Mid-Chem Unit Test Outline Review

Test Structure and Assessment Guidelines

• The Mid-Chem Unit Test is a cumulative term examination emphasizing Chapter 2 (The Periodic Table of Elements) while incorporating foundational science units studied throughout the term.
• Test question formats include:
  • Multiple choice
  • Matching
  • Reading diagrams (periodic table groups, periodic table symbols, Bohr diagrams, and measurement tools)
  • Drawing diagrams (Bohr models of atoms and ions)
  • Short answer questions
  • Lab application: Designing and drawing a data table skeleton based on a provided laboratory procedure.
• Periodic table reference conditions provided during testing:
  • The periodic table includes the stepped "staircase" line separating metals from non-metals.
  • The periodic table will not include a legend distinguishing metals, non-metals, and metalloids.
  • The periodic table will not display printed numbers for groups (columns) or periods (rows); these must be recognized and numbered independently.

Reading Periodic Table Symbols and Calculating Subatomic Particles

• Information displayed within a standard element tile on the periodic table includes:
  • Atomic Number: An integer representing the total number of protons found within the nucleus of an atom. The atomic number defines the identity of the element.
  • Element Symbol: A one- or two-letter abbreviation of the element name, starting with a capital letter (e.g., $\text{C}$, $\text{Na}$, $\text{Cl}$, $\text{Ba}$, $\text{Ar}$).
  • Atomic Mass: The weighted average mass of naturally occurring isotopes of that element, expressed in atomic mass units ($\text{amu}$).
  • Ion Charge (Combining Capacity): A signed superscript indicating the net electrical charge acquired when an atom loses or gains electrons during chemical reactions.
• Determining the quantity of subatomic particles in neutral atoms:
  • Number of Protons: Identical to the atomic number ($Z$):     $\text{Number of Protons} = \text{Atomic Number}$
  • Number of Electrons (in an uncharged neutral atom): Identical to the number of protons:     $\text{Number of Electrons} = \text{Number of Protons} = \text{Atomic Number}$
  • Number of Neutrons: Calculated by rounding the atomic mass to the nearest whole integer to obtain the mass number ($A$), then subtracting the atomic number ($Z$):     $\text{Number of Neutrons} = \text{Mass Number} - \text{Atomic Number}$

Organization of the Periodic Table

• Periods:
  • The horizontal rows on the periodic table, numbered sequentially from $1$ to $7$ from top to bottom.
  • The period number corresponds directly to the total number of electron shells (energy levels) occupied by electrons in an atom of that element.
• Groups (Families):
  • The vertical columns on the periodic table, numbered from $1$ to $18$ from left to right.
  • Elements within the same group share identical valence electron counts and exhibit similar physical and chemical behaviors.
• The Staircase Divider:
  • A distinct zigzag line that begins between boron ($\text{B}$) and aluminum ($\text{Al}$) and extends diagonally downward toward polonium ($\text{Po}$) and astatine ($\text{At}$).
  • It serves as the physical boundary separating metallic elements from non-metallic elements.

Metals, Non-Metals, and Metalloids (Semi-Metals)

