Atomic Structure, Matter, and Fundamental Chemical Laws: Study Notes for Introductory Chemistry

Atomic Structure

Introduction to Atomic Structure

Atomic structure is the foundation of chemistry, describing how atoms are composed and how their properties determine the behavior of matter. Understanding atomic structure is essential for grasping chemical reactions, bonding, and the periodic table.

• Atom: The smallest unit of an element, consisting of a nucleus (protons and neutrons) and electrons.

• Atomic number (Z): The number of protons in the nucleus of an atom. (atomic number of the element)

• Mass number (A): The sum of protons and neutrons in the nucleus. (can also say that it is molar mass of the element because the sum of neutron + protons will equal to the mass)

Subatomic Particles

Atoms are made up of three main subatomic particles, each with distinct properties.

Particle	Symbol	Mass (u)	Charge	Location in Atom
Proton	p+	1	+1	Nucleus
Neutron	n	1	0	Nucleus
Electron	e-	1/1837	-1	Outside nucleus

Isotopes

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

• Stable isotopes: Do not undergo radioactive decay (e.g., 12C and 13C for carbon).

• Radioactive isotopes: Undergo radioactive decay (e.g., 14C for carbon).

  • How do I identify stable vs unstable (radioactive) isotopes?

    • For lighter elements that have an atomic number of <20 use the proton/neutron ratio. Stable atoms have ratio closer to 1:1.

    • For heavier elements 20>, stability will require more number of neutrons than protons, so the ratio will increase ex: 1:2.

      • but if an atom has way too few or way too much neutrons in comparison to protons, it is most likely radioactive.

Example: Hydrogen has three isotopes:

• Protium (1H): 1 proton, 0 neutrons (proton/neutron ratio 1/0 = stable because 0 is closer to 1)

• Deuterium (2H): 1 proton, 1 neutron (1:1 stable)

• Tritium (3H): 1 proton, 2 neutrons (1:2 radioactive)

Nuclear Symbol Notation

Nuclear symbols are used to represent isotopes:

• $^A_ZX$ where X = element symbol, A = mass number, Z = atomic number

Example: $^{16}_8O$ (oxygen-16), $^{18}_8O$ (oxygen-18)

Matter and Energy

Definitions and Properties

• Matter: Anything that has mass and occupies space.

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

• Density: The mass per unit volume of a substance. $\text{Density} = \frac{\text{Mass}}{\text{Volume}}$ (move on)

Example: A box of rocks is less dense than a box of leaves of the same size because rocks have more mass per unit volume.

Measurement and Problem Solving

Scientific Notation

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

• Example: 602,200,000,000,000,000,000,000 can be written as $6.022 \times 10^{23}$(move on)

Approaching Calculations & Word Problems

1. Identify what is known.

2. Determine what you want to find.

3. List additional information or conversion factors needed.

4. Check if the answer makes sense.

Scientific Laws

Conservation of Mass

The law of conservation of mass states that during a chemical change, matter is neither created nor destroyed.

• Example: Decomposition of water: $\text{Mass of water} = \text{Mass of oxygen} + \text{Mass of hydrogen}$ If 85.0 g of water decomposes to produce 75.5 g of oxygen, the mass of hydrogen produced is: $85.0\,\text{g} - 75.5\,\text{g} = 9.5\,\text{g}$

Definite Proportions

A compound always contains the same elements in certain definite proportions, regardless of the source.

• Example: Copper carbonate always has the same composition, whether found in minerals or synthesized in a lab.

Atoms and Elements

The Periodic Table

The periodic table organizes elements by increasing atomic number and groups elements with similar chemical properties.

• Atomic number (Z): Number of protons

• Atomic mass: Weighted average mass of isotopes

• Chemical symbol: One- or two-letter abbreviation for each element

Moles and Chemical Composition

The Mole and Avogadro's Number

The mole is a counting unit in chemistry, defined as the number of atoms in 12.011 grams of carbon-12.

• Avogadro's Number: $6.022 \times 10^{23}$ entities per mole

Molar Mass

Molar mass is the mass, in grams, of one mole of a substance.

• Formula: $\text{Molar mass} = \sum \text{atomic masses of elements in compound}$

• Example: Water (H2O): $2 \times 1.008 + 15.999 = 18.015\,\text{g/mol}$

Problem Solving: Mass, Atomic Ratios, and Moles

• To convert between mass, moles, and number of atoms/molecules, use: $\text{Moles} = \frac{\text{Mass}}{\text{Molar Mass}}$ $\text{Number of entities} = \text{Moles} \times 6.022 \times 10^{23}$

• Example: How many moles of CO2 are in 150.0 g? $\text{Moles} = \frac{150.0\,\text{g}}{44.01\,\text{g/mol}} = 3.41\,\text{moles}$

Electricity and the Atom

Electrolysis and Discovery of Electrons

• Electrolysis: Chemical reactions caused by electricity, leading to the discovery of electrons.

• Electron: Small, negatively charged particle in atoms.

Ions

Definition and Types

• Ion: An atom or group of atoms with a charge.

• Anion: Negative ion (gained electrons).

• Cation: Positive ion (lost electrons).

Ions in the Environment

• Ocean acidification: CO2 + H2O → H2CO3 → H+ + HCO3-

• Cation retention in soils: Positive-negative charges attract, positive-positive and negative-negative charges repel.

Radioactivity and Nuclear Chemistry

Radioactivity

Radioactivity is the spontaneous emission of radiation or particles from unstable elements.

• Radioactive isotopes: Experience radioactive decay and are not stable.

• Half-life: Time for half the radioactive atoms to decay into a different isotope.

Types of Radioactivity

Name	Greek Letter	Mass (u)	Charge
Alpha particles	α	4	2+
Beta particles	β	1/1837	1-
Gamma rays	γ	0	0

Radioactive Decay Example

• Carbon-14 decay: Undergoes beta decay (a neutron turns into a proton, loses an electron), forming nitrogen-14.

Summary of Key Concepts

• Conservation of Mass and Definite Proportions are fundamental chemical laws.

• Moles and Molar Mass are essential for quantifying chemical substances.

• Atoms are composed of protons, neutrons, and electrons; isotopes differ in neutron number.

• Ions, isotopes, and radioactivity are crucial for understanding chemical and environmental processes.

Additional info: These notes expand on brief slide points to provide a comprehensive, self-contained study guide for introductory chemistry students.