Midterm Study Guide
Chapter 1 what is Chemistry?
The Student should able to:
o Define Chemistry
o Define Matter
o Explain the differences between pure substances and mixtures and give
examples
o Classify mixtures as homogenous or heterogeneous mixtures
o Explain the difference between elements and compounds
o Discuss the difference between chemical and physical changes
o Discuss the difference between chemical and physical properties
o Explain what is meant by state of matter and give the three most common
states matter.
o List terms associate with states of matter (e.g. melting, freezing,
condensation, sublimation)
o Determine the signs of chemical reaction.
o Discuss the scientific method including, knowing the steps
o Discuss the laws of conservation of energy and mass
Chemistry is the study of matter and the changes it undergoes.
1. What is matter? Anything that occupies space and has mass, i.e., anything
tangible.
a.
Types of Matter:
1
i.
Mixtures – proportions of components vary, separated into
components by physical means.
1.
Homogeneous – the same throughout (solutions and
alloys.)
2.
Heterogeneous- not the same from place to place
ii. Pure Substances – only one type of matter is present, fixed,
unvarying composition.
1. Elements - the simplest pure substances.
a.
A substance composed of only 1 type of “atom”,
smallest possible “piece” of an element.
b. Cannot be divided into simpler substances or
broken down by physical or chemical means.
c. Represented by chemical symbols
d. Organized into the Periodic Table of the Elements.
2. Compounds-Pure substances composed of more than one
kind of element,therefore, more than one kind of atom.
a. Can be broken down into simpler compounds or
elements by chemical means only.
b.
Transformation of Matter
i. Physical changes
1.
do not turn the substance into a different substance.
2. are reversible!
3. often involve change of “state” :solid  liquid  gas.
2
ii. Chemical changes
1.
turn the substance into a different substance (or
substances).
2.
are irreversible!
3.
called “chemical reactions”
iii. Physical properties-is any aspect of an object or substance that can
be measured or perceived without changing its identity.
iv. Chemical Property-a property that can only be observed by
changing the type of substance.
2. The Scientific Method –A way of solving problems or answering questions:
a. Steps:
i.
Make observations, accumulate facts, collect data.
1. Optional: formulate a law to summarize the observations
and data.)
2.
Devise a tentative hypothesis (or a “model”) to try to
explain these observations.
3. Design and run experiments to test the hypothesis.
a. A hypothesis (or model) that has never yet been
disproven becomes a theory.
3. Energy-The ablitlty to do work
a. Work-cause a change or move an object
b. Law of Conservation of Energy-Energy can be neither created or
destroyed in ordinary changes (not nuclear), it can only change form.
3
c. Law of Conservation of Mass-Mass can not be created or destroyed in
ordinary (not nuclear) changes.
i. All the mass can be accounted for
Chapter 2 – The Numerical Side of Chemistry
The Student should able to:
o Discuss the difference between precision and accuracy.
o Indicate whether a given data set is precise, acurate, both, or neither.
o Determine the number of significant figures in a numerical value.
o Convert any number in normal notation to scientific notation and vice versa
o Perform arithmetic operations(addition, subtraction, multiplication and division)
and express the answer with the correct number of significant figures.
o List all of the Greek prefixes used with SI and metric units.
o Perform conversions between units, including conversions within the metric
system, conversions within the English system and conversions between the
English and metric systems.
o Perform conversions between the three temperature scales
o Use the density equation to solve for the density, mass, or volume when the other
two quantities are known
o Use unit analysis to solve quantitative problems
o Perform algebraic manipulations to solve algebraic equations
o Define energy and give the definations of the two common units in which energy
is expressed.
4
o Calculate the amount of heat (in joules or calories), temperature, mass or specific
heat of a substance.
I. Accuracy and Precision
A. Mean (average) = sum of the measurements
number of measurements
B. Accuracy
1. Tells how close to the true value a measurement or the mean of a set of
measurements lies.
2. Percent Error – a quantitative measure of accuracy.
a. % Error = │True value – Experimental value │ x 100
True value
C. Precision
1. Tells how close to each other a set of measurements lie.
II. Significant Figures (Sig. figs.)
A. Definitions
1. Exact (counted ) numbers.
2. Measured numbers
B. Zeroes and significant numbers – learn the rules!
C. Rounding numbers – let’s all be “on the same page.”
III. Scientific Notation – also called exponential notation.
A. Decimal form -- scientific notation. 1098.45 - 1.09845 x 103
B. Scientific notation -- decimal form. 4.55 x 10-4 - 0.000455
5
IV. Doing Math with Sig Figs and Scientific Notation.
A. Multiplying and dividing sig figs.
B. Adding and Subtracting sig figs.
C. Multiplying exponential numbers.
D. Dividing exponential numbers.
V. Units of Measure
A. English System – still widely used in USA
B. Metric System - begun shortly after French Revolution, used today by most
countries in the world.
1.Systeme International d’Unites - 1960 - (SI units).
a. Length – meter (m),
b. Mass – kilogram (kg)
c. Time – second (s)
d. Temperature – Kelvin (K)
e. Amount of substance-Mole (mol)
2. Units commonly used in chemistry – centimeters (cm), millimeters
(mm), grams (g), Celsius (◦C), liter (L), (mL).
3. Prefixes – you need to memorize these!
4. Temperature conversions - no need to memorize these equations!
a. Fahrenheit and Celsius temperature scales.