• Classification by Location:
  • Metals: Located to the left and across the vast majority of the periodic table (occupying Groups $1$ through $12$, plus the lower left portions of Groups $13$ through $16$), excluding hydrogen.
  • Non-Metals: Clustered in the upper right-hand corner of the periodic table, with the single exception of hydrogen located in Group $1$.
  • Metalloids (Semi-Metals): Positioned directly adjacent to both sides of the staircase line (typically including boron, silicon, germanium, arsenic, antimony, tellurium, and polonium; aluminum is classified as a metal).
• Physical Properties versus Chemical Properties:
  • Physical Property: A characteristic of a substance that can be observed, identified, or measured without changing the chemical identity or composition of the substance.
  • Qualitative Physical Properties: Non-numerical descriptions based on sensory observation (e.g., color, texture, odor, state of matter, malleability, ductility, luster, clarity).
  • Quantitative Physical Properties: Measured physical attributes expressed with numerical values and appropriate units (e.g., density in $\text{g/cm}^3$, melting point in $^\circ\text{C}$, boiling point in $^\circ\text{C}$, mass in $\text{g}$, solubility).
  • Chemical Property: A characteristic describing the ability or inability of a substance to undergo a specific chemical change and interact with other substances to create entirely new materials (e.g., flammability, toxicity, reactivity with water, reactivity with oxygen/corrosion, reactivity with acids).
• General Properties of Metals:
  • Physical: Lustrous (shiny), malleable (can be hammered or rolled into flat sheets without shattering), ductile (can be drawn out into thin wires), excellent thermal and electrical conductors, high melting points, and high densities; all are solids at room temperature except mercury ($\text{Hg}$), which is liquid.
  • Chemical: Readily lose valence electrons to form positively charged ions (cations); react readily with oxygen and acids.
• General Properties of Non-Metals:
  • Physical: Dull in appearance, brittle when in solid form (shatter under pressure), poor thermal and electrical conductors (effective insulators), lower densities and melting points compared to metals; exist in all three states of matter at room temperature (e.g., solid carbon and sulfur, liquid bromine, gaseous oxygen and nitrogen).
  • Chemical: Tend to gain or share electrons during reactions to achieve stable electron configurations, forming negatively charged ions (anions) or covalent bonds.
• General Properties of Metalloids (Semi-Metals):
  • Exhibit intermediate properties that blend characteristics of both metals and non-metals.
  • Solid at room temperature; can appear shiny like metals but fracture brittly like non-metals.
  • Function as electrical semiconductors, meaning their electrical conductivity increases with temperature, making them critical in electronic components.

Chemical Families

• Alkali Metals:
  • Group Number: Group $1$ (excluding hydrogen).
  • Location: The leftmost vertical column of the periodic table.
  • Valence Electrons: Possess exactly $1$ valence electron in the outermost shell.
  • Ion Charge: Form ions with a $+1$ charge by shedding their single valence electron.
  • Main Chemical Properties: Highly reactive, metallic elements; soft enough to be sliced with a butter knife; silver-grey luster; react violently with water to form hydrogen gas ($\text{H}_2$) and alkaline solutions; must be stored submerged under mineral oil to prevent spontaneous reaction with atmospheric oxygen and water vapor.
• Alkaline Earth Metals:
  • Group Number: Group $2$.
  • Location: The second vertical column from the left.
  • Valence Electrons: Possess $2$ valence electrons in the outermost shell.
  • Ion Charge: Form ions with a $+2$ charge by losing both valence electrons.
  • Main Chemical Properties: Highly reactive metals (though slightly less reactive than alkali metals); harder, denser, and have higher melting points than Group $1$ metals; burn with distinct, intensely bright colored flames (frequently utilized in fireworks and flares).
• Halogens:
  • Group Number: Group $17$.
  • Location: The second vertical column from the right.
  • Valence Electrons: Possess $7$ valence electrons in the outermost shell.
  • Ion Charge: Form ions with a $-1$ charge by acquiring one electron to complete their valence shell.
  • Main Chemical Properties: The most reactive non-metallic family; exist as diatomic molecules in their pure elemental form ($\text{F}_2$, $\text{Cl}_2$, $\text{Br}_2$, $\text{I}_2$); highly toxic and corrosive; react vigorously with metals to produce salts (halides).
• Noble Gases:
  • Group Number: Group $18$.
  • Location: The far right vertical column of the periodic table.
  • Valence Electrons: Possess a full valence shell consisting of $8$ valence electrons (with the exception of helium, which possesses a full shell of $2$ electrons).
  • Ion Charge: Possess an ion charge of $0$; do not naturally form ions under standard conditions.
  • Main Chemical Properties: Chemically inert and virtually unreactive; exist as colorless, odorless, non-flammable monatomic gases at room temperature; emit characteristic glowing colors when subjected to high-voltage electric currents.

