called 1s.
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The second principal energy level (n = 2) has two sublevels, 2s and 2p. the second principal energy level has four
orbitals: 2s, 2px, 2py, and 2pz.
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The third principal energy level (n = 3) has three sublevels. 3s, 3p, and 3d. Thus the third principal energy level has
nine orbitals (one 3s, three 3p, and five 3d orbitals).
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The fourth principal energy level (n = 4) has four sublevels, 4s, 4p, 4d, and 4f. The fourth principal energy level,
then, has 16 orbitals (one 4s, three 4p, five 4d, and seven 4f orbitals).
Electron Configurations
 In an atom, electrons and the nucleus interact to make the most stable arrangement possible.
 The ways in which electrons are arranged in various orbitals around the nuclei of atoms are called electron
configurations.
 Three rules—the aufbau principle, the Pauli exclusion principle, and Hund’s rule—tell you how to find the
electron configurations of atoms
1. Aufbau Principle
o aufbau principle, electrons occupy the orbitals of lowest energy first.
o The orbitals for any sublevel of a principal energy level are always of equal energy.
o within a principal energy level the s sublevel is always the lowest-energy sublevel.
o range of energy levels within a principal energy level can overlap the energy levels of another principal
level.
2. Pauli Exclusion Principle
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Pauli exclusion principle, an atomic orbital may describe at most two electrons.
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To occupy the same orbital, two electrons must have opposite spins
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A vertical arrow indicates an electron and its direction of spin (↑ or ↓).
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An orbital containing paired electrons is written as
3. Hund’s Rule
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Hund’s rule states that electrons occupy orbitals of the same energy in a way that makes the number of
electrons with the same spin direction as large as possible.
o three electrons would occupy three orbitals of equal energy as follows:
o A convenient shorthand method for showing the electron configuration of an atom
o write the energy level and the symbol for every sublevel occupied by an electron.
o indicate the number of electrons occupying that sublevel with a superscript.
o Example: oxygen= 1s22s22p4.
o Note that the sum of the superscripts equals the number of electrons in the atom.
B. Exceptional Electron Configurations
Chapter 6: The Periodic Table
 Chemists used the properties of elements to sort them into groups.
 In the modern periodic table, elements are arranged in order of increasing atomic mass.
 Periodic law stated that when the elements are arranged in order of increasing atomic number, there is a
periodic separation of their physical and chemical properties.
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Three classes of elements are metals, nonmetals, and metalloids.
Metals are good conductors of heat and electric current
Metals are solid at room temperature (except mercury).
Metals are ductile (can be drawn into wires) and Malleable (hammered in to sheets)
Nonmetals are poor conductors of heat and electric current.
Solid nonmetals tend to be brittle.
Most are gasses at room temperature
Metalloids (or semi metals) have properties similar to metals and non metals.
Periodic Trends: as you move across a period (left to right) the elements get less metallic; as you move down a
group the elements get more metallic. Francium is most metallic element.
Classifying the Elements
 Elements can be sorted in to noble gases, representative elements, transition metals, or inner transitions
metals based on their electron configurations.
 Noble gases are elements in Group 8A and are called inert gases because they rarely take part in a reaction.
the s and p sublevels are completely filled with electrons.
 Halogens are elements in Group 7A
 Representative elements are in groups 1A through 7A. They display a wide range of physical and chemical
properties. The s and p sublevels of representative elements are not filled.
 Transition metals are in D-block and are usually displayed in the main body of a periodic table.
 Inner transition metals are below the main body of the periodic table. These elements are characterized by f
orbital’s that contain electrons.
Periodic trends
Atomic radius is one half of the distance between the nuclei of two atoms of the same elements when the atoms are
joined.
 In general, atomic size increases from top to bottom within a group and decreases from left to right across a
period.
Ions are an atom or groups of atoms that has a positive or negative charge.
 Positive and negative ions from when electrons are transferred between atoms.
Cation is an atom or groups of atoms with a positive charge
 Cations are always smaller than the atoms from which they form.
 Metals always form cations
Anion is an atom or groups of atoms with a negative charge
 Anions are always larger than the atoms from which they form.
