Atoms and Elements Study Notes

Atoms are the smallest identifiable units of an element.
There are approximately 91 different naturally occurring elements; over 20 synthetic elements have been created.

Law of Conservation of Mass
- Formulated by Antoine Lavoisier.
- States that matter cannot be created or destroyed during a chemical reaction.
- If the masses of all reactants are known, the total mass of products will equal the total mass of reactants.
- Supports the idea that matter comprises tiny, indestructible particles.
Law of Definite Proportions
- Proposed by Joseph Proust in 1797.
- All samples of a given compound contain the same proportion of constituent elements, irrespective of their source or preparation method.
- Example: The decomposition of 18.0 g of water produces 16.0 g of oxygen and 2.0 g of hydrogen, reflecting an oxygen-to-hydrogen mass ratio.
Law of Multiple Proportions
- Established by John Dalton in 1804.
- If two elements form more than one compound, the different masses of one element that combine with a fixed mass of the other can be expressed as small whole-number ratios.
- Example: Atoms of element A combine with varying numbers of atoms of element B to create distinct compounds (e.g., AB1, AB2, AB3).

Dalton’s Atomic Theory consists of the following postulates:
1. Elements are composed of tiny, indestructible particles called atoms.
2. All atoms of a given element are identical in mass and properties, but may have differing masses due to isotopes.
3. Atoms combine in simple, whole-number ratios to form compounds.
4. Atoms of one element cannot transform into atoms of another element during chemical reactions; they rearrange bonds.

All atoms consist of three main subatomic particles:
- Protons: Positively charged, found in the nucleus.
- Neutrons: Neutral charge, also found in the nucleus.
- Electrons: Negatively charged, orbit the nucleus.
The charge of protons and electrons is equal in magnitude but opposite in sign, while neutrons carry no charge.

Atomic Number (Z): The number of protons in an atom's nucleus, defining the element. It is unique for each element and corresponds to their position in the periodic table.
Mass Number (A): The sum of an atom's protons and neutrons, represented as:
$A = p + n$ .

Isotopes are variants of the same element with the same number of protons but different numbers of neutrons.
Natural Abundance: The percentage of each isotope in a naturally occurring sample of an element remains relatively constant.
Isotopes are typically denoted as the element symbol followed by the mass number (e.g., $^{32}S$ , $^{33}S$ ).

Exercises include tables mapping atomic numbers, mass numbers, and symbols, as well as additional conceptual practice.

In 1869, Dmitri Mendeleev created the periodic table by noting periodic properties in elements when arranged by increasing atomic mass. Now, elements are organized by increasing atomic number.

Metals: Good conductors of heat/electricity; malleable and ductile; lose electrons during chemical reactions.
Nonmetals: Poor conductors; tend to gain electrons in reactions; 17 total including solids (C, P, S), liquids (Br), and gases (H, He, N, O, among others).
Metalloids: Exhibit properties of both metals and nonmetals; lie along the dividing line in the periodic table.

Comprised of columns (groups or families) reflecting similar properties and rows (periods) for sequential elements.
The periodic table has 18 groups and 7 periods, with main-group elements having predictable properties.

Noble Gases (Group 8A): Unreactive elements (He, Ne, Ar, Kr, Xe).
Alkali Metals (Group 1A): Highly reactive metals (Li, Na, K, Rb, Cs).
Alkaline Earth Metals (Group 2A): Moderately reactive metals.
Halogens (Group 7A): Reactive nonmetals like fluorine (gas), bromine (liquid), and iodine (solid).

In chemical reactions, atoms can lose or gain electrons to form ions:
- Cations: Positively charged ions (e.g., $Na^+$ ).
- Anions: Negatively charged ions (e.g., $F^-$ ).
Main-group metals generally lose electrons to achieve a noble gas configuration, resulting in cations, while nonmetals tend to gain electrons to form anions.

Exercises related to identifying ions based on atomic structure and symbol representation are provided.

Atomic mass (often referred to as atomic weight) is the weighted average mass of the isotopes of an element based on natural abundance.

Formula for calculating atomic mass:
Atomic ext{ } mass = igg( ext{fraction of isotope 1} imes ext{mass of isotope 1} igg) + igg( ext{fraction of isotope 2} imes ext{mass of isotope 2} igg) + …
Example provided for chlorine with two isotopes, where the atomic mass is calculated using specific fractions and masses of isotopes.

Exercises involve calculating averages based on isotopic compositions and their respective abundances.

The mole (abbreviated mol) is the quantity of substances containing Avogadro’s number of particles, $6.02214 imes 10^{23}$ .
Definition presented: 1 mole of any substance equals $6.02214 imes 10^{23}$ units (e.g., atoms, molecules).
Example: 1 mol of carbon ( $C$ ) = 12 g = $6.022 imes 10^{23}$ carbon atoms.

This conversion mimics simple dozen conversions and utilizes Avogadro’s number as the conversion factor for atoms.

Molar mass represents the mass of 1 mole of an element, which equals the atomic mass in grams/mole.

To find the number of atoms:
1. Determine the mass of the sample.
2. Convert mass to moles using the molar mass.
3. Convert moles to atoms using Avogadro’s number.