Atomic Structure and Subatomic Particles

• Characteristics of Subatomic Particles:
  • Proton:
  • Relative Charge: Positive one ($+1$).
  • Location: Bound tightly inside the central atomic nucleus.
  • Relative Mass: Approximately $1\,\text{amu}$.
  • Neutron:
  • Relative Charge: Neutral / zero charge ($0$).
  • Location: Bound tightly inside the central atomic nucleus alongside protons.
  • Relative Mass: Approximately $1\,\text{amu}$.
  • Electron:
  • Relative Charge: Negative one ($-1$).
  • Location: Occupies the electron shells / energy levels in the vast space surrounding the nucleus.
  • Relative Mass: Extremely small, approximately $\frac{1}{1836}\,\text{amu}$ (considered negligible when determining atomic mass).
• Electrical Neutrality of an Atom:
  • An atom is electrically neutral (has an overall net charge of $0$) because the quantity of positively charged protons in its nucleus is exactly equal to the quantity of negatively charged electrons orbiting the nucleus.
  • The positive charge from each proton ($+1$) is canceled out by the equal and opposite negative charge of each electron ($-1$):     $\text{Net Charge} = (\text{Number of Protons} \times (+1)) + (\text{Number of Electrons} \times (-1)) = 0$

Atomic Models: Bohr and Lewis Dot Structures

• Electron Shell Capacity Pattern ($2, 8, 8, 18$):
  • Electrons populate defined energy levels (shells) moving outward from the nucleus.
  • Maximum electron capacity per shell:
  • Shell 1 (innermost shell): $2$ electrons
  • Shell 2: $8$ electrons
  • Shell 3: $8$ electrons (for elements up to calcium, $Z=20$)
  • Shell 4: $18$ electrons
• Bohr Models of Neutral Atoms:
  • Diagrammatic representation featuring a central circle representing the nucleus, inside which the exact counts of protons and neutrons are written (e.g., "$p^+ = 6$", "$n^0 = 6$").
  • Concentric circular rings are drawn around the nucleus to depict each electron shell.
  • Electrons are plotted as dots along these rings, filling inner shells to capacity before populating outer shells. In shells with up to $8$ electrons, dots are placed individually at four positions (top, bottom, left, right) before pairing up.
• Bohr Models of Ions:
  • Constructed identically to the atomic Bohr model, but electrons are either added (for non-metals) or removed (for metals) to achieve a full outermost valence shell.
  • The completed diagram is enclosed within large square brackets ($[\dots]$), and the resulting overall ion charge is written outside the bracket in the upper right corner (e.g., $[\dots]^{2+}$, $[\dots]^{3-}$).
• Lewis Dot Structures (Electron Dot Diagrams):
  • A simplified structural representation displaying only the element's chemical symbol and its valence electrons.
  • The chemical symbol represents the atomic nucleus and all inner (core) electron shells.
  • Dots representing valence electrons are placed around the four sides of the symbol, placed singly before pairing.
• Valence Electrons:
  • Shell Location: Always located in the outermost occupied electron shell of an atom.
  • Periodic Table Pattern: For main-group elements (Groups $1, 2, 13\text{--}18$), the number of valence electrons corresponds to the group number's ones digit:
  • Group $1$: $1$ valence electron
  • Group $2$: $2$ valence electrons
  • Group $13$: $3$ valence electrons
  • Group $14$: $4$ valence electrons
  • Group $15$: $5$ valence electrons
  • Group $16$: $6$ valence electrons
  • Group $17$: $7$ valence electrons
  • Group $18$: $8$ valence electrons (except helium, which has $2$)
  • Importance: Valence electrons govern the atom's chemical properties, reactivity, bonding type, and combining capacity with other elements.
  • Connection to Ion Charge: Atoms lose or gain valence electrons to reach a full stable outer shell; the number of valence electrons lost determines positive charge, while the number gained determines negative charge.

