Chemistry 1120 Exam 1 Study Guide: Matter, Measurement, Atomic Structure, and Nuclear Chemistry

Matter, Changes, and Classification

Chemical vs. Physical Changes

Understanding the distinction between chemical and physical changes is fundamental in chemistry. Chemical changes result in the formation of new substances, while physical changes alter the form or appearance without changing the substance's identity.

• Chemical Change: Produces new substances (e.g., rusting of iron).

• Physical Change: No new substances formed (e.g., melting ice).

• Example: Burning wood (chemical); dissolving sugar in water (physical).

Classification of Matter

Matter can be classified as elements, compounds, or mixtures based on composition and combination.

• Element: Pure substance consisting of one type of atom (e.g., O for oxygen).

• Compound: Substance formed by chemical combination of elements in fixed proportions (e.g., H2O).

• Mixture: Physical combination of substances; proportions can vary (e.g., air).

• Example: Salt water is a mixture; sodium chloride is a compound.

Phases of Matter

There are three common phases: solid, liquid, and gas.

• Solid: Definite shape and volume.

• Liquid: Definite volume, shape of container.

• Gas: No definite shape or volume.

Chemical Symbols and Formulas

Chemical symbols represent elements; formulas show the number and types of atoms in compounds.

• Example: Na = sodium; H2O = water.

• Writing Formulas: Use subscripts to indicate atom counts (e.g., CO2).

Measurement, Units, and Calculations

Metric Relationships and Conversions

Accurate measurement and unit conversion are essential for chemical calculations.

• Metric Relationships: 1 g = 1000 mg; 1 kg = 1000 g; 1 kg = 2.2 lb.

• Temperature Conversion: $\text{\degree C} = (\text{\degree F} - 32) \times \frac{5}{9}$

• Density Formula: $\text{Density} = \frac{\text{Mass}}{\text{Volume}}$

• Rearranged: $\text{Mass} = \text{Density} \times \text{Volume}$; $\text{Volume} = \frac{\text{Mass}}{\text{Density}}$

• Medication Calculations: Number of tablets = Ordered dose / Dose per tablet.

• Example: If ordered dose is 500 mg and each tablet is 250 mg, number of tablets = 2.

Significant Figures and Measurement

Counting Significant Figures

Significant figures reflect the precision of a measurement.

• Rules: All nonzero digits are significant; zeros between nonzero digits are significant; trailing zeros in a decimal are significant.

• Example: 52.0 mL has three significant figures.

• Precision: Use the measuring device's precision when recording and reporting results.

• Note: Decimal places are not the same as significant figures.

Atomic Structure and the Periodic Table

Subatomic Particles

Atoms consist of protons, neutrons, and electrons.

• Proton: +1 charge, ~1 amu.

• Neutron: 0 charge, ~1 amu.

• Electron: -1 charge, very small mass.

• Atomic Number: Number of protons; identifies the element.

• Neutral Atom: Number of electrons = number of protons.

• Mass Number: $\text{Mass Number} = \text{Protons} + \text{Neutrons}$

• Neutrons: $\text{Neutrons} = \text{Mass Number} - \text{Atomic Number}$

Isotopes and Notation

Isotopes are atoms of the same element with different numbers of neutrons.

• Isotope Notation: $^{A}_{Z}\text{Element}$, where A = mass number, Z = atomic number.

• Example: $^{14}_{6}\text{C}$ is carbon-14.

Periodic Table Organization

The periodic table is organized into groups (columns) and periods (rows).

• Groups: Vertical columns; elements in the same group have similar properties.

• Periods: Horizontal rows.

• Example: Sodium (Na) and potassium (K) are in the same group.

Electron Configurations and Periodic Trends

Electron Configurations

Electron configuration describes the arrangement of electrons in an atom.

• Notation: 1s, 2s, 2p, 3s, 3p, 4s, etc.

• s Sublevel: Holds 2 electrons; p Sublevel: Holds 6 electrons.

• Example: Oxygen: 1s2 2s2 2p4

• Valence Electrons: Electrons in the outermost shell; determine chemical reactivity.

Periodic Trends: Atomic Size

Atomic size changes across the periodic table.

• Trend: Atomic size decreases from left to right across a period.

• Reason: Increased nuclear charge pulls electrons closer.

Radioactivity and Nuclear Chemistry

Radioisotopes and Decay

Radioisotopes are unstable isotopes that undergo radioactive decay.

