Chemistry 1st Trimester Final Exam Review Material
Chapter 2 – Matter and Change
2.1 – Properties of Matter
 Properties used to describe matter can be classified as extensive or intensive.
 Every sample of a given substance has identical intensive properties because every
sample has the same composition.
 Three states of matter are solid, liquid, and gas.
 Physical changes can be classified as reversible or irreversible.
2.2 – Mixtures
 Mixtures can be classified as heterogeneous mixtures or as homogeneous mixtures, based
on the distribution of their components.
 Differences in physical properties can be used to separate mixtures.
2.3 – Elements and Compounds
 Compounds can be broken down into simpler substances by chemical means, but
elements cannot.
 If the composition of a material is fixed, the material is a substance. If the composition
may vary, the material is a mixture.
 Chemists use chemical symbols to represent elements, and chemical formulas to
represent compounds.
2.4 – Chemical Reactions
 During a chemical change, the composition of matter always changes.
 Four possible clues to chemical change include a transfer of energy, a change in color,
the production of a gas, or the formation of a precipitate.
 During any chemical reaction, the mass of the products is always equal to the mass of the
reactants.
Vocabulary:
Chemical change
Chemical Property
Chemical Reaction
Chemical symbol
Compound
Distillation
Element
Extensive Property
Filtration
Gas
Heterogeneous Mixture
Homogeneous Mixture
Intensive Property
Law of Conservation of
Mass
Liquid
Mass
Mixture
Phase
Physical change
Physical Property
Precipitate
Product
Reactant
Solid
Solution
Substance
Vapor
Volume
Chapter 5 – Electrons in Atoms
5.1 – Models of the Atom
 Rutherford’s planetary model couldn’t explain the chemical properties of elements.
 Bohr proposed that electrons move only in specific circular paths, or orbits, around the
nucleus.
Chemistry 1st Trimester Final Exam Review Material
5.2 – Electron Arrangement in Atoms
 Three rules – the aufbau principle, the Pauli-exclusion principle, and Hund’s rule – tell
you how to find the electron configurations of atoms.
Vocabulary:
Atomic Orbital
Aufbau Principle
Electron Configurations
Energy Levels
Hund’s Rule
Pauli exclusion principle
Chapter 6 – The Periodic Table
6.1 – Organizing the Elements
 Chemists used the properties of elements to sort them into groups.
 Menmdeleev arranged the elements in his periodic table in order of increasing atomic
mass.
 In the modern periodic table, elements are arranged in order of increasing atomic number,
The elements within a group in the table have similar properties.
 Three classes of elements are metals, non-metals, and metalloids.
6.2 – Classifying the Elements
 The periodic table displays the symbols and names of elements, along with information
on the structure of their atoms.
 Elements can be sorted into noble gases, representative elements, transition metals, or
inner transition metals based on their electron configurations.
 The periodic table can be divided into s, p, d, and f blocks that correspond to the highest
occupied sublevels in atoms of elements.
6.3 – Periodic Trends
 In general, atomic size increases from top to bottom within a group and decreases from
left to right across a period.
 Positive and negative ions form when electrons are transferred between atoms.
 First ionization energy tends to decrease from top to bottom within a group and increase
from left to right across a period.
 Cations are always smaller than the atoms from which they form. Anions are always
larger than the atoms from which they form.
 In general, electronegativity values decreases from top to bottom within a group. For
representative elements, the values tend to increase from left to right across a period.
 Trends in atomic size, ionization energy, ionic size, and electronegativity can be
explained by variations in atomic structure. The increase in nuclear charge within groups
and across periods explains many trends. Within groups an increase in shielding has a
significant effect.
Vocabulary:
Alkali Metals
Alkaline Earth Metals
Anion
Atomic Radius
Cation
Electronegativity
Halogens
Inner Transition Metal
Ion
Ionization Energy
Metalloids
Metals
Noble Gases
Nonmetals
Periodic Law
Representative Elements
Transition Metal
Chemistry 1st Trimester Final Exam Review Material
Chapter 7 – Ionic and Metallic Bonding
7.1 – Ions
 To find the number of valence electrons in an atom of a representative element, simply
look at its group number.
 Atoms of the metallic elements tend to lose their valence electrons, leaving a complete
octet in the next-lowest energy level. Atoms of some nonmetallic elements tend to gain
electrons to achieve a complete octet.
 An atom’s loss of valence electrons produces a positively charged cation.
 The gain of electrons by a neutral atom produces a negatively charged anion.
7.2 – Ionic Bonds and Ionic Compounds
 Although they are composed of ions, ionic compounds are electrically neutral.
 Most ionic compounds are crystalline solids at room temperature, and they generally
have high melting points. Ionic compounds can conduct an electric current when melted
or dissolved in water.
7.3 – Bonding in Metals
 The valence electrons of metal atoms can be modeled as a sea of electrons.
 Metal atoms are arranged in very compact and orderly patterns.
 Alloys are important because their properties are often superior to those of their
component elements.
Vocabulary:
Alloys
Chemical Formula
Electron Dot Structure
Formula Unit
Halide Ion
Ionic Bonds
Ionic Compounds
Metallic Bonds
Octet Rule
Valence Electron
Chapter 8 – Covalent Bonding
8.1 – Molecular Compounds
 Molecular compounds tend to have relatively low melting and boiling points.
 A molecular formula shows how many atoms of each element a molecule contains.
8.2 – The Nature of Covalent Bonding
 Electron sharing occurs so that atoms attain the configurations of noble gases.
 An electron dot structure shows the shared electrons of a covalent bond by a pair of dots.
 Atoms form double or triple bonds by sharing two or three pairs of electrons.
 In a coordinate covalent bond, the shared electron pair comes from a single atom.
 A large bond dissociation energy corresponds to a strong covalent bond.
 In ozone, the bonding of oxygen atoms is a hybrid of the extremes represented by the
resonance forms.
 The octet rule is not satisfied in molecules with an odd number of electrons, and in
molecules where an atoms has less, or more, than a complete octet of valence electrons.
8.3 – Bonding Theories
 Just as an atomic orbital belongs to a particular atom, a molecular orbital belongs to a
molecule as a whole.
 According to VSEPR theory, the repulsion between electron pairs causes moledular
shapes to adjust so that the valence-electron pairs stay as far apart as possible.
Chemistry 1st Trimester Final Exam Review Material

Orbital hybridization provides information about both molecular bonding and molecular
shape.
8.4 – Polar Bonds and Molecules
 When different atoms bond, the more electronegative atom attracts electrons more
strongly and acquires a slight negative charge.
 Polar molecules between oppositely charged metal plates tend to become oriented with
respect to the positive and negative plates.
 Intermolecular attractions are weaker than either an ionic or covalent bond.
 Melting a network solid requires breaking covalent bonds throughout the solid.
Vocabulary:
Bond Dissociation Energy
Bonding Orbital
Covalent Bond
Coordinate Covalent Bond
Diatomic Molecule
Dipole
Dipole Interactions
Dispersion Forces
Double Covalent Bond
Hybridization
Hydrogen Bonds
Molecular Compound
Molecular Formula
Molecular Orbital
Molecule
Network Solids
Nonpolar Covalent Bond
Pi Bond
Polar Bond
Polar Covalent Bond
Polar Molecule
Polyatomic Ion
Resonance Structure
Sigma Bond
Single Covalent Bond
Structural Formula
Tetrahedral Angle
Triple Covalent Bond
Unshared Pair
VSEPR Theory