Noble Gas Stability and Ion Formation

• Noble Gas Stability:
  • Noble gases are extremely stable and unreactive because their outermost valence shells are completely filled (a full octet of $8$ electrons, or a stable duet of $2$ electrons for helium).
  • A full valence shell represents a low-energy, highly favorable thermodynamic state, meaning noble gas atoms have no tendency to donate, accept, or share electrons.
• Ion Formation:
  • An ion is an atom or group of atoms that has acquired an electrical charge through the loss or gain of one or more electrons.
  • When an atom loses electrons, it forms a positively charged ion called a cation.
  • When an atom gains electrons, it forms a negatively charged ion called an anion.
• Connection Between Elements and Ion Type:
  • Metals have $1$, $2$, or $3$ valence electrons; they tend to lose these electrons during chemical reactions, forming positive cations (e.g., $\text{Na}^+$, $\text{Mg}^{2+}$, $\text{Al}^{3+}$).
  • Non-metals have $5$, $6$, or $7$ valence electrons; they tend to gain electrons from metals, forming negative anions (e.g., $\text{N}^{3-}$, $\text{O}^{2-}$, $\text{Cl}^-$, $\text{F}^-$).
• Isoelectronic Configuration with Nearest Noble Gas:
  • When atoms form ions, they lose or gain electrons until their valence electron arrangement is identical to that of the nearest noble gas on the periodic table (isoelectronic arrangement).
  • Example: A neutral sodium atom ($11$ protons, $11$ electrons) loses $1$ electron to become $\text{Na}^+$ ($10$ electrons), matching the electron configuration of neon ($10$ electrons).
  • Example: A neutral chlorine atom ($17$ protons, $17$ electrons) gains $1$ electron to become $\text{Cl}^-$ ($18$ electrons), matching the electron configuration of argon ($18$ electrons).
• Calculating the Charge on an Ion:
  • The electrical charge of an ion is calculated as the sum of the positive charges of its protons and the negative charges of its electrons:     $\text{Ion Charge} = (\text{Number of Protons}) - (\text{Number of Electrons})$

Cumulative Review: Lab Safety, WHMIS, and HHPS

• WHMIS (Workplace Hazardous Materials Information System):
  • A standardized Canadian system of safety pictograms, labels, and safety data sheets (SDS) designed to communicate the hazards of industrial and laboratory chemicals.
  • WHMIS Pictograms include:
  • Exploding Bomb: Explosive and self-reactive hazards.
  • Flame: Flammable materials that can ignite easily.
  • Flame Over Circle: Oxidizers that supply oxygen and intensify fires.
  • Gas Cylinder: Gases stored under high pressure.
  • Corrosion: Materials causing severe skin burns, eye damage, or metal corrosion.
  • Skull and Crossbones: Acute toxicity that can cause rapid death or severe poisoning upon short-term exposure.
  • Health Hazard: Chronic toxicity causing serious long-term health effects (carcinogens, respiratory sensitizers, reproductive toxins).
  • Exclamation Mark: Immediate less severe health effects (skin/eye irritation, acute toxicity category 4).
  • Biohazardous Infectious Materials: Biohazards containing infectious organisms (viruses, bacteria, parasites).
  • Environment: Aquatic environmental hazards.
• HHPS (Hazardous Household Products Symbols):
  • Safety warnings applied to household products, composed of two elements: a border indicating the level of danger and a central symbol identifying the specific hazard.
  • Danger Level Frames:
  • Octagon (Stop sign shape): Danger (highest hazard level).
  • Diamond: Warning (moderate hazard level).
  • Inverted Triangle (pointing down): Caution (lowest hazard level).
  • Hazard Content Pictograms:
  • Poisonous: Represented by a skull and crossbones.
  • Flammable: Represented by a fire/flame.
  • Explosive: Represented by an exploding object/burst.
  • Corrosive: Represented by a skeletal hand inside a container.
• Laboratory Safety Regulations:
  • Always wear approved safety goggles and lab aprons.
  • Know the locations and operation of safety equipment: eyewash station, safety shower, fire extinguisher, fire blanket, and broken-glass disposal container.
  • Never taste or directly inhale chemical vapors; utilize the "wafting" technique to detect odors safely.
  • Proper dilution protocol for acids: Always add acid to water slowly while stirring, never water to acid ("do as you oughta, add acid to water").