• Half-life: Time required for half of a radioactive sample to decay.

• Successive Half-lives: After each half-life, the remaining amount is halved.

• Calculation: Number of half-lives = Total time / Half-life.

• Medical Applications: Iodine-131 (hyperthyroidism, thyroid cancer), Iodine-123 (thyroid imaging), Technetium-99m (diagnostic imaging).

• Short Half-lives: Useful for diagnostics; minimize patient exposure.

Nuclear Particles and Nuclear Equations

Types of Radiation

Nuclear reactions emit different types of radiation.

• Alpha (α): Helium nuclei; decreases mass number by 4, atomic number by 2.

• Beta-minus (β⁻): Electron emission; increases atomic number by 1, mass number unchanged.

• Gamma (γ): Electromagnetic radiation; no change in mass or atomic number.

Balancing Nuclear Equations

Nuclear equations must conserve mass number and atomic number.

• Balanced Equation: Total mass and atomic numbers are equal on both sides.

• Identify Missing Particles: Use conservation laws to find unknowns.

• Example: $^{14}_{6}\text{C} \rightarrow ^{14}_{7}\text{N} + \beta^{-}$

Fission vs. Fusion

Nuclear fission splits nuclei; fusion combines nuclei.

• Fission: Splitting a heavy nucleus (e.g., uranium).

• Fusion: Combining light nuclei (e.g., hydrogen in the Sun).

Atomic Mass and Isotopic Abundance

Weighted Average Atomic Mass

The atomic mass listed on the periodic table is a weighted average of an element's isotopes.

• Calculation: Convert percent abundance to decimal.

• Formula: $\text{Average Atomic Mass} = \sum (\text{Isotope Mass} \times \text{Fractional Abundance})$

• Example: If isotope A is 70% (0.70) and isotope B is 30% (0.30): $\text{Average} = (\text{Mass}_A \times 0.70) + (\text{Mass}_B \times 0.30)$

• Result: Average mass is closer to the more abundant isotope.

Phase at a Given Temperature

Determining Phase

The phase of a substance depends on its melting and boiling points relative to the temperature.

• Below Melting Point: Solid phase.

• Between Melting and Boiling Points: Liquid phase.

• Above Boiling Point: Gas phase.

• Example: Water at 25°C is liquid (melting point 0°C, boiling point 100°C).

High-Priority Skills and Study Checklist

Essential Skills

• Distinguish chemical vs. physical changes

• Classify matter as element, compound, or mixture

• Write chemical symbols and formulas

• Perform metric and temperature conversions

• Calculate density and medication/tablet dosages

• Determine significant figures

• Identify protons, neutrons, electrons, and their properties

• Calculate neutrons from mass and atomic numbers

• Interpret isotope notation

• Use periodic table groups and periods

• Write and interpret electron configurations

• Know atomic size trends

• Identify alpha, beta-minus, and gamma emissions

• Balance nuclear equations

• Calculate remaining amount after multiple half-lives

• Understand medical uses of radioisotopes

• Distinguish fission from fusion

• Calculate weighted-average atomic mass

• Determine phase from melting and boiling points

Summary Table: Types of Matter

Type	Description	Example
Element	Pure substance, one type of atom	Oxygen (O)
Compound	Two or more elements chemically combined	Water (H2O)
Mixture	Physical combination, variable composition	Air

Summary Table: Types of Nuclear Radiation

Type	Symbol	Effect on Nucleus	Example
Alpha	α	Mass number -4, atomic number -2	$^{238}_{92}\text{U} \rightarrow ^{234}_{90}\text{Th} + \alpha$
Beta-minus	β-	Atomic number +1, mass number unchanged	$^{14}_{6}\text{C} \rightarrow ^{14}_{7}\text{N} + \beta^{-}$
Gamma	γ	No change	$^{60}_{27}\text{Co} \rightarrow ^{60}_{27}\text{Co} + \gamma$

Summary Table: Metric Conversions

Conversion	Relationship
Grams to milligrams	1 g = 1000 mg
Kilograms to grams	1 kg = 1000 g
Kilograms to pounds	1 kg = 2.2 lb
Temperature	$\text{\degree C} = (\text{\degree F} - 32) \times \frac{5}{9}$

Additional info: Tables and examples have been expanded for clarity and completeness. All formulas are provided in LaTeX format as required.