Atomic Structure, Moles, and Scientific Laws: Introduction to Chemistry Study Notes

How Matter Behaves

Scientific Laws: Conservation of Mass

The law of conservation of mass states that during a chemical change, matter is neither created nor destroyed. This fundamental principle underlies all chemical reactions and ensures that the total mass of reactants equals the total mass of products.

• Definition: In any chemical reaction, the mass of the products equals the mass of the reactants.

• Example: Decomposition of water: If 85.0 g of water is decomposed to produce 75.5 g of oxygen, the remaining mass (9.5 g) must be hydrogen.

Equation:

$\text{Total mass of reactants} = \text{Total mass of products}$

Scientific Laws: Definite Proportions

The law of definite proportions states that a chemical compound always contains the same elements in the same proportions by mass, regardless of the source or method of preparation.

• Definition: The composition of a compound is fixed and does not vary.

• Example: Copper carbonate always contains copper, carbon, and oxygen in the same mass ratio, no matter where it is found.

Atomic Theory

Atomic theory is the scientific theory that matter is composed of discrete units called atoms. This concept is foundational to all of chemistry.

• Atoms: The smallest unit of an element that retains the properties of that element.

• Elements: Substances made of only one type of atom.

Counting Atoms: The Mole and Avogadro's Number

The Mole

The mole is the SI unit for the amount of substance. It allows chemists to count atoms, molecules, or ions by weighing them.

• Definition: One mole contains $6.022 \times 10^{23}$ entities (Avogadro's number).

• Example: 1 mole of carbon-12 atoms has a mass of 12.011 grams and contains $6.022 \times 10^{23}$ atoms.

Molar Mass

The molar mass is the mass, in grams, of one mole of a substance. It is numerically equal to the sum of the atomic masses of the elements in a compound.

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

• Example: The molar mass of water ($\mathrm{H_2O}$) is $2 \times 1.008 + 15.999 = 18.015$ g/mol.

Conversions Using the Mole

To convert between mass, moles, and number of particles, use the following relationships:

• Number of moles: $\text{moles} = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}}$

• Number of particles: $\text{number of particles} = \text{moles} \times 6.022 \times 10^{23}$

The Periodic Table

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

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

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

• Atomic mass: Weighted average mass of the isotopes of an element.

Problem Solving: Mass, Atomic Ratios, and Moles

When solving problems involving mass, moles, and atomic ratios, follow a systematic approach:

• Identify: What is known, what is unknown, and what is needed.

• Set up conversions: Use molar mass and Avogadro's number as conversion factors.

• Check your answer: Does it make sense in the context of the problem?

Example: How many moles of CO2 are in 150.0 g of CO2?

$\text{Moles of CO}_2 = \frac{150.0\ \text{g}}{44.01\ \text{g/mol}} = 3.41\ \text{moles}$

Atomic Structure

Subatomic Particles

Atoms are composed of three main subatomic particles: protons, neutrons, and electrons.

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

Atomic Number and Mass Number

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

• Mass number (A): Total number of protons and neutrons in the nucleus.

• Example: Silicon (Si) has atomic number 14 and mass number 28.

Isotopes

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

• Example: Hydrogen has three isotopes: protium ($^1$H), deuterium ($^2$H), and tritium ($^3$H).

• Carbon isotopes: $^{12}$C (99%), $^{13}$C (1%) are stable; $^{14}$C is radioactive.

Nuclear Symbol Notation

The nuclear symbol is used to represent isotopes:

$^{A}_{Z}X$

• X: Element symbol

• A: Mass number

• Z: Atomic number

Example: $^{16}_{8}O$ (oxygen-16), $^{18}_{8}O$ (oxygen-18)

Radioactivity and Nuclear Chemistry

Radioactivity

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

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

• Half-life: The time required for half of the radioactive atoms in a sample to decay.

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

• Alpha decay: Emission of an alpha particle (2 protons, 2 neutrons).

• Beta decay: A neutron turns into a proton, emitting an electron.

• Gamma decay: Emission of high-energy photons.

Example: Carbon-14 Decay

• Carbon-14 undergoes beta decay: one neutron turns into a proton, and an electron is emitted.

• This transforms carbon-14 into nitrogen-14.

Ions and Their Importance

Definition of Ions

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

• Anion: Negatively charged ion (gained electrons).

• Cation: Positively charged ion (lost electrons).

Applications of Ions

• Ocean acidification: CO2 dissolves in water, forming carbonic acid, which dissociates into hydrogen and bicarbonate ions, increasing acidity.

• Cation retention in soils: Soil particles attract and retain positively charged ions (cations), which is important for nutrient availability in plants.

Summary Table: Subatomic Particles

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

Key Takeaways

• Conservation of mass and definite proportions are foundational scientific laws in chemistry.

• The mole and Avogadro's number allow chemists to count atoms and molecules by weighing them.

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

• Radioactivity involves the spontaneous emission of particles or energy from unstable nuclei.

• Ions play crucial roles in chemical reactions, environmental processes, and biological systems.