 Nonmetals always form anions
Ionization energy is the energy required to remove the first electron from an atom
 First ionization energy tends to decrease from top to bottom within a group and increase from left to right
across a period.
Electronegativity is the ability of an atom of an element to attract electrons when the atom is in a compound.
 Electronegativity values decrease from top to bottom within a group, and values tend to increase from left ot
right across a period.
 The most electronegative element is Fluorine.
Valance Electrons are the electrons in the outer energy level of an atom
 The number of valance electrons increases for the representative elements as you move across a period
 The number of valance electrons stays the same as you move down a family.
Ch 7 Ionic and Metallic Bonding
Valence Electrons
 Elements within each group of the periodic table behave similarly because they have the same number of valence
electrons.
 Valence electrons are the electrons in the highest occupied energy level
 The number of valence electrons determine the chemical properties
 To find the number of valence electrons in an atom of a representative element, look at the group number.
 Valence electrons are usually the only electrons used in chemical bonds
 Electron dot structures (Lewis dot structure) are diagram that show the valence electrons at dots.
The Octet Rule
 Octet rule stated that in forming compounds atoms tend to achieve the electron configuration of a noble gas.
 An octet is a set of eight electrons
 Atoms of metals tend to lose their valence electrons leaving a complete octet in the next-lowest energy level.
 Atoms of some nonmetals tend to gain electrons or to share electrons with another nonmetals to achieve a
complete octet.
Formation of Cations
 An atom’s loss of valence electrons produces a cation, or a positively charged ion.
 The most common cations are those produced by the loss of valence electrons from metal atoms.
 Example: Na: 1s2 2s2 2p6 3s1 loss of 1 electrons forms
Na+1: 1s2 2s2 2p6
 Some ions formed by transition metals do not have noble-gas electron configuration and are exceptions to the octet
rule.
Formation of Anions
 The gain of negatively charged electrons by neutral atom produces an anion.
 The name of an anion typically ends in –ide.
 Example: Chlorine: 1s2 2s2 2p5 gain one electron
Chloride ion: 1s2 2s2 2p6
Chapter 8 Molecular geometry
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Covalent Bond are atoms held together by sharing electrons.
Molecule is a neutral group of atoms joined together by covalent bonds.
Diatomic molecule is a molecule consisting of two atoms
Molecular compound is a compound composed of molecules.
Molecular compounds tent to have relatively lower melting and boiling points that ionic
compounds
Most are gases or liquids at room temperature, and most molecular compounds are composed of two or
more nonmetals.
Molecular formula is the chemical formula of a molecular compound
A molecular formula shows how many atoms of each element a molecule contains
In covalent bonds, electrons sharing usually occur so that atoms attain the electron configuration
of noble gases.
In covalent bonds elements usually acquire a total of eight electrons (an octet) by sharing electrons.
Single covalent bond is when atoms are held together by sharing a pair of electrons
An electron dot structure such as H:H represents the shared pair of electrons of the covalent bond
by two dots.
Structural formula represents covalent bonds by dashes and shows the arrangement of covalently bonded
atoms.
Unshared pair is a pair of valence electrons that is not shared between the atoms
The oxygen atom has two unshared pair of electrons and two single covalent bonds.
Atoms form double or triple covalent bonds if they can attain a noble gas structure by sharing two
pairs or three pairs of electrons.
Double covalent bond is a bond that involves two shared pairs of electrons
Triple covalent bond is a bond that involves three shared pairs of electrons
Chapter 9 Nomenclature
Chemistry Polyatomic Ions
+1 CHARGE
ion
name
-1 CHARGE
ion
name
-2 CHARGE
ion
name
-3 CHARGE
ion
name
NH4+1 ammonium C2H3O2-1
acetate
CO32-
carbonate
PO33- phosphite
H3O+1 hydronium ClO3-1
chlorate
CrO42-
chromate
PO43- phosphate
ClO2-1
chlorite
Cr2O72- dichromate
CN-1
cyanide
O22-
peroxide
OH-1
hydroxide
SiO32-
silicate
ClO-1
hypochlorite
SO42-
sulfate
NO3-1
nitrate
SO32-
sulfite
NO2-1
nitrite
O2-2
peroxide
ClO4-1
perchlorate
MnO4-1
permanganate