Chapter 1: Matter and Its Transformations

1.2 Matter

Matter is anything that has mass and occupies space. It can be classified based on its composition and properties.

• Pure Substances: Materials with a fixed composition and distinct properties. Examples include elements (e.g., O2) and compounds (e.g., H2O).

• Mixtures: Physical combinations of two or more substances. They can be separated by physical means.

• Homogeneous Mixtures (Solutions): Mixtures with uniform composition throughout (e.g., saltwater).

• Heterogeneous Mixtures: Mixtures with non-uniform composition (e.g., sand in water).

• Elements: Substances that cannot be broken down into simpler substances by chemical means.

• Compounds: Substances composed of two or more elements chemically combined in fixed proportions.

1.3 Matter and Its Physical Transformation

Matter exists in three primary phases:

• Solid: Definite shape and volume; particles are closely packed.

• Liquid: Definite volume but no definite shape; particles are less tightly packed than in solids.

• Gas: No definite shape or volume; particles are far apart and move freely.

1.4 Matter and Its Chemical Transformation

Chemical transformations involve changes in the composition of matter, resulting in the formation of new substances.

• Chemical Change: A process in which one or more substances are converted into different substances.

1.5 How Science is Done – The Scientific Method

The scientific method is a systematic approach to investigation:

1. Observation

2. Hypothesis formation

3. Experimentation

4. Analysis and conclusion

Chapter 2: Numbers and Calculations in Chemistry

2.1 Precision vs. Accuracy

• Precision: The closeness of repeated measurements to each other.

• Accuracy: The closeness of a measurement to the true or accepted value.

2.2 Uncertainty in Measurement

• All measurements have some degree of uncertainty, often due to instrument limitations.

• Uncertainty should be reported according to the method specified by your instructor.

2.3 Zeros and Significant Figures

• Significant Figures: Digits in a measurement that are known with certainty plus one estimated digit.

• Zeros may or may not be significant depending on their position.

• Counted quantities and conversion factors are considered to have infinite significant figures.

2.4 Scientific Notation

• Used to express very large or very small numbers in the form .

• Example:

2.5 Calculations with Significant Figures and Scientific Notation

• When multiplying or dividing, the result should have as many significant figures as the measurement with the fewest significant figures.

• When adding or subtracting, the result should have as many decimal places as the measurement with the fewest decimal places.

2.7 Density: A Useful Physical Property

• Density (d): The mass per unit volume of a substance.

• Where m is mass and V is volume.

• Units: g/cm3 or g/mL for solids and liquids; g/L for gases.

2.8 Unit Analysis (Dimensional Analysis)

• Method for converting between units using conversion factors.

• Set up calculations so that units cancel appropriately.

2.10 Quantifying Energy

• Energy: The capacity to do work or transfer heat.

• Specific Heat (c): The amount of heat required to raise the temperature of 1 gram of a substance by 1°C.

• Where q is heat (J), c is specific heat (J/g·°C), m is mass (g), and is the change in temperature (°C).

• Calorimeter: A device used to measure the heat absorbed or released during a chemical or physical process.

Chapter 3: Atomic Theory and the Periodic Table

3.1 Dalton’s Atomic Theory

• All matter is composed of indivisible atoms.

• Atoms of the same element are identical; atoms of different elements are different.

• Atoms combine in simple whole-number ratios to form compounds.

• Atoms are rearranged in chemical reactions, not created or destroyed.

Laws Related to Atomic Theory

• Law of Conservation of Matter: Matter is neither created nor destroyed in a chemical reaction.

• Law of Constant Composition: A given compound always contains the same proportion of elements by mass.

Law of Constant Composition Calculations

• Calculate the mass percent of each element in a compound using:

3.3 The Nucleus

• The nucleus is the dense, positively charged center of the atom, containing protons and neutrons.

3.4 The Structure of the Atom

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

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

• Neutrons:

• Protons: Equal to atomic number (Z).

• Electrons: In a neutral atom, equal to the number of protons.

• Atomic Mass: The weighted average mass of the isotopes of an element.

Determining Subatomic Particles

• Given mass number and atomic number, calculate protons, neutrons, and electrons.

• For ions, adjust the number of electrons according to the charge.

3.5 The Law of Mendeleev and the Law of Octets

• Law of Mendeleev: Elements are arranged in order of increasing atomic mass, showing periodic recurrence of properties.

• Law of Octets: Atoms tend to gain, lose, or share electrons to achieve eight electrons in their valence shell.

3.6 The Modern Periodic Table

• Elements are arranged by increasing atomic number.

• Groups (Columns): Labeled IA to VIIIA; elements in a group have similar chemical properties.

• Periods (Rows): Indicate the principal energy level of the valence electrons.

3.7 Ions and First Ionization Energy

• Ionization Energy: The energy required to remove an electron from a gaseous atom.

• Anion: A negatively charged ion (gains electrons).

• Cation: A positively charged ion (loses electrons).

• To determine the number of subatomic particles in ions, adjust the electron count according to the charge.

Chapter 4: Quantum Theory and Atomic Structure

4.2 A New Kind of Physics: Quantum Physics

• Quantum physics describes the behavior of matter and energy at the atomic and subatomic levels.

• Classical physics cannot explain phenomena such as atomic spectra and the stability of atoms.

4.3 The Bohr Theory of Atomic Structure

• Bohr proposed that electrons orbit the nucleus in specific energy levels (shells).

• Electrons can move between energy levels by absorbing or emitting energy.

• Bohr diagrams represent electrons in discrete shells around the nucleus.

• For ions, adjust the number of electrons in the diagram according to the charge.

4.4 Periodicity and Line Spectra Explained

• Periodicity: The recurring trends in properties of elements across periods and groups in the periodic table.

• Valence Shell: The outermost electron shell of an atom; determines chemical reactivity.

• Line spectra arise from electrons transitioning between energy levels, emitting or absorbing specific wavelengths of light.

4.5 Subshells and Electron Configuration

• Electrons occupy subshells (s, p, d, f) within principal energy levels.

• Electron Configuration: The arrangement of electrons in an atom's orbitals.

• Shorthand notation uses the previous noble gas to simplify configuration.

• Example: Sodium (Na): [Ne] 3s1

4.7 Compound Formation and the Octet Rule

• Octet Rule: Atoms tend to form compounds in ways that give them eight valence electrons, achieving a noble gas configuration.

• Predict the formula of a compound formed between a representative metal and a nonmetal by balancing charges.

• Example: Sodium (Na, +1) reacts with chlorine (Cl, -1) to form NaCl.