GOB Chemistry Study Guide: Matter, Measurement, and Atomic Structure (Chapters 2–4)

Chapter 2: Matter and Measurements

Scientific Notation

Scientific notation is a method for expressing very large or very small numbers in the form a × 10n, where a is a number between 1 and 10, and n is an integer.

• Very large numbers: Move the decimal to the left and use a positive exponent. Example: 5,600,000 = $5.6 \times 10^{6}$

• Very small numbers: Move the decimal to the right and use a negative exponent. Example: 0.00042 = $4.2 \times 10^{-4}$

Significant Figures (Sig Figs)

Significant figures are the digits in a measurement that are known with certainty plus one estimated digit. They reflect the precision of a measurement.

• Counting significant digits:

  • All nonzero digits are significant.

  • Zeros between nonzero digits are significant.

  • Leading zeros are not significant.

  • Trailing zeros after a decimal point are significant.

• Examples:

  • 0.00450 (3 sig figs)

  • 1200 (2 sig figs, unless specified otherwise)

  • 1.230 (4 sig figs)

Mathematical Operations and Significant Figures

• Addition/Subtraction: The answer should have the same number of decimal places as the measurement with the fewest decimal places.

• Multiplication/Division: The answer should have the same number of significant figures as the measurement with the fewest significant figures.

• Example: $2.34 \times 1.2 = 2.8$ (rounded to 2 sig figs)

Conversion Factors and Unit Conversions

Conversion factors are ratios used to express a quantity in different units. They are derived from equivalence statements.

• Example: 1 inch = 2.54 cm, so $\frac{1\ \text{in}}{2.54\ \text{cm}}$ or $\frac{2.54\ \text{cm}}{1\ \text{in}}$ can be used as conversion factors.

• Dimensional analysis: A systematic approach to problem-solving that uses conversion factors to move from one unit to another.

Metric System and Prefixes

The metric system uses prefixes to indicate multiples or fractions of base units.

• Kilo- (k): $10^{3}$

• Deci- (d): $10^{-1}$

• Centi- (c): $10^{-2}$

• Milli- (m): $10^{-3}$

• Micro- (\mu): $10^{-6}$

• Nano- (n): $10^{-9}$

• Mega- (M): $10^{6}$

• Giga- (G): $10^{9}$

Dimensional Analysis in Healthcare (Dosage Calculations)

Dimensional analysis is used to calculate medication dosages, converting between units as needed.

• Example: "You are to give ampicillin 25 mg/kg to a child with a mass of 67 lb. If stock on hand is 250 mg/capsule, how many capsules should be given?"

  • Convert pounds to kilograms: $1\ \text{kg} = 2.205\ \text{lb}$

  • Calculate total dose: $25\ \text{mg/kg} \times \text{mass in kg}$

  • Divide by mg per capsule to find number of capsules.

Density

Density is the mass of a substance per unit volume. It is used as a conversion factor between mass and volume.

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

• Units: g/mL, g/cm3, kg/L, etc.

• Application: Used to identify substances and convert between mass and volume.

Chapter 3: Atoms, Elements, Compounds, and Mixtures

Atoms: The Fundamental Units

An atom is the smallest unit of matter that retains the properties of an element.

• Structure:

  • Nucleus: Contains protons (positive charge) and neutrons (neutral).

  • Electrons: Negatively charged particles moving around the nucleus in energy levels (shells).

• Examples: Hydrogen (H), Carbon (C), Oxygen (O).

Elements

An element is a pure substance made of only one type of atom. Each element is represented on the periodic table.

• Examples: Oxygen gas (O2), Gold (Au).

Molecules

A molecule is formed when two or more atoms bond together chemically.

• Molecules of elements: O2 (oxygen gas), N2 (nitrogen gas).

• Molecules of compounds: H2O (water), CO2 (carbon dioxide).

• Properties: Molecules have characteristics different from the individual atoms.

Compounds

A compound is a substance formed when atoms of two or more different elements combine in fixed proportions.

• Examples: Water (H2O), Carbon dioxide (CO2).

• Key Point: Compounds always have a definite ratio of atoms.

Mixtures

A mixture is a physical combination of substances where each keeps its own properties.

• Homogeneous mixtures (solutions): Uniform composition (e.g., saltwater, air).

• Heterogeneous mixtures: Not uniform (e.g., sand and water, salad).

Summary Table: Compounds vs. Mixtures

Feature	Compound	Mixture
Type of change	Chemical	Physical
Composition	Fixed ratio	Variable ratio
Properties	New, different from elements	Original properties retained
Separation	Chemical methods needed	Physical methods possible
Example	H2O, CO2, NaCl	Air, saltwater, salad

Hierarchy of Building Blocks

• Atom → makes up Element

• Atoms combine → form Molecules/Compounds

• Ions → charged versions of atoms/molecules

• Mixtures → physical blends of substances

Chapter 4: Atomic Structure and Subatomic Particles

What Are Atoms?

Atoms are the smallest units of matter that still retain the properties of an element. They are the basic building blocks of all substances.

• Example: A single atom of oxygen (O) makes up oxygen gas when combined with other oxygen atoms (O2).

Subatomic Particles

Atoms are composed of three main subatomic particles:

• Protons

  • Charge: Positive (+1)

  • Mass: About 1 atomic mass unit (amu)

  • Location: In the nucleus (center of the atom)

  • Role: The number of protons equals the atomic number, which identifies the element. Example: Carbon has 6 protons, so its atomic number is 6.

• Neutrons

  • Charge: Neutral (no charge)

  • Mass: About 1 amu (similar to protons)

  • Location: In the nucleus

  • Role: Help stabilize the nucleus; different numbers of neutrons create isotopes. Example: Carbon-12 vs. Carbon-14.

• Electrons

  • Charge: Negative (−1)

  • Mass: Almost negligible (1/1836 of a proton)

  • Location: Move around the nucleus in energy levels (shells)

  • Role: Involved in chemical bonding; the number of electrons usually equals the number of protons in a neutral atom.

Calculating Subatomic Particles

• Atomic number (Z): Number of protons in the nucleus.

• Mass number (A): Total number of protons and neutrons. $A = Z + \text{number of neutrons}$

• Number of neutrons: $\text{Number of neutrons} = A - Z$

• Number of electrons: In a neutral atom, equals the number of protons.

Isotopes

Isotopes are atoms of the same element with different numbers of neutrons, resulting in different mass numbers.

• Example: Carbon-12 (6 protons, 6 neutrons), Carbon-14 (6 protons, 8 neutrons)

Role of Electrons in Chemical Bonding

• Electrons in the outermost shell (valence electrons) are involved in forming chemical bonds.

• Atoms may share (covalent bond) or transfer (ionic bond) electrons to achieve stable electron configurations.