Know the mp of ice and bp of water on both scales
28 ºC = ? ºF
ºF = (ºC x 9/5) + 32
95.2 ºF = ? ºC ºC = ( ºF - 32)(5/9)
6
b. Kelvin = Celsius + 273
VI. Unit Analysis.
A. Let the units be your guide – methodical, stepwise problem-solving.
B. Conversion Factors
C. Conversions among different units:
1. English ----> English
2. Metric ----> Metric - memorize prefixes.
3. English ----> Metric, Metric ----> English - conversions will be given
on test.
VII. Density
a. units are g/mL, g/cc3, g/cm3, g/L
VIII. Rearranging equations.
VIIII. Quantifying Energy
a. Specific heat
a. amount of heat needed to raise 1 g of a substance by 1ºC.
b. units are given in calories / g ºC or Joules / g ºC.
c. q= m*s*ΔT
7
Chapter 3: The Evolution of Atomic Theory
The Student should able to:
o State the law of conservation of matter in your own words.
o State the law of multiple proportions in your own words
o List Daltons postulates
o Explain why a couple of Dalton’s postulates are not exactly true.
o Draw and label Thomson’s experiement
o Draw a diagram of Thomson’s plum pudding model of the atom.
o Explain what alpha particles are and discuss their role in developing the
Rutherford model of the atom.
o Draw and label Rutherford’s apparatus
o State the meaning and indicate the abbreviations for the atomic number and the
mass number of an element. Indicate how these terms are related to the number
of subatomic particles in an atom of the element.
o Write a full atomic symbol for any isotope.
o Determine the number of protons, electrons and nuetrons in a given isotope
o Calculculate the atomic mass of an element from the masses and relative
percentages of the isotopes of the element.
o Explain the difference between groups and periods
o Know characterisitcs of the families
o Know which elements belong to which family
o Determine the number of valence electrons for a given family
8
o Describe how atomic radii and ionization energy change from left to right across a
period and from top to bottom down a group
o Exaplain what an ion is and describe the differences between anions and cations
1. Dalton’s Atomic Theory
a. Laws leading to the Formulation of the Theory
i. Law of conservation of matter
ii. Law of multiple proportions
b. Daltons five postualtes
2. Development of a model for atomic structure
i. Discovery of the existence of the electron, proton and neutron
ii. Properties of the electron, Proton and Neutron
iii. Thomsson’s experiemnt
1. Know the who, what, where and why
2. Know how to draw and label Thomson’s experiement
3. The Nucleus
a. Rutherfords Experiment
1. Know the who, what, where and why
2. Know how to draw and label Rutherford’s experiement
4. The Structure of the Atom
a. Atomic number and mass number
b. Isotopes
c. Meaning and calculation of atomic mass
9
5. The Periodic Table
a. Groups and Periods
b. Classification of elements
i. Properties of the individual families
c. Periodic Trends
i. Atomic Radius
ii. Ionization Energy
iii. Boiling/Melting point
iv. Ionic Radius
v. Anion and Cation formation
Chapter 4: The Modern Model of the Atom
The Student should able to:
o Explain the meaning of the term quantized.
o Explain the meaning of the term primary quantum number (n).
o Indicate the maximum number of electrons in each shell
o Describe how the valence electrons are related to chemical properties of the
elements.
o Calculate the energy and determine the color of the light emitted when an atom in
an excited state returns to the ground state.
o Use the octet rule to predict the formula of the compound that two elements will
form
o Write the electron configuration for any element. (both ways)
10
o Determine the number of valence electrons in an atom from its electron
configuration.
o Use the electron configuration to determine the atomic number, the element name
and the group number of any element
o Explain what orbitals are and indicate what happens to the size of an orbital as the
value of n increase
o Draw the shapes of the s and p orbitals
o Identify the parts of a wave
o Explain what is meant by wavelength and frequency.
o Know the electromagnetic spectrum
o Solve for wavelength, frequency, speed and energy
I.
Seeing the Light
a. Properties of Light-Speed, wavelength and energy
b. Mathematical relationship between speed, wavelength and Energy
c. Electromagnetic spectrum
II.
The Bohr Theory
a. Electron orbits
b. Principal quantum number
c. Maximum capacity of orbits
III.
Periodicity
a. Valence Electrons
b. Electron configurations
i. Ground state and excited state
11
ii. Long and Short
IV.
Subshells
a. Types of subshells
b. Maximum number of electrons each subshell holds
V.
Compound Formation and the Octet Rule
a. Chemistry of the representative elements
b. Predicting the formula of Ionic compounds
Chapter 5: Chemical Bonding and Nomenclature
The Student should able to:
o Define molecule, and discuss the relationship (if any) of the physical and
chemical properties of molecules to the properties of the atoms from which they
are made
o Explain what covalent bonds are and why they are formed
o Draw a Lewis dot diagram for any covalent molecule, indicating the number of
valence electrons on each atom in the diagram
o Discuss the role of the octet rule in determining when and how many covalent
bonds will form
o Draw Lewis dot diagrams for the atom of any main group element
I.
Molecules-What are they? Why are they
a. Definiation of molecule
b. Differences in properties of compounds and elements
II.
Holding molecules together-The Covalent Bond
a. Defination of a covalent bond
12
b. Valence electron role in bond formation
c. The octet rule
III.
Molecules, Dot Structures and the Octet Rule
a. Determination of valence electrons in neutral atoms
b. Predicting the number of bonds an atom will form
c. Lewis Dot diagrams for elements
i. Drawing dot diagrams of the elements
ii. Paired electrons and unpaired electrons
13