Cumulative Review: The Scientific Method and Experimental Data Tables

• Stages of the Scientific Method:
  • Question / Problem: Identifying a specific, testable inquiry regarding an observable phenomenon.
  • Hypothesis: A testable, falsifiable predictive explanation for an observation, structured in an "If… then… because…" framework.
  • Experimental Design and Variables:
  • Independent Variable (IV): The variable deliberately manipulated or altered by the experimenter.
  • Dependent Variable (DV): The variable measured or observed to evaluate the effect of changing the independent variable.
  • Controlled Variables (Constants): All other environmental and physical factors kept strictly identical across all experimental trials to ensure a fair test.
  • Data Collection: Systematic recording of qualitative observations and quantitative measurements.
  • Analysis and Conclusion: Evaluating patterns in data to determine whether the results support or refute the initial hypothesis.
• Designing a Laboratory Data Table Skeleton:
  • Must include a clear, descriptive, underlined title identifying the independent and dependent variables being evaluated.
  • Independent variable placed in the left-hand column, with controlled increments.
  • Dependent variable placed across columns to the right, subdivided into individual trial columns (Trial 1, Trial 2, Trial 3) and a calculated average column.
  • Column headers must include the variable name and metric units of measurement (e.g., "Time ($\text{s}$)", "Mass ($\text{g}$)", "Volume ($\text{cm}^3$")).
  • Clean, ruled borders separating rows and columns, leaving interior data cells empty for real-time experimental data entry.

Cumulative Review: Scientific Skills, Tools, and Measurement

• Scientific Notation:
  • A mathematical format for expressing very large or very small numbers:     $a \times 10^b$     where the coefficient $a$ is a real number satisfying $1 \le |a| < 10$, and the exponent $b$ is an integer.
• Metric System Conversions:
  • Utilizes standard base units: meters ($\text{m}$) for length, liters ($\text{L}$) or cubic centimeters ($\text{cm}^3$) for volume, and grams ($\text{g}$) for mass.
  • Common metric prefixes and conversion factors relative to the base unit:
  • Kilo- ($\text{k}$): $10^3 = 1000$
  • Hecto- ($\text{h}$): $10^2 = 100$
  • Deka- ($\text{da}$): $10^1 = 10$
  • Base unit: $10^0 = 1$
  • Deci- ($\text{d}$): $10^{-1} = 0.1$
  • Centi- ($\text{c}$): $10^{-2} = 0.01$
  • Milli- ($\text{m}$): $10^{-3} = 0.001$
• Reading Measurement Tools:
  • Volume (Graduated Cylinder):
  • Liquids form a curved surface called a meniscus.
  • Measurements must be read at eye level from the center bottom of the concave meniscus.
  • Mass (Triple Beam Balance or Electronic Scale):
  • Ensure the balance is zeroed (tared) before placing items on the pan.
  • Read rider indicators starting from the largest mass beam to the smallest.
  • Length (Ruler):
  • Align the object with the zero mark (not the edge of the ruler).
  • Estimation Rule for Analog Instruments:
  • Record all known, calibrated values from the tool's printed scale markings, plus one final estimated digit representing one decimal place beyond the instrument's smallest marked increment.
• Significant Digits (Significant Figures):
  • Rules for determining significant figures in a measurement:
  • All non-zero digits are significant.
  • Captive zeros (zeros between non-zero digits) are significant.
  • Leading zeros (zeros preceding non-zero digits) are not significant; they act as place holders.
  • Trailing zeros in a number containing an explicit decimal point are significant; trailing zeros without a decimal point are ambiguous and generally considered non-significant.

