Study Guide: Matter, Measurement, Atoms, and Covalent Bonding

Matter, Measurement, and Atoms

Pure Substances vs. Mixtures

Understanding the distinction between pure substances and mixtures is fundamental in chemistry.

• Pure Substances: Materials with a fixed composition; can be elements (e.g., O2) or compounds (e.g., H2O).

• Mixtures: Combinations of two or more substances that retain their individual properties; can be homogeneous (solutions) or heterogeneous.

• Example: Salt water is a mixture; table salt (NaCl) is a pure substance.

Periodic Table Regions and Group Names

The periodic table is organized into regions and groups with distinct properties.

• Regions: Metals, nonmetals, metalloids.

• Group Names: Alkali metals (Group 1), alkaline earth metals (Group 2), halogens (Group 17), noble gases (Group 18).

• Example: Sodium (Na) is an alkali metal; chlorine (Cl) is a halogen.

Scientific Notation

Scientific notation is used to express very large or small numbers concisely.

• Format: $a \times 10^n$ where a is a number between 1 and 10, and n is an integer.

• Example: 0.00045 = $4.5 \times 10^{-4}$

Significant Figures

Significant figures reflect the precision of a measured quantity.

• Rules: All nonzero digits are significant; zeros between nonzero digits are significant; leading zeros are not significant; trailing zeros in a decimal are significant.

• Example: 0.0300 has three significant figures.

Mathematical Operations with Significant Figures

When performing calculations, report answers with the correct number of significant figures.

• Addition/Subtraction: Result should have the same decimal places as the least precise measurement.

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

Metric Prefixes and Multiples

Metric prefixes are used to denote multiples of units.

• Kilo (k): $10^3$

• Deci (d): $10^{-1}$

• Centi (c): $10^{-2}$

• Milli (m): $10^{-3}$

• Micro (\mu): $10^{-6}$

• Example: 1 kilometer = 1000 meters

Unit Conversion

Unit conversion is essential for expressing measurements in different units.

• Set-up: Use conversion factors to cancel units.

• Example: To convert 5.0 cm to mm: $5.0 \text{ cm} \times \frac{10 \text{ mm}}{1 \text{ cm}} = 50 \text{ mm}$

Temperature Scales

Three common temperature scales are Fahrenheit (°F), Celsius (°C), and Kelvin (K).

• Conversion Equations:

  • $T_{\text{C}} = \frac{5}{9}(T_{\text{F}} - 32)$

  • $T_{\text{F}} = \frac{9}{5}T_{\text{C}} + 32$

  • $T_{\text{K}} = T_{\text{C}} + 273.15$

• Example: 25°C = $25 + 273.15 = 298.15$ K

Subatomic Particles of the Atom

Atoms are composed of three main subatomic particles.

• Protons: Positive charge, located in nucleus.

• Neutrons: No charge, located in nucleus.

• Electrons: Negative charge, orbiting nucleus.

• Example: Carbon-12 has 6 protons, 6 neutrons, 6 electrons.

Atomic Number and Mass Number

Atomic number and mass number are key identifiers for elements.

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

• Mass Number (A): Total number of protons and neutrons.

• Example: For $^{12}_{6}C$, Z = 6, A = 12.

Isotopes

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

• Definition: Same atomic number, different mass number.

• Representation: $^{A}_{Z}X$ format.

• Example: $^{14}_{6}C$ is an isotope of carbon.

Calculating Subatomic Particles in Isotopes

To determine the number of protons, neutrons, and electrons in an isotope:

• Protons: Equal to atomic number (Z).

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

• Electrons: Equal to protons in a neutral atom.

• Example: $^{17}_{8}O$: 8 protons, 9 neutrons, 8 electrons.

Valence Electrons for Main-Group Elements

Valence electrons are the outermost electrons involved in chemical bonding.

• Definition: Electrons in the outermost shell.

• Determination: For main-group elements, the group number indicates the number of valence electrons.

• Example: Oxygen (Group 16) has 6 valence electrons.

Covalent Bonding and Molecules

Covalent Bonds

Covalent bonds involve the sharing of electrons between atoms.

• Definition: A bond formed when two atoms share one or more pairs of electrons.

• Example: In H2O, each hydrogen shares electrons with oxygen.

Binary Covalent Compounds: Names and Formulas

Binary covalent compounds consist of two nonmetal elements.

• Naming: Use prefixes to indicate the number of atoms (mono-, di-, tri-, tetra-, penta-, etc.).

• Example: CO2 is carbon dioxide; N2O is dinitrogen monoxide.

Lewis Dot Symbols for Main-Group Elements

Lewis dot symbols represent valence electrons as dots around the element symbol.

• Representation: Each dot corresponds to a valence electron.

• Example: For oxygen: O with six dots.

Drawing Lewis Structures

Lewis structures show how atoms share electrons in molecules.

• Steps:

  1. Count total valence electrons.

  2. Arrange atoms and connect with single bonds.

  3. Distribute remaining electrons to satisfy octet rule.

• Example: Lewis structure for water (H2O): Oxygen in center, two hydrogens bonded, two lone pairs on oxygen.

Metric Prefix	Symbol	Multiplier
Kilo	k	$10^3$
Deci	d	$10^{-1}$
Centi	c	$10^{-2}$
Milli	m	$10^{-3}$
Micro	μ	$10^{-6}$

Additional info: Academic context and examples were added to expand brief points into full explanations and to ensure the notes are self-contained for exam preparation.