Cumulative Review: States of Matter and Kinetic Molecular Theory

• The Particle Model of Matter:
  • All matter is composed of extremely tiny particles.
  • All particles of a single pure substance are identical; particles of different substances are different.
  • Particles possess empty space between them.
  • Particles are in constant, continuous motion.
  • Attractive forces exist between particles; these forces become stronger as particles move closer together.
• Kinetic Molecular Theory (KMT) (p. 18–20):
  • Temperature is a direct measurement of the average kinetic energy of the particles within a substance.
  • Heat addition increases particle kinetic energy and velocity; cooling decreases particle kinetic energy and velocity.
• Characteristics of the Three States of Matter:
  • Solid:
  • Definite shape and definite volume.
  • Particles packed tightly together in fixed, orderly positions.
  • Particle motion is limited to vibrational motion about fixed points.
  • Strong intermolecular attractive forces.
  • Liquid:
  • Indefinite shape (takes the shape of its container) and definite volume.
  • Particles remain closely packed but have enough kinetic energy to slide, slip, and rotate past one another.
  • Particle motion includes rotational, vibrational, and translational movement.
  • Gas:
  • Indefinite shape and indefinite volume (expands to completely fill any container).
  • Particles separated by vast empty distances relative to their size.
  • Particles possess high kinetic energy, moving in rapid, straight-line, random paths until colliding with each other or the container walls.
  • Intermolecular forces between particles are negligible.
• Changes of State (Phase Transitions):
  • Melting (Liquefaction): Solid transitioning to liquid (endothermic, energy absorbed).
  • Freezing (Solidification): Liquid transitioning to solid (exothermic, energy released).
  • Evaporation / Boiling (Vaporization): Liquid transitioning to gas (endothermic).
  • Condensation: Gas transitioning to liquid (exothermic).
  • Sublimation: Solid transitioning directly to gas without passing through the liquid phase (endothermic).
  • Deposition: Gas transitioning directly to solid without passing through the liquid phase (exothermic).

Cumulative Review: Evolution of Atomic Theory

• John Dalton's Atomic Theory (Billiard Ball Model):
  • Proposed that all matter is made of indestructible, indivisible, solid spherical particles called atoms.
  • Stated that all atoms of a given element are identical in mass and properties, and atoms combine in simple whole-number ratios to form compounds.
• J.J. Thomson's Atomic Model (Plum Pudding / Raisin Bun Model):
  • Discovered the electron using the cathode ray tube experiment, proving that atoms are divisible.
  • Modeled the atom as a sphere of diffuse positive charge embedded with negatively charged electrons throughout, resembling raisins in a pudding.
• Ernest Rutherford's Nuclear Model:
  • Conducted the Gold Foil Experiment (shooting alpha particles at thin gold foil; most passed straight through while a small fraction deflected at extreme angles).
  • Conclusions: The atom is mostly empty space; almost all of its mass and all of its positive charge is concentrated in a tiny, dense, central core called the nucleus.
• Niels Bohr's Model (Planetary Model):
  • Modified Rutherford's model by proposing that electrons do not orbit randomly.
  • Stated that electrons travel only in specific, fixed circular pathways or energy levels (shells) at set distances from the nucleus.
  • Electrons jump between shells by absorbing or emitting discrete packets of energy (photons).
• Current Atomic Model: Quantum Mechanical Model / Cloud Theory Advancement (p. 32–33):
  • Modern atomic theory replaced Bohr's circular orbits with the concept of the "electron cloud" (orbitals).
  • Grounded in the Heisenberg Uncertainty Principle, which states that it is impossible to determine simultaneously both the exact position and velocity of an electron.
  • Electrons do not travel in fixed orbits; instead, they exist in three-dimensional mathematical probability regions around the nucleus known as electron clouds, where there is a high statistical probability of locating an electron.